general chemistry: chapter 1 slide 1 of 57 c h e m 1 0 0 general chemistry lecturer: prof. dr....

General Chemistry: Chapter 1Slide 1 of 57

C H E M 1 0 0

GENERAL CHEMISTRY

Lecturer: Prof. Dr. Tamerkan Özgen


Prentice-Hall © 2005

Chapter 1: Chemistry : Matter and Measurement

General ChemistryFouth Edition

Hill • Petrucci • McCreary • Perry


Contents

1.1 Chemistry: Principles and Applications

1.2 Getting Started: Some key terms

1.3 Scientific measurements

1.4 Precision and Accuracy in Measurements

1.5 A Problem-Solving Method

1.6 Further Remarks on Problem Solving



In their daily life people have always practiced chemistry

making drugs to fight disease;

computer chips;

pesticides to protect humans, animals and crops;

fertilizers to grow more food;

fibers for clothes;

building materials for housing;

replacing worn out body parts.



Glazing of pottery;

smelting of ores to produce metals;

tanning of leather;

dyeing of fabrics;

making of cheese, wine, beer and soap,

From petroleum, motor fuels and thousands of chemicals used in the

manufacture of plastics, synthetic fabrics, pharmaceuticals, and pesticides.

From coal thousands of chemicals for various uses



Modern chemical knowledge is also needed

to understand the processes that control life and

to understand and control processes in the environment,

such as the formation of smog and the destruction of

stratospheric ozone.

Early chemical knowledge consisted of the “how to” of chemistry

Modern chemical knowledge answers the “why” as well as the

“how to” of chemical change.


1.2 Getting Started. Some Key terms

Chemistry: Study of the composition, structure, and properties

of matter and of changes that occur in matter.

Matter: Occupies space, has mass and inertia

Atoms: The smallest distinctive units in a sample of matter

Molecules: Larger units in which two or more atoms are joined

together

Composition: Types of atoms and relative proportions of diferent

atoms in a sample of matter

ex. H2O, 11.19% H and 88.81% O

Properties: Distinguishing features, ( mainly physical and

chemical properties)


Properties of Matter

A physical property of matter displays without changing its

composition. Ex: color, brittleness, malleability, ductility.

A chemical property is a characteristic shown by a sample

of matter as it changes in composition.

In a chemical change, or chemical reaction, one or more

kinds of matter are converted to new kinds of matter with

different compositions (change in composition).

Ex. Zn + HCl giving ZnCl2 and H2 ,

Na + H2O giving NaOH and H2

HCOOH (formic acid) decomposing to CO and

H2O


States of Matter

a) solid b) liquid c) gas

Animations

1: Change of State

2: Physical Properties of Halogens



Chemical Symbol

A chemical symbol is a one- or two-letter symbol derived from

the name of an element.

Most symbols are based on English names; a few are based on

the latin name of the element. (eg. He, Ba, Na, Hg)

The first letter of a chemical symbol is always capitalized and the

second is never capitilized.

Chemists show compounds by combinations of chemical

symbols called chemical formulas.


Chemical Symbol

To represent a particular atom we use the symbolism:

A= mass number Z = atomic number

To represent an ion

Symbols


35 Cl

1727

Al13

56 Fe

26


Classification of Matter


Classifying Matter

Figure 1.3

Atoms Molecules

make up

ALL MATTER

which exists as

Substances Mixtures

which may be

Elements Compounds Heterogeneous

which may be

Homogeneous

Animations

3: Mixtures and Compounds

4: Dissolution of NaCl in Water



The Scientific Method

The ancient Greeks developed methods using knowledge

and mathematics.

They started with certain basic assumptions. Then by

deduction method certain conclusions logically followed.

Ex. if a = b and b = c, then a = c.

Greek philosopher Aristotle assumed four fundamental

substances: air, earth, water and fire. All other

materials, he believed, were formed by the combination of

these four elements. For many years this was accepted.



