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Mohd Faisol Mansor [email protected]

TABLE OF CONTENTS

CHAPTER 1 : INTRODUCTION TO CHEMISTRY Page 1

CHAPTER 2 : THE STRUCTURE OF THE ATOM Page 5

CHAPTER 3 : CHEMICAL FORMULAE & EQUATIONS 2

CHAPTER 4 : PERIODIC TABLE OF ELEMENTS Page 54

CHAPTER 5 : CHEMICAL BONDS Page 79

CHAPTER 6 : ELECTROCHEMISTRY Page 100

CHAPTER 7 : ACIDS & BASES Page 130

CHAPTER 8 : SALTS Page 148

CHAPTER 9 : MANUFACTURED SUBTSANCES Page 175

: IN INDUSTRY

mohd faisol mansor/chemistry form 4/chapter 1

1

CHAPTER 1 INTRODUCTION TO CHEMISTRY

Scientific Method : systematic method used to solve problems in Science.

Chemistry

Is the study of the composition,

structure, properties & interactions of

matter.

Uses of Chemical

Salt (NaCl) Vinegar (CH3COOH)

Tiles/Cement (CaCO3)

Careers

Doctor Chemical Engineer

Geologist Biochemist Pharmacist

Chemical Industry

Petroleum Food

Chemical Substance Electronic Agriculture

Contribution of Chemical Industries

Financial Job opportunity Development of

country

Making Observation

Making Inference

Identifying Problem

Making Hypothesis

Identifying Variables

Controlling Variables

Planning an Experiment

Collecting Data

Making Conclusion

Writing Report

Interpreting Data


2


3

Classify each of the following element into their group.

Hydrogen, H Oxygen, O Zinc, Zn

Chlorine, Cl Potassium, K Silver, Ag

Sodium, Na Carbon, C Phosphorus, P

Magnesium, Mg Sulphur, S Fluorine, F

Aluminium, Al Lithium, Li Barium, Ba

Copper, Cu Bromine, Br Lead, Pb

Neon, Ne Argon, Ar Calcium, Ca

Helium, He Iron, Fe Iodine, I

Metal Atom Nonmetal Atom


4

Classify each of the following element/compound into their group.

Sodium chloride, NaCl Hydrogen gas, H2

Chlorine gas, Cl2 Lead(II) oxide, PbO

Aluminium oxide, Al2O3 Silver bromide, AgBr

Carbon dioxide, CO2 Naphthalene, C10H8

Sugar, C6H12O6 Copper(II) iodide, CuI2

Oxygen gas, O2 Magnesium oxide, MgO

Zinc chloride, ZnCl2 Methane, CH4

Ethanol, C2H6O Water, H2O

Ionic Compound (Metal Atom + Nonmetal Atom)

Molecule (Nonmetal Atom + Nonmetal Atom)


5

CHAPTER 2 THE STRUCTURE OF THE ATOM

Matter

Anything that occupies space & has

mass.

Type of Particles

Atom Molecule

Ion

Atom

Smallest particle of an element that can

take place in chemical reaction.

Ion

Is a positive charged or negative charged

particle.

Element

Substance that consists of only one

type of atom.

Compound

Substance that contains two or more elements

that are chemically bonded together.

Physical State

Solid Liquid Gas

Subatomic Particles

Electron Proton

Neutron

Diffusion Is a random movement

between different particles from high

concentration to lower concentration.

Molecule

A group of two or more atoms which are

chemically bonded together.

Matter is made up of tiny & discrete particle.

Particle vibrate, moving & collide with each other.

Particles move randomly. There are forces between

the particle. Stronger the forces, particle close to each other.

Higher the temperature, higher the kinetic energy of particle.

Kinetic Theory of Matter


6

Example: Formation of WATER

Atom

+ Atom

Molecule (ELEMENT)

Atom

+ Atom

Molecule (COMPOUND)

Cation

+ Anion

Ionic Compound + + - -

+

Oxygen Atom

Hydrogen Atom

Water Molecule

A Glass of Water

Atom, Molecule & Ion


7

Diffusion in Solid: Test tube is filled with a hot jelly and copper(II) sulphate crystal.

Diffusion in Liquid: A beaker is filled with water & potassium manganate (VII).

Diffusion in Gas: A gas jar is filled with few drop of bromine liquid.

Difussion Rate:

Reason:

Difussion Rate:

Reason:

Difussion Rate:

Reason:


8

States Particles

Arrangement

Forces between

particles Particles Motion

SOLID

Packed closely in

orderly manner

Held together

very strong

Vibrate and rotate at

fixed position

LIQUID

Packed closely not in

orderly arrangement

Strong but weaker

than solid

Vibrate, rotate and

move through liquid &

collide against each

other

GAS

Far apart from each

other

weak Vibrate, rotate and

move freely

Solid

Liquid

Gas

Shape :

Ability to be

compressed:

Shape :

Ability to be

compressed:

Shape :

Ability to be

compressed:


9

Change in the State of Matter

Changes in the Heat Energy


10

Heating of Naphthalene Cooling of Naphthalene

1. Why is solid naphthalene, C10H8 not heated directly with Bunsen burner?

2. Why is water bath used to heat the naphthalene?

3. During the cooling of naphthalene, explain why

a) the boiling tube must place in a conical flask?

b) the naphthalene must stirred continuously?

4. What happens to the temperature of naphthalene during

a) melting? Give a reason.

b) freezing? Give a reason.

5. The melting point of sugar is 184oC. The melting point of sugar cannot be

determined using same apparatus. Why? What apparatus can be used

instead?


11

The Heating Curve of Naphthalene

* label the melting point of the naphthalene in the diagram above as MP.

Melting Point

Is the temperature at which a solid changes into

a liquid at a particular pressure.

Physical State

AB =

BC =

CD =

Temperature (oC)

Time (min)

A

B C

D

Why the temperature increase from A to B?

When solid is heated,

heat energy is

________________.

This cause particles to

_________ kinetic energy

and vibrate __________.

Why the temperature remains constant at BC?

Heat energy ___________

by the particles is used to

overcome the force

between the particles so

that the solid turn into

liquid.

Why the temperature increase from C to D?

When liquid is heated,

heat energy is

________________.



and move __________.

Also known as

latent heat of

fusion.


12

The Cooling Curve of Naphthalene

* label the freezing point of the naphthalene in the diagram above as FP.

Freezing Point

Is the temperature at which a liquid changes

into a solid at a particular pressure.

Physical State

EF =

FG =

GH =

Why the temperature decrease from E to F?

When liquid is cooled,

heat energy is

________________.



and move __________.

Why the temperature remains constant at FG?

Heat energy ___________

to the surroundings is

exactly same as the heat

energy released by

particle to form the

forces to become a solid.

Why the temperature decrease from G to H?

When solid is cooled,

heat energy is

________________.



and vibrate __________.

Temperature (oC)

Time (min)

E

F G

H

How to avoid supercooling?


13

The Atomic Structure

[Draw the atomic model & briefly explain]

(1) John Dalton

(1766 – 1844)

(2) J.J. Thomson

(1856 – 1940)

(3) Ernest Rutherford

(1871 - 1937)

(4) Neils Bohr

(1885 - 1962)

(5) James Chadwick

(1891 - 1974)


14

Subatomic Particles of an Atom

PROTON

Symbol : p

Relative electric

charge : +1

Relative mass: 1

NEUTRON

Symbol : n

Relative electric

charge : 0

Relative mass: 1

ELECTRON

Symbol : e

Relative electric

charge : -1

Relative mass: 0.0005

NUCLEUS

Proton + Neutron

NUCLEON NUMBER =

Proton + Number of

Number Neutron

NEUTRAL ATOM

Number of proton

is equal to the

number of electron.

Definition

Nucleon number is

the total number of

proton and neutrons

in its atom.

Definition

Proton number is the

number of protons in

its atom.

Atom Q has a proton number

of 11 and a nucleon number of

23. How many protons,

electrons and neutrons are

present in the atom?


15

Symbol of Elements

Element Symbol Element Symbol Element Symbol Element Symbol

Hydrogen

Carbon

Sodium

Sulphur

Helium

Nitrogen

Magnesium

Chlorine

Lithium

Oxygen

Aluminium

Argon

Beryllium Fluorine Silicon Potassium

Boron

Neon

Phosphorus

Calcium

Na

23

11

Also represented

by sodium-23

A fluorine atom has 9 protons and

10 neutrons. What is the proton

number and nucleon number of

the atom? Represent the atom in

the form of .

Bromine-80 has 45 neutrons.

What are the proton number

and nucleon number of the

atom? Represent the atom

in the form of .


16

ISOTOPES

Atoms of the same element has same proton number but

different number of neutron

Element Nucleon

Number

Proton

Number

No. of

Neutron

No. of

Electron

Hydrogen-1 1

Hydrogen-2

Hydrogen-3

Carbon-12 6

Carbon-13

Carbon-14

Oxygen-16 8

Oxygen-17

Oxygen-18

Chlorine-35 17

Chlorine-37

Bromine-80 35

Bromine-81

USES OF ISOTOPES

To treat

cancer.

To determine

the age of

artifact.

To detect leaks in

underground

petroleum pipe.

To detect the

location of a

tumor in the brain.

To determine the

rate of absorption

of fertilisers by

plants.

Why isotopes element has similar chemical

properties?


17

The Electronic Structure of an Atom

The last shell filled

with electrons

known as the

outermost

occupied shell.

Maximum number

of electron filled in

the shell:

1st = 2 electrons

2nd = 8 electrons

3rd = 8 electrons

4th = 8 electrons

For atoms with

more than 20

electrons, the third

shell can filled up

to 18 electrons.

Valence Electrons

Electrons in the

outermost

occupied shell.

Electron

Configuration

Chlorine has

proton number 17.

Write the electron

arrangement for

potassium & state the

valence electrons.


18

Element Proton

Number

Number of

Electron

Electron

Configuration

& Valence

Electron

Electronic Structure

Hydrogen

1

E.C =

V.E =

Helium

2

Lithium

3

Beryllium

4

Boron

5

Carbon

6

Nitrogen

7


19

Element Proton

Number

Number of

Electron

Electron

Configuration

& Valence

Electron


Oxygen

8

Fluorine

9

Neon

10

Sodium

11

Magnesium

12

Aluminium

13


20

Element Proton

Number

Number of

Electron

Electron

Configuration &

Valence

Electron


Silicon

14

Phosphorus

15

Sulphur

16

Chlorine

17

Argon

18


21

Element Proton

Number

Number of

Electron

Electron

Configuration &

Valence

Electron


Potassium

19

Calcium

20


22

CHAPTER 3 CHEMICAL FORMULAE & EQUATIONS

Relative Molecular Mass

Average mass of

one molecule

1/12 x mass of an

atom of carbon-12

Relative Atomic Mass

Average mass of

one atom of an

element

1/12 x mass of an

atom of carbon-12

Relative Molecular Mass (RMM)/Relative Formula Mass (RFM) can

be calculated by adding up the Relative Atomic Mass (RAM).

Water, H2O

RMM =

Sodium Chloride, NaCl

RFM =

Relative Formula Mass

Average mass of

one formula unit

1/12 x mass of an

atom of carbon-12

+ -

Example

1) Element mercury is 20 times

heavier than helium. Determine

the relative atomic mass of

element mercury if the relative

atomic mass of helium is 4.

