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SCHOLAR Study Guide

SQA CfE Higher ChemistryUnit 1: Chemical Changes andStructure

Authored by:Emma Maclean

Reviewed by:Diane Oldershaw

Previously authored by:Peter Johnson

Brian T McKerchar

Arthur A Sandison

Heriot-Watt University

Edinburgh EH14 4AS, United Kingdom.

First published 2014 by Heriot-Watt University.

This edition published in 2014 by Heriot-Watt University SCHOLAR.

Copyright © 2014 Heriot-Watt University.

Members of the SCHOLAR Forum may reproduce this publication in whole or in part foreducational purposes within their establishment providing that no profit accrues at any stage,Any other use of the materials is governed by the general copyright statement that follows.

All rights reserved. No part of this publication may be reproduced, stored in a retrieval systemor transmitted in any form or by any means, without written permission from the publisher.

Heriot-Watt University accepts no responsibility or liability whatsoever with regard to theinformation contained in this study guide.

Distributed by Heriot-Watt University.

SCHOLAR Study Guide Unit 1: SQA CfE Higher Chemistry

1. SQA CfE Higher Chemistry

ISBN 978-1-909633-20-9

Printed and bound in Great Britain by Graphic and Printing Services, Heriot-Watt University,Edinburgh.

AcknowledgementsThanks are due to the members of Heriot-Watt University's SCHOLAR team who planned andcreated these materials, and to the many colleagues who reviewed the content.

We would like to acknowledge the assistance of the education authorities, colleges, teachersand students who contributed to the SCHOLAR programme and who evaluated these materials.

Grateful acknowledgement is made for permission to use the following material in theSCHOLAR programme:

The Scottish Qualifications Authority for permission to use Past Papers assessments.

The Scottish Government for financial support.

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i

Contents

1 Reaction rates - collision theory 11.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31.3 Rate of reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41.4 Collision theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12

2 Reaction rates - reaction profiles 192.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212.2 Interpreting graphs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212.3 Activation energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 292.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342.5 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342.6 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35

3 Catalysis 393.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 413.2 Introduction to catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . 413.3 Catalyst mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 423.4 Catalysts in industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 433.5 Potential energy diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . 473.6 Catalysts and energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 563.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 633.8 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 643.9 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65

4 The Periodic Table 714.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 734.2 Arrangement of elements in the Periodic Table: Introduction . . . . . . . 734.3 History of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . 744.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . . . . . . . . 794.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 834.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 834.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 84

5 Bonding and structure 875.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 895.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

ii CONTENTS

5.3 Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 945.4 Covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 965.5 Ionic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1065.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1115.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1115.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112

6 Periodic Table trends 1176.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1196.2 Covalent radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1206.3 Ionisation energies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1236.4 Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1286.5 Summary of trends in the Periodic Table . . . . . . . . . . . . . . . . . . 1306.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1306.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1316.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132

7 Bonding continuum and polar covalent bonding 1397.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1417.2 Polar covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1417.3 Predicting bonding type using electronegativity . . . . . . . . . . . . . . 1447.4 Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1467.5 The bonding continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . 1517.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1537.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1537.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 154

8 Intermolecular forces 1598.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1628.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1628.3 London dispersion forces . . . . . . . . . . . . . . . . . . . . . . . . . . 1638.4 Hydrogen bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1688.5 Relating properties to bonding . . . . . . . . . . . . . . . . . . . . . . . 1718.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1748.7 Predicting solubilities from solute and solvent polarities . . . . . . . . . 1768.8 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1808.9 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1818.10 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 182

9 End of unit test 185

Glossary 195

Answers to questions and activities 1981 Reaction rates - collision theory . . . . . . . . . . . . . . . . . . . . . . 1982 Reaction rates - reaction profiles . . . . . . . . . . . . . . . . . . . . . . 2013 Catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2044 The Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2095 Bonding and structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2126 Periodic Table trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2177 Bonding continuum and polar covalent bonding . . . . . . . . . . . . . . 221

© HERIOT-WATT UNIVERSITY

CONTENTS iii

8 Intermolecular forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2249 End of unit test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 228


1

Topic 1

Reaction rates - collision theory

Contents

1.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2

1.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

1.3 Rate of reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

1.4 Collision theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6

1.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

1.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

1.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12

Prerequisite knowledge

Before you begin this topic, you should know:

• reactions can be monitored by measuring changes in concentration, mass andvolume of reactants and products;

• reactions can be monitored and graphs drawn and interpreted in relation to rate;

• the rates of reactions are affected by changes in concentration, particle size andtemperature);

• calculations can be carried out of the average rate of a chemical reaction from agraph of the change in mass or volume against time.

Learning Objectives

After studying this topic, you should be able to:

• state that reaction rates can be controlled by chemists;

• explain that if reaction rates are too low a manufacturing process will not beeconomically viable;

• explain that if reaction rates are too high there is a risk of thermal explosion;

• describe, using collision theory, the effects of concentration, pressure, surfacearea (particle size), temperature and collision geometry on reaction rates;

• calculate the relative rate of a reaction using the formula Rate = 1/t.

2 TOPIC 1. REACTION RATES - COLLISION THEORY

1.1 Prior knowledge

Test your prior knowledge

Q1: What three variables can be altered to change the rate of a reaction?

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Q2: Name two variables you could monitor to follow the course of a reaction.

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Q3: A student obtains the following set of results shown in the graph when carryingout a reaction with marble chips and dilute hydrochloric acid.

What is the average rate, in g s-1, of reaction between 60 and 120 seconds?

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TOPIC 1. REACTION RATES - COLLISION THEORY 3

1.2 Introduction

Chemical reactions occur at different rates depending on the nature of the reaction andthe conditions under which it occurs.

Fireworks are fast reactions Rusting is a slow reaction

Rate is very important to chemists. In industry, as in the lab, it is important tohave knowledge of how to control rates of reaction. If reaction rates are too low amanufacturing process will not be economically viable. However, if reaction rates aretoo high there is a risk of thermal explosion.

How can rate be measured and what is meant by rate anyway?

Think about this:

Imagine yourself having to travel 500 kilometres. There are three methods of travelavailable:

a) walking

b) fast car

c) rocket

In each case think about it and decide on a rough time it would take to complete thejourney, then work out the rate of travel.

The rate of travel in a), b) or c) will be measured in units of distance per time interval,possibly in these units:

a) km day-1

b) km hour-1

c) km second-1

In each case the rate is "per time interval". Rate is measured in units involving thereciprocal of time and if the time taken is low, the rate is high. If the time taken is highthe rate is low.



Rate is proportional 1t

The study of rates of reaction and the factors which influence the rate is known as"Kinetics".

Knowledge of how to control the rate of reaction is very important in the chemicalindustry.

1.3 Rate of reaction

During the course of a chemical reaction, reactants are being converted into products.Measurement of the rate of reaction involves measurement of the change in the amountof a reactant or product in a certain time. The rate of reaction changes as it progresses,being relatively fast at the start and slowing towards the end. What is being measuredis the average rate over the time interval chosen.

Reactions can be followed by measuring changes in concentration, mass and volumeor by using properties of reactants or products which change in step with concentrationsuch as pressure, conductivity, pH or colour intensity. If the change in concentration ismeasured then:

Average rate =change in concentration

change in time

The units of rate would be moles per second, (written as mol s-1). The units will, ofcourse, depend on which variable is being measured. In general:

Average rate =change in variable

change in time

It is worth mentioning at this stage that the unit of concentration of a solution ismeasured as moles per litre. In previous courses this may have been shortened tomol/� or even as mol/dm3 At this level the same unit will always be written as mol �-1.This means that a change in concentration will be measured as mol � -1 s-1.

Measuring rate

Measuring the speed of a car, or even a rocket, can be done directly, using aspeedometer. Measuring the rate of a reaction directly is more difficult since there is nosuch instrument as a reaction rate meter. A change in the amount of reactant orproduct over time can be used.

The online version of this activity shows a simulation of the reaction between zinc andhydrochloric acid. The volume of hydrogen produced is measured using a syringe.



When zinc reacts with hydrochloric acid, hydrogen gas is released. The volume of gasproduced can be recorded at regular time intervals. This figure above shows theapparatus used and the results.

Q4: What volume of hydrogen is released in the first 5 seconds?

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Q5: Calculate the rate of hydrogen production in the first 5 seconds (remember units)

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Q6: Calculate the rate of hydrogen production from 5 to 10 seconds in the reaction

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Q7: Explain the pattern of these two results.

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Q8: Calculate the average rate of reaction to 2 decimal places for the reaction fromthe start to the end of reaction. ( Hint: when has it ended?)

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1.4 Collision theory

All substances are made up of particles called atoms, ions or molecules, and theseparticles are constantly moving. The degree of movement depends upon the state ofthe substance. This is known as the "kinetic model" of matter. In any sample of solution,liquid or gas there is a range of kinetic energies known as an energy distribution.

The collision theory of reactions suggests that, for a chemical reaction to occur,particles must collide.

Simple collision is not enough however, as many collisions do not have sufficient kineticenergy to successfully rearrange the reactants to form new products. In many casesthe way the particles line up, sometimes called the collision geometry or orientation, iswrong and successful collision cannot occur.

Collision theory, based on the kinetic model of matter, provides an explanation for theeffect that various factors have on the rate of chemical reactions in terms of the numberof successful collisions which occur. Collision theory can be stated thus:

• particles must collide to react.

• not all collisions are successful.

• sufficient energy is needed.

• orientation must be correct.

In the picture shown above the single step formation of products in the reaction is asimplification. The reaction itself actually proceeds in a number of stages.

Earlier work in chemistry showed that the rate of a reaction can be influenced byconcentration of reactants, particle size and temperature. How can collision theoryexplain these effects?

A suitable reaction which will be used to study this over the next three activities is thereaction of hydrochloric acid with a sample of calcium carbonate (chalk).



Key Point

The rates of reactions are affected by changes in concentration, particle size andtemperature and the collision theory can be used to explain these effects.

Collisions and concentration

The online version of this activity is an animation of the reaction between hydrochloricacid and calcium carbonate at two different concentrations of acid, followed by questionstaken from the animation.

Look at the pictures showing the result of collisions between two different concentrationsof hydrochloric acid and calcium carbonate, both after 10 seconds of reaction. Thehydrochloric acid is represented as a large sphere and the calcium carbonate as a smallsphere. Products of the reaction are shown as a combination of a small sphere and alarge sphere.Answer the questions.

Q9: How many successful collisions have occurred using the 1mol �-1 acid after 10seconds?

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Q10: What is the rate of reaction with calcium carbonate and 1 mol � -1 acid expressedas successful collisions per second?

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Q11: What is the rate of reaction with calcium carbonate and 2 mol � -1 acid expressedin the same units?

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Q12: Predict what the rate might be if 4 mol �-1 acid were used in the same experiment.

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Q13: Look at the 2 mol �-1 animation again. Compare the successful collisions in thefirst five seconds period with the second period. Is this as expected? Why does it occur?

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Key Point

In many reactions the rate of reaction is directly proportional to the concentrationof reactant, but there is no simple way to predict the relationship in advance.

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Collisions and particle size

The online version of this activity is an animation of the reaction between hydrochloricacid and the same masses of calcium carbonate in two different particle sizes, followedby questions taken from the animation.

Look at the pictures showing the result of collisions between hydrochloric acid and twoparticle sizes of the same mass of calcium carbonate, both after 10 seconds of reaction(The large sample is the picture on the left). Products of the reaction are shown as acombination of a small sphere and a large sphere.Answer the questions.



Q14: Which sample of calcium carbonate has the greater surface area? ( Hint: comparethe number of particles on the outside edges of the solids which would have beenavailable before reaction)

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Q15: How many successful collisions have occurred with the large sample after 10seconds?

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Q16: What is the rate of reaction with the large sample expressed as successfulcollisions per second?

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Q17: What is the rate of reaction with the small sample expressed in the same units?

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Q18: Is the rate constant over the ten second period? Is this as expected? Explain whythis occurs.

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Key Point

In many reactions involving solid reactant the rate of reaction is raised if thesurface area of the reactant is increased. This is the same as saying the particlesize is decreased.

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Collisions and temperature

The online version of this activity is an animation of the reaction between hydrochloricacid and the same masses of calcium carbonate at two different temperatures, T◦C andT + 10◦C followed by questions taken from the animation.

Look at the pictures showing the result of collisions between hydrochloric acid andcalcium carbonate at temperature T◦C, and then at T + 10◦C both after 10 secondsof reaction. Products of the reaction are shown as a combination of a small sphere anda large sphere. Answer the questions.



Q19: How many successful collisions have occurred at T◦C after 10 seconds?

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Q20: What is the rate of reaction at T◦C expressed as successful collisions per second?

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Q21: What is the rate of reaction at T +10◦C expressed in the same units?

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Q22: What happens to the rate when the temperature is increased by 10 ◦C?

a) stays the sameb) halvesc) doublesd) quadruples

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Q23: Would you expect the rate to be constant over the ten second period? Why doesit change?

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Key Point

In many reactions a rise in temperature of 10◦C causes the rate of reaction toapproximately double.

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1.5 Summary

Summary

• Rates of chemical reactions can be controlled by chemists.

• If reaction rates are too low a manufacturing process will not beeconomically viable.

• If reaction rates are too high there is a risk of thermal explosion.

• The rates of reactions are affected by changes in concentration, particlesize and temperature and the collision theory can be used to explain theseeffects.

• The relative rate of a reaction can be calculated using the formula Rate =1/t.

1.6 Resources• Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,

ISBN 978-1444167528



1.7 End of topic test

End of topic 1 test

This end of topic test is available online. If you do not have access to the internet, hereis a paper version.

Q24:

For any chemical, the temperature is a measure of the:

A) average kinetic energy of all the particles.

B) minimum kinetic energy required before reaction occurs.

C) average kinetic energy of the particles which react.

D) activation energy.

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Q25:

Which of the following is not a factor which affects the rate of a reaction?

A) Time taken for reaction to complete

B) Concentration of reactants

C) Collision geometry

D) Kinetic energy of reactants

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Q26:

When marble pieces and hydrochloric acid are reacted, carbon dioxide is evolved. Thecurves below, showing the mass of the reaction vessel, were obtained under differentconditions.



The change in form of the rate curve from P to Q would be obtained by:

A) increasing the concentration of the acid.

B) decreasing the volume of the acid.

C) increasing the particle size of the marble.

D) decreasing the temperature of the reactants.

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Q27: Calculate the average rate of reaction in g s-1 for curve Q over the first 50 secondsof reaction.

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Q28: Calculate the average rate of reaction in g s-1 for curve Q over the 250 secondsshown on the graph.

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Q29:

Magnesium was added to 1.0 mol l-1 sulfuric acid. The volume of hydrogen gas releasedwas plotted as curve C.



Which curve shows the results of a repeat experiment using the same volume of 0.5 moll-1 sulfuric acid?

A) curve A

B) curve B

C) curve D

D) curve E

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Q30:

Which curve shows the results of a repeat experiment using double the volume of 0.5mol l-1 sulfuric acid?

A) curve A

B) curve B

C) curve D

D) curve E

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Q31:

The same reaction was carried out at four different temperatures. The table shows thetimes taken.



Temperature(◦C) 20 30 40 50

Time (s) 60 30 14 5

These results show that:

A) a small rise in temperature results in a large increase in reaction rate.

B) the activation energy increases with increasing temperature.

C) the rate of reaction is directly proportional to the temperature.

D) the reaction is endothermic.

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Q32:

A reaction was carried out under different conditions. The curves shown were obtainedwhen copper carbonate was reacted with acid, carbon dioxide being produced.

The change from P to Q could be brought about by:

A) decreasing the mass of copper carbonate.

B) increasing the particle size of the copper carbonate.

C) decreasing the particle size of the copper carbonate.

D) increasing the concentration of the acid.

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Q33:

Using the apparatus shown, a 6.075 g block of magnesium was added to 100 cm 3 of 1.0mol �-1 hydrochloric acid and the change in mass noted at regular time intervals.

Which of the following graphs would be drawn from the results?

A)



B)

C)

D)

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Q34: Calculate the number of moles of magnesium added to the acid.

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Q35: Calculate the number of moles of hydrochloric acid present at the start of thereaction.



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Q36: A vitamin C solution which is acidic slowly breaks down any sugar present.The sugar concentration measured over a number of hours was found to bedecomposing at an average rate of 0.0002 mol l -1 min-1.If the starting concentration in a fresh solution was found to be 0.5 mol l-1, what wouldbe its sugar concentration after 5 hours?

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19

Topic 2

Reaction rates - reaction profiles

Contents

2.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21

2.2 Interpreting graphs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21

2.3 Activation energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29

2.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34

2.5 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34

2.6 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35



• reactions can be monitored and graphs plotted and interpreted (National 4, Unit1);

• when chemical reactions occur, they are often accompanied by a change in energy(National 4, Unit 1);

• reactions can be exothermic or endothermic. This is dependent on the overallenergy change taking place (National 4, Unit 1);

• there are four factors that affect the rate of reaction: temperature, concentration,catalyst and surface area (National 3, Unit 1);

• describe using collision theory effects of concentration, pressure, surface area(particle size), temperature and collision geometry on reaction rates (Higher, Unit1, Topic 1).

Learning Objectives

After studying this topic, you should know:

• how to interpret rate graphs;

• that a potential energy diagram can be used to show the energy pathway for areaction;

• enthalpy change is the energy difference between products and reactants. It canbe calculated from a potential energy diagram;

20 TOPIC 2. REACTION RATES - REACTION PROFILES

• the enthalpy change has a negative value for exothermic reactions and a positivevalue for endothermic reactions;

• the activated complex is an unstable arrangement of atoms formed during areaction, at the maximum of the potential energy barrier;

• the activation energy is the energy required by colliding particles to form anactivated complex;

• the activation energy can be calculated from potential energy diagrams;

• temperature is a measure of the average kinetic energy of the particles of asubstance (revision; Higher, Unit 1, Topic 1);

• the activation energy is the minimum kinetic energy required by colliding particlesbefore reaction may occur;

• energy distribution diagrams can be used to explain the effect of changingtemperature on the kinetic energy of particles;

• the effect of temperature on reaction rate can be explained in terms of an increasein the number of particles with energy greater than the activation energy.


TOPIC 2. REACTION RATES - REACTION PROFILES 21

2.1 Prior knowledge


Q1: During the course of a combustion reaction, energy in the form of heat is givenout. What type of reaction is this?

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Q2: During the course of a decomposition reaction, energy in the form of heat is takenin. What type of reaction is this?

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Q3: Collision theory provides an explanation for the effect that various factors have onthe rate of chemical reactions. For successful collisions to occur, what must happen?(Name two things.)

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2.2 Interpreting graphs

Investigations which measure the change in a property, such as concentration, duringthe progress of a chemical reaction often have results displayed as a graph or a table. Inthe reaction between zinc and hydrochloric acid the total volume of hydrogen given offwas measured against time and the graph which resulted showed a steep slope whichtailed off to a level gradient as the reaction finished and the volume of hydrogen releasedstayed constant. Revisit the experiment above if you are unsure about this point.

It would also be possible to follow the progress of the same chemical reaction betweenzinc and hydrochloric acid by measuring the "mass of beaker and contents" againsttime using the apparatus shown in Figure 2.1. The graph which results is shown inFigure 2.2.

Figure 2.1: Zinc and hydrochloric acid mass experiment

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Graphs which use the same axes to plot results for change in concentration of bothreactant and product on the same graph may also be met. Figure 2.2 shows such agraph.

Figure 2.2: Reactant and product on shared axes

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Q4: Name the chemical which escapes from the apparatus in Figure 2.1.

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Q5: What is the mass loss in grams in Figure 2.1 during the first 4 minutes? (To onedecimal place.)

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Q6: Calculate the rate of fall in mass in g min-1 (Figure 2.1).

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Q7: Which of these units could also have been used to describe the rate of fall in massin Figure 2.1?

a) g-1 minb) g-1 min-1 min-1

c) s min-1

d) g s-1

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Q8: How long does it take before the reaction in Figure 2.2 reaches constantconcentrations of reactant and product?

a) 0.6 sb) 7 sc) 20 sd) 40 s

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Q9: Which of these statements is true in Figure 2.2?

a) After 40 seconds there is more reactant than product.b) The rate of production of product over 30 seconds is greater than the rate of fall of

reactant.c) The rate of production of product over 30 seconds is less than the rate of fall of

reactant.d) The rate of production of product over 30 seconds is the same as the rate of fall of

reactant.

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Graphs which use the same axes and place the results for different experiments inwhich the concentration or temperature are varied from one experiment to the next arecommon and show up how that variable affects the reaction progress.



Measuring rate at different concentrations

The online version of this activity is a simulation of a chemical reaction repeated atdifferent concentrations of acid. The data collected, involving changes of volume, isgraphed against time giving multiple graphs on the same axes.

When all three concentrations have been plotted, the graph shown below can be usedto answer the following questions.

The three features of each line which give information about the reaction, and arecommonly met in examination questions are:

• the steepness of the slope;

• the point at which the line becomes horizontal;



• the final level of the horizontal line.



Q10: Which graph shows the steepest slope?

a) 0.5 mol �-1

b) 1.0 mol �-1

c) 2.0 mol �-1

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Q11: Which graph shows the fastest rate of reaction?

a) 0.5 mol �-1

b) 1.0 mol �-1

c) 2.0 mol �-1

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Q12: How long did it take (in seconds) for the 2.0 mol �-1 reaction to finish?

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Q13: What volume of hydrogen (in cm3) in the 2.0 mol �-1 reaction after 30 seconds?

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Q14: Compare the 1.0 mol �-1 graph with the 2.0 mol �-1 graph in three areas:

• the steepness of the slope;

• the point at which the line becomes horizontal;

• the final level of the horizontal line.

Explain the differences.

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Q15: Predict the final level (in cm3 to 1 decimal place) of the horizontal line in the 0.5mol �-1 graph.

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In each of the graphs showing how a variable affects the reaction progress it is theslope of the graph which gives an indication of the average rate of reaction. The rate ofreaction changes as it progresses, being relatively fast at the start and slowing towardsthe end.

Average rate is measured as:

Average rate =change in variable

change in time

In the graphs developed above a comparison of rate of reaction of 0.5 mol � -1, 1.0 mol�-1 and 2.0 mol �-1 acid can be determined by looking at how much hydrogen has beengiven off in the same time span.



If the volume of hydrogen given off after 5 seconds is measured, the question "howdoes the rate of reaction depend on the concentration of acid?" can be answered.

The table of results (Table 2.1) taken from the graphs is shown below. The resultinggraph of rate of reaction is plotted against concentration of hydrochloric acid(Figure 2.3). It is clearly a straight line showing that, in this case, the rate is directlyproportional to the concentration of hydrochloric acid. (Remember however, that thereis no simple way to predict the relationship in advance and each concentration/raterelationship must be investigated experimentally to determine any relationship.)

Table 2.1: Rate versus concentration of hydrochloric acid

hydrochloric acidconcentration/mol �-1

volume ofhydrogen/cm3

rate/(cm3 of H2 s-1)

0.5 6.5 1.3

1.0 13.0 2.6

2.0 26.0 5.2

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Figure 2.3

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Q16: Predict the volume of hydrogen in cm3 given off in 5 seconds in a similarexperiment using 1.5 mol �-1 hydrochloric acid.

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Q17: Predict the average rate in cm3 s-1 in a similar experiment using 4.0 mol �-1

hydrochloric acid.



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Many investigations into how concentration of solutions or temperature changes affectthe rate of a reaction involve "clock reactions". The time taken for a colour change toappear may be an experiment you have seen. Sometimes a cross drawn on paper isplaced below the reaction vessel. The time taken for it be obscured is called the endpoint. In each case the time taken to reach the same stage in the reaction is beingmeasured.

The time is converted into rate by calculating the reciprocal of time:

Rate =1t

Conversions between time, rate and concentration or temperature are common.