Scientists often begin by making observations and then formulating a

hypothesis. A hypothesis is an explanation or prediction concerning

some phenomenon.

Scientists test a hypothesis through a carefully controlled procedure

called an experiment.

Scientific data are obtained during experiments.

Further experiments may be carried to refine these data.

Large collections of data may be summerized to identfy patterns which

are called scientific laws.


1.3 Scientific Measurements

• Scientists worldwide use a common system of measurement,

called Systeme Internationale d’Unites (International System

of Units) which is in short written as SI .

• This system was adopted in 1960, and is

• A modernized version of the metric system established in France

in 1791.

• In modern scientific work, all measured quantities can be

expressed in terms of the seven base units listed in the next slide.

• An essential aspect of SI is the use of exponential (powers of ten)

notation for numbers.


Units


Units

Derived Quantities

Force Newton, kg m s-2

Pressure Pascal, kg m-1 s-2

Energy Joule, kg m2 s-2

Other Common Units

Length Angstrom, Å, (10-8 cm)

Volume Litre, L, (10-3 m3)

Energy Calorie, cal, (4.184 J)

Pressure 1 Atm = 101.325 kPa

1 Atm = 760 mm Hg = 760 Torr


Wrong units

The Gimli Glider


EDMONTON

• The Gimli Glider is the nickname of an Air Canada aircraft which was

involved in a notable aviation incident.

• On 23 July 1983, Air Canada Flight 143, a Boeing 767-200 jet, ran

completely out of fuel at 41,000 feet (12,500 m) altitude, about halfway

through its flight from Montreal to Edmonton via Ottawa. The crew was

able to glide the aircraft safely to an emergency landing at Gimli Industrial

Park Airport, a former airbase at Gimli, Manitoba.

• The subsequent investigation revealed corporate failures and a chain of

minor human errors which combined to defeat built-in safeguards.

• In addition, fuel loading was miscalculated through misunderstanding of

the recently adopted metric system which replaced the Imperial system.


http://en.wikipedia.org/wiki/Air_Canada

http://en.wikipedia.org/wiki/Aviation

http://en.wikipedia.org/wiki/Boeing_767

http://en.wikipedia.org/wiki/Montreal

http://en.wikipedia.org/wiki/Edmonton,_Alberta

http://en.wikipedia.org/wiki/Ottawa

http://en.wikipedia.org/wiki/Gliding

http://en.wikipedia.org/wiki/Emergency_landing

• At the time of the incident, Canada was converting to the metric

system. As part of this process, the new 767s being acquired by Air

Canada were the first to be calibrated for the new system, using litres

and kilograms instead of gallons and pounds. All other aircraft were

still operating with Imperial units (gallons and pounds).

• For the trip to Edmonton, the pilot calculated a fuel requirement of

22,300 kilograms (49,000 lb).

• A dripstick check indicated that there were 7,682 litres (1,690 imp gal;

2,029 US gal) already in the tanks.


• In order to calculate how much more fuel had to be added, the crew

needed to convert the quantity in the tanks to a weight, subtract that

figure from 22,300 and convert the result back into a volume.

• A litre of jet fuel weighs 0.803 kg, so the correct calculation was:

7682 litres x 0.803 = 6169 kg

22300 kg – 6169 kg = 16131 kg

16131 kg ÷ 0.803 = 20088 litres


• Between the ground crew and flight crew, however, they arrived at an

incorrect conversion factor of 1.77 the weight of a litre of fuel in pounds.

This was the conversion factor provided on the refueller's paperwork

and which had always been used for the rest of the airline's imperial-

calibrated fleet.

• Their calculation produced:

7682 litres x 1.77 = 13597 'kg'

22300 kg – 13597 'kg' = 8703 kg

8703 kg ÷ 1.77 = 4916 litres

• They had 22,300 pounds on board instead of 22,300 kg of fuel which was

only a little over 10,000 kg, which was less than half the amount

required to reach their destination. The amount they loaded (4916 litres)

was not enough either.