2) The relative atomic mass of helium,

nitrogen and sulphur is 4, 14, and 32

respectively.

a) How many times is one atom of

sulphur heavier than one atom of

helium.

b) Calculate the number of atoms of

helium that have the same mass as two

atoms of nitrogen.


23

1. Calculate relative molecular mass of the following element or compound.

a) Oxygen gas, O2

b) Chlorine gas, Cl2

c) Carbon dioxide, CO2

d) Ammonia, NH3

e) Iodine gas, I2

f) Sulphur dioxide, SO2

g) Sugar, C6H12O6

h) Ethanol, C2H6O

2. Calculate relative formula mass of the following compound.

a) Magnesium oxide, MgO

b) Potassium iodide, KI

c) Calcium carbonate, CaCO3

d) Copper(II) nitrate, Cu(NO3)2

e) Aluminium oxide, Al2O3

f) Zinc Sulphate, ZnSO4

g) Hydrated magnesium sulphate,

MgSO4.7H2O

h) Hydrated copper(II) sulphate,

CuSO4.5H2O


Example


24

The Mole, Number of Particles, Mass & Volume of Substances.

MOLE

Amount of substance

that contains as many

particles as the number

of atoms in exactly 12 g

of carbon-12

NUMBER OF

PARTICLES

One mole of

substance contains

6.02 x 1023 particles.

MASS OF

SUBSTANCES

The mass of one

mole of the

substance equal to

the mass of 6.02 x

1023 particles.

VOLUME OF GAS

One mole of any gas

always has the same

volume under the

same temperature &

pressure.

Avogadro Constant

NA = 6.02 x 1023

mole = no of particles

NA

No of particles

= mole x NA

Molar Mass

= RAM/RMM/RFM

mole = Mass

Molar Mass

Mass = mole x MM

Molar Volume

1) Room Condition

= 24 dm3 mol -1

2) At S.T.P

= 22.4 dm3 mol -1

mole = Volume

Molar Volume

Volume = mole x MV

Unit conversion

1 dm3 = 1000 cm3


25

1. A closed glass bottle contains 0.5 mol of oxygen gas, O2.

a) How many oxygen molecules, O2 are there in the bottle?

b) How many oxygen atoms are there in the bottle?

2. Find the number of moles of hydrogen gas, H2 containing

a) 3.01 x 1024 hydrogen molecule, H2

b) 6.02 x 1023 hydrogen atoms.

3. Find the number of moles of molecules in a sample containing 9.03 x 1023

molecules of carbon dioxide, CO2.

4. A sample contains 6.02 x 1025 molecule of water. How many moles of water

are there in the sample?

Example 1


26

5. A container contains 1.806 x 1023 oxygen molecules, O2. A sample of 0.5 mol

of oxygen gas, O2 is added to the container. How many molecules are there

altogether in the container?

6. Calcium is needed for the formation of bones and teeth. How many calcium

ions are there in a serving of cereal that contains 0.007 mol of calcium ions?

7. A beaker contains 0.1 mol of zinc chloride, ZnCl2

a) Calculate the number of moles of chloride ions in the beaker.

b) Find the total number of ions in the beaker.


27

1. What is the mass of

a) 0.1 mol of magnesium? [RAM: Mg, 24]

b) 2.408 x 1023 atoms of magnesium? [RAM: Mg, 24 ; NA = 6.02 x 1023]

2. How many moles of molecules are there in 16 g of sulphur dioxide gas, SO2?

[RAM: O, 16 ; S, 32]

3. How many chloride ions are there in 27.2 g of zinc chloride, ZnCl2?

[RAM: Cl, 35.5 ; Zn, 65 ; NA = 6.02 x 1023]

Example 2


28

4. What is the mass of carbon that contains 6.02 x 1023 carbon atoms?

5. What is the mass of

a) 0.01 mol of ammonia gas, NH3?

b) 6.02 x 1024 nitrogen molecules, N2?

6. How many moles of molecules are there in 2.8 g of carbon monoxide, CO?


29

1. What is the volume of 1.2 mol of ammonia gas, NH3 at STP?

[Molar volume: 22.4 dm3 mol-1]

2. How many moles of ammonia gas, NH3 are present in 600 cm3 of the gas

measured at room conditions? [molar volume: 24 dm3 mol-1]

3. Calculate the volume of the following gases.

a) 0.3 mol of oxygen gas, O2, at room condition.

b) 4 mol of helium gas measured at STP.

4. Calculate the number of moles of 48 dm3 of chlorine gas, Cl2, at room

condition.

Example 3


30

1. What is the volume of 12.8 g of oxygen gas, O2, in cm3, at STP?

[RAM: O, 16 ; Molar volume: 22.4 dm3 mol-1]

2. How many molecules of carbon dioxide, CO2, are produced when 120 cm3 of

the gas is released during chemical reaction between an acid and a

carbonate at room conditions?

[Molar volume: 24 dm3 mol-1 ; NA = 6.02 x 1023]

3. What is the mass of 0.6 dm3 of chlorine gas, Cl2 at room condition?

[RAM: Cl, 35.5 ; Molar volume = 24 dm3 mol-1]

4. 3 dm3 of an unknown gas has a mass of 6.0 g at room conditions. Find the

molar mass of the gas.

Example 4


31

CHEMICAL FORMULAE

A representation

of a chemical

substance using

letters and

subscript numbers.

Magnesium Nitrate

Mg(NO3)2

Water

H2O

Empirical Formula

The simplest

number ratio of

atoms in the

compound.

Molecular Formula

The actual number

of atoms that are

present in the

compound.

Molecular Formula = (Empirical Formula)n

Compound Molecular

Formula

Empirical

Formula n

Water H2O

Ethene CH2 2

Glucose C6H12O6

RMM of Molecular Formula

RMM of Empirical Formula

= n

The empirical formula of a compound

is CH2. Its relative molecular mass is 42.

Find its molecular formula. [RAM: H, 1 ; C, 12]

Copper(II) Oxide

Magnesium Oxide

[state the number of particles consist in the substance above]


32

To determine Empirical Formulae of Magnesium Oxide

To determine Empirical Formulae of Copper(II) Oxide

1. Why is the magnesium ribbon cleaned with

sand paper before used?

2. Name the white fumes produced.

3. State the reason:

a) covering the crucible with its lid as soon as

the magnesium start burning.

b) raising the lid of the crucible at intervals

during heating.

c) heating, cooling & weighing are repeated

until constant mass is obtained.

4. Why is it important not to let any white fumes

escape from the crucible?

1. Why do we start off with copper(II) oxide instead of allowing copper to react with

oxygen in the air in this experiment?

2. How do you test that the air in the tube has been removed completely?

3. Explain what will happen if we burn excess hydrogen gas without removing the air

completely in combustion tube?

4. Why we need to continue the flow of hydrogen gas after the heating of copper(II)

oxide?

5. Why do we need to repeat heating, cooling and weighing until constant mass is

obtained?


33

a) A sample of aluminium oxide contains 1.08 g of aluminium and 0.96 g of

oxygen. What is the empirical formula of this compound? [RAM: O, 16 ; Al, 27]

Element Al O

Mass of Element (g)

Number of Mole

Ratio of Mole

Simplest Ratio

Empirical Formula of Aluminium Oxide =

b) 0.20 g of calcium reacts with fluorine to give 0.39 g of calcium fluoride. Find

the empirical formula of the calcium fluoride produced. [RAM: F, 19 ; Ca, 40]

Example 1


34

c) Find the empirical formula of a compound that consists of 32.4% of sodium,

22.6% of sulphur and 45.0% of oxygen. [RAM: O, 16, Na, 23 ; S, 32]

d) 60 g of aluminium sulphide contains 38.4 g of sulphur. Find the empirical

formula of the compound. [RAM: Al, 27 ; S, 32]


35

a) Butane has empirical formula of C2H5 and relative molecular mass of 58. Find

its molecular formula.

b) Ethanoic acid is an important ingredient of vinegar. The empirical formula of

this acid is CH2O. Given that its molar mass is 60 g mol-1, find its molecular

formula.

c) 6.24 g of element X combines with 1.28 g of oxygen to produce a compound

with an empirical formula of X2O. What is relative atomic mass of X?

[RAM: O, 16]

Example 2


36

d) Element Y react with oxygen to produce a compound with molecular

formula YO3. Given that the mass of 1 mol of the compound is 80 g.

Determine the relative atomic mass of element Y.

e) Determine the percentage composition by mass of water in hydrated

copper(II) sulphate, CuSO4.5H2O. [RAM: H, 1 ; O, 16 ; S, 32 ; Cu, 64]

f) Due to its high nitrogen content, urea, CO(NH2)2 is commercially used as

fertilizers. Calculate the percentage composition by mass of nitrogen in urea,

CO(NH2)2. [RAM: H, 1 ; C, 12 ; N, 14 ;O, 16]


37

Metal

Atom

Nonmetal

Atom

Cation

(+ve ion)

Anion

(-ve ion)