Problem

The graph shows a rate of reaction plotted against concentration of a reactant.

Figure 2.4: Rate versus concentration graph

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a) How long would it take to complete the reaction if the reactant concentration is 0.4mol �-1?

b) At what concentration was the reaction completed in 10 seconds?



Solution

a) From the graph, at 0.4 mol �-1 the rate = 0.2 s-1.

Rate =1

t

0.2 =1

t

t =1

0.2t = 5 seconds

b) Since the rate is given by:

Rate =1

t

Rate =1

10Rate = 0.1

From the graph this concentration = 0.2 mol l - 1

2.3 Activation energy

Collision theory states that colliding particles must have "sufficient energy" to react.

Some reactions start as soon as the reactants are mixed. A neutralisation reactionbetween hydrochloric acid and sodium hydroxide solutions would react immediately onmixing. The colliding particles, in this case, must have sufficient energy to react.

Some reactions require an input of energy to start. For example, a firework burns wellin air only after lighting. The particles in the firework had insufficient energy to react atroom temperature.

The activation energy is the minimum kinetic energy required by colliding particlesbefore reaction will occur. This is often thought of as an "energy barrier" which has tobe overcome. The kinetic model of matter shows the energy distribution of particles ina gas, liquid or solution. The average kinetic energy of the particles is measured as"temperature". Only those molecules with a kinetic energy greater than the activationenergy (shown as EA or as Ea) are capable of successful collision.



Figure 2.5: Kinetic energy distribution graph

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These questions refer to the four points on the distribution graph.

Q18: Which point represents particles with the lowest kinetic energy?

a) Ab) Bc) Cd) D

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Q19: Which point represents the least number of particles?

a) Ab) Bc) Cd) D

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Q20: The average kinetic energy of particles determines the "temperature". Which pointrepresents particles nearest to the average kinetic energy.

a) Ab) Bc) Cd) D

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Q21: Which point represents particles with sufficient kinetic energy to reactsuccessfully?

a) Ab) Bc) Cd) D

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The effect of temperature on kinetic energy

Temperature is a measure of the average kinetic energy of the particles of a substance.An increase in temperature changes the energy distribution and increases the averagekinetic energy.

Controlling the rate - temperature & kinetic energy

The online version of this activity is a simulation followed by questions which investigatetemperature as a measure of the average kinetic energy of the particles of a substance.Energy distribution diagrams are described and these are used to explain how anincrease in temperature increases the number of particles with energy greater than theactivation energy

These graphs show the kinetic energy distribution at T◦C and T +10◦C of a group ofparticles involved in a chemical reaction.Compare the graphs and consider the effect of increasing the temperature by 10◦C.



Q22: How has the increased temperature affected the value of the activation energy?

a) Increasedb) Decreasedc) Not changed

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Q23: How has the increased temperature affected the average kinetic energy of theparticles.

a) Increasedb) Decreasedc) Not changed

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Q24: At the increased temperature, which point or points represent particles withsufficient kinetic energy to react successfully?

a) Db) C and Dc) B,C and Dd) A,B,C and D

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Q25: What does the area to the right of EA in the right hand graph represent?

a) Particles with sufficient energy to react.b) Particles with insufficient energy to react.c) Particles with the average kinetic energy.d) Particles with the same kinetic energy.

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Q26: How has the 10◦C increase in temperature affected the number of particles withenergy greater than the activation energy?

a) Roughly the same.b) Roughly halved.c) Roughly doubled.

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Remember, the Chemical Industry is not in existence to manufacture chemicals: like anyother industry it exists to create wealth and wealth can only be created if it can makeprofits.

For this reason, an understanding of rate is very important to chemists. In industry, asin the lab, it is important to have knowledge of how to control rates of reaction.

If reaction rates are too low a manufacturing process will not be economically viable.However, if reaction rates are too high there is a risk of thermal explosion.

In recent years, a number of thermal explosions at chemical plants have made the news;including West Thurrock, Essex and The Imperial Sugar Company, USA.



Reactions increase their rate at higher temperatures because a higher proportion ofthe molecules involved have energy in excess of the activation energy and moresuccessful collisions can occur. It is observed that a 10◦C rise is responsible for anapproximate doubling of rate in many reactions.

Energy distribution diagrams can be used to explain how an increase in temperatureincreases the kinetic energy of particles and therefore increases the number of particleswith energy greater than the activation energy.

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The effect of light energy on kinetic energy

With some chemical reactions light energy is absorbed and can provide particles withenergy sufficient to raise their kinetic energy above the value needed to cross theactivation energy barrier. Photosynthesis and photography provide two commonexamples of photochemical reactions.

(This subject will be considered further in a later topic on skin care products.)

There are many examples of this type of reaction taking place in the atmosphere. Highenergy ultraviolet radiation which would be harmful to humans is absorbed by ozone(O3) molecules in the upper atmosphere.

The ozone layer which protects us is fragile however, as the ozone can be broken downby certain chemicals. Moves have taken place to reduce the use of these chemicalsand protect the ozone layer.

Key Point

• Temperature is a measure of the average kinetic energy of the particles ofa substance and activation energy is the minimum kinetic energy requiredby colliding particles before reaction can occur.

• Energy distribution diagrams can be used to explain how an increase intemperature or, in some chemical reactions the energy from light, increasesthe number of particles with energy greater than the activation energy.



2.4 Summary

Summary

• Graphs which use the same axes and place the results for differentexperiments in which the concentration or temperature are varied from oneexperiment to the next are common and show up how that variable affectsthe reaction progress.

• Temperature is a measure of the average kinetic energy of the particles ofa substance.

• Activation energy is the minimum kinetic energy required by collidingparticles before reaction can occur.

• Energy distribution diagrams can be used to explain how an increase intemperature or, in some chemical reactions the energy from light, increasesthe number of particles with energy greater than the activation energy(Ea/EA).

• Reactions increase their rate at higher temperatures because a higherproportion of the molecules involved have energy in excess of the activationenergy and more successful collisions can occur.

• It is observed that a 10◦C rise is responsible for an approximate doubling ofrate in many reactions.

• The effect of temperature on reaction rate can be explained in terms of anincrease in the number of particles with energy greater than the activationenergy.


ISBN 978-1444167528




End of topic 2 test


Q27: A piece of phosphorus ignites when touched with a hot wire, whereas magnesiumribbon needs strong heating before it will burn. What can you deduce about theactivation energies of the two reactions?

a) The activation energy if less for magnesium than for phosphorus.b) The activation energy if less for phosphorus than for magnesium.c) The activation energies are the same for the reactions.d) No information about activation energy can be deduced.

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Q28: a) A student carries out a reaction and then repeats the experiment at thetemperature ten degrees higher. What is the likely effect on rate of reaction?

a) No effectb) The reaction rate halvesc) The reaction rate doublesd) The reaction rate triples

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Q29: b) Why is this?

a) The activation energy has been decreased.b) The activation energy has increased.c) Fewer particles have energy greater than the activation energy.d) More particles have energy greater than the activation energy.

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Q30: Activation energy is:

a) the minimum kinetic energy required by colliding molecules for a reaction to occur.b) the maximum kinetic energy colliding molecules can reach for a reaction to occur.c) the temperature at which a reaction will take place.d) the temperature at which rate of reaction will be greatest.

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Q31: The activated complex is:

a) the initial stage of reaction when the reactants are first put together.b) an intermediate stage at the top of the activation energy barrier.c) the maximum kinetic energy colliding molecules can reach for a reaction to occur.d) the final stage of a reaction when the products have been formed.

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Q32: The activation energy can be calculated using:

a) time taken for reaction to complete.b) temperature the reaction occurs at.c) potential energy diagrams.d) the activated complex.

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Q33: The graph below shows how the concentrations of both a reactant and a productchange during the course of a reaction.

The average rate of reaction over the first 15 seconds is:

a) 0.0025 mol �-1 s-1

b) 0.000 mol �-1 s-1

c) 0.025 mol �-1 s-1

d) 0.025 s-1 mol �-1

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Q34: View how the correct answer to the above question is calculated.

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Q35: This graph shows the effect of temperature on the rate of a reaction.

How long, in seconds, did the reaction take to complete at 40◦C?

a) 35.00 secondsb) 0.066 secondsc) 47.00 secondsd) 31.25 seconds

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Q36: View how the correct answer to the above question is calculated.

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Q37: Explain in terms of the collision theory why increasing the temperature increasesthe reaction rate.

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Q38: A reaction between A and B forms product C. At a temperature of 40 ◦C thereaction is complete after 20 seconds.At what temperature would the rate be expected to be four times as fast?

A) 50 ◦C

B) 60 ◦C

C) 80 ◦C

D) 160 ◦C



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Q39: A reaction was carried out at four different temperatures. The table below showsthe time taken for the reaction to occur.

Temp (◦C) 20 30 40 50Time (seconds) 60 30 14 5

The results show that:

a) the activation energy increases with increasing temperature.b) a small rise in temperature results in a large increase in reaction rate.c) the rate of the reaction is directly proportional to the temperature.d) the reaction is endothermic.

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39

Topic 3

Catalysis

Contents

3.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

3.2 Introduction to catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

3.3 Catalyst mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42

3.4 Catalysts in industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43

3.4.1 Examples of catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44

3.5 Potential energy diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47

3.5.1 What produces this energy change? . . . . . . . . . . . . . . . . . . . . 51

3.5.2 The thermochemical equation . . . . . . . . . . . . . . . . . . . . . . . . 53

3.6 Catalysts and energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 56

3.6.1 The activated complex . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57

3.6.2 How catalysts work . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 60

3.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63

3.8 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 64

3.9 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65



• there are four factors that affect the rate of reaction: temperature, concentration,catalyst and surface area (National 3, Unit 1);

• adding a catalyst will speed up a reaction (National 3, Unit 1);

• reactions can be exothermic or endothermic. This is dependent on the overallenergy change taking place (National 4, Unit 1);

• using collision theory the effects of concentration, pressure, surface area (particlesize), temperature and collision geometry on reaction rates can be understood(Higher, Unit 1, Topic 1);

• a potential energy diagram can be used to show the energy pathway for a reaction(Higher, Unit 1, Topic 1);

• enthalpy change is the energy difference between products and reactants. It canbe calculated from a potential energy diagram (Higher, Unit 1, Topic 1);

40 TOPIC 3. CATALYSIS

• the enthalpy change has a negative value for exothermic reactions and a positivevalue for endothermic reactions (Higher, Unit 1, Topic 1);

• the activation energy can be calculated from potential energy diagrams (Higher,Unit 1, Topic 1);

Learning Objectives

At the end of this topic, you should know or be able to:

• the enthalpy change for a reaction has a negative value for exothermic reactionsand a positive value for endothermic reactions;

• the activated complex is an unstable arrangement of atoms formed at themaximum of the potential energy barrier, during a reaction;

• the activation energy is the energy required by colliding particles to form anactivated complex;

• the activation energy can be calculated from potential energy diagrams;

• a catalyst provides an alternative reaction pathway with a lower activation energy;

• a potential energy diagram can be used to show the effect of a catalyst onactivation energy;

• explain how catalysis works and learn about adsorption;

• define the terms activated complex and activation energy and hence explain howcatalysts work in terms of potential energy diagrams.


TOPIC 3. CATALYSIS 41

3.1 Prior knowledge


Q1: For any chemical, the temperature is a measure of the:

a) average kinetic energy of all the particles;b) minimum kinetic energy required before reaction occurs;c) average kinetic energy of the particles which react;d) activation energy.

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Q2: During the course of a decomposition reaction, energy in the form of heat is takenin. What type of reaction is this?

a) Exothermicb) Endothermic

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Q3: What is activation energy?

a) The minimum kinetic energy required by colliding molecules for a reaction to occur.b) The maximum kinetic energy colliding molecules can reach for a reaction to occur.c) The temperature at which a reaction will take place.d) The temperature at which rate of reaction will be greatest.

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3.2 Introduction to catalysis

Catalysts are amongst the most important chemicals in the world around us.

Catalysts are essential in:

• the production of materials such as plastics;

• the production of fuels to power our transport;

• the removal of pollutants produced by transport;

• the production of medicines and food;

• the functioning and health of our bodies.

It is said that 90% of all manufactured items use catalysts at some stage of theirproduction.

A catalysts is a substance which alters the rate of a reaction, but is chemicallyunchanged at the end of the reaction. They can be recovered at the end of the reactionchemically unchanged. They are neither reactants nor products and do not appear inthe chemical equation. Often, the catalyst used in a reaction is shown above the arrowin the equation, as in the Haber Process.



Figure 3.1

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Catalysts allow chemical reactions to occur more quickly at lower temperatures and soreduce energy costs.

3.3 Catalyst mechanism

Hydrogenation of ethene

Often the catalyst is in a different state from the reactants. An example of this is theaddition of hydrogen gas to ethene gas using a solid nickel catalyst.

Addition of hydrogen to ethene

The atoms within the metal are fully bonded to their neighbours whereas the atoms onthe surface have 'spare' bonds and are capable of forming weak bonds with the reactantmolecules. This is known as adsorption. (This should not be confused with absorptionwhich involves a substance penetrating inside a solid - like water being soaked up by asponge.) When the molecules leave the surface of the catalyst, the process is known adesorption.

Q4: Reactant molecules bond to the surface of the metal. This is known as;

a) Absorptionb) Desorptionc) Adsorptiond) Activation

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Q5: What happens to the bonds within the molecules when they attach to the surface?

a) They break.b) They get weaker.c) They get stronger.d) They stay the same.

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Q6: When the product molecules leave the surface of the catalyst, what change hastaken place to the surface of the catalyst?

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Q7: Which of these statements is false?

a) The catalyst speeds up the reaction by weakening the bonds in the reactantmolecules.

b) The larger the surface area of the catalyst the more effective it will be.c) The catalyst speeds up the reaction without taking part.d) The catalyst speeds up the reaction without being used up.

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Enzymes

Enzymes are biological catalysts that are responsible for the chemical processes thatoccur in all living organisms, from the simplest bacteria to the most complex mammals.They are proteins and are the target for many modern pharmaceutical products used tocontrol diseases.

Enzymes are studied in more detail in the topic on Proteins later in the course.

Key Point

Many catalysts work by adsorbing reactant molecules onto their surface.

3.4 Catalysts in industry

Many important industrial processes involve catalysis.

Figure 3.2: Examples of industrial catalysis

Process Use Catalyst

Haber Production of ammonia IronContact Production of sulfuric acid Vanadium (V) oxide

Oswald Production of nitric acid PlatinumHardening of oils Production of margarine NickelCatalytic cracking Production of alkenes Aluminium oxide

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3.4.1 Examples of catalysts

Catalysis by cobalt(II) chloride solution

The oxidation of aqueous tartrate ions by hydrogen peroxide solution below is slow evenwhen the solution is heated to 60◦C. The reaction can be catalysed by adding a solutioncontaining cobalt(II) ions.

Figure 3.3: Oxidation of tartrate ions

5H2O2(aq) + C4H4O62-(aq) → 4CO2(g) + 6H2O(l) + 2OH-(aq)

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Catalysis by cobalt(II) chloride solution

The online version of this topic contains a video clip. If you do not have access to this,study the following diagram and answer the questions which follow.

Catalysis by cobalt(II) chloride

Q8: Which species is responsible for the pink colour at the beginning?

a) Cl-(aq) ionsb) Co2+(aq) ionsc) Co3+(aq) ionsd) tartrate ions

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Q9: Suggest a species that could be responsible for the green colour which appears?


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Q10: Which species acts as the catalyst?


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Q11: A student defines a catalyst as a substance which speeds up a reaction withouttaking part. Explain whether or not this is correct, with reference to the above reaction.

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Key Point

Cobalt(II) chloride speeds up the reaction between tartrate ions and hydrogenperoxide. Colour changes clearly show that the catalyst takes part in the reactionand is chemically unchanged at the end of the reaction.

Ozone layer destruction

Another example of catalysis is found in the destruction of the ozone layer by CFCs(chlorofluorocarbons). Ozone, O3, is constantly being formed in the upper atmospherewhere it provides a protective layer by absorbing a lot of the harmful ultraviolet radiationwhich could cause serious skin damage including skin cancer. Ozone slowly breaksdown by reacting with oxygen atoms to form oxygen molecules below.

The breakdown of ozone

This breakdown is accelerated by chlorine atoms below which are produced whensunlight in the upper atmosphere causes the CFCs to break down.

The catalysed breakdown of ozone



Q12: Which species is acting as a catalyst?

a) O3

b) O2

c) ClOd) Cl

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The Monsanto process

A very important industrial example of catalysis is the Monsanto process for thesynthesis of ethanoic acid from methanol.

The Monsanto process

The reaction takes place in solution using a soluble compound containing rhodium.

Industrial uses of catalysts

Throughout this topic, examples have been given of the use of catalysts in industry. Youdo not need to know them all but should be able to give some examples. Without theuse of catalysts most industrial processes would not be feasible. You may need to referback to the earlier sections of this topic or use the Internet or some of the text bookslisted in the section on Resources.

Q13:

Copy or get a photocopy of the following table. Use the word bank to help you completeit. You may need to refer back to the earlier sections of this topic or use some of thetext books listed in the section on Resources.



Ethanoic acid from methanol & CO

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3.5 Potential energy diagrams

Consider the exothermic reaction between zinc and sulfuric acid.

Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

After the reaction, the energy stored in the products is less than the energy originallystored in the reactants - that is, some of the chemical potential energy stored inreactants has been transferred to the surroundings. This transferred heat energy isdefined as the enthalpy change for this reaction.

Strictly, the change in energy which occurs during a chemical reaction consists of twocomponents: heat and work. We live on the surface of the Earth, at the bottom of a 'sea'of atmosphere which exerts pressure. In the reaction above, the gaseous hydrogenproduced will have to 'work' to displace the atmosphere. This will require use of someenergy, so the heat energy released to the surroundings will be less than the total energychange. In this reaction, for 1 mole of zinc, the heat (enthalpy) change is 152.4 kJ andthe work against the atmosphere 2.5 kJ, the overall energy change being 154.9 kJ.

Chemists often carry out reactions at constant pressure, hence the use of 'enthalpychange', the heat change at constant pressure.



Let's consider another familiar exothermic reaction, the burning (oxidation) of methanein a flame. You see this every time you use natural gas.

These changes are often shown as a potential energy diagram. An example is shownin Figure 3.4 for methane oxidation.

Figure 3.4: Methane oxidation

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This diagram shows the potential energy of the reactants, at the start of the reaction. Italso shows the potential energy of the products, at the end of the reaction. The arrowshows the energy given to the surroundings.

The easiest measurement to make is the enthalpy change, ΔH, when the reactionoccurs. This is defined as the difference in enthalpy of the products and reactants bythe equation below. The units of energy and enthalpy are joules (J).

Enthalpy change

ΔH is negative for an exothermic reaction. The energy of the products is less than thatof the reactants.



Exothermic reaction

The potential energy diagram for an endothermic reaction, the decomposition ofcalcium carbonate

is shown below.

Calcium carbonate decomposition

In this case the energy of the products is greater than that of the reactants.

Q14: So what will happen to the temperature of the surroundings?

a) It will increase.b) It will not change.c) It will decrease.

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Q15: What will the value of ΔH be?

a) Negativeb) Zeroc) Positive

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This is shown below.

Endothermic reaction

In order to distinguish endothermic and exothermic reactions the + and - sign arealways used when giving ΔH values.

Q16:

The equation for the "water gas" reaction is

Figure 3.5: Water gas reaction

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Note that this equation states that the carbon is in the solid state (s), and the othermaterials are gases (g).

This reaction absorbs energy from its surroundings. What type of reaction is this?

a) Fastb) Exothermicc) Explosived) Endothermic

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Q17: What will be the sign of ΔH?

a) Positiveb) Has no sign.c) Negatived) Insufficient information to know.

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Q18: Draw a potential energy diagram for this reaction, and compare it with the answerat the back of the book. .

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But how much energy is involved in this reaction? In the case of the water gas reaction(Figure 3.5) 121 kJ is absorbed for each mole of carbon and steam used. In otherwords ΔH = +121 kJ mol-1.

The full potential energy diagram for the water gas reaction (see below).

Reaction progress

H/k

J

-111

-232

ΔH = +121 kJ mol-1

ReactantsC(s) + H2O(g)

ProductsCO(g) + H2(g)

Water gas potential energy diagram

Key Point

A potential energy diagram can be used to show the energy changes for areaction.The enthalpy change is the energy difference between products and reactants.The enthalpy change can be calculated from a potential energy diagram.The enthalpy change has a negative value for exothermic reactions and a positivevalue for endothermic reactions.

3.5.1 What produces this energy change?

The chemical energy in molecules is stored in the bonds. When a reaction occurs thebonds in the reactant molecules are broken and the atoms rearrange to form new bondswith different energies in the products.

For example, think of making ammonia from nitrogen and hydrogen.



The balanced chemical equation is

N2 + 3H2 2NH3

You can think of this reaction as first breaking the bonds in the nitrogen and hydrogenmolecules. This stage requires an input of energy.

Then these atoms rearrange and combine to form new N - H bonds in ammonia. Thisstage releases energy.

The enthalpy change will be given by the total energy in new bonds minus the energypresent in the original bonds.

It is important to remember that reactions do not normally proceed by such a route, butthe enthalpy change can be calculated as if they did.



3.5.2 The thermochemical equation

It is usual to express this information in a thermochemical equation, such as that forthe water gas reaction below (Figure 3.6).

Figure 3.6: Thermochemical equation for the water gas reaction

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Three points should be noted.

1. The enthalpy value quoted in the balanced equation is measured in kilojoules permole (kJ mol-1).

Look at Figure 3.6 again. The 121 kJ of energy is absorbed when one mole of solidcarbon and one mole of steam react; two moles of each reacting would absorb 242 kJ,and so on. (This makes sense when you think that burning 2 kg of carbon would giveout twice the energy from burning 1 kg.)

2. The equation always contains the state symbols of reactants and products.

The enthalpy of reactions will change if the states are different.

Look carefully at these two equations for the reaction of hydrogen and oxygen.

(N.B. the equation has been written for forming 1 mole of water. The "1/2" before theoxygen means 1/2 mole of reactant not 1/2 a molecule of oxygen.)

The extra enthalpy in the second case is due to the heat released when the watervapour condenses to liquid water.

3. Since the enthalpy change is independent of the route of the reaction, the enthalpychange for the reverse reaction will be equal in value, but of opposite sign.



For the reaction below:

Q19:

So the enthalpy change for the reverse reaction:

CO2(g) → CO(g) + 1/2O2(g)

will be?

a) ΔH = -566 kJ mol-1

b) ΔH = -283 kJ mol-1

c) ΔH = 0 kJ mol-1

d) ΔH =+283 kJ mol-1

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Q20: The equation for calculating the enthalpy change (ΔH) for a reaction is:

a) ΔH = Hproducts + Hreactants

b) ΔH = Hproducts - Hreactants

c) ΔH = Hreactants + Hproducts

d) ΔH = Hreactants - Hproducts

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Q21: A ΔH value of - 500 kJ mol-1 would indicate that the reaction is;

a) Fastb) Reversiblec) Exothermicd) Endothermic

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Q22:

The potential energy diagram for the combustion of one mole of ethene is:

What is the enthalpy change for this reaction?

a) ΔH = +52 kJ mol-1

b) ΔH = -1362 kJ mol-1

c) ΔH = -1310 kJ mol-1

d) ΔH = -1414 kJ mol-1

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Q23:

The equation for the combustion of ethanol is

How many kJ of energy will be released to the surroundings if 5 moles of ethanol areburned?

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Q24:

Given that

what would be the enthalpy change for the reaction below? (Hint: don't forget the sign.)

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Q25: The Ostwald process for producing nitric acid involves the oxidation of ammonia(NH3) to form nitric oxide (NO) and water. Write a balanced equation for this reaction.