EDMONTON


Wrong units

The Gimli Glider


Cesium 133 is the element most commonly

chosen for atomic clocks.

Cesium atomic clocks employ a beam of

cesium atoms. The clock separates cesium

atoms of different energy levels by magnetic

field.

To turn the cesium atomic resonance which

is 9,192,631,770 Hz (Hz= cycles/second),

into an atomic clock, it is necessary to

measure one of its transition or resonant

frequencies accurately.

Cesium atomic clock


How the NIST-F1 Cesium Fountain Clock Works

FIGURE 1: A gas of cesium atoms enters the clock's vacuum chamber. Six lasers slow the movement of the atoms, cooling them to near absolute zero and force them into a spherical cloud at the intersection of the laser beams.

FIGURE 2: The ball is tossed upward by two lasers through a cavity filled with microwaves. All of the lasers are then turned off.

FIGURE 3: Gravity pulls the ball of cesium atoms back through the microwave cavity. The microwaves partially alter the atomic states of the cesium atoms.

FIGURE 4: Cesium atoms that were altered in the microwave cavity emit light when hit with a laser beam. This fluorescence is measured by a detector (right). The entire process is repeated until the maximum fluorescence of the cesium atoms is determined. This point defines the natural resonance frequency of cesium, which is used to define the second.


Mass and Weight

Mass (m) is the quantity of matter in an object. In SI the standard of

mass is 1 kilogram (kg). This is a large quantity for most applications

in chemistry. More commonly chemists use the unit gram (g).

Weight (W) is the force of gravity (g) on an object. It is directly

proportional to mass. The mathematical expression is

W = m . g

An object has a fixed mass, but its weight may change.

Thus, an object that weighs 100.0 kg in St. Petersbugh, Russia,

weighs only 99.6 kg in Panama (about 0.4% less).

the same object would weigh only about 17 kg on the moon.


A modern electronic

balance

Laboratory weighing is one

of the most reliable

measurements we can make.

The mass of an object

weighing a few grams can be

determined to six or seven

significant figures.


Volume


Temperature

a) Freezing point

b) Boiling point


Relative Temperatures


Uncertainties

• A) Systematic errors.– Instrumental errors. (Thermometer constantly 2°C too low)– Error in method. (using not suitable method)– Personal errors. (Limitation in reading a scale)

• B) Random errors

- Source not known and can not be detected

• Precision– Reproducibility of a measurement.

• Accuracy– How close to the real value.


1.4 Precision and Accuracy in Measurements

• The precision of a set of measurements refers to how closely

individual measurements agree with one another.

• The precision is good (or high) if each of the measurements in a

set is close to the average of the set.

• The accuracy of a set of measurements refers to how close the

average of the set comes to the true, or most probable, value.

• Measurements of high precision are more likely to be accurate

than those of poor precision, but even high precise measurements

are sometimes inaccurate.


Rules for Significant Figuresin Calculations

KEY POINT: A calculated quantity can be no more

precise than the least precise data used in the

calculation

Analogy: a chain is only as strong as its

weakest link

… and the reported result should reflect this fact


Significant Figures


Significant Figures

Counting Sign. Fig.: from left and from first non-zero digit.

Numbers

6.29 g0.00348 g9.0 1.0 10-8

100 g = 3.14159

SignificantFigures

322

bad notationvarious

3


Significant figures

Multiplying and dividing :

Use the fewest significant figures.

0.01208 0.236

= 0.0512 or

= 5.12 10-2

= 0.051186

Adding and subtracting :

Use the number of decimal places in the number with thefewest decimal places.

1.14 0.611.67613.416 13.4


Significant figures

Rounding off :

3rd digit is increased if4th digit 5

Report to 3 significant figures.

10.235 12.4590 19.75 15.651

10.212.519.815.7


Significant Figuresin Calculations

Multiplication and Division:

the reported results should have

no more significant figures than

the factor with the fewest

significant figures

1.827 m × 0.762 m = ?