IONIC

COMPOUND

Ionic

Formulae

Silver

Chlorine

Silver

Chloride

Ionic

Formulae

Zinc

Zn

Oxygen

O

Zn2+

O2-

Zinc Oxide

ZnO

Ionic

Formulae


38

Charge Name of Cation Formula of Cation

1+

Hydrogen ion H +

Lithium ion Li +

Sodium ion Na +

Potassium ion K +

Silver ion Ag +

Ammonium ion NH4 +

2+

Barium ion Ba 2+

Calcium ion Ca 2+

Magnesium ion Mg 2+

Zinc ion Zn 2+

Copper(II) ion Cu 2+

Iron(II) ion Fe 2+

Lead(II) ion Pb 2+

3+ Aluminium ion Al 3+

Iron(III) ion Fe 3+

Charge Name of Anion Formula of Anion

1-

Hydroxide ion OH -

Chloride ion Cl -

Fluoride ion F -

Bromide ion Br -

Iodide ion I -

Nitrate ion NO3 -

2-

Oxide ion O 2-

Sulphate ion SO4 2-

Carbonate ion CO3 2-

3- Phosphate ion PO4 3-


39

Name of Cation Formula of

Cation

Hydrogen ion

Lithium ion

Sodium ion

Potassium ion

Silver ion

Ammonium ion

Barium ion

Calcium ion

Magnesium ion

Zinc ion

Copper(II) ion

Iron(II) ion

Lead(II) ion

Aluminium ion

Iron(III) ion


Cation

Hydrogen ion

Magnesium ion

Barium ion

Potassium ion

Iron(II) ion

Ammonium ion

Zinc ion

Aluminium ion

Lithium ion

Iron(III) ion

Sodium ion

Calcium ion

Silver ion

Copper(II) ion

Lead(II) ion

Name of Anion Formula of

Anion

Hydroxide ion

Chloride ion

Fluoride ion

Bromide ion

Iodide ion

Nitrate ion

Oxide ion

Sulphate ion

Carbonate ion

Phosphate ion


Anion

Phosphate ion

Bromide ion

Oxide ion

Carbonate ion

Iodide ion

Chloride ion

Sulphate ion

Nitrate ion

Hydroxide ion

Fluoride ion


40


Cation

Hydrogen ion

Lithium ion

Sodium ion

Potassium ion

Silver ion

Ammonium ion

Barium ion

Calcium ion

Magnesium ion

Zinc ion

Copper(II) ion

Iron(II) ion

Lead(II) ion

Aluminium ion

Iron(III) ion


Anion

Hydroxide ion

Chloride ion

Fluoride ion

Bromide ion

Iodide ion

Nitrate ion

Oxide ion

Sulphate ion

Carbonate ion

Phosphate ion


Anion

Hydroxide ion

Chloride ion

Fluoride ion

Bromide ion

Iodide ion

Nitrate ion

Oxide ion

Sulphate ion

Carbonate ion

Phosphate ion


Cation

Hydrogen ion

Lithium ion

Sodium ion

Potassium ion

Silver ion

Ammonium ion

Barium ion

Calcium ion

Magnesium ion

Zinc ion

Copper(II) ion

Iron(II) ion

Aluminium ion

Iron(III) ion

Lead(II) ion


41


Cation

Hydrogen ion

Lithium ion

Sodium ion

Potassium ion

Silver ion

Ammonium ion

Barium ion

Calcium ion

Magnesium ion

Zinc ion

Copper(II) ion

Iron(II) ion

Lead(II) ion

Aluminium ion

Iron(III) ion


Anion

Hydroxide ion

Chloride ion

Fluoride ion

Bromide ion

Iodide ion

Nitrate ion

Oxide ion

Sulphate ion

Carbonate ion

Phosphate ion


Anion

Hydroxide ion

Chloride ion

Fluoride ion

Bromide ion

Iodide ion

Nitrate ion

Oxide ion

Sulphate ion

Carbonate ion

Phosphate ion


Cation

Hydrogen ion

Lithium ion

Sodium ion

Potassium ion

Silver ion

Ammonium ion

Barium ion

Calcium ion

Magnesium ion

Zinc ion

Copper(II) ion

Iron(II) ion

Aluminium ion

Iron(III) ion

Lead(II) ion


42

Construct the chemical formula for each of the following ionic compound.

a)Magnesium hydroxide

b) Silver iodide c) Potassium Bromide

d) Zinc nitrate

e) Sodium carbonate f) Aluminium oxide

g) Copper(II) iodide

h) Iron(II) sulphate i) Magnesium oxide

j)Calcium carbonate

l) Ammonium phosphate m) Sodium hydroxide

n) Zinc bromide

o) Lead(II) nitrate p) copper(II) sulphate

Example


43

1. Write the formula of the following substances.

a) Potassium iodide =

b) Magnesium oxide =

c) Carbon dioxide =

d) Copper(II) oxide =

e) Lead(II) bromide =

f) Calcium chloride =

g) Hydrochloric acid =

h) Copper(II) sulphate =

i) Hydrogen gas =

j) Water =

k) Sulphuric acid =

l) Silver chloride =

m) Potassium nitrate =

n) Calcium carbonate =

o) Aluminium oxide =

p) Oxygen gas =

q) Oleum =

r) Ammonium sulphate =

s) Sodium chloride =

t) Zinc oxide =

Example

1) Nitric acid =

2) Lead(II) iodide =

3) Copper(II) nitrate =

4) Zinc sulphate =

5) Iron(II) chloride =

6) Iron(III) chloride =

7) Chlorine gas =

8) Potasium nitrate =

9) Silver nitrate =

10) Magnesium bromide =

11) Zinc chloride =

12) Sodium hydroxide =

13) Ammonia =

14) Iron(II) sulphate =

15) Lead(II) oxide =

16) Carbon monoxide =

17) Magnesium sulphate =

18) Ammonium nitrate =

19) Potassium hydroxide =

20) Lithium oxide =


44




c) Carbon dioxide =






i) Hydrogen gas =

j) Water =

k) Sulphuric acid =





p) Oxygen gas =

q) Oleum =



t) Zinc oxide =

1) Nitric acid =



4) Zinc sulphate =



7) Chlorine gas =


9) Silver nitrate =


11) Zinc chloride =


13) Ammonia =







20) Lithium oxide =

Example


45




c) Carbon dioxide =






i) Hydrogen gas =

j) Water =

k) Sulphuric acid =





p) Oxygen gas =

q) Oleum =



t) Zinc oxide =

1) Nitric acid =



4) Zinc sulphate =



7) Chlorine gas =


9) Silver nitrate =


11) Zinc chloride =


13) Ammonia =







20) Lithium oxide =

Example


46

Name the following ionic compound by using their IUPAC name.

Ionic

Formula Name

Ionic

Formula Name

NaCl KI

MgO BaSO4

Cu(NO3)2 CaCO3

Al2O3 FeCl3

ZnCl2 LiOH

CuO FeSO4

AgNO3 NaOH

MgBr2 ZnO

PbSO4 PbI2

Example

NaBr

Sodium Bromide

Br -

Bromide

Na +

Sodium


47

CHEMICAL EQUATIONS

a) Qualitative Aspect

For each equation, identify the reactant(s), product(s) and the state of each

of them. Then, balance the equation.

a) H2 (g) + O2 (g) H2O (l)

b) CuO (s) + HCl (aq) CuCl2 (aq) + H2O (l)

c) Cl2 (g) + NaBr (aq) NaCl (aq) + Br2 (l)

d) Mg (s) + HCl (aq) MgCl2 (aq) + H2 (g)

K(S) + H2O(l) KOH(aq) + H2(g)

Reactant

Product

Meaning: Solid Potassium react with water liquid to give

potassium hydroxide solution and hydrogen gas

p/s: 1) Able to classify reactant and product.

2) Able to balance the equation.

Example 1


48

Write a balanced equation for each of the following reactions.

a) Carbon monoxide gas + oxygen gas Carbon dioxide gas

b) Hydrogen gas + nitrogen gas Ammonia gas

c) Aluminium + iron(III) oxide Aluminium oxide + iron

d) Ammonia gas react with oxygen gas to yield nitrogen monoxide gas and

water.

e) Silver nitrate solution is added to calcium chloride solution. Silver chloride

precipitate and calcium nitrate solution are produced.

f) When solid zinc carbonate is heated, it decomposes into zinc oxide powder

and carbon dioxide gas.

Example 2


49

1. Construct balanced chemical equations:

a) Magnesium react with oxygen will produce magnesium oxide.

b) Sodium metal react with chlorine gas will produce sodium

chloride.

c) Potassium oxide react with water will produce potassium

hydroxide.

d) Lithium metal react with water will produce lithium hydroxide

and hydrogen gas.

e) Zinc metal react with water will produce zinc oxide and

hydrogen gas.

f) Calcium carbonate react with hydrochloric acid will produce

calcium chloride, water and carbon dioxide.

g) Hydrochloric acid react with sodium hydroxide will produce

sodium chloride and water.

h) Potassium oxide react with nitric acid will produce potassium

nitrate and water.

Example 3


50

i) Iron metal react with chlorine gas will produce iron(III) chloride.

j) Magnesium metal react with nitric acid with produce

magnesium nitrate and hydrogen gas.

k) Zinc metal dissolved in copper(II) chloride will produce zinc

chloride and copper metal.

l) Chlorine gas react with potassium bromide will produce

potassium chloride and bromine gas.

m)Copper(II) carbonate when heated will produce copper(II)

oxide and carbon dioxide.

n) Lead(II) nitrate when heated will produce lead(II) oxide,

nitrogen gas and oxygen gas.

o) Potassium iodide react with lead(II) nitrate will produce lead(II)

iodide and potassium nitrate solution.

p) Sodium hydroxide react ammonium chloride will produce

sodium chloride, water and ammonia gas.

q) Zinc metal react with hydrochloric acid will produce zinc

chloride and hydrogen gas.

r) Magnesium oxide react with sulphuric acid will produce

magnesium sulphate react with water.


51

b) Quantitative Aspect

1. Copper(II) oxide, CuO reacts with aluminium according to the following

equation.

3CuO (s) + 2Al (s) Al2O3 (s) + 3Cu (s)

Calculate the mass of aluminium required to react completely with 12 g of

copper(II) oxide, CuO. [RAM: O, 16 ; Al, 27 ; Cu, 64]

2H2 (g) + O2 (g) 2H2O(l)

2 molecule

Or

2 mol

Note: The coefficient in the reaction tell the exact proportions of

reactant and product in chemical reaction.

Example

1 molecule

Or

1 mol

2 molecule

Or

2 mol


52

2. A student heats 20 g of calcium carbonate, CaCO3 strongly. It decomposes

according to the equation below.

CaCO3 (s) CaO (s) + CO2 (g)

a) If the carbon dioxide produced is collected at room conditions, what is its

volume?

b) Calculate the mass of calcium oxide, CaO produced.

[RAM: C, 12 ; O, 16 ; Ca, 40 ; Molar volume = 24 dm3 mol-1]


53

3. Hydrogen peroxide, H2O2 decomposes according to the following equation.

2H2O2 (l) 2H2O (l) + O2 (g)

Calculate the volume of oxygen gas, O2 measured at STP that can obtained

from the decomposition of 34 g of hydrogen peroxide.

[RAM : H, 1 ; O, 16 ; Molar volume = 22.4 dm3 mol-1]

4. 16 g of copper(II) oxide, CuO is reacted with excess methane, CH4. Using the

equation below, find the mass of copper that is produced.

4CuO (s) + CH4 (g) 4Cu (s) + CO2 (g) + 2H2O (l)

[RAM : H, 1 ; C, 12 ; O, 16 ; Cu, 64]


54

CHAPTER 4 PERIODIC TABLE OF ELEMENTS


55

PERIODIC TABLE


56

Historical Development of Periodic Table

Johann Dobereiner

(1780 – 1849)

divided element

according their

similar chemical

properties

Antoine Lavoiser

(1743 – 1794)

first scientist classify

substances.

John Newlands

(1837 – 1898)

arranged element

in order of

increasing atomic

mass

Lothar Meyer

(1830 – 1895)

showing that the

properties of the elements

formed a periodic

pattern against their

atomic masses.

Dmitri Mendeleev

(1834 – 1907)

arranged the elements in

order of increasing atomic

mass and

grouped them

according to similar

chemical properties.

Henry J.G. Moseley

(1887 – 1915)

Studied the X-ray

spectrum of elements.

He rearranged the

elements in order of

increasing proton number.

Concluded that the

proton number should

be the basis for the

periodic change of

chemical properties

instead of the atomic

mass.

Leading to the modern

Periodic Table.


57

Arrangement of Element in the Periodic Table

Element in the

Periodic Table are

arranged in an

increasing order of

proton number

ranging 1 to 118.

Element with

similar chemical

properties are

placed in the same

vertical column

18 vertical column

called a group

7 horizontal rows

called a period

GROUP

The number of

valence electron in

an atom decides the

position of the group

of an element.

GROUP

Valence electron 1

and 2

GROUP

Valence electron 3

until 8

PERIOD

Equal to the number

of shell occupied

with electrons in its

atom

Example:

An atom of element X has a proton

number of 15. In which group and period

we can find element X in Periodic Table.


58

1. Element D has a proton number 19. Where is element D located in Periodic

Table?

2. An atom of element E has 10 neutrons. The nucleon number of element E is 19.

In which group and period of element E located in the Periodic Table?