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Q26: Given that the enthalpy of the reactants (as written in the balanced equation) is-184 kJ and the enthalpy of the products is -1090 kJ, calculate the enthalpy change forthis reaction in kJ. Remember the sign.

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Q27: What would be the enthalpy change per mole of ammonia oxidised? Give youranswer in kJ mol-1 to 1 decimal place.

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Q28: When 102 g of ammonia is oxidised, how many kJ of heat have to be removed tomaintain a constant temperature?

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3.6 Catalysts and energy

The effect of a catalyst on a reaction rate can be explained in terms of energy.

Potential energy diagrams are used to show the energy pathway of a chemical reaction,see below.

You will study these diagrams in much more detail in the next topic in this unit.

Q29: Which diagram shows an exothermic reaction, one in which the products haveless energy than the reactants, so heat is released to the surroundings?

a) Ab) B

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Q30: Enthalpy change is the difference in energy between products and reactant. Whatis the approximate enthalpy change in diagram A?

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Q31: Enthalpy change is the difference in energy between products and reactant.Which of these could be the enthalpy change in B?

a) +1040 kJb) -1040 kJc) +890 kJd) -890 kJ

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3.6.1 The activated complex

In considering energy changes in reactions, only the starting point and finishing pointare considered. Energy has to be put in to break the old bonds before energy can bereleased when the new bonds form. As a result, the potential energy diagram for anexothermic reaction has the shape shown below.

Figure 3.7: Potential energy diagram

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There is an energy barrier between the reactants and products. Reactant moleculesmust have enough energy when they collide to get over this barrier.

The maximum point, X, in the graph above corresponds to the activated complex. Thisis a very unstable arrangement of atoms in which the old bonds are half broken and thenew bonds are half formed. This species can either break down to reform the reactants(fall back down the same side of the hill) or break down to form the products (fall downthe other side of the hill).



Reaction profile

Reaction profile with activated complex

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The activation energy (EA), is the energy required by the colliding molecules to formthe activated complex, i.e. it is the height of the barrier. The activation energy can becalculated from potential energy diagrams.

Q32: Which of these could be the activation energy of the reaction shown in Figure 3.7?

a) -180 kJb) +180 kJc) -110 kJd) +110 kJ

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Q33: Which of these could be the activation energy of the reaction shown above?

a) -22 kJb) -55 kJc) +55 kJd) +90 kJ

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Q34: Explain what will happen to the rate of the reaction if the activation energy isdecreased.

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Earlier, the activation energy was introduced as the minimum kinetic energy requiredby colliding molecules before reaction can occur, as shown in the energy distributiondiagram below.



Figure 3.8: Energy distribution graph

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Q35: Which letter shows the molecules with least energy?

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Q36: Which letter shows molecules with enough energy to react?

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Q37: What will happen to the number of molecules with enough energy to react if theactivation energy is decreased?

a) The number of molecules with enough energy will decrease.b) The number of molecules with enough energy will stay the same.c) The number of molecules with enough energy will increase.

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The two definitions of activation energy are quite compatible. As colliding molecules getcloser together, they begin to repel (potential energy increases) and slow down (kineticenergy decreases). Kinetic energy is converted into potential energy. If the collidingmolecules have enough kinetic energy, then enough potential energy will be producedto form the activated complex and so cause a reaction. A successful collision will havetaken place.

Key Point

The activated complex is an unstable arrangement of atoms formed during areaction at the maximum of a potential energy barrier. The activation energyis energy required by colliding molecules to form the activated complex. Theactivation energy can be calculated from potential energy diagrams.



3.6.2 How catalysts work

How catalysts work

The following diagrams show two images taken from an activity which can be viewed inthe online version of this topic.

Without a catalyst

With a catalyst

Q38: Without a catalyst, which of the following is true?

a) High activation energy, lots of successful collisions.b) High activation energy, few successful collisions.c) Low activation energy, lots of successful collisions.d) Low activation energy, few successful collisions.

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Q39: With a catalyst, which of the following is true?

a) High activation energy, lots of successful collisions.b) High activation energy, few successful collisions.c) Low activation energy, lots of successful collisions.d) Low activation energy, few successful collisions.

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Q40: What effect does a catalyst have on the reaction rate?

a) It speeds up the reaction by adding extra energy.b) It speeds up the reaction by making the reaction exothermic.c) It speeds up the reaction by lowering the activation energy.d) It speeds up the reaction by raising the activation energy.

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Q41: What effect does a catalyst have on the enthalpy change?

a) Increases it.b) No effect.c) Decreases it.

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The catalyst provides an alternative pathway for the reaction. often the catalyst providesa surface on which the reaction occurs. The bonds in the reactant molecules areweakened when the molecules are adsorbed onto the surface. This pathway has alower activation energy and so the reaction is faster.

Mechanism of catalysis

The online version of this Topic contains an activity which shows how a catalyst mightspeed up the gas phase reaction between two reactants A and B, as discussed earlier.The equations which follow show the possible steps.

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In the uncatalysed reaction, A2 molecules and B2 molecules must collide with sufficientenergy to form the activated complex above.



In the catalysed reaction, step 1 involves a lower energy collision between the catalyst(Y) and an A2 molecule. In step 2, the intermediate, A2Y, then undergoes another lowenergy collision with B2 to form the products. Y is regenerated in step 2 and is free torepeat the process.

A potential energy diagram can be drawn for the process.

Figure 3.9: PE diagram with catalyst

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Q42: Which letter corresponds to the activation energy of the catalysed formation ofAB?

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Q43: Which arrow shows the activation energy of the uncatalysed decomposition ofAB?

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Q44: Which arrow (Figure 3.9) shows the enthalpy change for the catalysed formationof AB?

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Q45: What effect has the catalyst had on the activation energy of the reverse reaction?

a) Increases itb) Decreases itc) No effect

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Q46: By considering the equations in Figure 3.9, which letter in Figure 3.9 shows thestage in the uncatalysed reaction at which you are most likely to find A2Y + B2?

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Q47:

By considering the equations, in Figure 3.9 which letterin Figure 3.9 shows the stage in the catalysed reactionat which you are most likely to find the speciesopposite?

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Q48: Draw a possible structure for the species that exists at point Z (Figure 3.9).

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Key Point

Catalysts speed up reactions by providing an alternative pathway which has alower activation energy. A potential energy diagram can be used to show theeffect of a catalyst on activation energy.

3.7 Summary

Summary

• In general, the lower the activation energy the faster the reaction.

• Catalysts are substances which speed up chemical reactions without beingused up in the process. They are widely used in industrial processes.

• Catalysts allow chemical reactions to occur more quickly at lowertemperatures and so reduce energy costs.

• Catalysts work by the adsorption of reactant molecules onto the surfaceof the catalyst with consequent weakening of the bonds in the reactantmolecules.

• Catalytic converters are fitted to cars to catalyse the conversion ofpoisonous carbon monoxide and oxides of nitrogen to carbon dioxide andnitrogen. Cars with catalytic converters can only use 'lead-free' petrol toprevent poisoning of the catalyst.

• Enzymes catalyse the chemical reactions which take place in the living cellsof plants and animals. They are also widely used in industrial processes.

• Cobalt(II) chloride speeds up the reaction between tartrate ions andhydrogen peroxide. Colour changes clearly show that the catalyst takespart in the reaction and is chemically unchanged at the end of the reaction.



Summary Continued

• The activated complex is an unstable arrangement of atoms formed duringa reaction at the maximum of a potential energy barrier.

• A potential energy diagram can be used to show the energy pathway for areaction.

• The enthalpy change, which can be calculated from the potential energydiagram, is the energy difference between products and reactants.

• The enthalpy change has a negative value for exothermic reactions, whichcause heat energy to be released to the surroundings.

• The enthalpy change has a positive value for endothermic reactions, whichcause absorption of heat energy from the surroundings.

• The activation energy is energy required by colliding molecules to form theactivated complex.

• The activation energy can be calculated from potential energy diagrams.

• Catalysts speed up reactions by providing an alternative pathway which hasa lower activation energy.

• A potential energy diagram can be used to show the effect of a catalyst onactivation energy.


ISBN 978-1444167528




End of topic 3 test


Q49:

In the (red) shaded area on the graph:

A) molecules always have energy above the activation energy.

B) there are no molecules with kinetic energy above the average.

C) there are no molecules with energy above the activation energy.

D) collisions between molecules are always successful in forming products.

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Q50: Activation energy is:

a) the minimum kinetic energy required by colliding molecules for a reaction to occur.b) the maximum kinetic energy colliding molecules can reach for a reaction to occur.c) the temperature at which a reaction will take place.d) the temperature at which rate of reaction will be greatest.

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The progress of reactions can be followed by energy diagrams.



Q51:

Which graph represents the catalysed version of the reaction in diagram F?

i A

ii B

iii C

iv D

v E

vi F

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Q52:

Which graph represents the reaction with the highest activation energy?

i A

ii B

iii C

iv D

v E

vi F

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Q53:

Which diagram represents the reaction with an enthalpy change of -200 kJ mol -1?

i A

ii B

iii C

iv D

v E

vi F

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Q54:

What is an activated complex?

A) A very unstable intermediate is formed when bonds within the reactant moleculesbegin to break and new bonds begin to form.

B) The energy required for a reaction to take place.

C) When the surface of a catalyst forms weak bonds with reacting molecules.

D) A compound containing a transition metal.

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Use the diagram to answer the following questions.

Q55: What type of reaction does this graph represent?

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Q56: What is the activation energy value for the forward reaction?

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Q57: What is the activation energy value for the reverse reaction?

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Q58: What is the value of the enthalpy change for the forward reaction?

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Q59: What is the value of the enthalpy change for the reverse reaction?

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Q60:

a) Name the energy changes represented by the numbered intervals on the energydiagram shown.

b) The reaction shown is exothermic, how can you tell this from the shape of thegraph?

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71

Topic 4

The Periodic Table

Contents

4.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73

4.2 Arrangement of elements in the Periodic Table: Introduction . . . . . . . . . . . 73

4.3 History of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74

4.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . . . . . . . . . . . . 79

4.4.1 Melting point and boiling point . . . . . . . . . . . . . . . . . . . . . . . 79

4.4.2 Atomic size . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 81

4.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 83

4.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 83

4.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 84



• atoms contain protons, neutrons and electrons each with a specific charge, massand position within the atom. The number of protons defines an element and isknown as the atomic number (National 4, Unit 1);

• learners should have knowledge of: sub-atomic particles, their charge, mass andposition within the atom, the structure of the Periodic Table, groups, periods andatomic number (National 5, Unit 1);

• all matter is made of atoms. When a substance contains only one kind of atom itis known as an element (National 4, Unit 1);

• elements are arranged in the Periodic Table in order of increasing atomic number;elements with similar chemical properties are grouped together (National 4, Unit1);

• or be familiar with the seven diatomic elements (National 5, Unit 1);

• elements can be categorised as metals and non-metals (National 4, Unit 1);

Learning Objectives

At the end of this topic, you should know that:

• elements are arranged in the Periodic Table in order of increasing atomic number;

72 TOPIC 4. THE PERIODIC TABLE

• the Periodic Table allows chemists to make accurate predictions of physicalproperties and chemical behaviour for any element based on its position;

• there are periodic variations in the densities, melting points and boiling points ofthe elements across a Period and down a Group.


TOPIC 4. THE PERIODIC TABLE 73

4.1 Prior knowledge


Q1: The number of protons in an atom determines what?

a) Atomic Numberb) Mass Numberc) Name of Elementd) Charge of atom

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Q2: Name the seven diatomic elements.

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Q3: Complete the following table;

Particle Charge Mass LocationProtonNeutronElectron

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Q4: What is a row in the Periodic Table is called?

a) Setb) Rowc) Groupd) Period

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Q5: What is a column in the Periodic Table is called?

a) Setb) Columnc) Groupd) Period

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4.2 Arrangement of elements in the Periodic Table:Introduction

As more and more elements were discovered and their properties investigated, chemistsshowed a natural desire to simplify the study of the elements by organising themaccording to similarities in their chemical behaviour. Eventually this resulted in themodern Periodic Table. If you understand how the Periodic Table is constructed, youwill realise that it contains a huge amount of information stored in a very compact form.This makes it a vital resource for any chemist.



4.3 History of the Periodic Table

Background Information

In 1800, about 33 elements were known but there was no obvious pattern or relationshipbetween them.

By 1830, a further 20 or so elements had been discovered and some similarities inproperties within small groups of elements was recognised. The German chemist,Johan Wolfgang Dobereiner, made a tentative connection between chemical behaviourand the atomic masses of certain groups of elements, each containing three elements,which he called 'triads'.

Figure 4.1: Triads

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In each case, the atomic mass of the central element was approximately the mean ofthe other two.

The next significant development occurred in 1866 when the English chemist, JohnNewlands, published a paper on the 'Law of Octaves'.

Newlands' octaves

The on-line version of this topic contains two activities which explain how Newlandsorganised the elements.

Figure 4.2 shows how Newlands organised the first 14 elements. Study the diagramcarefully and then answer the questions which follow.

Figure 4.2



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Questions about Newlands' arrangement of the first 14 elements.

Q6: What property did Newlands use to put the elements in order?

a) Name (i.e. alphabetical order)b) Colourc) Atomic massd) Atomic number

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Q7: What property is shared by the elements in the first column?

a) Both are gases.b) Both are reactive metals.c) None - they are completely different.

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Q8: What property is shared by the elements in the second column (i.e. Li and Na)?

a) Both are gases.b) Both are reactive metals.c) None - they are completely different.

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Q9: Why do you think Newlands referred to these as 'octaves' ? (Hint - think aboutmusical scales.)

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Soh-fah so good! Now look at the second diagram (Figure 4.3) then answer thequestions. The next seven elements have been added.

Figure 4.3

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Q10: At first sight, the pattern seems to continue. For how many of the further elementsdoes the pattern work?

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Q11: Can you think of a reason why chromium, manganese and iron do not fit with theelements immediately above them?

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When Newlands arranged the elements in order of increasing atomic mass, similarchemical properties were repeated with every eighth element but this only worked forthe first 17 elements.

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Newlands' ideas were subjected to ridicule and it was even suggested unkindly that hewould get better agreement if he arranged the elements in alphabetical order. However,he was on the right lines.

The Modern Periodic Table

The major breakthrough, indeed one of the most important advances in all chemistry,was provided by Dmitri Ivanovitch Mendeleev in 1869. Like Newlands, he organised theelements in order of increasing atomic mass but there were other important differences.

Mendeleev's Periodic Table

Mendeleev organised the first 45 of the 62 elements known at that time. The diagramsbelow (Figure 4.4) show his final arrangement.

Study the tables carefully and then answer the questions which follow.

Figure 4.4: Constructing Mendeleev's table



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Q12: Unlike Newlands, Mendeleev left spaces (shown as *) in his table. Why?

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Q13: Explain why Mendeleev swapped the positions of the elements iodine andtellurium.

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Q14: Name the Group of elements which is found in the modern Periodic Table but isabsent from both Newlands' and Mendeleev's tables.

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Key Point

Mendeleev organised the elements in order of increasing atomic mass (mostly) inconjunction with similar chemical properties, leaving gaps for elements yet to bediscovered.

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Figure 4.5

The online version of this topic contains a further activity that shows the relationshipbetween Mendeleev's table and the modern Periodic Table.

Here is our modern Periodic Table.

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Mendeleev was so confident that his table was correct that he predicted the propertiesof some of the undiscovered elements. His predictions proved to be very accurate whenthese elements were finally isolated, providing startling vindication of his theory.

The modern Periodic Table can be shown in a variety of different ways. The usual formis the one used in the recommended data booklet. The resources at the end of this topic



contain addresses for a number of websites which contain interactive Periodic Tables.One of the most useful is http://www.webelements.com/.

Q15: In the modern Periodic Table, the elements are not arranged in order of increasingatomic mass. What is used instead?

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4.4 Trends and patterns (periodicity)

When the elements are arranged in order of increasing atomic number, many propertiesvary in a regular way. As you move across a Period from left to right, a pattern emerges.A similar pattern appears on crossing the next Period. Properties which behave in thisway are said to be periodic. Periodicity is the regular recurrence of similar elementproperties.

Throughout this topic, we will concentrate on the properties of the main group elementsand in particular on the elements in Periods 2 and 3. Much of this work involves thecollection of data, presentation of data in graphical form and the interpretation of suchgraphs.

4.4.1 Melting point and boiling point

Figure 4.6: Period 2 and 3 (mp and bp)

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http://www.webelements.com/


Melting points and boiling points are also periodic properties. Melting points and boilingpoints depend on the strength of the forces which exist between the particles whichmake up a substance.

Q16: Write the name of the element in Period 2 which has the strongest forces betweenits atoms in the solid state.

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Q17: Write the name of the Period 3 element which is the easiest to boil?

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Q18: Of all the elements shown in the data booklet, which element which has theweakest forces between its atoms.

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Again a more obvious trend can be seen on descending a Group.

Figure 4.7: Alkali metals (mpt and bpt)

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Q19: Write the name of the alkali metal which has the strongest forces between itsatoms.

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Q20: Write the name of the alkali metal which has the weakest forces between itsatoms.



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Q21: What happens to the forces between the atoms on descending Group 1?

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The variation in melting point and boiling point will be explained when the bonding andstructure of the first 20 elements is considered later in this Topic.

Key Point

There are periodic variations in the densities, melting points and boiling points ofthe elements across a Period and down a Group.

4.4.2 Atomic size

Explanation

The size of an atom is determined by the amount of space taken up by the electronsand so must be connected to the electron arrangement of the atom. The electronarrangement is itself a periodic property. The online version of this topic contains anactivity to explain the trends in covalent radii across a Period and down a Group. If youdo not have access to the online version, use the following diagrams (Figure 4.8) toanswer the questions.

Figure 4.8



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Q22: Which of the following provides the best reason for the increase in covalent radiuson going down a Group?

a) The number of protons increases.b) The number of electrons increases.c) The number of electron shells increases.d) The number of neutrons increases.

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Q23: Which of the following provides the best reason for the decrease in covalent radiuson going from left to right across a Period?

a) The number of electrons increases.b) The number of electron shells increases.c) The number of neutrons increases.d) The number of protons increases.

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Key Point

• The covalent radius decreases across a Period because the increase innuclear charge attracts the electrons more strongly.

• The covalent radius increases on going down a Group as the number ofoccupied electron shells increases.

4.5 Summary

Summary

• Elements are arranged in the Periodic Table in order of increasing atomicnumber.

• The Periodic Table allows chemists to make accurate predictions of physicalproperties and chemical behaviour for any element based on its position.

• There are periodic variations in the densities, melting points and boilingpoints of the elements across a Period and down a Group.


ISBN 978-1444167528




End of topic 4 test


Q24:

Menedeleev is famous for producing the Periodic Table on which the modern version isbased.

Which of the following statements is true?

a) Mendeleev organised the elements in order of their atomic number.b) Mendeleev left gaps because some elements did not fit the pattern of reactivity.c) Mendeleev left gaps for elements which had not yet been discovered.d) Mendeleev swapped some elements round so that their atomic masses fitted the

pattern.

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Q25:

Which of the following statements about the Periodic Table are true?

a) There is a steady increase in melting point across a period from left to right.

b) There is a steady decrease in density on going down Group 1.

c) There is a steady decrease in atomic size across a period from left to right.

d) There is a decrease in first ionisation energy on going down Group 0.

e) There is a decrease and then an increase in boiling point on crossing a period fromleft to right.

f) There is an increase in electronegativity on going down Group 7.

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Q26:

What does the number of protons in an atom determine?

a) Atomic Number.b) Mass Number.c) Name of Element.d) Charge of atom.

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Q27:

Name the seven diatomic elements.

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Q28:

What is Group 1 in the Periodic Table called?

a) The halogens.b) The alkali metals.c) The noble gases.d) Transition metals.

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Q29:



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Q30:



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Q31:

Elements in the same group in the Periodic Table have the same:

a) number of occupied energy shells;b) density;c) number of outer electrons;d) number of protons.

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Q32:

Where in the Periodic Table are non-metals found?

a) The top.b) The bottom.c) The right hand side.d) The left hand side.

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87

Topic 5

Bonding and structure

Contents

5.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 89

5.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

5.3 Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 94

5.3.1 Metallic structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95

5.4 Covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 96

5.4.1 Covalent molecular structures . . . . . . . . . . . . . . . . . . . . . . . . 98

5.4.2 Covalent network structures . . . . . . . . . . . . . . . . . . . . . . . . . 102

5.5 Ionic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106

5.5.1 Ionic lattice structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . 109

5.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111

5.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111

5.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112



• elements can be categorised as metals and non-metals (National 4, Unit 1);

• experimental procedures are required to confirm the type of bonding present in asubstance (National 5, Unit 1);

• metallic bonding can explain the conductivity of metals (National 5, Unit 3);

• covalent compounds form when non-metal atoms form covalent bonds by sharingtheir outer electrons (National 4, Unit 1);

• covalent molecular compounds have low melting and boiling points. As a result,they can be found in any state at room temperature (National 4, Unit 1);

• in a covalent bond, the shared pair of electrons is attracted to the nuclei of the twobonded atoms (National 5, Unit 1);

• more than one bond can be formed between atoms leading to double and triplecovalent bonds (National 5, Unit 1);

• covalent substances can form either discrete molecular or giant network structures(National 5, Unit 1);

88 TOPIC 5. BONDING AND STRUCTURE

• diagrams show how outer electrons are shared to form the covalent bond(s) in amolecule and the shape of simple two-element compounds (National 5, Unit 1);

• covalent molecular substances have low melting and boiling points due to onlyweak forces of attraction between molecules being broken (National 5, Unit 1);

• giant covalent network structures have very high melting and boiling pointsbecause the network of strong covalent bonds must be broken (National 5, Unit 1);

• when there is an imbalance in the number of positive protons and electrons theparticle is known as an ion (National 5, Unit 1);

• ionic bonds are the electrostatic attraction between positive and negative ions.Ionic compounds form lattice structures of oppositely charged ions (National 5,Unit 1);

• ionic compounds have high melting and boiling points because strong ionic bondsmust be broken in order to break down the lattice. Dissolving also breaks downthe lattice structure (National 5, Unit 1);

• ionic compounds have high melting and boiling points. As a result, they are foundin the solid state at room temperature (National 4, Unit 1);

• ionic compounds form when metal atoms join to non-metal atoms by transferringelectron(s) from the metal to the non-metal. The resulting charged particles arecalled ions and an ionic bond is the attraction of the oppositely charged ions(National 4, Unit 1);

• ionic compounds conduct electricity, only when molten or in solution due to thebreakdown of the lattice resulting in the ions being free to move (National 5, Unit1).

Learning Objectives

At the end of this topic, you should know that the first 20 elements in the Periodic Tablecan be categorised according to bonding and structure:

• metallic Li, Be, Na, Mg, Al, K, Ca;

• covalent molecular H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60);

• covalent network B, C, Si (diamond, graphite and the element silicon and silicondioxide) monatomic (noble gases).

You should be able to explain each of these types of substance in terms of bonding andstructure.


TOPIC 5. BONDING AND STRUCTURE 89

5.1 Prior knowledge


Q1: Covalent bonding involves:

a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) gaining electrons.

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Q2: Metals can conduct electricity in their solid state because of their:

a) ions.b) positive cores.c) lattice structure.d) delocalised electrons.

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Q3: Ionic bonding involves:

a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) protons only.

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Q4: One property of covalent networks is that they:

a) have high melting and boiling points.b) have low melting and boiling points.c) are soluble in water.d) conduct electricity when molten or in solution.

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Q5: One property of ionic substances is that they:

a) are all white in colour.b) have low melting and boiling points.c) are insoluble in water.d) conduct electricity when molten or in solution.

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5.2 Introduction

Most elements are found on Earth as compounds.

Q6: Can you think of any elements that are found free, not as compounds?