EOS

0.762 has 3 sig figs so the

reported answer is 1.39 m2


A Problem-Solving Method

Chemistry problems usually require calculations,

and yield quantitative (numerical) answers

For example,1 inch = 2.54 cm

EOS

The unit-conversion method is useful

for solving most chemistry problems –

the focus here is on “unit equivalents”


1.5 A Problem-Solving Method

Much of the study of chemistry involves solving

problems.

The Unit-Conversion Method

In chemical calculations frequently one unit is converted

to another unit.

In these calculations a conversion unit is used to convert

one unit to another.

By using conversion units stepwisely (usually more than

one step), calculations can be performed easily and

correcly.


Conversion units

are obtained

from tables


Conversion

What is the mass of a cube of osmium in grams,

that is 1.25 inches on each side?

Have volume, need density= 22.48g/cm3 (from table)


Two Examples

How many cm are in 26 inches?

26 in × cmin

2.541

= 66 cm? cm =


Two Examples

EOS

How many mm are in 1.25 foot?


Density

d = m/V

kg/m3 g/cm3 g/mL

Although the mass of an object remains constant as

the temperature is raised, the volume generally

increases – the object expands – and therefore its

density decreases.

It is usually necessary to state the temperature at

which a density is measured.


Density is the ratio of mass to volume:

md = –––

VDensity can be used as a conversion factor.

For example, the density of methanol is 0.791 g/mL;

therefore, there are two conversion factors, each equal

to one:

Density: A Physical Property and Conversion Factor

0.791 g methanol–––––––––––––– and 1 mL methanol

1 mL methanol––––––––––––––0.791 g

methanol


Measuring Density of Solids

A comparison of some densities

The liquids are as follows: Hexane (d = 0.66 g/cm3), which does not mix with water, floats on the water. Water (d = 1.00 g/cm3), which does not mix with chloroform, floats on the chloroform. Chloroform (d = 1.48 g/cm3) floats on the mercuryMercury (d = 13.6 g/cm3) at the bottom. Wood, a material of variable composition, has a range of densities. The densities listed here are representative for three types of wood. Balsa wood (d = 0.11 g/cm3) has such a low density that it floats on hexane; padouk wood (d = 0.86 g/cm3) floats on water but not on hexane; and ebony wood (d = 1.2 g/cm3) sinks in water but floats on chloroform. Liquid mercury (d = 13.6 g/cm3). is so dense that copper (d = 8.94 g/cm3) floats on it.


Measuring volumes of irregular shapes

The volume of the

irregularly shaped

objects can be

measured by water

displacement in a

graduated cylinder.


Determining body volume

It is easy to measure a person's mass (weight), but what about a person's volume? When submerged in water, a body displaces its own volume of water. The difference in the person's weight in air and when submerged in water equals the mass of displaced water. This mass, divided by the density of water, yields the volume of displaced water and--with an appropriate correction for the volume of air in the lungs and gases in the intestine--the person's volume (see Problem 103).

Notes:This method is used for any irregular object.


Mass off floating wood

The floating block of

wood displaces water

from the container.

The mass of the water

displaced is equal to the

mass of the wood object.


A pycnometer can be used to

determine the densities of liquids

very precisely.

The exact volume of the pycnometer

must be determined by first

determining the mass of the empty,

dry pycnometer, then filling it with

water of a known temperature and

determining the mass of the water

and pycnometer. Using the formula: V

= Mwater/Dwater, the volume can be

calculated.

The pycnometer is again dried, then

filled with the liquid. The mass of the

liquid, divided by the known volume,

gives the density of the liquid.

Density of Liquids

Diff: 9.9570 gVol: 9.9750 ml

Diff: 8.7680 g d: 0.8790 g/ ml


Separations

general chemistry: chapter 1 slide 1 of 57 c h e m 1 0 0 general chemistry lecturer: prof. dr....

Documents

o slide

general chemistry lecturer

key terms chemistry

various uses slide

tamerkan zgen slide

chemical properties

chemical change

sample of matter molecules