3. An atom of element G has 3 shell occupied with electrons. It is placed in

Group 17 of the Periodic Table. What is the electron arrangement of atom G?

4. An atom Y is located in Group 18 and period 2 of the periodic Table. What is

the electron arrangement and proton number of atom Y?

5. Element R has a proton number of 11. Its atom has 6 neutrons. In which group

and period can you find element R in the Periodic Table?


Exercise


59

The advantage of grouping elements in the Periodic Table

1. Helps us to study the element systematically especially their

physical and chemical properties.

2. Element with a same number valence electron is place in the

same group because they have the same chemical properties.

3. It could be used to predict undiscovered properties.

4. To relate the characteristic of an element with atomic structure


60

Physical Properties of Group 18

GROUP 18

consist of helium,

neon, argon,

krypton, xenon

and radon

GROUP 18

Known as

NOBLE GAS

Exist as

monoatomic gas.

GROUP 18

NOBLE GAS

COLOUR SOLUBILITY

ELECTRIC

CONDUCTIVITY

MELTING &

BOILING POINT

DENSITY


61

The inert properties of G18

All noble gas are inert which means unreactive.

Noble gas are inert because the outermost occupied shell are full.

This electron arrangements are very stable.

Helium have 2 valence electron, this electron arrangement is called

duplet electron arrangement whereas other noble gas have eight

valence electron which called octet electron arrangement.

p/s: All other element try to achieve noble gas electron arrangement to

become stable.

The Physical Properties of G18

The physical properties vary down the group.

physical properties vary down the group are related to atomic size.

Melting and boiling point of an element increase when going

down the group because the atomic sizes increase, the forces

attraction (Van der Waals forces) becomes stronger. Thus more

heat energy is required to overcome the forces.

The strength of Van der Waals forces propotional to the size of

particle


62

Uses of Noble Gas Elements

Fill light bulb laser for repair retina,

photographic flash lamps

fill airships and weather

balloons

treatment cancer

advertising light and

television tubes

Car bulb, use in bubble chamber

[ State the element of group 18 ]


63

1. Table below shows the electron arrangements of atoms of elements P, Q,

and R.

Element Electron Arrangement

P 2.8

Q 2.8.18.8

R 2.8.18.32.18.8

a) Arrange the element in ascending order of boiling points. Give reasons for

your answer.

b) Elements P, Q, and R are chemically unreactive. Why?

2) What is the meaning of

a) Duplet electron arrangement

b) Octet electron arrangement

Exercise


64

Physical Properties of Group 1

GROUP 1

Known as Alkali

Metals.

Have valence

electron of 1.

Consist of elements

such as lithium,

sodium, potassium,

rubidium, caesium

and francium.

GROUP 1

ALKALI

METALS

HARDNESS APPEARANCE

ELECTRIC

CONDUCTIVITY

MELTING

POINT

DENSITY


65

A melting point of a metal indicates the strength of its metallic bonding in its

structure.

Generally, the strength of metallic bonding is directly proportional to the number

of valence electron per atom divided by the radius of a metal.

A metallic bond can be defined as the electrostatic force between the

positively charged metallic ions and the ‘sea’ of electrons.

Electropositivity is the measurement of an atom to release an electron and form

positive ion

When going down the group, what happen to the electropositivity of the

element? Why?

Element Atomic Size

Melting /

Boiling

point

Density Hardness

Lithium

Sodium

Potassium

Rubidium

Caesium

Francium


66

Chemical Properties of Group 1 Elements

Exercise :

1. Why does Group 1 element have similar chemical properties?

2. Why is paraffin oil used to store metals such as sodium and potassium?

3. Element in Group 1 has similar chemical properties but differ in reactivity.

Why?

React vigorously with water to produce alkaline metal hydroxide solution

and hydrogen gas, H2.

2Li (s) + 2H2O 2LiOH (aq) + H2 (g)

Burn in oxygen gas rapidly to produce solid metal oxide

4Li (s) + O2 (g) 2 Li2O (s)

Solid metal oxide dissolve in water to form alkaline metal hydroxide

Li2O (s) + H2O 2LiOH (aq)

Alkaline metal burn in chlorine gas, Cl2 to form white solid metal chloride

2Na (s) + Cl2 2NaCl (aq)


67

To investigate the chemical properties of Lithium, Sodium & Potassium

[ Discussion ]

1. Why are the experiments involving sodium and potassium demonstrated by

your teacher and not carried out by the students?

2. Write the chemical equations for the reactions of lithium, sodium and

potassium with

a) Water, H2O

Lithium :

Sodium :

Potassium :

b) Oxygen gas, O2

Lithium :

Sodium :

Potassium :

[ Draw the diagram ]

1) The reaction of alkali metals with water

2) The reaction of alkali metals with O2 gas

[ Draw the diagram ]


68

3. Write the chemical equations for the reactions between the products from

the combustion of each alkali metal with water.

Lithium :

Sodium :

Potassium :

1. A student performs two experiments to study the reaction of alkali metal with

water.

Experiment Metal used Observation

1 Sodium Sodium moves rapidly and randomly on the

water surface and emits ‘hiss’ sounds.

2 Lithium Lithium moves slowly on the water surface.

a) Write the chemical equation for the reaction between sodium and lithium

with water.

Lithium :

Sodium :

b) Between reactions of Experiment 1 and 2, which is more vigorous? Explain

your answer from the point of electron arrangement.

c) Explain why sodium and lithium show similar chemical properties?

d) Calculate the volume of hydrogen gas produced if 2.3 g of sodium is used.

[RAM: Na, 23 ; molar volume: 24 dm3 mol-1]

Exercise


69

GROUP 17

consist of fluorine,

chlorine, bromine,

iodine and

astatine.

GROUP 17

Known as

HALOGENS.

Have valence

electron of 7.

GROUP 17

Elements are

poisonous.

Exist as diatomic

molecules.

Element Symbol of

Atom

Symbol of

Molecule

Physical

state at

room

condition

Colour

Fluorine

Chlorine

Bromine

Iodine

Astatine

Physical State

HALOGENS

Melting & Boiling

Point Electric

Conductivity

Density

Smell


70

Electronegativity is an ability of atom to receive an electron to become

negative charged ions.

All halogen have seven valence electron. Their atoms always gain one electron

to form an ion with a charge of –1, in order to achieve stable octet electron

arrangement.

Chemical Properties of Group 17 Elements

Element Atomic

Size

Melting /

Boiling

point

Density Electronegativity

Fluorine

Chlorine

Bromine

Iodine

Astatine

Halogen react with water to produce two acids

Cl2 + H2O HCl + HOCl

Halogen react with metal to produce metal halides

2Fe (s) + 3Cl2 (g) FeCl3 (s)

Halogen react with sodium hydroxide solution to form sodium

halide and sodium halate (I) and water

I2(s) + 2NaOH (aq) NaI (aq) + NaOI (aq) + H2O (l)


71

To Investigate the Chemical Properties of Halogens

[Draw the Apparatus]

SECTION A : Reaction with Water

1) Chlorine

2) Bromine 3) Iodine

SECTION B : Reaction with Iron

1) Chlorine


SECTION C : Reaction with Sodium Hydroxide

1) Chlorine



72

[Discussion]

1. Name the products formed when chlorine, bromine, and iodine

react with water.

Chlorine :

Bromine :

Iodine :

2. Write the chemical equations for the reaction of chlorine,

bromine and iodine with water.

Chlorine :

Bromine :

Iodine :

3. What is the function of soda lime in SECTION B?


react with iron.

Chlorine :

Bromine :

Iodine :

5. Write the chemical equation for the reaction of chlorine,

bromine, and iodine with iron.

Chlorine :

Bromine :

Iodine :


73


react with sodium hydroxide solution.

Chlorine :

Bromine :

Iodine :

5. Write the chemical equation for the reaction of chlorine,

bromine, and iodine with sodium hydroxide solution.

Chlorine :

Bromine :

Iodine :

6. Describe the changes in reactivity of Group 17 elements when

going down the group. Explain your answer.


74

1. Table below shows several halogen elements with their proton numbers.

Halogen Proton Number

X 9

Y 17

Z 35

a) State the group that the halogen elements belong to in the Periodic Table.

b) i) Write the electron arrangement of elements X and Y.

ii) From your answers in (b)(i), deduce the period of elements X and Y.

c) Draw the electron arrangement of element Y.

d) State the changes of properties of halogens down the group from X to Z in

terms of:

i) Atomic radius

ii) Electronegativity

iii) Melting point and boiling point

Exercise


75

e) i) Iron glow brightly when reacting with element Y. Write the chemical

reaction to represent this reaction.

ii) Predict the observation for the reaction between iron and element Z.

iii) Between Y and Z, which is more reactive? Explain why.

f) i) Determine the elements X, Y and Z.

ii) State the colour for each elements X, Y, and Z.

iii) Why the physical state of halogen differ when down the group?


76

Comparison & Similarities between Group 1 and Group 17

Properties Group Down the Group

Atomic

Size

1

17

Melting &

Boiling

Point

1

17

Density 1

17

Reactivity 1

17


77

Elements in a Period 3

Element Sodium Magnesium Aluminium Silicon Phosphorus Sulphur Chlorine Argon

Symbol

Proton Number

Electron

Arrangement

Metal Properties

Physical State

[RC]

Atomic Radius

Electronegativity

Property of the

oxide

Note: the atomic radius ______________ and electro negativity ______________

due to the increasing nuclei attraction on the valence electrons.

(nuclei attraction force increase with the increase of proton number)

The proton number ____________ by one unit.

All atom of the element have __________ shells occupied by electron.

The number of valence electron ____________ from 1 to 8.

All element exist as __________ except chlorine and argon.

The atomic radius of element ________________.

The electronegativity of element ______________.

The oxides of element in Period 3 change from basic to acidic properties,

therefore the metallic properties decrease across the period. On the other

hand non-metallic properties of the elements increase.


78

Transition Elements

Transition elements are elements from Group 3 to Group 12

All the elements are metals, usually solids with shiny surfaces, ductile,

malleable and have tensile strength.

Have high melting and boiling point as well as high density.

Good conductors of heat and electricity.

Three special characteristics of transition element

Show different oxidation number in their compound.

Form coloured ions or compound.

Act as catalysts.


79

CHAPTER 5 CHEMICAL BONDS

Atom of other

element that have

less than eight

valence electron

are not stable

Almost all chemical

substances exist as

compounds in nature

except inert gases

and other stable

element (such as

gold and silver).

Less stable atom

will tend to release,

accept or share

electron to

achieve the stable

electron

arrangement

of an inert gas.

All other elements

combine together to

achieve the stability by

forming duplet or octet

electron arrangement by

i) The transfer of electron

ii) Sharing of electron

Two types of chemical

bonds formed:-

i) ionic bonds

ii) covalent bond

Covalent Bond

Ionic Bond


80

Ionic Bond

Ionic bond formed

when metal atom

transfer electrons to

non-metal atom to

form ionic

compound.

Formation of

Cation

Formation of

Anion

Metal atom from

group 1,2 and 13

tend to released all

their valence

electrons.

Non-Metal atom

from group 15, 16

and 17 tend to

accept the

electrons.

Example:

Draw the formation

of sodium ion.