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Comparing the elements iron and tantalum

Sometimes ores can be simple compounds (e.g. haematite, an oxide of iron, Fe2O3);some are complex minerals (e.g. columbite, containing tantalum, (Fe,Mn)(Ta,Nb)2O6).

Haematite ("Haematite", by By Sailko, is licensed under CC BY 3.0)

Columbite is a black mineral group that is an ore of niobium ("Columbite", byRob Lavinsky, is licensed under CC BY 3.0)


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http://commons.wikimedia.org/wiki/User:Sailko

http://creativecommons.org/licenses/by/2.0/deed.en

http://commons.wikimedia.org/wiki/File:Columbite-75444.jpg

http://www.irocks.com

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Iron is also abundant in the Earth's crust (5%); tantalum is rare (1 - 2 parts per million).

The extraction of iron from its ores is an established, well-understood process; theextraction of tantalum is complex, mainly because it occurs with the very similar metal,niobium.

Q7: Niobium is used for a very special purpose in space rockets. Can you find outwhat it is for?

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Iron is used for a vast variety of purposes; tantalum is almost exclusively used to makehigh-performance capacitors in electronic equipment, for example, mobile phones.

Despite these contrasts, you should always remember that all the 'chemicals' wehumans use have to be extracted from the finite resources present in the Earth.

A variety of chemical processes are used to extract elements from their compounds,but the choice depends on the bonding and structure of the materials containing theelement.

Bonding and structure in elements and simple compounds

Everything we see around us is made from fewer than 100 different types of atomschemically bonded in various ways to produce a multitude of different molecules.

The different types of chemical bonding determine the structure that elements andcompounds adopt. In turn the structure, together with the size of intermolecular forcesmainly determines the physical and chemical properties possessed by these materials.

Bonding and structure will be studied in this topic. Intermolecular forces (the forces thatexist between molecules) and properties are studied in a later topic.

Examples of some substances with different properties, depending on different bondingand structures are shown in the images below.



Iron screws - a typical metal Water - a volatile liquid

Nitrogen dioxide - a brown gas Sodium chloride - a white crystallinesolid

No atoms, except those of the noble gases, exist in isolation under normal conditions.They all interact in one way or another to form more stable structures.

Atoms consist of a tiny positively charged nucleus, with different numbers of electrons,in shells of increasing energy levels, around it.

The activity below shows the final electronic structure of a sodium atom.



Construction of a sodium atom

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Construction of a sodium atom

The illustration shows that in the sodium atom there is a single electron in theoutermost shell.

In the case of the atoms of noble gases all shells are completely filled with electrons.This arrangement is particularly stable, which makes noble gases unreactive.

Atoms of other elements combine in ways which try to achieve this stable noble gasarrangement of electrons, in other words, to become isoelectronic with a noble gas.

Q8: Look at the Periodic Table in the SQA data booklet (page 2). Which noble gas hasan electronic structure closest to chlorine?

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Q9: What must happen to a chlorine atom to get the noble gas electronicarrangement?

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Q10: Look again at the diagram of a sodium atom. What must happen to it for it to havea noble gas electronic structure?

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Q11: What will the sodium atom become?

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Q12: Which noble gas has the same electronic structure as a sodium ion?

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5.3 Metallic bonds

Metals occur to the left and centre of the Periodic Table (see Topic 4). When atomsof metals come together the most stable condition is for them to "pool" their outerelectrons and become, in effect, a regular arrangement of fixed metal ions in a "sea"of delocalised electrons shared by all the ions.

For example, in metallic sodium each atom will donate an electron to the delocalisedpool, and will achieve the stable sodium ion structure (isoelectronic with neon).

Metallic bonds

The next diagram shows the structure of metallic sodium.

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Such an array of positively charged ions would normally split apart by mutualrepulsions, but the influence of the negative electrons holds the particles in thestructure so well that most metals are hard solids with high melting and boiling points.

Q13: There are exceptions. Can you think of soft metals, and one that is liquid undernormal conditions?

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5.3.1 Metallic structures

In metallic bonding the delocalised electrons are able to migrate freely throughout themetal (making them good conductors of electricity), but the positive ions have a regular3D structure known as a lattice.

Part of the structure of the giant lattice for copper is shown below.

Copper lattice

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Q14: The electrons in a metallic bond are said to be:

a) loose.b) ionised.c) delocalised.d) hard.

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Q15: What do the ions in a metallic bond form into?

a) Gelb) Gridc) Bard) Lattice

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Q16: Suggest a reason why aluminium is a better conductor of electricity than sodium.

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Key Point

Metallic bonding is the electrostatic force of attraction between positively chargedions and delocalised outer electrons.A metallic structure consists of a giant lattice of positively charged ions anddelocalised outer electrons.

5.4 Covalent bonds

You should already have studied covalent bonds, which are the most common type ofbond. This section revises your knowledge.

Covalent bonds form when atoms share two electrons, enabling both atoms tocomplete their valency shells.

An on-line activity shows two chlorine atoms forming a chlorine molecule.

The activity below illustrates two separate chlorine atoms and the sharing of electronsto form a chlorine molecule.

Covalent bond formation

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This sharing of two electrons (one from each atom) to complete an outer shell ofelectrons is called a covalent bond.



In most cases the outer shell will contain 8 electrons, but for hydrogen only two arepresent.

The figure below shows the electrostatic forces in a hydrogen molecule. The positivelycharged nuclei will repel each other, as will the negatively charged electrons; but theseforces are more than balanced by the attraction between the nuclei and electrons.

Attractive and repulsive forces in a hydrogen molecule

When atoms require more than one electron to complete their outer shell, they canshare two or three electrons to make a double or triple covalent bond, or share withmore than one other atom.

The diagrams show the shared electrons in a molecule of oxygen (O2) and ammonia(NH3). Notice the easier way to show a covalent bond with a line for each bonding(shared) pair of electrons.

Covalent bonding in oxygen and ammonia

The dark pairs of electrons in above figure are the bonding electrons, sharedbetween the atoms. The light pairs of electrons, completing the shells around theatoms, are called lone pairs.



Q17: How many electrons are involved in the double bond in an oxygen molecule?

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Q18: How many lone pairs are there in an ammonia molecule?

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Q19: Which two words complete this sentence?A covalent bond is formed when two atoms —— a pair of electrons, so that each canachieve a noble gas ——— configuration.

a) share, protonb) share, electronc) donate, atomicd) donate, ionic

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Q20: The dominant attractive force in a covalent bond is between:

a) the positively charged nuclei.b) the negatively charged shared electrons.c) negatively charged nuclei and positively charged shared electrons.d) positively charged nuclei and negatively charged shared electrons.

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Q21: Draw the structure of iodine chloride. Look at the SQA data booklet page 2 to getthe electron structures of the atoms.

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Key Point

Atoms in a covalent bond are held together by electrostatic forces of attractionbetween positively charged nuclei and negatively charged shared electrons.

5.4.1 Covalent molecular structures

Most covalent substances (for example all those mentioned in the previous section) existas discrete molecules. There are strong covalent bonds binding the atoms together inthe molecule, but much weaker forces between these molecules. You will study theseforces in detail later (refer to the 'Intermolecular Forces' Topic). Consequently manysmall covalent molecules are gases (e.g. fluorine, oxygen, nitrogen, carbon dioxide,sulfur dioxide etc.). Larger molecules make structures which are liquids or low meltingpoint solids. (Think of candle wax, with molecular mass around 500, but a low meltingpoint about 70◦C.)

The molecules of nitrogen, oxygen, and the halogens consist of two atoms, they arediatomic. (You always write them N2, O2 etc.)

The non-metals phosphorus and sulfur form larger molecules as described in the nextactivities.



White phosphorus

White phosphorus

Phosphorus (in the common form of "white" phosphorus) forms P4 molecules, witheach phosphorus atom at the corner of a tetrahedron. A model is shown below.

Tetrahedron of four atoms in white phosphorus

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Red phosphorus

Red phosphorus

Phosphorus can also form a much less reactive (and less toxic) form, "red" phosphoruswhich is used in making matches. Its structure consists of chains of these tetrahedra.The structure is shown below.

These P4 tetrahedra are quite stable, so that it is usual to write "P4" in equations wherephosphorus is involved. It is even retained in some compounds - the oxides aremolecules of formula P4O6 and P4O10.

Red phosphorus

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Sulfur structure

Molecules of sulfur (in the normal solid state as rhombic sulfur) consist of puckeredrings of eight sulfur atoms, written S8. A model showing two of these rings see below.

Rings of eight sulfur atoms

In this case, using S8 in equations would be rather cumbersome, so that sulfur inequations is written "S".

Allotropes of sulfur

In addition to the eight-membered rings in the rhombohedral, α-sulfur found in common'flowers of sulfur', there are many other allotropes of sulfur. At temperatures above95◦C the S8 rings pack to form monoclinic crystals of β-sulfur.

Sulfur can also form rings and chains with any number of S atoms from 2 to 20.S2 is a gaseous form, an analogue of O2; the blue colour of burning sulfur is due to theemission of light by S2 molecules produced in the flame.

S3 is cherry red and has a structure like ozone, O3.

N.B. The recently discovered molecular structures of carbon (the fullerenes) arediscussed in the next section, after the other common forms of carbon.

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5.4.2 Covalent network structures

In some elements and compounds the covalent bonds are not limited only to those withinthe molecules, but all the atoms are held to others by strong covalent bonds.

An example of this is the structure of diamond. It consists solely of a network of carbonatoms held together by covalent bonds. Each carbon atom makes covalent bonds tofour different carbon atoms, which each bond to three more carbons, and so on,binding the whole structure together, as shown below (Diamond).

This covalent network structure has the characteristic properties of hardness and ahigh melting point, both due to the strong bonding throughout the structure.

Diamond

Arrangement of carbon atoms in diamond

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Q22: How many pairs of electrons are around each carbon atom?

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Q23: What type of bond holds the carbon atoms together in diamond?

a) Single covalent.b) Double covalent.c) Triple covalent.

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Silicon dioxide

Look at the model below, which compares the structure of silica (quartz, SiO2) withdiamond. Again, this consists of a network of covalent bonds, with each silicon atombonded to four oxygen atoms which in turn bond to another silicon atom, to form astrongly-bound network. Quartz is hard and has a high melting point, like diamond.

Silica and diamond

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Graphite structure

Carbon has another interesting structure, graphite. In this case carbon atomscovalently bond with only three other atoms to form a sheet of hexagonal rings, asshown below.

The distance between carbon atoms within these sheets is 142 pm.



Q24: This is less than the bond length in diamond (154 pm) and suggests that the bondstrength is:

a) weaker.b) the same.c) stronger.

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Q25: In these sheets, each carbon atom is bonded to only three others. How manybonding electrons surround each carbon atom in this case?

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Graphene

It is only recently that single layers of carbon atoms as found in graphite have beenisolated in sufficient quantities to study their properties. The Nobel Prize for physics for2010 was awarded to Andre Geim and Konstantin Novoselov for work on graphene, asit was called, done in Manchester University in 2004. They found that these isolatedsingle layers have very unusual electrical and physical properties, which have startedresearch into new electronic applications.

Graphite

The carbon atoms have four electrons in their outer shell. Only three are used to bondwith three other atoms, leaving a fourth electron. These sheets are layered togetherand the carbon atoms donate their extra electron to a delocalised pool which holds thesheets in a metallic type bonding. The structure is shown in the model below.

Layers of carbon atoms in graphite

The distance between these sheets is 335 pm which indicates a weak bonding. Theseweaker forces produce a material which is greasy (as the layers slip over each other)and can be used as a lubricant, and the delocalised electron pool makes graphite agood conductor of electricity.

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Fullerenes

Another form of carbon (discovered in the 1980s) exists as covalent molecularstructures. These are the fullerenes.

The most stable of these has 60 carbon atoms arranged in 5- and 6-memberedrings forming a large sphere, looking like a football, see the figure below. It iscalled buckminsterfullerene after the architect (Robert Buckminster Fuller) who createsgeodesic domes resembling the structure found in these carbon molecules.

Buckminsterfullerene

These fullerenes have interesting properties, which are currently being researched. Forexample, the "buckyball" above, will dissolve in benzene to form a red solution.

Q26: Why do you think buckminsterfullerene is soluble when diamond and graphite arenot?

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Long "nanotubes" of carbon have tensile strengths 50 to 100 times that of steel! Astructure is shown below.



Why are carbon nanotubes so strong? The diagram shows only a small part of a tube.When very many of these long molecules are combined, the result consists of verylarge molecules of carbon atoms bonded to each other with strong bonds, similar tothose in sheets of graphite. These bonds are much stronger than metallic bonds insteel.

Key Point

A covalent molecular structure consists of discrete molecules held together byintermolecular forces (see next Topic).A covalent network structure consists of a giant lattice of covalently bondedatoms.

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5.5 Ionic bonds

Polar covalent bonds (we will be studying 'Polar covalent bonding' in a later Topic) areformed when the atoms involved in a bond have different attractions for the bondingelectrons. When two atoms have a large difference in their attraction for electrons (e.g.sodium and chlorine), it is most energetically favourable for an electron in the metal tobe donated completely to the non-metal. This is the concept of electronegativity and willbe fully explained in the next topic. Both then achieve a noble gas structure as shown inthe activity below.



Q27: How many outer electrons does a chlorine atom have? (Answer with a number)

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Q28: When a sodium atom loses its outer electron to a chlorine atom, how many outerelectrons do both ions then have?

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Q29: Why is this special?

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The original atoms change to ions (the metal forms a positive ion, the non-metal anegative ion) which are held together by electrostatic attraction.

You can think of covalent, polar covalent and ionic bonds as forming a spectrum: bondsbecome more polar, then ionic as the difference in electronegativity increases.

The extent of covalent or ionic character depends mainly on the difference inelectronegativity. The properties of the different bonds are summarised in the tablebelow.



Bond Covalent Polar covalent IonicElectrons Equally shared Unequally shared Totally transferred

ChargeDistribution

None Partial +ve and -veFully charged +ve

and -ve ions

Other ionic compounds

Metals in Group 2, for example calcium, need to lose two electrons from their atoms toachieve a stable noble gas structure. This is reflected in the relatively low first andsecond ionisation energies. These elements will require two chlorine atoms to receivethese two electrons, and maintain electrical neutrality, so that the ratio of metal to nonmetal is 1:2 (e.g. [Ca2+][ Cl-]2)

Metals from other groups in the Periodic Table can form ions with 3+ charge, and othernon metals ions with 2- and 3- charges. Higher charges than these are rare, and whenelements form ionic bonds the ratio of ions is such that +ve and -ve charges alwaysbalance.

Further consideration of the nature of the ions produced by different elements will bediscussed in a later section.

Q30: Ionic bonds are formed between atoms which have a ——– difference inelectronegativity.

a) zerob) smallc) larged) 0.2

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Q31: Caesium chloride is an ionic solid consisting of:

a) negative caesium ions and negative chloride ions.b) negative caesium ions and positive chloride ions.c) positive caesium ions and negative chloride ions.d) positive caesium ions and positive chloride ions.

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Q32: An ionic crystal consists of M2+ and X3- ions. What will be the ratio of ions (M2+ :X3-) in the solid?

a) 1 : 1b) 2 : 3c) 3 : 2d) 3 : 3

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5.5.1 Ionic lattice structures

The mutual electrostatic attraction of positive and negative ions is completely nondirectional (unlike covalent bonds which are highly directed). When solid ioniccompounds form, a positive ion will be surrounded by several negative ions which inturn will attract more positive ions. This process results in the formation of a lattice ofregularly arranged ions all held together by electrostatic forces. These forces act equallyin all directions , so that no one ion is attracted particularly to any other specific ion. Forexample, in sodium chloride there are no moleculesformed. The ionic lattice for sodiumchloride is shown below.

Sodium chloride lattice

This extensively bonded ionic solid is relatively hard, though not as hard as thecovalent network solids, but is brittle. (Think of how easily salt (sodium chloride)crystals can be crushed.)

Rock salt cellar



Q33: What are Caesium chloride (CsCl) crystals formed from?

a) Covalent molecules.b) A covalent lattice.c) An ionic lattice.d) Ionic molecules.

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Now that we have looked at the three types of bonding in elements and compounds,see the table below roughly compares their strengths. The table also gives a figure forthe weaker bonds between molecules, which will be studied in detail in the next topic.

Bond type Strength/kJ mol-1

Metallic 80 - 600

Covalent 100 - 500

Ionic 100 - 450

Between covalent molecules 1 - 30

Strengths of bonds

You can see that metallic bonds have the greatest range of strengths. (This is sensibleif you think of hard metals like tungsten and chromium, and soft metals like sodium.)

The strengths of covalent and ionic bonds are similar, and much greater than the bondswhich hold covalent molecules together. (These are the intermolecular forces, whichare studied in the next topic.)

Key Point

Ionic bonding is the electrostatic force of attraction between positively andnegatively charged ions.An ionic structure consists of a giant lattice of oppositely charged ions.



5.6 Summary

Summary

• Metallic bonding is the electrostatic force of attraction between positivelycharged ions and delocalised outer electrons.

• A metallic structure consists of a giant lattice of positively charged ions anddelocalised outer electrons.

• Atoms in a covalent bond are held together by electrostatic forces ofattraction between positively charged nuclei and negatively charged sharedelectrons.

• A covalent molecular structure consists of discrete molecules held togetherby intermolecular forces.

• A covalent network structure consists of a giant lattice of covalently bondedatoms.

• Ionic bonding is the electrostatic force of attraction between positively andnegatively charged ions.

• An ionic structure consists of a giant lattice of oppositely charged ions.

• Elements can be categorised into four classes according to their bondingand structure.

– Metallic

– Covalent molecular

– Covalent network

– Monatomic


ISBN 978-1444167528




End of topic 5 test

This end of topic test is only available online.

Q34: In which of the following compounds do both ions have the same electronarrangement as argon?

a) Calcium sulfideb) Magnesium oxidec) Sodium sulfided) Calcium bromide

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Q35: Which element is a solid at room temperature and consists of discrete molecules?

a) Sulfurb) Neonc) Silicond) Boron

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Q36: Graphite, a form of carbon, conducts electricity because it has:

a) pure covalent bonding.b) delocalised electrons.c) metallic bonding.d) van der Waals' bonding.

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Q37: Which of the following can be applied to lithium but not to carbon?

a) Covalentb) Metallicc) Made up of discrete molecules.d) Made up of diatomic molecules.e) Gasf) Solid

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Q38:

Which of the following can be applied to both fluorine and phosphorus?

a) Covalent

b) Metallic

c) Made up of discrete molecules.

d) Made up of diatomic molecules.

e) Gas



f) Solid

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Q39: Identify the element which exists as a covalent network solid.

a) Boronb) Chlorinec) Nitrogend) Phosphoruse) Sodiumf) Sulfur

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Q40:

Identify the two elements which exist as discrete covalent molecular solids.

a) Boron

b) Chlorine

c) Nitrogen

d) Phosphorus

e) Sodium

f) Sulfur

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Q41:

Identify the two elements which react to form a compound with the most ionic character.

a) Boron

b) Chlorine

c) Nitrogen

d) Phosphorus

e) Sodium

f) Sulfur

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Q42: Which of the elements is most likely to have a covalent network structure?

ElementMelting Point(K)

Boiling Point(K)

Density (gcm-3)

Conductionwhen solids

A 317 553 1.82 NoB 933 2740 2.7 YesC 1683 2628 2.32 NoD 387 457 4.93 No

a) Ab) Bc) Cd) D

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Q43: Which type of structure is found in a substance melting at 771 K which conductselectricity when molten, but not when solid?

a) Ionicb) Covalent (network structure)c) Metallicd) Covalent (discrete molecules)

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Q44: Identify the covalent network substance.

a) NH4Cl (s)b) CH3OH (l)c) C6H14 (l)d) SO2 (g)e) Na2CO3 (s)f) SiO2 (s)

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Q45:

Identify the two substances which are ionic.

a) NH4Cl (s)

b) CH3OH (l)

c) C6H14 (l)

d) SO2 (g)

e) Na2CO3 (s)

f) SiO2 (s)

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Q46:

Diamond and graphite are forms of carbon with very different properties. Graphite canmark paper, is a lubricant and is a conductor of electricity. Diamond has none of theseproperties.

Draw a diagram to show the structure of diamond.

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Q47:

Diamond and graphite are forms of carbon with very different properties. Graphite canmark paper, is a lubricant and is a conductor of electricity. Diamond has none of theseproperties.

Why is graphite an effective lubricant?

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Q48:

Boron nitride can form a similar structure to graphite. The boron and nitrogen atomsalternate throughout the structure as shown.

Why is this substance a non-conductor, while graphite conducts electricity?

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Q49: Suggest why the bonds between the layers in boron nitride are stronger than thebonds between the layers in graphite.

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Q50: Which of the following structures is never found in compounds?

a) Ionicb) Monatomicc) Covalent Moleculard) Covalent Network

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117

Topic 6

Periodic Table trends

Contents

6.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 119

6.2 Covalent radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120

6.2.1 The trends in covalent radius . . . . . . . . . . . . . . . . . . . . . . . . 121

6.3 Ionisation energies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 123

6.3.1 Ionisation energies activity . . . . . . . . . . . . . . . . . . . . . . . . . 125

6.3.2 Explanation of the trends in first ionisation energy . . . . . . . . . . . . 125

6.4 Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 128

6.5 Summary of trends in the Periodic Table . . . . . . . . . . . . . . . . . . . . . . 130

6.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 130

6.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131

6.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132


Before you begin this topic, you should be able to state that:

• atoms contain protons, neutrons and electrons each with a specific charge, massand position within the atom. The number of protons defines an element and isknown as the atomic number (National 4, Unit 1);

• learners should have knowledge of: sub-atomic particles, their charge, mass andposition within the atom, the structure of the Periodic Table, groups, periods andatomic number (National 5, Unit 1);

• all matter is made of atoms. When a substance contains only one kind of atom itis known as an element (National 4, Unit 1);

• elements are arranged in the Periodic Table in order of increasing atomic number;elements with similar chemical properties are grouped together (National 4, Unit1);

• you should also be familiar with the seven diatomic elements (National 5, Unit 1);



118 TOPIC 6. PERIODIC TABLE TRENDS


• more than one bond can be formed between atoms leading to double and triplecovalent bonds (National 5, Unit 1);



• when there is an imbalance in the number of positive protons and electrons theparticle is known as an ion (National 5, Unit 1).

Learning Objectives

By the end of this topic, you should be able to state that:

• the covalent radius is a measure of the size of an atom;

• the trends in covalent radius across periods and down groups can be explained interms of the number of occupied shells, and the nuclear charge;

• the trends in ionisation energies across periods and down groups can be explainedin terms of the atomic size, nuclear charge and the screening effect due to innershell electrons;

• atoms of different elements have different attractions for bonding electrons;

• electronegativity is a measure of the attraction an atom involved in a bond has forthe electrons of the bond;

• electronegativity values increase across a period and decrease down a group;

• electronegativity trends can be rationalised in terms of nuclear charge, covalentradius and the presence of 'screening' inner electrons.


TOPIC 6. PERIODIC TABLE TRENDS 119

6.1 Prior knowledge

Prior knowledge: Periodic Table trends

Q1: Complete the following table:

Particle Charge Mass LocationProton

NeutronElectron

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Q2: Elements in the same group in the Periodic Table have the same:

a) number of occupied energy shells.b) density.c) number of outer electrons.d) number of protons.

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Q3: Elements in the same period in the Periodic Table have the same:

a) number of occupied energy shells.b) density.c) number of outer electrons.d) number of protons.

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Q4: What determines the number of protons in an atom?

a) Atomic numberb) Mass numberc) Name of elementd) Charge of atom

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a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) gaining electrons.

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6.2 Covalent radius

The size of individual atoms is difficult to measure since the size is determined by thespace taken up by the constantly moving electrons. However, the distance between thenuclei of atoms in the solid state can be measured by a technique called X-ray diffraction.Consequently, the size of an atom is usually described in terms of its covalent radius.This is defined as half the distance between the nuclei of two bonded atoms of theelement (see figure below).