81

1. Draw the formation of the following cations:

a) Potassium ion

b) Magnesium ion

c) Aluminium ion

Exercise


82

2. Draw the formation of the following anions:

a) Chloride ion

b) Oxide ion

c) Nitride ion


83

Formation of Ionic Compound

Example: Formation of Sodium Chloride, NaCl

1. Metal atom and non-metal atom are not stables.

2. To achieve stability, metal atom will combine with nonmetal atom by transfer of

electrons.

3. Metal atoms will donate electrons to form cation.

4. Non-metal atom will accept electrons to form anion.

5. Both ions already achieve stable duplet/octet electron arrangement.

6. The formation of cation & anion will create strong electrostatic force between the ions.

7. Ionic bond will pull cation & anion together in crystal lattice form.

8. The compound formed called as ionic compound.

X X+ + e

Y + e Y-


84

1. Explain the formation of ionic compound below:

a) Lithium fluoride

b) Magnesium oxide

Exercise


85

c) Calcium chloride

d) Aluminium oxide


86

2. Draw the formation of the following ionic compound:

a) Lithium fluoride

b) Magnesium oxide


87

c) Calcium chloride

d) Aluminium oxide


88

Ionic Equations

Equation that represent the formation of ionic compounds are

known as ionic equation.

Example:

a) Formation of sodium chloride, NaCl

i) Chemical Equation :

ii) Half-ionic Equation :

b) Formation of Magnesium oxide, MgO

i) Chemical Equation :

ii) Half-ionic Equation :

1. Write an ionic equation of the following compound

a) Lithium fluoride

b) Magnesium chloride

c) Aluminium oxide

Exercise


89

1. Atom X and Y each have proton numbers of 3 and 8. What is

the ionic compound formula formed between atoms X and Y?

2. Complete each of the following table:

Atom Proton

Number

Electron

Arrangement

Ionic

Formula Atom

Proton

Number

Electron

Arrangement

Ionic

Formula

Compound

Formula

A 3 B 9

1 C 11 D 8

2 E 12 F 17

3 G 20 H 9

4 I 19 J 17

5 K 13 L 8

6 M 13 N 17

Exercise


90

Covalent Bond

Covalent bond is the

chemical bond

formed through the

sharing of electron

between two or more

non metal atom to

form covalent

compound.

Three types of covalent

bonds:

single covalent bond

( sharing one pair of e )

double covalent bond

( sharing two pairs of e )

triple covalent bond

( sharing three pairs of e )

Single Covalent Bond

Example:

Draw the formation of chlorine gas.


91

Double Covalent Bond

Example:

Draw the formation of oxygen gas.

Triple Covalent Bond

Example:

Draw the formation of nitrogen gas.


92

1. Draw the formation of the following compound.

a) water

b) Carbon dioxide

Exercise


93

c) Ammonia

b) Tetrachloromethane , CCl4.


94

Determine the Formula of Covalent Compound

Guideline:

1. State the electron configuration of atoms.

- Make sure electron valence for both atoms is either 4, 5, 6, and 7.

2. Determine the number of electrons needed to achieve stability.

3. Write the number of electron needed to achieve stability at the

below right corner of each atom.

4. Cross the number.

Example:

If atom P has 8 protons and atom Q has 9 protons, determine the

formula of the covalent compound formed.


95

1. Atoms K and S each have a proton number of 6 and 8

respectively. What is the formula of the covalent compound

which is formed by K and S?

2. Complete the table below to show the formulae of compounds

which are formed.

Atom Proton

number

Electron

config. Atom

Proton

number

Electron

config.

Compound

formula

A 6 2.4 B 9 2.7 AB4

C 6

D 16

E 7

F 1

G 1

H 8

I 6

J 17

K 6

L 8

Exercise


96

Comparison between the formation of the ionic bond and the

covalent bond

IONIC BOND COVALENT BOND

Similarity

Differences

Formation

Particles

Force of

Attraction


97

The following figure compares and contrasts the properties

of ionic compound and covalent compound

IONIC COMPOUND COVALENT COMPOUND

PROPERTIES

Physical State

Melting &

Boiling point

Solubility

Electric

Conductivity


98

1. (a) Table 1.1 shows the proton number of three elements, X, Y, and

Z. The letters used do not represent the actual symbols of the

elements.

Element Proton Number

X 6

Y 12

Z 17

Table 1.1

i) Write the electron arrangement of:

Atom Y : _______________________________________________

The ion of Z : ___________________________________________

ii) Write the formula of the compound formed between

elements Y and Z.

________________________________________________________

iii) Element X reacts with element Z to form a covalent

compound with a formula XZ4. State two physical

properties of this compound.

________________________________________________________

________________________________________________________

i) Draw the electronic structure of the compound XZ4.

Exercise


99

(b) Table 1.2 shows some physical properties of two compounds, U

and V.

Compound Melting pt

(oC)

Boiling pt

(oC)

Solubility in

water

Solubility in

organic

solvent

U 800 1 420 Soluble Insoluble

V - 95 86 Insoluble Soluble

Table 1.2

i) State the physical state of the following compound at room

condition.

U : _______________________________________________________

V : _______________________________________________________

ii) State the type of compound for U.

__________________________________________________________

iii) Explain why melting point and boiling point of compound U is

higher than V?

___________________________________________________________

___________________________________________________________

___________________________________________________________

___________________________________________________________


100

CHAPTER 6 ELECTROCHEMISTRY

Electrolytes are

substances that can

conduct electricity when

they are in molten state

and aqueous solution.

This due to the present of

free moving ions in the

electrolytes.

Electrolysis is a process

whereby compounds in

molten or aqueous state

are broken down into

their constituent

elements by passing

electricity through them.

Non-electrolytes are

substances that can not

conduct electricity when

they are in all state. This

because non-electrolyte

exist as molecule which

means contain no ions.

Example

2NaCl (l) 2Na (s) + Cl2 (g)

Sodium Metal

Chlorine Gas


101

Electrolytic Cell (molten state)

a) Electrode attach to positive terminal (battery) =

b) Electrode attach to negative terminal (battery) =

c) Anion (negative ion) discharged at electrode =

Anion will _______________ electrons.

d) Cation (positive ion) discharged at electrode =

Cation will ______________ electrons.

e) Electron flow from ___________________ to ___________________

f) Electrolytic Cell will change the _______________ energy to

________________ energy.

[ Draw the apparatus of electrolysis molten sodium chloride ]


102

Product of Electrolysis

Ion

discharged

at Cathode

Observation Half-equation Product

All ion metal

except

Copper ion

Copper ion

Hydrogen

ion

Ion

discharged

at Anode

Observation Half-equation Product

Oxide ion

Chloride ion

Bromide ion

Iodide ion

Hydroxide

ion

Gas Test

Hydrogen

gas

Oxygen gas

Chlorine gas


103

Electrolysis Molten Lead(II) Bromide

1. State the ion consists in the electrolyte.

2. Which electrode is

a) Cathode =

b) Anode =

3. Which ion will be discharged at

a) Cathode =

b) Anode =

4. State the observation at

a) Cathode =

b) Anode =

5. State the product formed at

a) Cathode =

b) Anode =

6. Write the half equation at

a) Cathode =

b) Anode =

7. Draw the electron flow on the diagram above.

P

Q


104

1. State the ion present in the following electrolyte. Predict the

products from the electrolysis of some molten compound and

write the ionic equation involved.

a) Magnesium oxide

b) Copper(II) chloride

c) Lead(II) iodide

Exercise


105

2. State the meaning of the following terms.

a) Anode b) Cathode c) Electrolysis

3. A molten oxide, R2O3 is electrolysed using carbon electrodes.

a) Draw a labeled diagram to show the set-up of apparatus for

the electrolysis.

b) What ions are present in the electrolyte? Write the formulae for

the ions present in the electrolyte.

c) State the ions move to each of the electrodes during

electrolysis.


106

d) Write half equation of the reaction at each of the electrodes.

e) Name the substances formed at each of the electrodes.

f) Label the flow of electron in the diagram (a).


107


108

Electrolysis of Aqueous Solution

Aqueous solution consists of four types of ions. Two ions from the

compound and two ions from the water.

Example:

Molten sodium chloride Sodium chloride solution

Generally, there are 3 factors that may influence the selective of ions

during electrolysis of an aqueous solution

1. Position of ions in the electrochemical series

2. Concentration of ions in the electrolytes

3. Types of electrodes used in the electrolysis


109

Position of ions in the electrochemical series (ECS)

The ions that are lower in the ECS will selected to be discharged.



a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


[ Draw the apparatus of electrolysis sodium chloride solution ]


110

1.

For the electrolysis of copper(II) sulphate solution,

(a) State all the ions that are present in the electrolyte.

(b) State the ions in (a) which discharged to the

i) anode :

ii) cathode :

(c) Write a half equation for the reaction at the

i) anode :

ii) cathode :

(d) The blue colour of copper(II) sulphate solution fades if the

electrolysis is carried for a long period of time. Explain why.

Exercise

Carbon electrode

Copper(II)

sulphate solution


111

2.

For the electrolysis of dilute sulphuric acid,

a) State all the ions that are present in the electrolyte

b) State the ion in (a) which discharged to

i) anode

ii) cathode

c) Write half equation for the reaction at the

i) anode

ii) cathode

d) Explain why the concentration of dilute sulphuric acid increases

gradually during the electrolysis

3. Base on the answer 1(c) and 2(c), name the process that occur

at the

a) anode

b) cathode

Dilute sulphuric

acid

Carbon electrode


112

Concentration of ions in the electrolytes

If the concentrations of particular ions are high, the ion is selectively

discharged



a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


[ Draw the apparatus of electrolysis concentrated sodium chloride solution ]


113

1.

Diagram above show the apparatus set up for the experiments of

electrolysis using two different concentration of hydrochloric acid.


i) Experiment A :

ii) Experiment B :

Exercise

Carbon

electrode

Dilute Hydrochloric

acid solution

Concentrated

Hydrochloric acid

solution

Experiment A

Carbon

electrode

Experiment B


114

b) State the ion in (a) which discharged to anode and cathode in

i) Experiment A :

ii) Experiment B :

c) Write half equation for the reaction at the anode and cathode in

i) Experiment A :

ii) Experiment B :

d) State the observation occur at cathode and anode in

i) Experiment A :

ii) Experiment B :

e) State the product formed at cathode and anode in

i) Experiment A :

ii) Experiment B :


115

Types of electrodes used in the electrolysis

If using the active electrode at anode, ions that are present in the

electrolytes are not discharge. Instead the active electrode will

corrodes and dissolves in the electrolytes.



a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


a) Cathode =

b) Anode =


[ Draw the apparatus of electrolysis silver chloride solution using silver electrodes ]


116

1.

Diagram above show the apparatus set up for the experiments of

electrolysis using two different electrodes immersed in copper(II)

sulphate solution.


i) Experiment A :

ii) Experiment B :

Exercise

Carbon

electrode

Copper(II) sulphate

solution

Experiment A

Copper

plate

Experiment B

Copper(II) sulphate

solution


117

b) State the observation occur at anode and cathode in

i) Experiment A :

ii) Experiment B :

c) Write half equation for the reaction at the anode and cathode in

i) Experiment A :

ii) Experiment B :

d) Explain the observation on the colour of copper(II) sulphate

solution in

i) Experiment A :

ii) Experiment B :


118

Electrolysis in Industry

Most common

application:

i) Extraction of metal

ii) Purification of metal

iii) Electroplating

1) Extraction of Metal

Extraction of aluminium

from aluminium oxide.