Figure 6.1: Definition of covalent radius

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In the diagram above, the distance between the nuclei is shown as 2r and so the covalentradius in each case is r.

Covalent radius - relative sizes

The covalent radius is another periodic property (see below).

Figure 6.2: The covalent radius

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Q6: What happens to the size of the atoms on crossing a Period from left to right?

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Q7: What happens to the size of the atoms on descending a group?

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Q8: Can you suggest why are there no covalent radii quoted for the noble gases(Group 0)?

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6.2.1 The trends in covalent radius

The size of an atom is determined by the amount of space taken up by the electronsand so must be connected to the electron arrangement of the atom. The electronarrangement is itself a periodic property.

Use the following diagrams to explain the trends in covalent radius.

The trends in covalent radius

Figure 6.3: The trends in covalent radius



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Q9: Which of the following provides the best reason for the increase in covalent radiuson going down a group?

a) The number of protons increases.b) The number of electrons increases.c) The number of electron shells increases.d) The number of neutrons increases.

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Q10: Which of the following provides the best reason for the decrease in covalent radiuson going from left to right across a Period?

a) The number of electrons increases.b) The number of electron shells increases.c) The number of neutrons increases.d) The number of protons increases.

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Key Point

The covalent radius decreases across a Period because the increase in nuclearcharge attracts the electrons more strongly.The covalent radius increases on going down a group as the number of occupiedelectron shells increases.



6.3 Ionisation energies

The electrons orbiting the nucleus of an atom are held by the electrostatic attractionbetween the negative electrons and the positive nucleus. Some atoms lose electronsrelatively easily, whereas some lose electrons only with great difficulty. No atoms simplygive away electrons; to remove an electron always requires energy to overcome the forceof attraction between the electron and the nucleus. The electron lost always comes fromthe outer shell of electrons.

The energy required to remove an electron from an atom can be measured. It is normalto quote values not for an individual atom but for one mole of atoms.

Moles

At National 5 level, the mole was introduced as the gram formula mass of a substance(the formula mass expressed in grams). For an element, one mole is normally the gramformula mass, e.g.

• 12g of carbon is 1 mole

• 4g of helium is 1 mole

• 40g of calcium is 1 mole

These three quantities are different masses but they have one very important featurein common: they all contain the same number of atoms. This arises because calciumatoms are ten times as heavy as helium atoms while carbon atoms are three times asheavy as helium atoms.

This number is also known as one mole which refers to a particular number of particles(this will be investigated more fully in Unit 3). In fact, one mole of a substance containsa huge number of particles.

Ionisation energy

The first ionisation energy (IE) is defined as the energy required to remove one moleof electrons from one mole of gaseous atoms (one electron from each atom). The unitsused are kilojoules per mole (kJ mol-1).

Consider the element carbon. The value quoted for the 1st ionisation energy of carbonis 1090 kJ mol-1. In other words, 1090 kJ of energy is required to remove one electronfrom each atom in one mole (12 g) of gaseous carbon.

This can be represented in the following way:

• the first ionisation energy for an element E refers to the reaction

E(g) → E+(g) + e-;

• the second ionisation energy refers to the reaction

E+(g) → E2+(g) + e-.



Figure 6.4: First ionisation energy of carbon

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The C+ ion is smaller than the C atom because the remaining five electrons in the ionstill feel the full attractive force from six protons and so they are more tightly held. As aresult, it is even more difficult to remove the next electron.

Figure 6.5: Second ionisation energy of carbon

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The removal of one mole of electrons from one mole of gaseous C+ ions is known as thesecond ionisation energy of carbon. This has a value of 2360 kJ mole -1. The figurebelow shows equations representing the first four ionisation energies of carbon.

Figure 6.6: Successive ionisation energies

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6.3.1 Ionisation energies activity

Ionisation energies activity

Use values from the data booklet to plot a graph of first ionisation energy against atomicnumber for the first 20 elements.

Q11: Is the first ionisation energy a periodic property?

a) Yesb) No

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Q12: From the shape of the graph, explain your answer to the previous question.

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Q13: Name the group of elements which appears at the peaks in the graph.

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Q14: Name the group of elements which have the lowest ionisation energies.

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Q15: What is the general trend in first ionisation energy, going across a Period from leftto right?

a) A steady increase.b) A steady decrease.c) An increase followed by a decrease.d) A decrease followed by an increase.

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Q16: What is the general trend in first ionisation energy, going down a group?

a) A steady increase.b) A steady decrease.c) An increase followed by a decrease.d) A decrease followed by an increase.

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Q17: Confirm your answer to the previous question by listing the the first ionisationenergies of the alkali metals in order of atomic number.

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6.3.2 Explanation of the trends in first ionisation energy

Going down a group

In general, the first ionisation energy decreases on going down a group. For thisexercise we will concentrate on Group 1 since there is only one outer electron.

The higher the ionisation energy the more difficult it is to remove the electron, i.e. thestronger are the attractive forces between the electron and the nucleus. The strength ofthis force of attraction will depend on:



1. the size of the nuclear charge

2. the distance between the electron and the nucleus

3. the number of other electrons between the electron and the nucleus (i.e. thenumber of inner-shell electrons). The inner electrons cause a screening effectwhich prevents the outer electron from feeling the full effect of the nuclear charge.

First ionisation energy

Consider the electron arrangements of lithium, sodium and potassium (see figurebelow).

Figure 6.7: Electron arrangements of alkali metals

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Now discuss the following points with a partner, a group or even your tutor:

• What effect will an increase in nuclear charge have on the ionisation energy?

• What effect will an increase in covalent radius have on the ionisation energy?

• What will happen to the number of inner shell electrons on going down a group?

• What effect will this have on the ionisation energy?

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After discussion, you should be able to answer the following question. Write a detailedanswer on paper before revealing the display answer.

Q18: Explain fully why the first ionisation energy decreases on going down Group 1.You should mention nuclear charge, covalent radius and screening effect.

(There would be about three marks allocated to this type of "explain" question.)

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Going across a Period

The first ionisation energy increases on going from left to right across a Period. Thesame factors, as described above, will be involved in explaining this trend. Concentrateon Period 2 for this part.

Discuss the same points with a partner as before. Then answer the following question.

Q19: Explain fully why the first ionisation energy increases on going across Period 2from left to right.

(There would be about three marks allocated to this type of "explain" question.)

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Q20: What happens to the nuclear charge?

a) It decreases.b) It increases.c) It stays the same.

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Q21: What effect will this have on the first ionisation energy?

a) It increases.b) It decreases.c) It stays the same

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Q22: What happens to the number of inner shell electrons?


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Q23: What happens to the screening effect?


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Q24: As a result, what happens to the size of the atoms?


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• Use your data book to calculate the energy required to carry out the followingreactions; Al(g) ——> Al+(g) + e- Mg(g) ——> Mg2+(g) + 2e- -



• Why is there such a large increase in the energy required to remove a fourthelectron from aluminium compared to removing the first, second or third electrons?

Key Point

The first ionisation energy is the energy required to remove one mole of electronsfrom one mole of gaseous atoms. The second and subsequent ionisationenergies refer to the energies required to remove further moles of electrons.First ionisation energies increase across a Period and decrease down a group.This can be explained in terms of atomic size, nuclear charge and the screeningeffect due to inner shell electrons.

6.4 Electronegativity

The first ionisation energy involves the removal of electrons from gaseous atoms and sois a measure of how strongly an isolated atom holds on to its outermost electrons. In theworld around us, isolated atoms are very rare and the vast majority of atoms are foundbonded to one or more other atoms. How and why atoms bond together is the basis ofChemistry and electrons are the fundamental particles involved in bonding.

In all theories of bonding, different types of bond arise because atoms of differentelements have different attractions for electrons. Atoms which tend to attract theelectrons within a bond are said to be electronegative.

An atom with a high electronegativity will tend to attract bonded electrons towardsit whereas an atom with a low electronegativity will have a very weak attraction forelectrons. There will be a 'tug of war' between the different atoms for the electrons.

Electronegativity

In the figure below, the darker shading shows the electrons being pulled towards themore electronegative atom.

Figure 6.8: Electronegativity

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There are several different methods for estimating electronegativity values. The onemost commonly used is that devised by double Nobel Prize winner, Linus Pauling. Heassigned a value to each of the elements most commonly found in bonds. Lithiumwas assigned a value of 1.0 while fluorine was assigned the highest value of 4.0. Theelectronegativity values produced by Pauling are quoted in the SQA data booklet (page11). Study these values carefully and look for any patterns.

Q25: What is the group number of the elements which have the lowestelectronegativities? Please answer using an integer / integers.

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Q26: What is the group number of the elements with the highest electronegativities?Please answer using an integer / integers.

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Q27: Type the name of the group for which no values are quoted.

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Q28: Can you suggest a reason why no values are quoted for these elements?

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In general, the electronegativity increases from left to right across a period. This isbecause as you move from left to right, there are more protons and so the nuclear chargeis increased. As you move down a group, atomic size is increasing so outer electrons arefurther from the positively charged nucleus. As a result, the electronegativity decreases.

Q29: Which of the following statements describes the trends in electronegativity valuesin the Periodic Table?

a) Electronegativity values increase on going from left to right and increase on goingdown a group.

b) Electronegativity values decrease on going from left to right and decrease on goingdown a group.

c) Electronegativity values increase on going from left to right and decrease on goingdown a group.

d) Electronegativity values decrease on going from left to right and increase on goingdown a group.

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Key Point

Electronegativity is a measure of the attraction an atom in a bond has for theelectrons of the bond. Electronegativity values increase across a Period anddecrease down a group.



6.5 Summary of trends in the Periodic Table

Summary of trends in the Periodic Table

Q30: Complete the table by selecting the appropriate arrow symbol.

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6.6 Summary

Summary

• The covalent radius decreases across a Period because the increase innuclear charge attracts the electrons more strongly.

• The covalent radius increases on going down a group as the number ofoccupied electron shells increases.

• The first ionisation energy is the energy required to remove one mole ofelectrons from one mole of gaseous atoms.

• The second and subsequent ionisation energies refer to the energiesrequired to remove further moles of electrons.

• First ionisation energies increase across a Period and decrease down agroup.

• This can be explained in terms of atomic size, nuclear charge and thescreening effect due to inner shell electrons.

• Electronegativity is a measure of the attraction an atom in a bond has forthe electrons of the bond.

• Electronegativity values increase across a Period and decrease down agroup.



Summary Continued

• This can be explained in terms of atomic size, nuclear charge and thescreening effect due to inner shell electrons.

• Electronegativity is a measure of the attraction an atom in a bond has forthe electrons of the bond.

• Electronegativity values increase across a Period and decrease down agroup.

6.7 Resources

Text

• Higher Chemistry for CfE, J Anderson, E Allan and J Harris, Hodder Gibson,ISBN 978-1444167528




End of topic 6 test

Q31: Menedeleev is famous for producing the Periodic Table on which the modernversion is based. Which of the following statements is true?

a) Mendeleev organised the elements in order of their atomic number.b) Mendeleev swapped some elements round so that their atomic masses fitted the

pattern.c) Mendeleev left gaps for elements which had not yet been discovered.d) Mendeleev left gaps because some elements did not fit the pattern of reactivity.

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Q32: Which of the following statements about the Periodic Table are true?

1. There is a steady decrease in density on going down Group 1.

2. There is an increase in electronegativity on going down Group 7.

3. There is a decrease in first ionisation energy on going down Group 0.

4. There is a decrease and then an increase in boiling point on crossing a period fromleft to right.

5. There is a steady increase in melting point across a period from left to right.

6. There is a steady decrease in atomic size across a period from left to right.

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Q33: Which property of the Group 1 elements could be represented by the followinggraph?

Li Na K Rb

Arbitrary Scale

a) Electronegativityb) Atomic sizec) Melting pointd) First ionisation energy

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Q34: Some information about four atoms A, B, C and D is presented below.

AB

CD

Covalent Radius

Number ofOccupiedElectronShells

A 3B 4C 3D 4

Which atom will have the largest nuclear charge?

a) Ab) Bc) Cd) D

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Q35: Calcium has a larger covalent radius than magnesium because calcium has:

a) a smaller first ionisation energy?b) a larger nucleus?c) a larger nuclear charge?d) more occupied electron shells?

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Q36: The difference in atomic size between sodium and chlorine is mainly due to thenumber of:

a) electron shells?b) neutrons?c) protons?d) electrons?

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Q37: Which of the following equations represents the first ionisation energy ofmagnesium?

a)Mg(s) → Mg+(g) + e−

b)Mg(g) → Mg2+(g) + 2e−

c)Mg(g) → Mg+(g) + e−

d)1

2Mg(g) → 1

2Mg2+(g) + e−

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Q38: Which of the following equations represents the first ionisation energy of fluorine?

a)F− (g) → F (g) + e−

b)1

2F2 (g) → F+ (g) + e−

c)F (g) + e− → F− (g)

d)F (g) → F+ (g) + e−

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Q39: Identify the statements which correctly describe a trend in the Periodic Table.

1. The covalent radius increases from Li to F.

2. The boiling point increases from Li to F.

3. The first ionisation energy decreases from Na to Cl.

4. The first ionisation energy decreases from Li to Cs.

5. Electronegativity decreases from Li to Cs.

6. The number of electrons in the outer shell increases from Li to Cs.

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Q40: The change in first ionisation energy from Li to F is mainly due to:

a) increasing number of electron shells?b) increased screening effect?c) increasing number of electrons?d) increasing nuclear charge?

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Q41: The bar graph shows the variation in the first ionisation energy with atomic numberfor sixteen consecutive elements in the Periodic Table. The element at which the bargraph starts is not specified.

Atomic Number

ZFirst Ionisation

Energy(kJ mol-1)

Element Z is identified in the bar graph. In which group of the table is element Z?

a) 1b) 3c) 5d) 6

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Q42: Which of the following graphs shows the variation in first ionisation energy ongoing from left to right across period 2?

a)

Atomic Numberb)

Atomic Number



c)

Atomic Numberd)

Atomic Number. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q43: Which of the following will not affect the first ionisation energy of an element?

a) Screening effectb) Atomic sizec) Atomic numberd) Atomic mass

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Q44: �� is a measure of the ability of an atom in a bond to attract thebonding electrons.

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Q45: �� has the highest attraction for bonding electrons.

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Q46: Atoms of different elements have different attractions for the electrons in a bond.This property shows periodic variation.

Which of the following shows the correct trends within the Periodic Table?

a)



b)

c)

d)

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Q47: Explain fully why the first ionisation energy increases on going across Period 2from left to right. Write a detailed answer on paper before revealing the display answer.(3 marks)

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Q48: Explain fully why the first ionisation energy decreases on going down Group 1.You should mention nuclear charge, covalent radius and screening effect. (3 marks)

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Q49: Why are no electronegativity values are quoted for the Group 0 elements?

a) They have a value of zero.b) They generally do not form bonds with other elements.c) They are unreactive non-metals.d) Their electronegativities are too high for the scale.

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139

Topic 7

Bonding continuum and polarcovalent bonding

Contents

7.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 141

7.2 Polar covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 141

7.3 Predicting bonding type using electronegativity . . . . . . . . . . . . . . . . . . 144

7.4 Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 146

7.4.1 Polar molecules and molecular polarity . . . . . . . . . . . . . . . . . . 148

7.5 The bonding continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151

7.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 153

7.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 153

7.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 154


Before you begin this topic, you should know that:






• the first 20 elements in the Periodic Table can be categorised according to bondingand structure:

– metallic (Li, Be, Na, Mg, Al, K, Ca)

– covalent molecular (H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60))

140 TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING

– covalent network (B, C (diamond, graphite), Si) monatomic (noble gases)

• electronegativity is a measure of the attraction an atom in a bond has for theelectrons of the bond;

• electronegativity values increase across a Period and decrease down a Group(Higher, Unit 1);

• this can be explained in terms of atomic size, nuclear charge and the screeningeffect due to inner shell electrons (Higher, Unit 1);

• electronegativity is a measure of the attraction an atom in a bond has for theelectrons of the bond (Higher, Unit 1);

• electronegativity values increase across a Period and decrease down a Group(Higher, Unit 1).

Learning Objectives

By the end of this topic, you should be able to state that:

• in a covalent bond, atoms share pairs of electrons;

• the covalent bond is a result of two positive nuclei being held together by theircommon attraction for the shared pair of electrons;

• polar covalent bonds are formed when the attraction of the atoms for the pair ofbonding electrons is different;

• delta positive and delta negative notation can be used to indicate the partialcharges on atoms, which give rise to a dipole;

• pure covalent bonding and ionic bonding can be considered as being at oppositeends of a bonding continuum with polar covalent bonding lying between these twoextremes;

• the larger the difference in electronegativities between bonded atoms, the morepolar the bond will be;

• if the difference is large then the movement of bonding electrons from the elementof lower electronegativity to the element of higher electronegativity is completeresulting in the formation of ions;

• compounds formed between metals and non-metals are often, but not alwaysionic.


TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING 141

7.1 Prior knowledge

Prior knowledge: bonding continuum and polar covalent bonding


a) transfer of electrons.b) a shared pair of electrons.c) delocalised electrons.d) gaining electrons.

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Q2: Why are no electronegativity values quoted for the group 0 elements?

a) They have a value of zero.b) They are unreactive non-metals.c) They generally do not form bonds with other elements.d) Their electronegativities are too high for the scale.

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Q3: Which of the following elements has the greatest attraction for the shared pair ofelectrons in a bond?

a) Fluorineb) Carbonc) Hydrogend) Chlorine

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Q4: Which of the following elements has the least attraction for the shared pair ofelectrons in a bond?

a) Fluorineb) Carbonc) Hydrogend) Chlorine

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7.2 Polar covalent bonds

Previously, you saw how the electronegativity of atoms is a measure of their attractionfor electrons in a bond. Only when the two atoms bonded by a covalent bond are thesame (e.g. Cl2 in chlorine gas, C-C in diamond) will the electrons be exactly sharedequally. In other cases there will be an unequal distribution of the electrons, the atomwith the higher electronegativity will have a greater share of the electrons.

Since electrons carry a negative charge, an unequal distribution will result in the bondhaving a partial negative charge (called "delta negative" and drawn δ-) where there is anexcess of electrons around the most electronegative atom, and a partial positive charge



(δ+) where there is a deficiency. This is not to be confused with a full charge as foundon ions.

Figure 7.1: Polar bond

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In the case of HCl (see figure above), chlorine, with electronegativity 3.0, becomesnegative and hydrogen, at 2.2, positive. The δ partial charge is about 0.17 of a fullcharge.

This type of bond is called a polar covalent bond (sometimes abbreviated to a polarbond). The greater the difference between the electronegativities of the atoms, thegreater the distortion of the electrons in the bond, and the greater the chargedistribution. This effect is shown below.

Electronegativity

Figure 7.2: Electronegativity and charge

N.B. The electron cloud in this image shows the situation with X more electronegativethan Y.



YX YX YX

The blue shading is intended to show where electrons are most likely to be found.The darker the colour the higher the probability of finding an electron.

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Use the table of electronegativities, where necessary, (SQA data booklet, page 11) toanswer the following questions.

Q5: Which molecule contains polar bonds?

a) Hydrogen (H2)b) Chlorine (Cl2)c) Hydrogen chloride (HCl)d) Nitrogen (N2)

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Q6: Which molecule contains the most polar bonds?

a) Hydrogen fluoride (HF)b) Hydrogen chloride (HCl)c) Hydrogen bromide (HBr)d) Hydrogen iodide (HI)

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Q7: Which molecule has the bonds that are most polar?

a) Hydrogen sulfide (H2S)b) Phosphine (PH3)c) Methane (CH4)d) Ammonia (NH3)

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The symbol can be used to show the direction of the dipole, the arrowpointing to the negative side.

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Direction of charge

Use the electronegativity values (page 11 of the data booklet) to help.

Q8: Complete the table by selecting the correct sign for the bond shown.

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Key Point

Polar covalent bonds occur when the atoms of the bond attract the bondingelectrons unequally causing the atoms to have partial positive and negativecharges.The polarity of a covalent bond depends on the difference in electronegativitybetween the bonded atoms, the most electronegative becoming more negative.

7.3 Predicting bonding type using electronegativity

One of the main factors determining the type of bond formed between two elements ina compound is the difference in electronegativity. As previously discussed,electronegativity is a periodic property, with increasing electronegativity as you moveacross a period from left to right, and decreasing electronegativity as you move down agroup.

The relationship between electronegativity and bond type is summarised below.



Figure 7.3: Relationship between electronegativity and bond type

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Electronegativity difference is only one predictor of bond type. For a more completedescription, the properties of the substances, as described in another topic, need to beconsidered.

Most covalent bonded compounds generally exist as discrete molecular structures.However, two important covalent network compounds are silicon dioxide (SiO2, ) andsilicon carbide (SiC, carborundum), which has a similar structure to diamond. Both ofthese are used as abrasives on account of their hardness.

Use the electronegativity values in the SQA data booklet to help you answer thesequestions.

Q9: Hydrogen, bromine and potassium have electronegativities of 2.2, 2.8 and 0.8respectively. What type of bonding would you expect:

• between two bromine atoms in a Br2 molecule?

• between hydrogen and bromine in HBr?

• between potassium and bromine in KBr?

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Q10: What type of bonding would you expect between the Group 1 metal caesium andthe Group 6 element sulfur?

a) Pure covalentb) Polar covalentc) Ionicd) Metallic

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Q11: What type of bonding would you expect between phosphorus and hydrogen inphosphine (PH3) ?

a) Pure covalentb) Polar covalentc) Ionicd) Metallic

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Q12: Which of the elements below will bond with oxygen, to produce a polar bond witha partial positive charge on the oxygen atom?

a) Carbonb) Fluorinec) Hydrogend) Lithium

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Key Point

The type of bonding between two atoms depends mainly on the difference inelectronegativity between the atoms.

When the difference is zero, the bond will be covalent. With a small difference, apolar covalent bond is likely. When the difference is large, fully charged ions areproduced and ionic bonding will be predicted.

7.4 Polar molecules

When two atoms of different electronegativity are bonded together by sharing electrons,one atom attracts the electrons more than the other and a polar bond results. In simplemolecules like hydrogen chloride and iodine chloride this polar bond has a permanentdipole.

Polar molecules are attracted to one another by forces called permanentdipole-permanent dipole interactions (see figure below) as well as theLondon dispersion attractions caused by the movement of electrons observed in thelast section.



Permanent dipole interactions

Figure 7.4: Permanent dipoles in iodine chloride

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How do permanent dipole-permanent dipole interactions compare to Londondispersion forces of attraction in terms of strength?

A comparison of strengths can be made by looking at the table below.

Figure 7.5: Boiling points of halogen containing compounds

A B CCl - Cl I - Cl Br - Br-35oC 97oC 59oC

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Any comparison of boiling points has to be between molecules of similar size andshape so that the London dispersion forces are similar.

Q13: Which molecule (A, B, or C) has permanent dipole-permanent dipole interactions?

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Q14: Which other molecule would have almost the same London dispersion forces?

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Q15: Which of the molecules has the strongest intermolecular forces?

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Key Point

Permanent dipole-permanent dipole interactions act in addition to Londondispersion electrostatic attractions between polar molecules and are strongerthan these attractions for molecules of equivalent size.

7.4.1 Polar molecules and molecular polarity

Not all covalent molecules with polar bonds result in polar molecules. Those moleculeswhich are highly symmetrical prove to be non-polar (see figure below).

Figure 7.6: Non-polar molecules

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In each of the two molecules shown, although there is a difference in electronegativityin the polar bonds, the charge is distributed around the central carbon atoms with thepositive and negative charges balancing out. The molecules have no overall dipole.

Notice that the central carbon on each of the molecules is shown with only one (δ+)sign for clarity, even though each atom has sufficient charge to balance out thenegative charge.