2) Purification of Metal

In purification:

The impure metal is made to be the

anode

The cathode is a thin layer of pure metal

3) Electroplating

Electroplating is a process to coat

one metal onto another metal.

The purposes of electroplating

onto metal are:-

i) Make it look more attractive

ii) more resistant to corrosion

In electroplating :

object to be electroplated

as the cathode

anode is the metal used for plating

Electrolyte is a solution of the

compound of the electroplating

metal

Copper nugget


119

1) Ion present in electrolyte =

2) Ion discharged

a) Cathode =

b) Anode =

3) Observation

a) Cathode =

b) Anode =

4) Half equation

a) Cathode =

b) Anode =

5)Function of cryolite, Na3AlF6

Extraction of Metal

[ Draw the apparatus of electrolysis for extraction of aluminium from aluminium oxide ]


120


2) Ion discharged

a) Cathode =

b) Anode =

3) Observation

a) Cathode =

b) Anode =

4) Half equation

a) Cathode =

b) Anode =

5) Colour changes of electrolyte

Purification of Metal

[ Draw the apparatus of electrolysis for purification of impure copper ]


121


2) Ion discharged

a) Cathode =

b) Anode =

3) Observation

a) Cathode =

b) Anode =

4) Half equation

a) Cathode =

b) Anode =

5) Colour changes of electrolyte

Electroplating

[ Draw the apparatus of electrolysis to electroplate key by using copper as electrode ]


122

Voltaic Cells

A simple voltaic cell can be made by dipping two different

types of metals in an electrolyte

Electron flow from one metal to another metal through the

connecting wire in the external circuit.

More electropositive metal will release electron, thus act as

the negative terminal. Less electropositive metal will accept

electron and act as the positive terminal.

Continuous flow of electron produces an electric current.

Simple zinc-copper

Voltaic

Zinc more reactive than

copper

Zinc will act as terminal

________________, and

copper will act as

terminal ______________.

Zinc will release electron

to form Zn 2+.

Half equation:

Cu 2+ ions from copper(II) sulphate solution receive

electron to form copper metal.

Half equation :

Overall equation:

The further the distance between the position of two

metals is in ECS the bigger the cell voltage.


123

Different Types of Voltaic Cells

Two types of voltaic cell:

1) Primary cells: non rechargeable cell

Example: Daniell cell, dry cells, alkaline cell

2) Secondary cells: rechargeable cells

Example: Lead-acid accumulator, Nickel-cadmium

Daniell Cell 1

1. Used salt bridge

Salt bridge contain inert

ions or salt that does not

react with electrolyte.

Example:

Daniell Cell 2

2. Used porous pot

Porous pot has fine pores

that allow ions flow through.

What is the function of salt

bridge and porous pot in

Daniell Cell?


124

1. For the simple voltaic cell that you see at the diagram

a) State how electricity was produced.

b) What are the chemical changes that occur at the

magnesium ribbon and the copper plate.

c) Write the half equation for the changes that occur at each

the electrode.

d) What is the direction of electron flow from terminal to

another through the external circuit.

Exercise

Mg

Cu

Magnesium sulphate

solution


125

Electrochemical series (ECS)

The electrochemical series (ECS) can be constructed by two

method:

a) The potential difference (voltage difference) between pairs of

metal.

b) The ability of metal to displace another metal from its salt solution.

Tendency of

metal to release

electrons to

form ions

increases

Tendency of cation

to receive

electrons to form

metals

increases


126

The potential difference (voltage difference) between pairs of metal.

The bigger the voltage value the further apart their position.

The metal act as negative terminal is placed at higher position

in electrochemical series (ECS).

How to determine the positive/negative terminal?

Example: The voltaic cells are constructed as shown in the figure. The

voltmeter reading of the cell I is 1.1 V while that of cell II is

2.5 V.

Arrange the metals in descending order in the

electrochemical series.

Cell 1

Cell 2

P

P

Q

R


127

The ability of metal to displace another metal from its salt solution

If the M can displace metal N from an aqueous N salt solution, then:

i) Metal M is more electropositive than metal N

ii) Metal M is placed at a higher position than metal N in the ECS

Example: Zinc and copper(II) sulphate sulphate solution

Observation :

Half-equation :

The Important of ECS

ECS can be used to determine:

The terminal of voltaic cell

The standard cell voltage

The ability of a metal to displace another metal from its

salt solution.


128

1. The diagram shows an electrolytic cell. The left section of the cell

(S) is a source of electricity to drive the right section (T) of the

cell.

a) State the change of energy in cell S

b) i) For cell S, state the positive terminal of the cell

ii) Explain your choice for b(i)

c) i) State what has happened at the negative terminal

ii) What process has happened in this electrode

Exercise

Copper Zinc Aluminium

Copper(II) Sulphate Zinc Sulphate

Aluminium

Sulphate

S T


129

d) Explain why the color of copper(II) sulphate remain unchanged

e) Determine the anode of cell T

f) Write down the half equation for the cathode in cell T

g) What will happened if the aluminium in cell S is replaced by copper


130

CHAPTER 7 ACIDS AND BASES

The role of water

In the presence of water an acid will ionise to form hydrogen ion and

alkaline will dissociate into hydroxide ions, OH-.

Therefore, water is essential for the formation of hydrogen ions, H+

that cause acidity and hydroxide ions, OH- that cause alkalinity.

Arrhenius Theory

An acid is a chemical

compound that

produces hydrogen ions,

H+ or hydroxonium ions

H3O+ when dissolve in

water.

A base defined as a

chemical substance

that can neutralise an

acid to produce a

salt and water.

An alkali is defined as

a chemical

compound that

dissolve in water to

produce hydroxide

ions, OH- .

Example

HCl (g) H2O H+ (aq) + Cl- (aq)

NaOH(s) H2O Na+(aq) + OH-(aq)


131

BASICITY OF AN ACID

Is the number of ionisable hydrogen atoms per molecule of an

acid.

Example: a) CH3COOH? (ethanoic acid) =

b) H2SO4? (sulphuric acid ) =

c) HNO3? (Nitric acid) =

d) H3PO4? (Phosphoric acid) =

STRENGTH OF ACID & ALKALI

Monoprotic Acid

Acid which produces

1 hydrogen ion when

one molecule of an

acid ionises in H2O.

Diprotic Acid

Triprotic Acid


132

STRONG ACID

Strong acid will dissociate or ionize

completely in water to produce

hydrogen, H+ ions.

Degree of dissociation is higher.

Thus, higher concentration of

hydrogen ions in aqueous acid

solution.

Therefore, low pH value of the acid

solution.

WEAK ACID

STRONG ALKALI

WEAK ALKALI


133

PHYSICAL PROPERTIES OF ACID & ALKALI

ACID

ALKALI pH Value Litmus paper

Taste Corrosive

pH Value Litmus paper

Corrosive Taste

Conduct

electricity

Conduct

electricity


134

CHEMICAL PROPERTIES OF ACID & ALKALI

Acids can react with,

i) bases to produce salts and water

eg :

ii) metal to produce salts and hydrogen gas

eg :

iii) metal carbonates to produce salts, carbon dioxide

and water

eg :

iv) alkali to produce salts and water (neutralization)

eg :

Alkali can react with,

i) acid to produce salts and water (neutralization)

eg :

ii) ammonium salt to produce salts, water and ammonia

gas

eg :


135

The concentration of acid and alkali

Concentration are measurement of the quantity of solutes dissolved

in a quantity of solvent.

1. The molarity of a bottle of nitric acid, HNO3 solution is 2.0 mol dm-3.

What is the concentration of the solution in g dm-3?

[RAM: H, 1 ; N, 14 ; O , 16]

Grams per dm3

( g dm-3)

Moles per dm3

( mol dm-3)

*known as Molarity Concentration

Grams per

dm3 ( g dm-3)

Moles per dm3

( mol dm-3)

Unit Conversion

× molar mass

÷ molar mass

Example


136

2. Calculate the molarity of a sodium sulphate, Na2SO4 solution with

a concentration of 28.4 g dm-3. [RAM: O, 16 ; Na, 23 ; S, 32]

Calculating Involving Concentration and Molarity

1. 5.00 g of copper (II) sulphate is dissolved in water to form 500

cm3 solution. Calculate the concentration of copper (II) sulphate

in g dm-3.

2. A 250 cm3 nitric acid solution contains 0.4 moles. Calculate the

molarity of the nitric acid.

No of mole = Molarity x Volume ( cm3)

1000

n = MV

1000


137

3. What is the mass of sodium carbonate required to dissolve in

water to prepare 200 cm3 solution contains 50 g dm-3.

4. Calculate the number of moles of ammonia in 150 cm3 of 2 mol

dm-3 aqueous ammonia.

5. Calculate the volume in dm3 of a 0.8 mol dm-3 sulphuric acid

that contains 0.2 mol.

6. 4.0 g sodium carbonate powder, Na2CO3 is dissolved in water

and made up to 250 cm3. What is the molarity of the sodium

carbonate solution. [RAM: C,12;O,16;Na,23]


138

7. Dilute hydrochloric acid used in the school laboratories usually

has a concentration of 2.0mol dm-3. Calculate the mass of

hydrogen chloride that found in 250 cm3 of the hydrochloric

acid? [RAM : H,1; Cl,35.5]

8. The concentration of a potassium hydroxide solution is

84.0 g dm-3. Calculate the number of moles of potassium

hydroxide present in 300 cm3 of the solution. [RAM: K,39,H,1 O,16]


139

9. Calculate the number of moles of hydrogen ions present in

200 cm3 of 0.5 mol dm-3 sulphuric acid.

Preparation of Standard Solution

A solution in which its concentration is accurately known is a

standard solution.

Preparation of a solution by dilution method

Adding water to a concentrated solution changes the concentration

of the solution but does not change the amount solutes ( number of

moles)of solution present in solution.

Moles of stock solution = moles of dilute solution

n1 = n2

M1V1 = M2V2

Add water Stock solution

Dilute solution


140

1. Find the volume of 2.0 mol dm-3 sulphuric acid, H2SO4 needed to

prepare 100 cm3 of 1.0 mol dm-3 sulphuric acid, H2SO4.

2. Calculate the volume of a concentrated solution needed to

prepare each of the following dilute solution:

a) 50 cm3 of 0.1 mol dm-3 sodium hydroxide, NaOH solution

from 2.0 mol dm-3 sodium hydroxide, NaOH solution.

b) 100 cm3 of 0.5 mol dm-3 potassium manganate(VII),

KMnO4 solution from 1.0 mol dm-3 potassium

manganate(VII), KMnO4 solution.

Exercise


141

3. Calculate the volume of 2.0 mol dm-3 sulphuric acid, H2SO4

needed to prepare 2.5 dm3 of 0.5 mol dm-3 of the same acid

solution.

4. Calculate the molarity of potassium hydroxide, KOH, solution if

200 cm3 of 2.0 mol dm-3 potassium hydroxide, KOH, solution is

added to 200 cm3 of water.

5. 60 cm3 of 0.5 mol dm-3 sodium hydroxide, NaOH, solution is

diluted with 30 cm3 of water. Calculate the molarity of the

solution produced.


142

The pH Values and Molarity

As the molarity of an acid increases, the pH value of the acid

decreases, however the pH value of an alkali increases when the

molarity of the alkali increases.

Neutralisation

Reaction of an acid

and a base that

produce salt and

water.

Acid-base titration

Titration is a very useful

laboratory technique in

which one solution is

used to analyse another

solution.

An acid of known

concentration is

carefully delivered from

burette to completely

neutralise a known

volume of an alkali in a

conical flask.

Acid-base indicator is

used to detect the end

of titration.

Eg: methyl orange,

phenolphthalein and

litmus.

The point at which the

colour of the solution

change is called the

end point.

Titration using

phenolphthalein


143

ACID-BASE

TITRATION

NEUTRALISATION


144

1. Write a balanced equation for the neutralization of each of the

following:

a) Sulphuric acid, H2SO4 and barium hydroxide, Ba(OH)2

solution.

b) Nitric acid, HNO3 and calcium hydroxide, Ca(OH)2 solution.

c) Ethanoic acid, CH3COOH and potassium hydroxide, KOH

solution.

Acid-base Indicator

Indicator Colour in

alkalis

Colour in neutral

solution

Colour in

acids

Methyl orange

phenolphthalein

litmus

The end-point of neutralisation also can determined by another two

method:-

i) Measurement of pH values by computer

ii) Measurement of electrical conductivity during titration

Example


145

Numerical Problem involving Neutralisation

Say the balance equation is

aA + bB product

which,

A = acid

a = no of mole of acid

B = base

b = no of mole of base

MAVA = a and MBVB = b

Therefore MAVA a

MBVB b

1. In an experiment, 25.0 cm3 of a sodium hydroxide solution of

unknown concentration required 26.50 cm3 of 1.0 mol dm-3

sulphuric acid to complete a reaction in titration. Calculate the

molarity of sodium hydroxide.

=

Exercise


146

2. What is the volume of 0.5 mol dm-3 sulphuric acid, H2SO4 needed

to neutralize 25.0 cm3 of 0.8 mol dm-3 ammonia, NH3 solution?

3. A sample of copper(II) oxide, CuO was found to completely

neutralize 100 cm3 of 0.5 mol dm-3 hydrochloric acid, HCl.

Calculate the mass of the sample. [RAM: O, 16 ; Cu, 64]


147

4. The volume of 0.15 mol dm-3 sulphuric acid, H2SO4 required to

completely neutralize 25.0 cm3 of potassium hydroxide, KOH

solution is 30.5 cm3. Calculate the molarity of the potassium

hydroxide, KOH solution.

5. A student dissolved 3.65 g of hydrogen chloride gas, HCl in water

to make 1.0 dm3 of solution. Calculate the volume of a 0.1 mol

dm-3 barium hydroxide, Ba(OH)2 solution required to completely

neutralize 25.0 cm3 of the acid solution. [RAM : H, 1 ; Cl, 35.5]


148

CHAPTER 8 SALTS

A salt is an ionic

substance produced

when the hydrogen

ion of the acid is

replaced by metal

ion or an ammonium

ion.

The salt consists of two

parts, cation from base

and anion from acid.

NaCl

NaOH

(Base)

HCl

(Acid)


149

Table of Salts

Complete the table below.

Metal ion Sulphate salt

(SO42-)

Chloride salt

(Cl-)

Nitrate salt

(NO3-)

Carbonate salt

(CO32-)

K+ K2SO4 KCl KNO3 K2CO3

Na+

Ca2+

Mg2+

Al3+

Zn2+

Fe2+

Sn2+

Pb2+

Cu2+

Ag+

NH4+

Ba2+

Based on the table above, mark the insoluble salt.


150

SOLUBLE & INSOLUBLE SALTS

SALT SOLUBLE INSOLUBLE

K+ , Na+ , NH4+ All soluble none

Nitrate salts

(NO3- )

All nitrate salts none

Chloride salts

(Cl-) All chloride salts

Lead (II) chloride,

PbCl2

Silver chloride, AgCl

Mercury chloride,

HgCl

Sulphate salts

(SO42-)

All sulphate salts

Lead (II) sulphate

Calcium sulphate

Barium sulphate

Carbonate

salts (CO32-)

Sodium carbonate,

Na2CO3

Potassium carbonate,

K2CO3

Ammonium carbonate,

(NH4)2CO3

All others carbonate

salts

Oxide salts

(O2-)

Sodium oxide, Na2O

Potassium oxide, K2O

Calcium oxide, CaO (slightly soluble)

All oxide salts

Hydroxide salts

(OH-)

Sodium hydroxide, NaOH

Potassium hydroxide, KOH

Calcium hydroxide,

Ca(OH)2

(slightly soluble)

All hydroxide salts


151

Preparation and Purification of Soluble Salts

Soluble salt can be prepared by the following ways:

1. Reaction between acid and alkali - preparation for

sodium, potassium and ammonium salts only.

Eg: HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Burette

HCl solution

NaOH solution

[write the step of preparation of the salts on the diagram]

NaCl solution

Evaporating dish

NaCl crystal salt

NaCl salt crystal


152

Preparation and Purification of Soluble Salts

Soluble salt can be prepared by the following ways:

1. Reaction between acid and metal oxide

Eg: HNO3(aq) + MgO(s) Mg(NO3)2 (aq) + H2O(l)

2. Reaction between acid and metal

Eg: H2SO4(aq) + Zn (s) ZnSO4(aq) + H2 (g)

3. Reaction between acid and metal carbonate

Eg: HCl(aq) + CaCO3(s) CaCl2(aq) + CO2(g) + H2O(l)

[write the step of preparation of the salts on the diagram]

heat heat

Copper(II) nitrate

Solution + excess

CuO

Copper(II) nitrate

solution

Evaporating Dish

Copper(II) nitrate solution

& crystal salt

Cu(NO3)2 crystal salt


153

Crystallization

is a process to crystallize

the soluble salts.

Recrystallization

process will carried out in

order to get pure and

more defined crystal

Physical characteristic of

crystals

Fixed geometrical shapes

such as a cuboids,

rhombic or prism

Flat surface, straight edges

and sharp angles.

Same shapes for same

substance but differ in

sizes

Fixes angle between two

neighbouring surfaces.


154

Preparation of insoluble salts

An insoluble salt is prepared through precipitation method.

Aqueous solution of two soluble salts are mixed to form

insoluble and soluble salt:

The reaction is called double decomposition.

Two solutions contain ions that make up the insoluble salts.

Eg: Preparation of lead(II) iodide salt by using lead(II)

nitrate solution and potassium iodide solution. [write the balance chemical equation]

Double Decomposition

[write the general equation of preparation insoluble salt]


155

Chemical and Ionic Equation

Chemical and ionic equation can be written for all reaction

That used to prepare salts.

Example: Formation of precipitate Barium Sulphate, BaSO4.

1. Chemical Equation:

BaCl2 (aq) + Na2SO4 (aq) BaSO4(s) + 2NaCl (aq)

2. Ionic Equation:

Ionic equation shows the ions take part in the reaction.

1. Change each of the following word equations to a balanced

chemical equation.

a) Sulphuric acid + zinc zinc sulphate + hydrogen gas.

b) Silver nitrate + potassium iodide silver iodide + potassium

nitrate

c) Nitric acid + chromium(III) hydroxide chromium(III) nitrate

+ water

Exercise


156

2. Complete the following equations.

a) HCl (aq) + NiO (s)

b) HNO3 (aq) + Ca(OH)2 (aq)

c) H2SO4 (aq) + MgCO3 (s)

3. Write an ionic equation for each reaction between the following

pairs of substances.

a) Sulphuric acid, H2SO4 and barium hydroxide, Ba(OH)2 solution

b) Ammonium chloride, NH4Cl solution and silver nitrate, AgNO3

solution.

c) Lead(II) nitrate, Pb(NO3)2 solution and copper(II) sulphate,

CuSO4 solution.

d) Iron(III) oxide, Fe2O3 and hydrochloric acid, HCl.


157

Constructing Ionic Equation using the Continuous Variation Method

Continuous variation method can be used to construct ionic

equation for the formation of insoluble salts.

Fixed volume of a reactant A is react with varying volumes

of a reactant B to determine the mole ratio of reactant A

that react completely with reactant B.

If x mol of reactant A with y mole of reactant B, than the

empirical formula for insoluble salt is A x B y.

Example:

2Fe3+ (aq) + 3CO32- (aq) Fe2(CO3)3 (s)

1 2 3 4 5 6 7 8


158


159

1. 6.0 cm3 of 0.2 mol dm-3 Xn+ solution reacts completely with 4.0

cm3 of 0.1 mol dm-3 Ym- solution to form a salt XmYn. Write the

ionic equation and hence determine the empirical formula of

the salt reaction.

2. 18.0 cm3 of 0.1 mol dm-3 solution of Px+ ions reacts completely

with 9.0 cm3 of 0.1 mol dm-3 solution of Qy- ions to form a salt

PyQx. Write the ionic equation and hence determine the

empirical formula of the salt in this reaction.

Example


160

Solving Problem Involving Calculation of Quantities of Reactants or

Product in Stoichiometric Reactions

Since the quantities of chemicals involved in a reaction are in term

of moles, the quantities of chemicals (volume, mass and number of

particles) must be converted to moles in calculation regarding

quantities of reactant and products.

1. Calculate the number of moles of aluminium sulphate

produced by the reaction of 0.2 mole of sulphuric acid with

excess aluminium oxide. [0.067 mol]

2. 2.0 g of sodium hydroxide reacts with excess sulphuric acid.

What is the mass of sodium sulphate produces

[RAM: H,1 ; O,16 ; Na,23 ; S,32] [ 3.55 g]

Exercise


161

3. What the volume of carbon dioxide gas evolved at s.t.p when

2.1 g of magnesium carbonate reacts with excess nitric acid.

[ RAM: C,12;O,16;Mg,24; s.t.p = 22.4 dm3] [ 560 cm3]

4. What is the mass of magnesium required to react with 20 cm3 of

2.0 mol dm3 hydrochloric acid to produce 120 cm3 of hydrogen

at temperature? [RAM: Mg,24 ; 1 mol = 24 dm3 at room temp.]