In molecules with polar bonds which are not symmetrical (see figure below) the dipolescannot cancel each other out. If the molecule has a permanent slight positive charge atone side and a negative charge at the other then it is a polar molecule. Bond polaritycan thus be predicted from electronegativity differences, and molecular polarity can bepredicted from electronegativity differences if we also take into account the shape ofthe molecule.

The symbol can be used to show the direction of the dipole, the arrowpointing to the negative side. The molecules shown in the figure below are both polarmolecules.



Figure 7.7: Polarity of molecules

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Predicting molecular polarity

Electronegativity values are used to predict the polarity of bonds, and the shapes ofmolecules are used to predict molecular polarity. You may find page 11 in the databooklet helpful.

Q16: Draw out this table, or photocopy the page. Complete the table by predicting thepolarities of the molecules and drawing the items in the word bank into the blank spaces.The dipole symbols can be repeated.

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Detecting polar molecules

Look at the picture showing the results of the experiment and answer the questions.

Figure 7.8

Water Paraffin

ChargedRods

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Q17: Which liquid is affected by the charged rod?

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Q18: Which liquid is polar?

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Q19: Which of these statements is correct when a jet of polar molecules passes acharged rod?

a) It is not affected by the rod.b) It is attracted by the rod.c) It is repelled by the rod.d) It is speeded up by the rod.

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Q20: Which of these statements is correct when a jet of non-polar molecules passes acharged rod?

a) It is not affected by the rod.b) It is attracted by the rod.c) It is repelled by the rod.d) It is speeded up by the rod.

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Key Point

An electrostatically charged rod can be used to detect the presence of polarmolecules in a liquid. Polar molecules are attracted to both a negative and positiverod.

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7.5 The bonding continuum

Covalent bonding usually involves elements close to each other in the Periodic Table.Ionic bonding generally involves metals from the left side, and non-metals from theopposite side of the table. Most chemical bonds are somewhere between the twoextremes and it is best to think of ionic and covalent bonding as being at opposite endsof a bonding continuum with varying degrees of polar covalent bonding lying betweenthe extremes (see figure below).

Figure 7.9: Bonding spectrum

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In most cases the bigger the difference in electronegativity between the atoms, themore polar the bond and the greater the ionic character. However, other factors make acontribution and care needs to be taken before jumping to conclusions. Identical atomsshare the electrons in a covalent bond equally. In the hydrogen molecule there is nocharge distribution. Non-identical atoms like hydrogen and chlorine attract the bondingelectrons unequally, since chlorine attracts electrons more strongly. A polar bondresults (see figure below).



Figure 7.10: Polar bond

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In some cases, the bond polarity results in a polar molecule. In certain symmetricalmolecules the polar nature of the bonds tends to cancel out (see figure below).

Figure 7.11: Polar and non-polar molecules

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In water, the molecule is asymmetrical, so the δ+ on the hydrogen atoms will add andthe molecule will have a resultant charge distribution; it will be a polar molecule.

The tetrachloromethane molecule is highly symmetrical, so the charge distribution onthe C-Cl bonds will cancel out and the molecule will be non-polar, despite having polarbonds.

Q21: Will trichloromethane (CHCl3, chloroform) be a polar or non-polar molecule?

a) Polarb) Non-polar

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7.6 Summary

Summary

• Polar covalent bonds occur when the atoms of the bond attract the bondingelectrons unequally causing the atoms to have partial positive and negativecharges.

• The polarity of a covalent bond depends on the difference inelectronegativity between the bonded atoms, the most electronegativebecoming more negative.

• Between pure covalent and pure ionic bonds there are polar covalent bonds.

• The type of bonding between two atoms depends mainly on the differencein electronegativity between the atoms.

– When the difference is zero, the bond will be covalent.

– With a small difference, a polar covalent bond is likely.

– When the difference is large, fully charged ions are produced and ionicbonding will be predicted.

• Permanent dipole-permanent dipole interactions act in addition to Londondispersion electrostatic attractions between polar molecules and arestronger than these attractions for molecules of equivalent size.

• Not all covalent molecules with polar bonds result in polar molecules.

• Molecules which are highly symmetrical tend to be non-polar.

• An electrostatically charged rod can be used to detect the presence of polarmolecules in a liquid. Polar molecules are attracted to both a negative andpositive rod.

• There is a complete range of bond types leading to a bonding spectrummainly based on electronegativity.

7.7 Resources• Higher Chemistry for CfE, J Anderson, E Allan and J Harris, Hodder Gibson,

ISBN 978-1444167528




End of topic 7 test

Q22: In which of the following compounds do both ions have the same electronarrangement as argon?

a) Calcium bromideb) Magnesium oxidec) Sodium sulfided) Calcium sulfide

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Q23: Which element is a solid at room temperature and consists of discrete molecules?

a) Sulfurb) Carbonc) Silicond) Boron

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Q24: Graphite, a form of carbon, conducts electricity because it has:

a) metallic bonding?b) pure covalent bonding?c) delocalised electrons?d) London dispersion forces?

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Q25: Which of the following can be applied to lithium but not to carbon?

a) Covalentb) Metallicc) Made up of discrete moleculesd) Made up of diatomic moleculese) Gasf) Solid

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Q26: Which of the following can be applied to both fluorine and phosphorus?

1. Covalent

2. Metallic

3. Made up of discrete molecules

4. Made up of diatomic molecules

5. Gas

6. Solid

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a) Boronb) Chlorinec) Nitrogend) Phosphoruse) Sodiumf) Sulfur

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Q28: Identify the two elements which exist as discrete covalent molecular solids.

1. Boron

2. Chlorine

3. Nitrogen

4. Phosphorus

5. Sodium

6. Sulfur

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1. Boron

2. Chlorine

3. Nitrogen

4. Phosphorus

5. Sodium

6. Sulfur

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Q30: In which molecule will the chlorine atom carry a partial positive charge (δ+)?

a) Cl-Fb) Cl-Brc) Cl-Cld) Cl-I

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Q31: Which of the elements is most likely to have a covalent network structure?

Element Melting point(K)

Boiling point(K)

Density ( gcm-3)

Conductionwhen solid

A 317 553 1.82 NoB 933 2740 2.7 YesC 1683 2628 2.32 NoD 387 457 4.93 No

a) Ab) Bc) Cd) D

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Q32: Which of the following chlorides is likely to have least ionic character?

a) LiClb) CsClc) BeCl2d) CaCl2

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Q33: Which type of structure is found in a substance melting at 771 K which conductselectricity when molten, but not when solid?

a) Covalent (network structure)b) Covalent (discrete molecules)c) Ionicd) Metallic

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Q35: Identify the two substances which are ionic.

1. NH4Cl (s)

2. CH3OH (l)

3. C6H14 (l)

4. SO2 (g)

5. Na2CO3 (s)



6. SiO2 (s)

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159

Topic 8

Intermolecular forces

Contents

8.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 162

8.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 162

8.3 London dispersion forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163

8.4 Hydrogen bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 168

8.5 Relating properties to bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . 171

8.5.1 Boiling point . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172

8.5.2 Density . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 173

8.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174

8.7 Predicting solubilities from solute and solvent polarities . . . . . . . . . . . . . 176

8.8 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 180

8.9 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 181

8.10 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 182



• the first 20 elements in the Periodic Table can be categorised according to bondingand structure:

– metallic Li, Be, Na, Mg, Al, K, Ca;

– covalent molecular H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60);

– covalent network B, C (diamond, graphite), Si;

– monatomic (noble gases).

• electronegativity is a measure of the attraction an atom in a bond has for theelectrons of the bond;


• this can be explained in terms of atomic size, nuclear charge and the screeningeffect due to inner shell electrons (Higher, Unit 1);

• electronegativity is a measure of the attraction an atom in a bond has for theelectrons of the bond (Higher, Unit 1);

160 TOPIC 8. INTERMOLECULAR FORCES


• in a covalent bond, atoms share pairs of electrons (National 4, Unit 1);

• the covalent bond is a result of two positive nuclei being held together by theircommon attraction for the shared pair of electrons (National 4, Unit 1);

• polar covalent bonds are formed when the attraction of the atoms for the pair ofbonding electrons is different (Higher, Unit 1);

• delta positive and delta negative notation can be used to indicate the partialcharges on atoms, which give rise to a dipole (Higher, Unit 1);

• pure covalent bonding and ionic bonding can be considered as being at oppositeends of a bonding continuum with polar covalent bonding lying between these twoextremes (Higher, Unit 1);

• the larger the difference in electronegativities between bonded atoms, the morepolar the bond will be (Higher, Unit 1).

Learning Objectives

At the end of this topic, you should know that:

• all molecular elements and compounds and monatomic elements condense andfreeze at sufficiently low temperatures. For this to occur, some attractive forcesmust exist between the molecules or discrete atoms;

• any ‘intermolecular’ forces acting between molecules are known as van der Waals’forces;

• there are several different types of van der Waals’ forces such as Londondispersion forces and permanent dipole: permanent dipole interactions, whichinclude hydrogen bonding;

• London dispersion forces are forces of attraction that can operate between allatoms and molecules;

• these forces are much weaker than all other types of bonding;

• London dispersion forces are formed as a result of electrostatic attraction betweentemporary dipoles and induced dipoles caused by movement of electrons in atomsand molecules;

• the strength of London dispersion forces is related to the number of electronswithin an atom or molecule;

• a molecule is described as polar if it has a permanent dipole. The spatialarrangement of polar covalent bonds can result in a molecule being polar;

• permanent dipole-permanent dipole interactions are additional electrostatic forcesof attraction between polar molecules;

• permanent dipole-permanent dipole interactions are stronger than Londondispersion forces for molecules with similar numbers of electrons;


TOPIC 8. INTERMOLECULAR FORCES 161

• bonds consisting of a hydrogen atom bonded to an atom of a stronglyelectronegative element such fluorine, oxygen or nitrogen are highly polar;

• hydrogen bonds are electrostatic forces of attraction between molecules whichcontain these highly polar bonds;

• a hydrogen bond is stronger than other forms of permanent dipole-permanentdipole interaction but weaker than a covalent bond;

• melting points, boiling points and viscosity can all be rationalised in terms of thenature and strength of the intermolecular forces which exist between molecules;

• by considering the polarity and number of electrons present in molecules, it ispossible to make qualitative predictions of the strength of the intermolecular forces;

• the melting and boiling points of polar substances are higher than the melting andboiling points of non-polar substances with similar numbers of electrons;

• the anomalous boiling points of ammonia, water and hydrogen fluoride are a resultof hydrogen bonding;

• boiling points, melting points, viscosity and solubility/miscibility in water areproperties of substances which are affected by hydrogen bonding;

• hydrogen bonding between molecules in ice results in an expanded structurewhich causes the density of ice to be less than that of water at low temperatures.



8.1 Prior knowledge


Q1: Which of these is a measure of the ability of an atom in a bond to attract thebonding electrons?

a) Ionisation Energyb) Periodicityc) Electronegativityd) Polarity

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Q2: Which of the following elements has the greatest attraction for the shared pair ofelectrons in a bond?

a) Fluorineb) Nitrogenc) Phosphorusd) Lithium

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Q3: As you move across the Periodic Table, the electronegativity:

a) stays the same.b) decreases.c) increases.d) decreases then increases.

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8.2 Introduction

Simple molecules like nitrogen N2, methane CH4 and water H2O, have their atoms heldtogether by covalent bonds within the molecule.

Bonds within a molecule are called intramolecular forces (intra: on the inside, as inintramuscular, intravenous).

Nitrogen and methane are both gases at room temperature, but can be made liquid bycooling to a very low temperature. Water is a liquid at room temperature and othercovalent molecules like candle wax are solid. Forces of attraction must exist betweenall these covalent molecules, or it would never be possible to have covalent molecularsolids at all. The molecules would fly apart and all discrete covalent molecules wouldbe gases even at the lowest temperatures possible.

The forces between molecules are called intermolecular forces (inter: between, as ininter-city, international).



This diagram shows both intermolecular and intramolecular forces in a sample of water.

Figure 8.1

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The way that the charge is distributed in a water molecule is not symmetrical. There ismore negative charge on one side than the other. It is said to have a dipole. TheGreek letter delta, δ, is used to show a "small amount".

Van der Waals' forces

Johannes Diderik van der Waals recognised that relatively weak were responsible forthe change of state from gas to liquid. He was awarded the Nobel prize for his work in1910.

There are several different forms of these forces and these will be discussed in thefollowing sections.

• London dispersion forces.

• Permanent dipole-permanent dipole interactions.

• Hydrogen bonding.

In this topic the forces of attraction caused by dipoles will be explored.

8.3 London dispersion forces

The strength of the between covalent substances can be estimated by looking at theboiling points. The lower the boiling point the weaker the forces must be, since it is thewhich have to be broken to change the liquid into gas.



Water

Steam

Heat energy added

Liquid to gas

The covalent diatomic elements like hydrogen, H2 gas and bromine, Br2, which is aliquid at room temperature, are completely non polar and would seem to have no dipoleat all. As well as these, even monatomic gases like helium and neon can be liquified.

How can intermolecular forces exist in these substances?

The work of Fritz London (in his 1930 paper) led to suggestions that an atom ormolecule could have a temporary dipole at a particular instant in time and that, if therewere other atoms or molecules nearby, this temporary dipole might affect them andinduce a dipole in the nearby atom or molecule. The resultant electrostatic attractionbetween the temporary dipole and the induced dipole might, although very weak, beable to hold the substance together.These forces of attraction are known as London dispersion forces.

Induced dipoles

The online version of this activity contains a simulation of the formation of temporarydipoles and induced dipoles in covalent substances.

This picture shows how the temporary dipole in a helium atom can exist at a singlebrief instant in time. This can induce a dipole in its neighbour. Hydrogen molecules areshown forming the same temporary dipole-induced dipole pair. The attractionsbetween the dipole pairs form the same type of bonds; London dispersion forces.



Temporary dipoles and induced dipoles

Q4: In the single brief instant in time when the picture was taken, and the electronmovement was paused the electron distribution is:

a) evenly spread.b) unevenly spread.

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Q5: This distribution of electrons causes

a) a permanent dipole.b) a temporary dipole.

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Q6: What effect does this have on the neighbouring particle?

a) The permanent dipole causes an induced dipole.b) The temporary dipole causes an induced dipole.c) The permanent dipole induces a temporary dipole.d) The temporary dipole induces a permanent dipole.

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Q7: The London dispersion forces of attraction formed can be said to be:

a) intramolecularb) intermolecular

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London dispersion forces of attraction can operate between all atoms and molecules.They are much weaker than other types of bonding, at around 2 kJ mol -1 compared tocovalent bonds of between 150-500 kJ mol-1. In many substances London dispersionforces are not large enough to influence the behaviour of substances. If they are theonly type of intermolecular force present however, as in the noble gases, they must beconsidered. One example of this relates the London dispersion forces to the size ofatoms or molecules.



This diagram shows the relative sizes of atoms and boiling points (◦C) of the first fournoble gases.

Noble gas boiling points

Q8: Name the noble gas shown which has the largest atom.

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Q9: Which noble gas shown has the highest boiling point?

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Q10: Which noble gas shown must have the strongest London dispersion forces holdingthe atoms together?

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Q11: Look at the relationship between the London dispersion forces and size. Which ofthese statements is true?

a) The size of the atom doesn't affect the London dispersion forces.b) The larger the atom the weaker the London dispersion forces.c) The larger the atom the stronger the London dispersion forces.d) The smaller the atom the stronger the London dispersion forces.

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The lowest temperature it is possible to reach is called Absolute Zero. This occurs at-273◦C.

Solid helium is changed into liquid helium at approximately -272◦C and liquid helium togas at -269◦C, both by breaking London dispersion forces of attraction.

Q12: What does this suggest about the strength of London dispersion forces?

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Molecular substances which have only London dispersion forces between themolecules i.e. intermolecular, also show a relationship between size and strength ofthe London dispersion forces. This table shows the boiling points of the first four alkanemolecules.



Name Formula Boiling point (◦C)

methane CH4 -164

ethane C2H6 -89propane C3H8 -42

butane C4H10 -1

Alkane boiling points

Q13:

Describe the relationship between size and boiling point and explain fully why thisoccurs. Write your answer on paper before consulting the answer at the back.

(There would be about three or four marks allocated to this type of "describe" and"explain" question and these require practice.) If you are in any way unsure about beingable to explain your answer fully, work through the next four questions before attemptingthe complete answer in the repeated question. If you feel confident, check your answerwith the display answer.

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Q14: Name the alkane shown which has the largest molecules.

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Q15: Name the alkane shown which has the highest boiling point.

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Q16: Which alkane shown must have the strongest London dispersion forces holdingthe molecules together?

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Q17: Look at the relationship between the London dispersion forces and size. Which ofthese statements is true?

a) The size of the molecule doesn't affect the London dispersion forces.b) The larger the molecule the weaker the London dispersion forces.c) The larger the molecule the stronger the London dispersion forces.d) The smaller the molecule the stronger the London dispersion forces.

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Q18:

Now try again to use these answers to write a response which would be worth threemarks. Here is the question again.

Describe the relationship between size and boiling point and explain fully why thisoccurs. Write your answer on paper before consulting the answer at the back.

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Key Point

London dispersion forces of attraction operate between all atoms and moleculesand are weaker than all other types of bonding.The strength of London dispersion forces is related to the size of the atoms ormolecules.

8.4 Hydrogen bonding

Hydrogen bonding is the name given to an intermolecular force which is actually a typeof permanent dipole-permanent dipole attraction. These hydrogen bonds areseparated from other examples of this type of interaction because they are unusuallystrong (about 15-25 kJ mol-1) as shown below.

For hydrogen bonding to occur the hydrogen atom involved needs to:

• be the positive end of a strong dipole (estimated from the difference inelectronegativity).

• have a small, highly electronegative atom on a neighbouring molecule.

Only the three elements fluorine, oxygen and nitrogen are considered able to satisfythese conditions.

Hydrogen bonding elements

Notice that the bonds called hydrogen bonds are the forces of electrostatic attractionBETWEEN the molecules which contain these highly polar bonds. i.e. the hydrogenbonds are the intermolecular forces.



Key Point

Bonds consisting of a hydrogen atom bonded to an atom of a stronglyelectronegative element such as fluorine, oxygen or nitrogen are highly polar andthe electrostatic attractions between molecules which contain these highly polarbonds are called hydrogen bonds.

Strength of hydrogen bonds

The online version of this activity contains an illustration of the formation of hydrogenbonds in a sample of water.

This picture shows the formation of hydrogen bonds in a sample of water, and thediagram below shows the boiling points and relative sizes of the first four hydrides ofthe Group 6 elements.

Strength of hydrogen bonds

A comparison of the strength of hydrogen bonds with other permanentdipole-permanent dipole interactions and London dispersion forces can be made byconsidering the data in the diagram below.

Boiling points of the Group 6 hydrides



Remember that London dispersion forces of attraction will operate between all atomsand molecules, and some of the hydride molecules will have permanentdipole-permanent dipole attractions since they are polar molecules.

There would be three or four marks allocated to the next "compare" and "explain"question using the evidence in the diagram. Try writing an answer on paper beforeconsulting the answer at the back.

Q19: Use the boiling point data to compare the intermolecular forces present in waterwith those present in the other Group 6 hydrides and explain fully how hydrogen bondscompare in strength to others present.

If you are unsure about being able to explain your answer, work your way through thenext four questions below.

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Q20: Name the type of weak intermolecular force of attraction which increases as sizeincreases.

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Q21: Name the substance which shows boiling point evidence which goes against thistrend.

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Q22: Name the substance which has the strongest intermolecular force of attractionholding the molecules together.

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Q23: How do the hydrogen bonds in water compare in strength to the Londondispersion forces present and the other permanent dipole-permanent dipoleinteractions present in some of the other molecules?

a) The hydrogen bonds are weaker than London dispersion forces and the otherpermanent dipole-permanent dipole interactions.

b) The hydrogen bonds are stronger than London dispersion forces and the otherpermanent dipole-permanent dipole interactions.

c) The hydrogen bonds are stronger than London dispersion forces but weaker thanthe other permanent dipole-permanent dipole interactions.

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Q24:

Now try again to use these answers to write a response. Here is the question again.

Use the boiling point data to compare the intermolecular forces present in water withthose present in the other Group 6 hydrides and explain fully how hydrogen bondscompare in strength to others present.

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A comparison of the strength of hydrogen bonds with intramolecular bonds like thecovalent bond between hydrogen and oxygen can be made by remembering that thehydrogen bond has a value of about 15-25 kJ mol-1 and the strength of the covalentO-H bond can be found in the data booklet (page 9). The table of "Mean BondEnthalpies" gives the value.

Q25: Which statement is true about hydrogen bonds in comparison to covalent bonds?

a) Covalent bonds are much stronger than hydrogen bonds.b) Hydrogen bonds are much stronger than covalent bonds.c) Covalent bonds are the same strength as hydrogen bonds.d) Hydrogen bonds are twice as strong as covalent bonds.

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Key Point

A hydrogen bond is stronger than other forms of permanent dipole-dipoleinteraction but weaker than a covalent bond.

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8.5 Relating properties to bonding

Intermolecular bonds are the relatively weak bonds which attract molecules to eachother. They are not as strong as the bonds, for example covalent bonds, that bind theatoms together into a molecule.

Q26: Which of the following properties of a substance are a reflection of the betweenthe molecules of the substance?

• Melting point

• Boiling point

• Density

• Viscosity

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Water is extremely important for life on Earth; in fact the search for extraterrestrial lifeinvolves looking for water, or evidence that water was present at some time. Water is thecommonest liquid on Earth and also one of the most unusual! The next sections exploresome properties of water from the point of view of hydrogen bonding.



8.5.1 Boiling point

The figure below shows the boiling points of the hydrides along the periods for groups 4to 7 elements.

Boiling points of hydrides.

Hydrides of non-metals

Q27: The trend in boiling points in group 4 from CH4 to SnH4 shows a regular increasewith formula mass. Is there any evidence of polar attractions between molecules?

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Q28:

Three of the second period hydrides have higher boiling points than expected bycomparison with the trends seen for other elements. (Remember that higher formulamass usually means higher boiling point.)

Which are the three hydrides?

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Q29: What is the main force between molecules in group 4 hydrides, from CH4 to SnH4?

a) Polar-polar attractionsb) Van der Waals attractionsc) Covalent bondsd) Ionice) Hydrogen bonds

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Key Point

The higher-than-expected boiling points of ammonia, water and hydrogenfluoride result from additional intermolecular forces, in this case, hydrogenbonds.

8.5.2 Density

Most substances contract when they are cooled and solids are normally denser thantheir own liquids. This causes the majority of solids to sink when placed in their ownliquid. The structure of ice, however, is very unusual because the solid is less densethen the liquid water.

As water freezes, the intermolecular hydrogen bonding spreads out the watermolecules into a strong "open" structure with large spaces in it (see below). Thismakes ice less dense and able to float on water.

The arrangement of water molecules in ice maximises the hydrogen bonding betweenthem and leads to an open structure. As it melts the structure collapses into the spaces

and the liquid becomes less dense.

The on-line version of this topic contains a three-dimensional representation of thestructure of ice at this point.



The fact that ice floats on water is good news for fish living in ponds. The ice forms onthe surface and the fish and other aquatic life can exist below this layer. The layer of ice

thermally insulates the water beneath.

A disadvantage of water expanding as it forms ice is that water trapped in pipes over acold spell causes the pipes to crack open as it freezes, resulting in leaks when it thaws.

Key Point

Hydrogen bonding between molecules in ice results in an expanded structurewhich causes the density of ice to be less than that of water at low temperatures.

8.6 Viscosity

The stronger the intermolecular forces are between molecules in a liquid, the moreviscous the liquid. The viscosity is a measure of how easily it flows (how thick orsyrupy a liquid is). The molecules have to be able to move past each other to flow.



Testing viscosity

Figure 8.2

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Q30: In which liquid does the marble fall fastest?

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Q31: Which chemical structure has the least hydrogen bonding?

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Q32: Which chemical structure has the most hydrogen bonding?

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Q33: If the marble in methanol took 1.5 seconds to reach the bottom, calculate the rateof descent of the marble through methanol in m s-1.

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Q34: Predict the rate of descent of a marble in an identical experiment usingethoxyethane.

a) 0.8 m s-1

b) 0.6 m s-1

c) 0.4 m s-1

d) 0.1 m s-1

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Q35: Assuming your answer to the last question is correct, how long would it take amarble to travel through the ethoxyethane?

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Viscosity is a term which applies to fluids (both liquids and gases), and since these arecovalent substances, being able to move past each other involves breaking the Londondispersion forces, dipole-dipole interactions or hydrogen bonds. Of these, the hydrogenbonds are the strongest.

Key Point

The viscosity of a liquid is related to the strength of the intermolecular bonding.

8.7 Predicting solubilities from solute and solvent polarities

When you add a spoon of sugar to your cup of coffee and stir the sugar dissolves. Infact sugar will dissolve in water to a considerable extent to produce treacle or 'Goldensyrup'.

However, if you added a spoon of sugar to a cup of petrol no amount of stirring wouldmake it dissolve.

A polar solvent is one whose molecules exhibit strong permanent dipoles. Because ofthis, polar solvents like water can generally dissolve both polar substances, like sucrose,and ionic solids, for example sodium chloride. The expression often used to describethis is:



Key Point

LIKE DISSOLVES LIKE

In the same way, because the intermolecular attractions are of a similar type, non-polarsolvents are more likely to dissolve non-polar substances: LIKE DISSOLVES LIKE.For example wax dissolved in white spirit in a furniture polish.

If however a mixture of polar and non-polar substances is used, no dissolving occurs.Opposites do not dissolve.

Polar and ionic substances tend to be insoluble in non-polar solvents. Non-polarsubstances tend to be insoluble in polar solvents. The attractions are not sufficientlystrong to allow dissolving.

When an ionic compound dissolves in water, the ions need to be separated from thelattice. The polar water molecules can sometimes pull the ions into solution andsurround them.

Salt dissolving



The way the water molecules direct themselves depends on the charge the ion carries.The ions are said to be hydrated.

Dissolving of sodium chloride can be shown in an equation.

Dissolving sodium chloride

A polar molecule like hydrogen chloride dissolves in water as the water "interacts" withthe molecules.

Hydrogen chloride dissolving



Hydrogen chloride dissolving can also be shown in an equation.

Hydrogen chloride dissolving (equation).

In a similar way, a polar substance such as sucrose (C12H22O11) containing several -OHgroups, can bond with polar water molecules.

Q36: Ammonia (NH3) is a gas which is very soluble in water, whereas nitrogen (N2) isalmost insoluble. Can you explain why?

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Key Point

Ionic compounds and polar molecular compounds tend to be soluble in polarsolvents such as water and insoluble in non-polar solvents. Non-polar molecularsubstances tend to be soluble in non-polar solvents and insoluble in polarsolvents.

Miscibility

When two liquids mix thoroughly with no visible boundary between them, they are saidto be miscible e.g. water and methanol. Both of these molecules have hydrogenbonds shown in Figure 8.3(a).

Miscibility arises when the intermolecular attractions between two types of substancesare fairly similar. They are then able to mix easily.

If one substance has different intermolecular attractions from the other, its moleculesstay grouped around each other and form a separate layer, e.g. water (which is hydrogenbonded), and tetrachloromethane (which is non-polar) shown in Figure 8.3(b).

Figure 8.3: Miscible and immiscible liquids



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Whether two liquids are miscible or immiscible can only be determined by experiment,but the general rule of LIKE DISSOLVES LIKE can be useful in trying to make a roughprediction.

8.8 Summary

Summary

• All molecular elements and compounds and monatomic elements condenseand freeze at sufficiently low temperatures. For this to occur, some attractiveforces must exist between the molecules or discrete atoms.

• Intermolecular forces acting between molecules are known as van derWaals’ forces.

• There are several different types of van der Waals’ forces such as Londondispersion forces and permanent dipole: permanent dipole interactions,which include hydrogen bonding.

• London dispersion forces are forces of attraction that can operate betweenall atoms and molecules.

• These forces are much weaker than all other types of bonding.

• They are formed as a result of electrostatic attraction between temporarydipoles and induced dipoles caused by movement of electrons in atomsand molecules.

• The strength of London dispersion forces is related to the number ofelectrons within an atom or molecule.

• Bonds consisting of a hydrogen atom bonded to an atom of a stronglyelectronegative element such fluorine, oxygen or nitrogen are highly polar.

• Hydrogen bonds are electrostatic forces of attraction between moleculeswhich contain these highly polar bonds.

• A hydrogen bond is stronger than other forms of permanent dipole-permanent dipole interaction but weaker than a covalent bond.

• Melting points, boiling points and viscosity can all be rationalised in termsof the nature and strength of the intermolecular forces which exist betweenmolecules.

• By considering the polarity and number of electrons present in molecules,it is possible to make qualitative predictions of the strength of theintermolecular forces.



Summary Continued

• The melting and boiling points of polar substances are higher than themelting and boiling points of non-polar substances with similar numbers ofelectrons.

• The anomalous boiling points of ammonia, water and hydrogen fluoride area result of hydrogen bonding.

• Boiling points, melting points, viscosity and solubility/miscibility in water areproperties of substances which are affected by hydrogen bonding.

• Hydrogen bonding between molecules in ice results in an expandedstructure which causes the density of ice to be less than that of water atlow temperatures.

• Ionic compounds and polar molecular compounds tend to be soluble in polarsolvents such as water and insoluble in non-polar solvents.

• Non-polar molecular substances tend to be soluble in non-polar solventsand insoluble in polar solvents.


ISBN 978-1444167528




End of topic 8 test

This end of topic test is only available online.

Q37: London dispersion forces of attraction are a result of:

a) an induced dipole causing an temporary dipole.b) a permanent dipole causing an permanent dipole.c) an induced dipole causing an permanent dipole.d) a temporary dipole causing an induced dipole.

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Q38: Which of these is the strongest?

a) Hydrogen bondsb) Covalent bondsc) London dispersion forcesd) Permanent dipole-permanent dipole interactions

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Q39: What is the connection between the size of molecules and the London dispersionforces of attraction?

a) The size makes no difference.b) The smaller the molecule the weaker the force.c) The smaller the molecule the stronger the force.d) The weakest force is between the largest molecules.

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Q40: Which of these molecules has the largest permanent dipole?

a) HFb) HClc) HBrd) HI

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Q41: Identify the trends which would occur in descending the elements of the halogengroup.

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Q42:

Look at this table showing boiling points for some substances.

Substance Boiling Point (K)Bonds (or Forces) Brokenat the Boiling Point

Sodium 610 MetallicNeon 27 (A)

Water 373 (B)

Which word(s) should be inserted at (A) ?

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Q43: Which word(s) should be inserted at (B)? (See the table above.)

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Q44: Identify the substance which has neither covalent or metallic bonding.

a) NH3 (l)b) CCl4 (l)c) He (g)d) Hg (l)e) H2 (g)f) HBr (l)

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Q45: Identify the substance which has hydrogen bonding between the molecules.


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Q46: Identify the substances which are polar molecules.

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Q47: Identify the substance which has polar bonds but is a non-polar molecule.


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Q48:

This diagram shows the structure of KEVLAR, which is strong because of theintermolecular bonding between neighbouring molecules.

Name the type of intermolecular bonding involved.

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Q49: Highlight the area of the diagram to show the position of one such bond.

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185

Topic 9

End of unit test

Contents

186 TOPIC 9. END OF UNIT TEST

End of unit 1 test

Q1: Ethanol is formed industrially by of ethene.

a) increases.b) decreases.

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Q2: When the enthalpy change has a positive sign, the reaction is:

a) exothermic.b) endothermic.c) large.d) small.

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Q3: Ionic bonding is most likely when the electronegativity difference between theelements is:

a) exothermic.b) endothermic.c) large.d) small.

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Q4: Which of the following exist as diatomic molecules?

a) Ethaneb) Potassium bromidec) Carbon monoxided) Neon

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Q5: When an atom X of an element in Group 2 reacts to become X2+:

a) the mass number of X increases.b) the charge on the nucleus increases.c) the atomic number of X decreases.d) the number of occupied energy levels decreases.

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Q6: The same reaction was carried out at four different temperatures. The table belowshows the time taken for the reaction to occur.Temp (◦C) 20 30 40 50Time (seconds) 60 30 14 5

The results show that:

a) the rate of the reaction is directly proportional to the temperature.b) a small rise in temperature results in a large increase in reaction rate.c) the reaction is endothermic.


TOPIC 9. END OF UNIT TEST 187

d) the activation energy increases with increasing temperature.

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The potential energy diagram below refers to the reversible reaction involving reactantsR and products P.

Q7: What is the value of the activation energy for the forward reaction (reactants toproducts)?

a) 10 kJ mol-1

b) 20 kJ mol-1

c) 40 kJ mol-1

d) 60 kJ mol-1

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Q8: What is the enthalpy change for the forward reaction (reactants to products)?

a) -30 kJ mol-1

b) -20 kJ mol-1

c) +20 kJ mol-1

d) +30 kJ mol-1

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Q9: Which of the following molecules may be described as polar?

a) C1 Be C1b) H C1c)

C1 C1d) C1

CC1 C1

C1

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Q10: The melting points of Group 7 elements increase on descending the groupbecause the �� increase:

a) mean bond energiesb) nuclear chargesc) covalent bond lengthsd) London dispersion forces

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Q11: The difference between the covalent radii of lithium and carbon is mainly due tothe difference in the:

a) number of electrons.b) number of protons.c) mass of each atom.d) number of neurons.

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Q12: In general, covalent substances have lower melting points than ionic substancesbecause:

a) covalent bonds have no electrostatic attractive forces.b) bonds between molecules are weaker than bonds between ions.c) covalent compounds are composed of non-metals which have low melting points.d) ionic bonds are stronger than covalent bonds.

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Q13: Carbon dioxide is a gas at room temperature while silicon dioxide is a solidbecause:

a) London dispersion forces are much weaker than covalent bonds.b) carbon dioxide contains double covalent bonds and silicon dioxide contains single

covalent bonds.c) carbon-oxygen bonds are less polar than silicon-oxygen bonds.d) the relative formula mass of carbon dioxide is less than that of silicon dioxide.

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The apparatus shown can be used to prepare iron(III) chloride.

Q14: On the diagram, click on a substance which contains metallic bonding.

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Q15: By considering the method of production, predict the type of bonding in iron(III)chloride.

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Q16: Which of the following describes the bonding in the chlorides across Period 3 fromNaCl to SCl2?

a) Ionic ⇒ polar covalentb) Ionic ⇒ polar covalent ⇒ pure covalentc) Polar covalent ⇒ ionicd) Ionic ⇒ pure covalent ⇒ polar covalent

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Q17: The elements in the third row of the Periodic Table are shown below.

Na Mg Al Si P S Cl Ar

Why does the atomic size decrease crossing the period from sodium to argon?

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Q18: Atoms of different elements have different attractions for bonded electrons. Whatterm is used as a measure of these differing attractions?

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Q19: Use a table of these values, from the SQA data booklet, to explain why nitrogenchloride contains pure covalent bonds.

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Q20: Atoms of different elements have different ionisation energies. Explain clearly whythe first ionisation energy of sodium is less than the first ionisation energy of lithium.

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Q21: The ability of an atom to form a negative ion is measured by its Electron Affinity.The Electron Affinity is defined as the energy change when one mole of gaseous atomsof an element combines with one mole of electrons to form gaseous negative ions. Writethe equation, showing state symbols, that represents the Electron Affinity of chlorine.

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Graph A shows the volume of hydrogen gas produced against time when an excess ofmagnesium is added to 50 cm3 of hydrochloric acid of concentration 1 mol l-1 at 20 ◦C.

Time (s)

Volu

me

of H

ydro

gen

A

B

Graph B was obtained when the reaction was repeated with excess magnesium andhydrochloric acid of the same concentration but at a different temperature.

Q22: Which reaction was at the higher temperature?

a) Ab) B

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Q23: How many cm3 of acid was required to produce graph B?

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The potential energy diagram below represents the decomposition of hydrogen iodide:2HI (g) ⇒ H2(g) + I2(g)

Q24: What is the value for the activation energy (EA) for the above reaction?

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Q25: What would be the enthalpy change (ΔH) for the reaction H2(g) + I2(g) ⇒ 2HI (g)?

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Q26: The use of a catalyst would �� the activation energy in part 1.

a) increaseb) decreasec) have no effect on

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Q27: The use of a catalyst would �� the enthalpy change in part 2.

a) increaseb) decreasec) have no effect on

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Q28: The enthalpy change for the reaction 1/2H2(g) + 1/2I2(s) ⇒ HI (g) is known as theenthalpy of formation of hydrogen iodide.

It differs from the above equation in two respects:

• it involves the formation of one mole of HI;

• the reactants are in their standard states, so iodine is a solid.

Given that the enthalpy of sublimation of iodine I2(s) ⇒ I2(g) has ΔH = + 62 kJ, calculatethe enthalpy of formation of hydrogen iodide.

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Q29: If the enthalpy of formation of hydrogen chloride is -92 kJ mol-1, predict a valuefor the enthalpy of formation of hydrogen bromide.

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Q30: Match the following chemical terms to their correct definitions:

DefinitionChemical term(randomised)

For a chemical reaction to occur, particles must collide. Enthalpy change

The minimum kinetic energy required by colliding particlesbefore a reaction will occur.

Activated complex

Intermediate stage at the top of the energy barrier whenreactants change to products.

Catalyst

The difference in potential energy between products andreactants.

Collision theory

A substance that alters the rate of a reaction without beingused up.

Activation energy

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Regular arrangement of positively charged ions surroundedby delocalised electrons.

Covalent bonding

The electrostatic force of attraction of a shared pair ofelectrons for two positive nuclei.

Metallic bonding

The electrostatic attraction between negative and positiveions.

Intramolecular forces

Weak bonds between molecules. Dipole

A slight negative or positive charge caused by the unevendistribution of electrons.


Bonds holding atoms together within molecules. Ionic bonding

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Half the distance between the nuclei of two covalentlybonded atoms of an element.

Ionisation energy

The strength of the attraction of an element for the electronsof its bonding electrons. Covalent radius

The energy required to remove an electron from a gaseousatom to form an ion with a single positive charge.

Electronegativity

The temperature at which a substance changes from a solidto a liquid.

Boiling point

The temperature at which a substance changes from aliquid to a gas.

Melting point

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Q33: Sort the forces of attraction into order of strength from strongest to weakest:

• Hydrogen bonding

• Permanent dipole - permanent dipole interactions

• London dispersion forces

• Covalent bonding

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Q34: Match up the start of the statement with its end, selecting 'increases' or'decreases' where appropriate.

• Going across a period, atomic size .......... increases / decreases.

• Going down a group, atomic size .......... increases / decreases.

• Going down a group, electronegativity .......... increases / decreases.

• Going across a period, electronegativity .......... increases / decreases.

• Going across a period, boiling point .......... increases / decreases.

• Going down a group, boiling point .......... increases / decreases.

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A student writes the following two statements. Both are incorrect.In each case explainthe mistake in the student's reasoning.

Q35: All ionic compounds are solids at room temperature. Many covalent compoundsare gases at room temperature. This proves that ionic bonds are stronger than covalentbonds.

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Q36: The formula for magnesium chloride is MgCl2 because, in solid magnesiumchloride, each magnesium ion is bonded to chloride ions.

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Q37: Although propane and ethanol have similar molecular masses the alkane is a gasat room temperature while the alcohol is a liquid. Explain why propane is a gas at roomtemperature whereas ethanol is a liquid. (This question is worth 3 marks)

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GLOSSARY 195

Glossary

Activated complex

the activated complex is a very unstable arrangement of atoms formed at themaximum of the potential energy barrier, during a chemical reaction

Activation energy

is the minimum kinetic energy required by colliding particles before reaction willoccur, since a high energy activated complex must be formed

Adsorption

adsorption occurs when molecules become bonded to the surface of a catalyst

Allotropes

one of two or more existing forms of an element. For example, graphite anddiamond are allotropes of carbon

Bonding electrons

are shared pairs of electrons from both atoms forming the covalent bond

Chemical bonding

is the term used to describe the mechanism by which atoms are held together

Chemical structure

describes the way in which atoms, ions or molecules are arranged

Collision theory

of reactions suggests that, for a chemical reaction to occur, particles must collide

Covalent bond

a covalent bond is formed when two atoms share electrons in their outer shell tocomplete the filling of that shell

Covalent radius

half the distance between the nuclei of two bonded atoms of an element

Delocalised

delocalised electrons, in metallic bonding, are free from attachment to any onemetal ion and are shared amongst the entire structure

Desorption

desorption occurs when the bonds between the molecules and the surface breakand the molecules leave the surface of the catalyst

Diatomic

molecules with only two atoms are described as diatomic (e.g. oxygen, O2, andcarbon monoxide, CO.)

Dipole

an atom or molecule in which a concentration of positive charges is separated froma concentration of negative charge


196 GLOSSARY

Electronegativity

a measure of the attraction that an atom involved in a bond has for the electronsof the bond

Enthalpy change

for a reaction is defined as the change in heat energy when 1 mole of reactant isconverted to product(s) at constant pressure, and has the symbol ΔH and units ofkJ mol-1

Fullerenes

are molecules of pure carbon constructed from 5- and 6-membered ringscombined into hollow structures. The most stable contains 60 carbon atoms ina shape resembling a football

Hydrogen bonds

are electrostatic forces of attraction between molecules containing a hydrogenatom bonded to an atom of a strongly electronegative element such as fluorine,oxygen or nitrogen, and a highly electronegative atom on a neighbouring molecule


are those which attract molecules together. They are weaker than chemical bonds

Intramolecular forces

are forces of attraction which exist within a molecule

Ionisation energy

the energy required to remove one mole of electrons from one mole of atoms inthe gaseous state

Isoelectronic

means having the same arrangement of electrons. For example, the noble gasneon, a sodium ion (Na+) and a magnesium ion (Mg2+) are isoelectronic

Lattice

a lattice is a regular 3D arrangement of particles in space. The term is applied tometal ions in a solid, and to positive and negative ions in an ionic solid

London dispersion forces

the forces of attraction which result from the electrostatic attraction betweentemporary dipoles and induced dipoles caused by movement of electrons in atomsand molecules

Lone pairs

are pairs of electrons in the outer shell of an atom which take no part in bonding

Miscible

fluids are fluids which mix with or dissolve in each other in all proportions

Periodicity

is the regular recurrence of similar properties when the elements are arranged inorder of increasing atomic number


GLOSSARY 197

Polar covalent bond

a covalent bond between atoms of different electronegativity, which results in anuneven distribution of electrons and a partial charge along the bond

Potential energy diagram

shows the enthalpy of reactants and products, and the enthalpy change during achemical reaction

Properties

of a substance are their physical and chemical characteristics. These are often areflection of the chemical bonding and structure of the material.

Thermochemical equation

states the enthalpy change for the reaction defined, with reactants and products inthe states shown

Viscosity

is the resistance to flow that is exhibited by all liquids


198 ANSWERS: TOPIC 1

Answers to questions and activities

1 Reaction rates - collision theory

Test your prior knowledge (page 2)

Q1: Any three from:

• concentration;

• particle size / surface area;

• temperature.

Q2: Any two from:

• mass of products;

• mass of reactants;

• mass of gas given off;

• concentration of reactants;

• concentration of products;

• volume of reactants;

• volume of products;

• pressure;

• conductivity;

• pH;

• colour Intensity.

Q3: 0.01 g s-1 (to 2 dec. pl.)Working:

Rate =Change in mass

Change in time

=0.7

60= 0.01 g s−1

Measuring rate (page 4)

Q4: 7.0 cm3

Q5: 1.4 cm3 s-1

Q6: 0.7 cm3 s-1

Q7: The rate of reaction is dropping as it progresses, being relatively fast at the startand slowing towards the end. This is because the reactants are being used up.

Q8: 0.52 cm3 s-1


ANSWERS: TOPIC 1 199

Collisions and concentration (page 7)

Q9: 3

Q10: 0.3

Q11: 0.6

Q12: The rate would be 1.2 collisions per second. It would seem that doubling theconcentration of acid has doubled the rate, and in some reactions this is correct; therate can be directly proportional to the concentration of a reactant. On the other hand,sometimes doubling the concentration of a reactant only increases rate a little, or evennot at all. There is no simple way to predict the relationship in advance and eachconcentration/rate relationship must be investigated experimentally to determine anyrelationship. It can be said however that for many reactions increase in concentrationincreases rate.

Q13: There are more successful collisions in the earlier part of the reaction, when thereare plenty of reactant particles available to react. As some react, there are less availablefor collision and therefore the rate of reaction falls. This is as expected as a consequenceof collision theory.

Collisions and particle size (page 8)

Q14: small

Q15: 6

Q16: 0.6

Q17: 0.8

Q18: There are more successful collisions in the earlier part of the reaction, whenthere are plenty of reactant particles available to react. As some react, there are feweravailable for collision and therefore the rate of reaction falls. This is as expected as aconsequence of collision theory.

Collisions and temperature (page 9)

Q19: 4

Q20: 0.4

Q21: 0.8

Q22: c) doubles

Q23: There are more successful collisions in the earlier part of the reaction, whenthere are plenty of reactant particles available to react. As some react, there are feweravailable for collision and therefore the rate of reaction falls. This is as expected as aconsequence of collision theory.



End of topic 1 test (page 12)

Q24: A: average kinetic energy of all the particles.

Q25: A: Time taken for reaction to complete.

Q26: A: increasing the concentration of the acid.

Q27: 0.2 g s-1

Q28: 0.08 g s-1

Q29: D: curve E

Q30: C: curve D

Q31: A: a small rise in temperature results in a large increase in reaction rate.

Q32: A: decreasing the mass of copper carbonate.

Q33: A: graph A

Q34: 0.25 moles

Q35: 0.1 moles

Q36: 0.44 mol l-1



2 Reaction rates - reaction profiles


Q1: Exothermic

Q2: Endothermic

Q3: Any 2 from:

• particles must collide to react;

• not all collisions are successful;

• sufficient energy is needed and orientation must be correct.

Answers from page 22.

Q4: hydrogen

Q5: 0.8

Q6: The rate of fall in mass is = 0.2 g min-1. Strictly speaking , since the massdecreases with time, this value should be reported as a negative value. In most caseshowever the phrase "rate of fall" will describe the direction adequately enough.

Q7: d) g s-1

Q8: c) 20 s

Q9: d) The rate of production of product over 30 seconds is the same as the rate offall of reactant.

Measuring rate at different concentrations (page 24)

Q10: c) 2.0 mol �-1

Q11: c) 2.0 mol �-1

Q12: 20

Q13: 26

Q14:

• The steepness of the slope for the 1.0 mol �-1 is less because the reaction rate islower (fewer acid particles).

• The point at which the 1.0 mol �-1 line becomes horizontal takes longer to reach(25 seconds rather than 20 seconds) because the reaction is slower (fewer acidparticles).

• The final level of the 1.0 mol �-1 horizontal line is lower by half because the acid ispresent in only half the quantity.

Q15: The 0.5 mol �-1 graph starts to become level from about 30 seconds onward andwould settle completely at a volume of 6.5cm3. This is predictable because 6.5cm3 ishalf the value of 13.0cm3 which is the final volume for the 1.0 mol �-1 graph.




Q16: 20 cm3 (1.5 mol �-1 hydrochloric acid is three times the concentration of 0.5 mol�-1 hydrochloric acid.)