[ 0.12 g]


162

Qualitative Analysis of Salts

General procedure of qualitative analysis

Colour & Solubility of the Salt

GREEN PALE GREEN

Iron(II) Sulphate, FeSO4

Iron(II) Nitrate, Fe(NO3)2

Iron(II) Chloride, FeCl2

BROWN YELLOW/BROWN [depend on concentration]

Iron(III) Sulphate, Fe2(SO4)3

Iron(III) Nitrate, Fe(NO3)3

Iron(III) Chloride, FeCl3

REDDISH-BROWN INSOLUBLE

Iron(III) Oxide, Fe2O3

GREEN INSOLUBLE

Copper(II) Carbonate, CuCO3

BLUE BLUE

Copper(II) Sulphate, CuSO4

Copper(II) Nitrate, Cu(NO3)2

Copper(II) Chloride, CuCl2

BLACK INSOLUBLE

Copper(II) Oxide, CuO

Observations on the physical

properties of salts Action of heat on salts

Tests for cations and

anions

Confirmatory tests for

cations and anions


163

Colour & Solubility of the Salt

WHITE COLOURLESS

Potassium Oxide, K2O

Sodium Oxide, Na2O

Calcium Oxide, CaO

WHITE INSOLUBLE

Magnesium Oxide, MgO

Aluminium Oxide, Al2O3

INSOLUBLE

Zinc Oxide, ZnO

BROWN [hot]

INSOLUBLE

Lead(II) Oxide, PbO

WHITE [cold]

YELLOW [hot]

YELLOW [cold]


164

Gas test

HCl acid

Carbon dioxide Sulphur dioxide

HCl acid

Sodium Carbonate,

Na2CO3

Sodium Sulphite, Na2SO3

Oxygen gas hydrogen gas

Carbon dioxide gas sulphur dioxide gas

chlorine gas ammonia gas

White fumes

Glass rod dipped into

concentrated HCl acid

Ammonia gas, NH3 Chlorine gas, Cl2

Red litmus paper

Wooden splinter

Oxygen gas, O2 Hydrogen gas, H2


165

EFFECT OF HEAT ON SALTS

Sodium carbonate & potassium carbonate are very stable. They do

not decompose on heating.

Carbonate Salt Metal Oxide + Carbon dioxide

Salts Chemical Equation

1. Calcium carbonate

2. Magnesium carbonate

3. Aluminium carbonate


1. Zinc carbonate

Carbonate salts

Heating

Heating

Carbonate salts [ white ]

residue [ white ]

Heating


residue [ yellow – hot ] [ white – cold ]


166

Carbonate Salt Metal Oxide + Carbon dioxide


1. Iron(III) carbonate


1. Copper(II) carbonate


1. Lead(II) carbonate

Heating

Carbonate salts [ brown ]

residue [ brown ]

Heating

Carbonate salts [ green ]

residue [ black ]

Heating


residue [ brown – hot ]

[ yellow – cold ]

Heating


167

Carbonate Salt Metal + Carbon dioxide + Oxygen gas


1. Mercury(II) carbonate


1. Silver carbonate


1. Aurum(II) carbonate

Heating


residue [ grey ]

Heating


residue [ shiny grey ]


residue [ golden yellow ]

Heating

Heating


168

Nitrate Salt Metal Oxide + Oxygen gas + Nitrogen dioxide


1. Calcium nitrate

2. Magnesium nitrate

3. Aluminium nitrate


1. Zinc nitrate

nitrate salts

Heating

Heating

nitrate salts [ white ]

residue [ white ]

Heating


residue [ yellow – hot ] [ white – cold ]


169

Nitrate Salt Metal Oxide + Oxygen gas + Nitrogen dioxide


1. Iron(III) nitrate


1. Copper(II) nitrate


1. Lead(II) nitrate

Heating

nitrate salts [ brown ]

residue [ brown ]

Heating

nitrate salts [ blue ]

residue [ black ]


residue [ brown – hot ]

[ yellow – cold ]

Heating

Heating


170

Nitrate Salt Metal + Nitrogen dioxide + Oxygen gas


1. Mercury(II) nitrate


1. Silver nitrate

Nitrate Salt Metal nitrite + Oxygen gas


1. Potassium nitrate

2. Sodium nitrate

Heating


residue [ grey ]

Heating


residue [ shiny grey ]

Heating

Heating

Heating


residue [ white ]


171

TEST FOR ANIONS

Salt needed to be

dissolved into water

first to produce

aqueous salt

solution.

Unknown aqueous salt solution

[ state the procedure ] [ state the procedure ] [ state the procedure ] [ state the procedure ]

CONCLUSION CONCLUSION CONCLUSION CONCLUSION

OBSERVATION OBSERVATION OBSERVATION OBSERVATION [ label the diagram ] [ label the diagram ] [ label the diagram ] [ label the diagram ]

[ state the anion ] [ state the anion ] [ state the anion ] [ state the anion ]

Ionic Equation Ionic Equation Ionic Equation Ionic Equation


172

TEST FOR CATIONS

Salt needed to be


(soluble salts) or in

dilute acid then

filtered (insoluble

salts) first to produce

aqueous salt solution.


No precipitate White precipitate Coloured precipitate

Green

Brown

Blue

Add NaOH drop by drop

Add NaOH drop by drop until excess

Dissolve in excess NaOH Insoluble in excess NaOH

* All coloured ions insoluble in excess NaOH


173

TEST FOR CATIONS

Salt needed to be


(soluble salts) or in

dilute acid then

filtered (insoluble

salts) first to produce

aqueous salt solution.


No precipitate White precipitate Coloured precipitate

Green

Brown

Blue

Add NH3 drop by drop

Add NH3 drop by drop until excess

Dissolve in excess NH3 Insoluble in excess NH3

Dissolved in excess NH3

Dark blue solution

* Fe2+ & Fe3+ ions insoluble in excess NH3


174

Confirmatory Test for Cation

Cation Procedure Observation

Fe 2+

Fe 3+

Pb 2+

NH4 +

1. Identify the aqueous solutions based on the test and observation given.

Type Observation Answer

a. Potassium thiocyanate,

KSCN, solution is added. Blood red solution formed.

b. Potassium iodide, KI, solution

is added.

Yellow precipitate is

formed.

c. Ammonia solution, NH3, is

added until excess.

Blue precipitate dissolve to

form dark blue solution.

d. A little hydrochloric acid is

added.

Effervescene occur and

lime water turn into chalky.

e. A sulphuric acid, iron(II)

sulphate solution and

concentrated sulphuric acid

is added.

Brown ring formed.

f. Ammonia solution is added

until excess.

White precipitate

dissolved.

Exercise

ZnCl2 NaNO3 PbCl2 CuSO4 FeCl3 K2CO3


175

CHAPTER 9 MANUFACTURED SUBSTANCES IN INDUSTRY

SULPHURIC ACID

The Uses of

Sulphuric Acid


176

POLLUTION of

Sulphur DIOXIDE

Formation of Acid Rain

Burning of Sulphur

Burning of Sulphur Dioxide

[ Cause by ]

[ Cause by ]

[ Effect ]

[ Explain the formation of acid rain & write the chemical equation ]

[ Explain the formation of acid rain & write the chemical equation ]


177

Effect of Acid Rain


178

Industrial Process in the Manufacture of Sulphuric Acid

Contact Process

Burn in

the air Stage 1

Stage 2

Stage 3 ( i )

Stage 3 ( ii )

Dissolves in

concentrated H2SO4

Dilute with water

Catalyst:

Temp. :

Pressure:

[ write the chemical equation ]





179

AMMONIA

The Uses of

ammonia

[ Label the uses of ammonia ]


180

Properties of

ammonia

White fumes

Glass rod dipped into

concentrated HCl acid

Ammonia gas, NH3


181

Industrial Process in the Manufacture of Ammonia

Haber Process

1 part 3 parts

Catalyst chamber

Catalyst :

Temperature :

Pressure :

The unreacted

mixture will flowed

back

[ Write the balance chemical equation ]

Condenser

[ state what happen in the condenser ]

[ state the final product ]


182

Preparation of Ammonium Fertilisers in the laboratory

Ammonia fertilizers can be prepared by using neutralization reaction between acid and

ammonia solution.

Eg: Preparation of ammonium sulphate salt

Burette

H2SO4 acid solution

Ammonia solution

Acid + Ammonia Ammonium salts H2SO4 + NH3 (NH4)2SO4

(NH4)2SO4 solution

Evaporating dish

(NH4)2SO4 crystal salt

(NH4)2SO4 crystal salt


183

Alloys

Physical

Properties

Of pure metal

Pure metal is soft and not strong.

Atom of pure metal have similar shape and size.

The particles in pure metal are arranged layer by layer and

easily slide along between each other.

Ductile

malleable

[ state the definition of the ductile and draw the particle arrangement ]

[ state the definition of the malleable and draw the particle arrangement ]


184

Aim of making

alloys

alloys

Alloy is a mixture of two or more elements

which is the major component is pure

metals.

Foreign elements either metal or

non-metal is added into pure metal.

The size of foreign elements either

smaller or bigger.

It will disrupt the orderly arrangement of

pure metal. Thus, the properties of pure metals improved.

[ draw the particle arrangement of alloy ]


185

Uses of Alloys & Their Compositions

STEEL

COPPER-NICKEL

BRASS

STAINLESS STEEL

BRONZE

DURALUMIN

PEWTER

[ state the components, properties & uses of alloy ]

Pure metal : 99% Iron Foreign element : 1% Carbon

- Strong - Hard - Withstands corrosion - To make bridge, vehicles, building & train tracks


186

Natural

polymers

Synthetic Polymers

Polymers are long chains of molecules made from combination of monomers by

polymerisation process.

Two types of polymer:

i) Natural polymer

ii) Synthetic polymer

Polymer that occurs naturally made by living organisms.

RUBBER

Polymer

Monomer

CELLULOSE

Polymer

Monomer

STARCH

Polymer

Monomer

PROTEIN

Polymer

Monomer

FAT

Polymer

Monomer

Amino acid

Isoprene


187

Synthetic

polymers

Polymer that are man-made polymer produced from

chemical compound through polymerisation.

Two types of polymerisation:

i) Addition polymerisation

ii) Condensation polymerisation

POLYTHENE

Polymer

Monomer

POLYPROPENE

Polymer

Monomer

PERSPEX

Polymer

Monomer

POLYVINYL CHLORIDE (PVC)

Polymer

Monomer

NYLON

Polymer

Monomer


188

Glass & ceramics

glass

Main component of glass is silica & silicon dioxide, SiO2.

Main

characteristics


189

Type of Glass

- Transparent

- Low melting point

- Easily shaped

- Easily broken

- Cannot withstand heat & chemical

reaction

SODA-LIME GLASS

Method of Productions

Compositions

Properties

Uses

- Very high melting point

- Not easy to change its shape

- Does not easily expand or shrink

with changes of temperature

- Transparent to ultraviolet ray

FUSED SILICA GLASS


Compositions

Properties

Uses


190

Type of Glass

- Very transparent

- Shiny

- High density

LEAD CRYSTAL GLASS


Compositions

Properties

Uses

- Withstand heat & chemical reaction

- High melting point

- Transparent to light & infrared ray

but not to ultraviolet ray

- Expand & shrink a little when

temperature changes

BOROSILICATE GLASS


Compositions

Properties

Uses


191

ceramics

Made from clay at very high temperature and the main

component is silicate.

Common

properties


192

Uses of

ceramics


193

USES OF COMPOSITE MATERIALS

Composite Materials

Produced from the combination of two or more different

compound such as alloys, metals, glass, ceramic & polymers.

SUPERCONDUCTOR

Compositions

Properties

Uses

REINFORCED CONCRETE

Compositions

Properties

Uses


194


FIBRE OPTICS

Compositions

Properties

Uses

FIBRE GLASS

Compositions

Properties

Uses

“Success is not the key to happiness. Happiness is the key to success.

If you love what you are doing, you will be successful.”

- Albert Schweitzer -

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195


CERAMIC GLASS

Compositions

Properties

Uses

PHOTOCHROMIC GLASS

Compositions

Properties

Uses

“ There are no secrets to success. It is the result of preparation, hard work, and learning from failure.”

- Colin Powell -

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