Q17: 10.4 cm3 s-1


Q18: a) A

Q19: d) D

Q20: b) B

Q21: d) D

Controlling the rate - temperature & kinetic energy (page 31)

Q22: c) Not changed

Q23: a) Increased

Q24: b) C and D

Q25: a) Particles with sufficient energy to react.

Q26: c) Roughly doubled.


Q27: b) The activation energy if less for phosphorus than for magnesium.

Q28: c) The reaction rate doubles

Q29: d) More particles have energy greater than the activation energy.

Q30: a) the minimum kinetic energy required by colliding molecules for a reaction tooccur.

Q31: b) an intermediate stage at the top of the activation energy barrier.

Q32: c) potential energy diagrams.

Q33: c) 0.025 mol �-1 s-1

Q34: Either curve can be usedRate = (change in concentration)/(change in time) (1/2 mark)= 0.375/15 (1/2 mark)= 0.025 (1/2) mol �-1 s-1 (1/2 mark)Allow ± 0.001 only. Two sig fig must be shown. (-1/2)



Q35: d) 31.25 seconds

Q36: At 40◦C, rate = 0.032 s-1 (1/2 mark)Time = 1/rate (1/2 mark)= 1/0.032 (1/2 mark)= 31.25 seconds (1/2 mark - units must be given.)

Q37: Increasing temperature increases kinetic energy (or speed of movement) ofmolecules (1/2 mark)The (potential) energy generated in collisions increases (collisions have more force) (1/2mark)More collisions will result in the activation energy being attained (1 mark)

Q38: B 60 ◦C

Q39: b) a small rise in temperature results in a large increase in reaction rate.



3 Catalysis


Q1: a) average kinetic energy of all the particles;

Q2: b) Endothermic

Q3: a) The minimum kinetic energy required by colliding molecules for a reaction tooccur.

Addition of hydrogen to ethene (page 42)

Q4: c) Adsorption

Q5: b) They get weaker.

Q6: none

Q7: c) The catalyst speeds up the reaction without taking part.

Catalysis by cobalt(II) chloride solution (page 44)

Q8: b) Co2+(aq) ions

Q9: c) Co3+(aq) ions

Q10: b) Co2+(aq) ions

Q11: This is a poor definition. It is correct to say that the Co2+(aq) ions do speed upthe reaction but it is wrong to say that they do not take part in the reaction. The pinkCo2+(aq) ions are changed into something green (probably Co3+(aq) ions) which arelater changed back into pink Co2+(aq) ions. So the Co2+(aq) ions can be recovered atthe end.


Q12: d) Cl



Industrial uses of catalysts (page 46)

Q13:


Q14: c) It will decrease.

Q15: c) Positive


Q16: d) Endothermic

Q17: a) Positive

Q18:




Q19: d) ΔH =+283 kJ mol-1

Q20: b) ΔH = Hproducts - Hreactants

Q21: c) Exothermic

Q22: d) ΔH = -1414 kJ mol-1

Q23: 6835

Q24: +150

Q25: 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)

Q26: -906

Q27: -226.5

Q28: 1359


Q29: b) B

Q30:

The enthalpy change (ΔH) is given by the expression:

ΔH = enthalpy of products − enthalpy of reactants

From the graph,

Enthalpy of reactants = −1200 kJ

Enthalpy of products = −630 kJ

So, ΔH = (−630) − (−1200) kJ

= +570 kJ

Q31: d) -890 kJ


Q32: d) +110 kJ

Q33: c) +55 kJ

Q34:

If the activation energy is less, the barrier will be lower and it will be easier for collidingmolecules to get over the barrier. There will be more molecules with enough energy tocollide successfully and so the reaction rate will increase.

In general, the lower the activation energy the faster the reaction.




Q35: A

Q36: D

Q37: c) The number of molecules with enough energy will increase.

How catalysts work (page 60)

Q38: b) High activation energy, few successful collisions.

Q39: c) Low activation energy, lots of successful collisions.

Q40: c) It speeds up the reaction by lowering the activation energy.

Q41: b) No effect.


Q42: X

Q43: 3

Q44: 4

Q45: b) Decreases it

Q46: W

Q47: X

Q48:

Point Z is at the top of the barrierfor the second step in thecatalysed reaction, i.e. itrepresents the activated complexfor the second step. The old bondswill be half broken and the newbonds half formed. A possiblestructure is shown opposite.


Q49: A: molecules always have energy above the activation energy.

Q50: a) the minimum kinetic energy required by colliding molecules for a reaction tooccur.

Q51: iv) D



Q52: v) E

Q53: iii) C

Q54: A. A very unstable intermediate is formed when bonds within the reactantmolecules begin to break and new bonds begin to form.

Q55: Exothermic

Q56: 10 kJ

Q57: 35 kJ

Q58: -25 kJ

Q59: +25 kJ

Q60: a)

1. Energy of reactants / Reactant molecules collide with correct geometry.

2. Energy of activated complex / The collision has enough energy to overcome theEa.

3. Activation energy of the forward reaction / Activated Complex: bonds in the reactantmolecules break and new bonds begin to form.

4. Enthalpy change of the forward reaction

b) The product molecule(s) has less energy than the reactants.



4 The Periodic Table


Q1: a) Atomic Number

Q2:

• Hydrogen,

• Nitrogen,

• Oxygen,

• Fluorine,

• Chlorine,

• Bromine,

• Iodine.

• (In any order)

Q3:

Particle Charge Mass LocationProton +1 1 NucleusNeutron 0 1 Nucleus

Electron -1 0Orbiting nucleus orin shells

Q4: d) Period

Q5: c) Group

Newlands' octaves (page 74)

Q6: c) Atomic mass

Q7: a) Both are gases.

Q8: b) Both are reactive metals.

Q9:

When listed in order of increasing atomic mass, similar properties appear with everyeighth element. If the elements are numbered in order, element 1 has properties incommon with element 8, element 2 has properties in common with element 9, etc.

This is similar to musical notation. Notes in a scale are described as the letters A to G.The note after G is A and so on. A to G is said to be one octave. The same applies ifthe notes are described as 'doh-ray-me-fah-soh- etc'.

Q10: 3

Q11: Chromium, manganese and iron are metals. The elements above them inNewlands' table are non-metals.



Mendeleev's Periodic Table (page 76)

Q12: Mendeleev believed that elements, as yet undiscovered, would fit in these spaces.

Q13: Iodine has similar chemical properties to bromine and so fits better in the columnwhich contains bromine.

Q14: noble gases


Q15: atomic number


Q16: carbon

Q17: argon

Q18: helium


Q19: lithium

Q20: caesium

Q21: Melting point and boiling point both decrease and so the forces must be gettingweaker.


Q22: c) The number of electron shells increases.

Q23: d) The number of protons increases.


Q24: c) Mendeleev left gaps for elements which had not yet been discovered.

Q25: c) There is a steady decrease in atomic size across a period from left to right andd) There is a decrease in first ionisation energy on going down Group 0.

Q26: a) Atomic Number.

Q27: Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine (In any order)

Q28: b) The alkali metals.



Q29: a) The halogens.

Q30: c) The noble gases.

Q31: a) number of occupied energy shells;

Q32: c) The right hand side.



5 Bonding and structure


Q1: a) a shared pair of electrons.

Q2: d) delocalised electrons.

Q3: b) transfer of electrons.

Q4: a) have high melting and boiling points.

Q5: d) conduct electricity when molten or in solution.


Q6:

You have probably thought of gold, nuggets of which are found in rocks and as goldparticles in some river beds. Platinum and other platinum group metals (rhodium,iridium, ruthenium and osmium) are also found native.

Sulfur is also found as yellow crystals in rocks and around volcanos. Copper is alsofound in elemental form, as well as in other ores.

gold sulfur native copper


Q7: Niobium is used in steel superalloys which can resist high temperatures and areused in jet engine components and heat shields.


Q8: argon



Q9: Gain an outer electron.

Q10: Lose the outer electron.

Q11: A positive sodium ion (Na+)

Q12: neon


Q13: The group 1 metals sodium and potassium are easily cut with a knife (which ismade from hard metals - iron and chromium). Mercury is liquid (melting point -39◦C)and gallium melts in the hand (at 30◦C).


Q14: c) delocalised.

Q15: d) Lattice

Q16: Each aluminium atom will ionise to produce three electrons and a Al3+ ion. Sodiumwill produce only one electron (to carry the electric current) from each atom.


Q17: 4

Q18: 1

Q19: b) share, electron

Q20: d) positively charged nuclei and negatively charged shared electrons.



Q21:


Q22: 4

Q23: a) Single covalent.

Graphite structure (page 103)

Q24: c) stronger.

Q25: 6

Fullerenes (page 105)

Q26: Diamond and graphite are network covalently bonded, so that it would require avery large input of energy to separate the atoms in a solution. "Buckyballs" are covalentmolecular bonded, with strong bonds binding the C60 atoms into a molecule, but muchweaker bonds between the molecules in the solid. These can be easily broken to forma solution.


Q27: 7

Q28: 8

Q29: Both ions have a stable noble gas electron configuration.




Q30: c) large

Q31: c) positive caesium ions and negative chloride ions.

Q32: c) 3 : 2


Q33: c) An ionic lattice.


Q34: a) Calcium sulfide

Q35: a) Sulfur

Q36: b) delocalised electrons.

Q37: b) Metallic

Q38: a) Covalent and c) Made up of discrete molecules.

Q39: a) Boron

Q40: d) Phosphorus and f) Sulfur

Q41: b) Chlorine and e) Sodium

Q42: c) C

Q43: a) Ionic

Q44: f) SiO2 (s)

Q45: a) NH4Cl (s) and e) Na2CO3 (s)

Q46:



Q47: The bonds between the layers of carbon in graphite are comparatively weak,which means that the layers can easily slide over each other providing lubrication.

Q48:

The carbon atoms in graphite have electrons not involved on the covalent bondsbetween atoms within the layers. These delocalised electrons permit conduction ofelectricity.

The boron atoms in boron nitride have no free electrons once they have formed threecovalent bonds, so there is no delocalised pool of electrons to conduct electricity.

Q49: The nitrogen atoms in the layers have two free electrons which can form strongcovalent bonds between adjacent layers in boron nitride.

Q50: b) Monatomic



6 Periodic Table trends

Prior knowledge: Periodic Table trends (page 119)

Q1:

Particle Charge Mass LocationProton +1 1 Nucleus

Neutron 0 1 Nucleus

Electron -1 0Orbitting Nucleus

or In Shells

Q2: c) number of outer electrons.

Q3: a) number of occupied energy shells.

Q4: a) Atomic number

Q5: a) a shared pair of electrons.


Q6: decreases

Q7: increases

Q8: The covalent radius is defined as half the distance between the nuclei of bondedatoms. Noble gases do not form bonds because they are so unreactive.


Q9: c) The number of electron shells increases.

Q10: d) The number of protons increases.

Ionisation energies activity (page 125)

Q11: a) Yes

Q12: The shape of the first part of the graph (from atomic number 3-10) is repeated forelements 11-18

Q13: noble gases

Q14: alkali metals

Q15: a) A steady increase.

Q16: b) A steady decrease.



Q17:

Metal 1st Ionisation energy / kJ mol-1

Lithium 526

Sodium 502

Potassium 425

Rubidium 409

Caesium 382


Q18:

The nuclear charge increases on going down a group which should make the outerelectron more difficult to remove. (This would get 1 mark.)

However, the covalent radius increases and so the outer electron is further away fromthe nucleus. The electrostatic forces get weaker as the distance between the chargesincreases. This should make the outer electron easier to remove. (1 mark)

On going down the group, there are more and more inner electrons which prevent theouter electron from experiencing the full effect of the nuclear charge. These innerelectrons increasingly screen the outer electron from the nucleus. Consequently, theouter electron becomes easier to remove as the atom gets bigger, i.e. the first ionisationenergy decreases on going down a group. (1 mark)


Q19:

On going from left to right across Period 2, the nuclear charge increases. The electronsare held more tightly making it more difficult to remove one of the outer electrons. (1mark)

Each additional electron goes into the second shell. This makes little difference to thescreening effect. Consequently, the increased nuclear charge pulls in the electrons andthe covalent radius decreases. (1 mark)

Because the outer electrons are closer to the nucleus, they are more strongly held andso the first ionisation increases across the Period. (1 mark)

Q20: b) It increases.

Q21: a) It increases.

Q22: c) It stays the same.

Q23: c) It stays the same.

Q24: a) It decreases.




Q25: 1

Q26: 7

Q27: noble gases

Q28: The noble gases do not generally form bonds with other elements.


Q29: c) Electronegativity values increase on going from left to right and decrease ongoing down a group.

Summary of trends in the Periodic Table (page 130)

Q30:


Q31: c) Mendeleev left gaps for elements which had not yet been discovered.

Q32: 3. and 6.

Q33: b) Atomic size

Q34: b) B

Q35: d) more occupied electron shells?

Q36: c) protons?

Q37: c)Mg(g) → Mg+(g) + e−

Q38: d)F (g) → F+ (g) + e−



Q39: 4. and 5.

Q40: d) increasing nuclear charge?

Q41: d) 6

Q42: b)

Atomic Number

Q43: d) Atomic mass

Q44: Electronegativity

Q45: Fluorine

Q46: d)

Q47: On going from left to right across Period 2, the nuclear charge increases. Theelectrons are held more tightly making it more difficult to remove one of the outerelectrons. (1 mark)

Each additional electron goes into the second shell. This makes little difference to thescreening effect. Consequently, the increased nuclear charge pulls in the electrons andthe covalent radius decreases. (1 mark)

Because the outer electrons are closer to the nucleus, they are more strongly held andso the first ionisation increases across the Period. (1 mark)

Q48: The nuclear charge increases on going down a group which should make theouter electron more difficult to remove. (1 mark)

However, the covalent radius increases and so the outer electron is further away fromthe nucleus. The electrostatic forces get weaker as the distance between the chargesincreases. This should make the outer electron easier to remove. (1 mark)

On going down the group, there are more and more inner electrons which prevent theouter electron from experiencing the full effect of the nuclear charge. These innerelectrons increasingly screen the outer electron from the nucleus. Consequently, theouter electron becomes easier to remove as the atom gets bigger, i.e. the first ionisationenergy decreases on going down a group. (1 mark)

Q49: b) They generally do not form bonds with other elements.



7 Bonding continuum and polar covalent bonding

Prior knowledge: bonding continuum and polar covalent bonding (page 141)

Q1: b) a shared pair of electrons.

Q2: c) They generally do not form bonds with other elements.

Q3: a) Fluorine

Q4: c) Hydrogen

Electronegativity (page 142)

Q5: c) Hydrogen chloride (HCl)

Q6: a) Hydrogen fluoride (HF)

Q7: d) Ammonia (NH3)

Direction of charge (page 144)

Q8:


Q9:

• Pure covalent - electronegativity difference 0; Br - Br

• Polar covalent - electronegativity difference 0.6; δ+H - Brδ-

• Ionic - electronegativity difference 2.0; K+ ...... Br-

Q10: c) Ionic

Q11: a) Pure covalent

Q12: b) Fluorine




Q13: B

Q14: C

Q15: B

Predicting molecular polarity (page 149)

Q16:

Detecting polar molecules (page 150)

Q17: water

Q18: water

Q19: b) It is attracted by the rod.

Q20: a) It is not affected by the rod.


Q21: a) Polar


Q22: d) Calcium sulfide

Q23: a) Sulfur

Q24: c) delocalised electrons?

Q25: b) Metallic

Q26: 1. Covalent and 2. Made up of discrete molecules

Q27: a) Boron

Q28: 4. Phosphorus and 6. Sulfur



Q29: 2. Chlorine and 5. Sodium

Q30: a) Cl-F

Q31: c) C

Q32: c) BeCl2

Q33: c) Ionic

Q34: f) SiO2 (s)

Q35: 1. NH4Cl (s) and 5. Na2CO3 (s)

Q36: b) CH3OH (l)



8 Intermolecular forces


Q1: c) Electronegativity

Q2: a) Fluorine

Q3: c) increases.

Induced dipoles (page 164)

Q4: b) unevenly spread.

Q5: b) a temporary dipole.

Q6: b) The temporary dipole causes an induced dipole.

Q7: b) intermolecular


Q8: krypton

Q9: krypton

Q10: krypton

Q11: c) The larger the atom the stronger the London dispersion forces.


Q12: The energy required to break forces which are only effective at such lowtemperatures is very small, and in fact, London dispersion forces of attraction are weakerthan all other types of bonding.


Q13:

As the alkane molecules in the family get bigger from methane to butane, the boilingpoint increases (butane at -1◦C is a higher temperature than methane at -164◦C). (1mark)

Since boiling point depends on the strength of the London dispersion forces, butanemust have stronger London dispersion forces than methane. (1 mark)

Therefore the larger the molecule the stronger the London dispersion forces. (1 mark)

Q14: butane

Q15: butane



Q16: butane

Q17: c) The larger the molecule the stronger the London dispersion forces.

Q18:

As the alkane molecules in the family get bigger from methane to butane, the boilingpoint increases (butane at -1◦C is a higher temperature than methane at -164◦C).(1mark)

Since boiling point depends on the strength of the London dispersion forces, butanemust have stronger London dispersion forces than methane. (1 mark)

Therefore the larger the molecule the stronger the London dispersion forces. (1 mark)

Strength of hydrogen bonds (page 169)

Q19: As size increases, London dispersion forces increase, so one might expectwater (the smallest) to have the lowest boiling point. However, the boiling point ofwater goes against the trend and is much higher than might be expected. Thisis because water molecules exhibit hydrogen bonding between the molecules, dueto the highly polar character of the water molecule. Although the other hydrideshave some permanent dipole-permanent dipole interactions, the hydrogen bondedpermanent dipole-permanent dipole interactions in water are stronger than other formsof permanent dipole-permanent dipole interactions and London dispersion forces, asevidenced by the high boiling point.

Q20: London dispersion forces

Q21: water

Q22: water

Q23: b) The hydrogen bonds are stronger than London dispersion forces and the otherpermanent dipole-permanent dipole interactions.

Q24: As size increases, London dispersion forces increase, so one might expectwater (the smallest) to have the lowest boiling point. However, the boiling point ofwater goes against the trend and is much higher than might be expected. Thisis because water molecules exhibit hydrogen bonding between the molecules, dueto the highly polar character of the water molecule. Although the other hydrideshave some permanent dipole-permanent dipole interactions, the hydrogen bondedpermanent dipole-permanent dipole interactions in water are stronger than other formsof permanent dipole-permanent dipole interactions and London dispersion forces, asevidenced by the high boiling point.

Q25: a) Covalent bonds are much stronger than hydrogen bonds.


Q26: All of the above.



Hydrides of non-metals (page 172)

Q27: No

Q28: NH3, H2O and HF

Q29: b) Van der Waals attractions

Testing viscosity (page 175)

Q30: methanol

Q31: methanol

Q32: glycerol

Q33: 0.66 m s-1

Q34: a) 0.8 m s-1

Q35: 1.25 seconds


Q36: Ammonia is a polar substance which can form hydrogen bonds with water;nitrogen is non-polar and unable to bond to water.


Q37: d) a temporary dipole causing an induced dipole.

Q38: b) Covalent bonds

Q39: b) The smaller the molecule the weaker the force.

Q40: a) HF

Q41: The London dispersion forces become stronger., The relative atomic massincreases.

Q42: The forces broken when neon boils are London dispersion forces.

Q43: Water has hydrogen bonding (a form of permanent-permanent dipole interaction)between the molecules.

Q44: c) He (g)

Q45: a) NH3 (l)

Q46: ETh1, ETh6

Q47: b) CCl4 (l)



Q48: Kevlar is strong because of the intermolecular bonding between neighbouringmolecules. This bonding is called hydrogen bonding.

Q49:



9 End of unit test

End of unit 1 test (page 186)

Q1: b) decreases.

Q2: b) endothermic.

Q3: c) large.

Q4: c) Carbon monoxide

Q5: d) the number of occupied energy levels decreases.

Q6: b) a small rise in temperature results in a large increase in reaction rate.

Q7: b) 20 kJ mol-1

Q8: a) -30 kJ mol-1

Q9: b) H C1

Q10: d) London dispersion forces

Q11: b) number of protons.

Q12: b) bonds between molecules are weaker than bonds between ions.

Q13: a) London dispersion forces are much weaker than covalent bonds.

Q14:

Heat

Iced water

FeCl3(g)

Fe(s)

Cl2(g)

r

Fe(s)

Q15: Polar covalent

Q16: b) Ionic ⇒ polar covalent ⇒ pure covalent

Q17: The number of protons increases OR greater nuclear charge OR greater nuclearattraction.



Q18: Electronegativity

Q19: Same electronegativity values.

Q20: Bigger atom so electron is further from the nucleus; inner electrons shield (screen)the outer electron from attraction of the nucleus.

Q21: Cl(g) + e- ⇒ Cl-(g)

Q22: b) B

Q23: 100 cm3

Q24: 190 kJ

Q25: -10 kJ

Q26: b) decrease

Q27: c) have no effect on

Q28: 26 kJ mol-1

Q29: -33 kJ mol-1

Q30:

Definition Chemical term

For a chemical reaction to occur, particles must collide. Collision theory

The minimum kinetic energy required by colliding particlesbefore a reaction will occur.

Activation energy

Intermediate stage at the top of the energy barrier whenreactants change to products.

Activated complex

The difference in potential energy between products andreactants.

Enthalpy change

A substance that alters the rate of a reaction without beingused up.

Catalyst



Q31:


Regular arrangement of positively charged ions surroundedby delocalised electrons.

Metallic bonding

The electrostatic force of attraction of a shared pair ofelectrons for two positive nuclei.

Covalent bonding

The electrostatic attraction between negative and positiveions.

Ionic bonding

Weak bonds between molecules. Intermolecular forces

A slight negative or positive charge caused by the unevendistribution of electrons.

Dipole

Bonds holding atoms together within molecules. Intramolecular forces

Q32:


Half the distance between the nuclei of two covalentlybonded atoms of an element.

Covalent radius

The strength of the attraction of an element for the electronsof its bonding electrons.

Electronegativity

The energy required to remove an electron from a gaseousatom to form an ion with a single positive charge.

Ionisation energy

The temperature at which a substance changes from a solidto a liquid.

Melting point

The temperature at which a substance changes from aliquid to a gas.

Boiling point

Q33:

1. Covalent bonding

2. Hydrogen bonding

3. Permanent dipole - permanent dipole interactions

4. London dispersion forces

Q34:

• Going across a period, atomic size decreases.

• Going down a group, atomic size increases.

• Going down a group, electronegativity decreases.

• Going across a period, electronegativity increases.

• Going across a period, boiling point decreases.

• Going down a group, boiling point increases.



Q35: Covalent bonds not being broken OR Intermolecular bonds that are breaking. Anyalternative wording that recognises that covalent bonds are not broken when covalentsubstances melt or boil.

Q36: Formula refers to the ratio of Mg2+:Cl- ions (in lattice) OR alternative wording iein the lattice there are twice as many chloride ions as magnesium ions OR Mg2+ ionssurrounded by > 2 Cl- ions OR Cl- surrounded by >1 Mg2+ Any reference to "chlorineions" is not acceptable.

Q37: As a general rule, award yourself a half mark, up to a maximum of three marksin total, for each point you made (NB this is a general rule only, there are no half marksawarded at Higher):

• Propane molecules are held together by weak intermolecular forces / ethanolmolecules are held together by strong intermolecular forces.

• The only intermolecular forces in propane are London Dispersion forces.

• These are weak forces are due to momentary displacement of electrons betweenatoms creating temporary dipoles.

• In response to these temporary dipoles an induced dipole can occur.

• The intermolecular forces in ethanol are hydrogen bonds.

• Hydrogen bonding arises because the O-H bond is highly polar (there is a largedifference in the electronegativities of O and H).

• The small positive dipole on H and small negative dipole on O strongly attract.

• This causes ethanol to have a higher boiling point than propane as more energy isneeded to overcome the stronger hydrogen bonds than the weak LDF's.


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