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Modul Chemistry Form 4 [email protected]

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CHAPTER 2: THE STRUCTURE OF ATOM

A MATTER

Activity 1

Fill in the blanks with suitable word(s) in the box given

1. Matter is made up of ………………………. and …………………… particles.

2. The tiny particles may be atoms ……………….. and ……………………….

3. An atom is the ………………… particle of an element that can …………………….. in a

chemical reaction.

4. A molecule is a group of two or more …………………… which are ……………………

bonded together.

5. An ion is a …………………………………. or negatively – charged particle

6. Diffusion occurs when particles of a substance move ……………… between the particles

of another substance.

7. Diffusion of matter occurs most rapidly in ………………… state, slower in

………………….. state and slowest in …………………….. state. This is due to the

different ……………………… and ………………………. of particles in the three states of

matter.

8. Matter consists of small particles that always collide among each other. The particles

move faster when energy is …………………… and the particles move slower when the

energy is ….………………..….

Learning Outcomes You should be able to:

describe the particulate nature of matter,

state the kinetic theory of matter,

define atoms, molecules and ions,

relate the change in the state of matter to the change in heat,

relate the change in heat to the change in kinetic energy of particles,

explain the inter-conversion of the states of matter in terms of kinetic theory of matter.

Ions tiny molecules discrete smallest take part positively-charged in released atoms gas chemically solid absorbed different arrangement movement liquid


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Activity 2

1. Complete the table below.

State of matter Solid Liquid Gas

Draw the

arrangement of

particles

Arrangement of

particles

The particles are

packed ……………..

together in an

…………………………

manner

The particles are

packed ……………….

together but not in

….……..……………….

………………………….

The particles are

…………….. apart from

each other and in

……………………………

motion.

Movement of

particles

Particles can only

…………………… and

………………….. about

their fixed positions

Particles can

………………………...,

….……… and…………

throughout the liquid.

Particles can

………………………….,

……………………… and

……………. freely

Attractive forces

between the

particles

Particles are attracted

by very ……………..

…………….. between

the partcles

Particles are held

together by strong

forces but ……………..

than the forces in solid

(moderately)

The attraction forces

between particles are

……………….… forces

Energy content of

particles

……………………..

……………………….

……………………………

2. Underline the correct word in the passage below.

When heat energy is supplied to particles in matter, its kinetic energy (increases

/decreases) and

the particles in matter vibrate ( faster/ slower). When matter loses heat energy, the

kinetic energy

of the particles (increases/decreases) and they vibrate ( faster/ slower).

3. State the change of matter for each conversion in the spaces provided.


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A …………………………… B …………….……………….. C……………..……………..

D……………………………. E …………………………….. F ……………………………

4. Complete the passage below by using the words given below. (solid, gas, boiling point, melting point, solid, gas, liquid, intermolecular, released, absorbed, overcome )

The temperature at which a ………………………………… completely changes to a liquid

is called ……………………………… Boiling point is the temperature at which

a…………………… changes into …………….. . During the boiling process, the

temperature remains constant because the heat energy is …………………… by the

particles and is used to …………………...………………… the ……………………………

forces between particles.

5. The graph below shows the change in temperature with time when a matter in solid state was heated.

Temperature/ OC

A

B

C

D

Time/s


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Based on the graph above, complete the table below.

Point States of matter Explanation in terms of energy change and movement of particles

A to B

Heating causes the particles to ……………………….. more

energy and vibrate ……………………….. The temperature of

the substance and the kinetic energy

…………………………………

B to C

Continuous heating does not cause the temperature of the

substance to increase. The energy absorbed is used to

…………………………… the forces of attraction between the

………………………. The constant temperature is called the

…………………..…………………………………………..

C to D

Continuous heating causes the temperature of the liquid to

…………………………… The particles move……………………

because their kinetic energy ………………………………………..

6. (a) The graph below shows the change in temperature with time when a matter in liquid

state is left to cool.

Based on the graph above, complete the table below.

Point State of matter Explanation

P to Q

As cooling continues, particles lose

their…………………………

and move ………………………. . The

temperature……………...

Q to R

The stronger bonds ……………………. during freezing

release energy. This energy released is the same as the

Temperature / OC

Time/s

P

Q

Q

A

R

B

A

A

SD

A


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energy ………………. to the surroundings during cooling.

Thus the temperature remains unchanged. This constant

temperature is called …………………………..

R to S

The ………………is cooled. The particles

vibrate……………….

as the temperature ……………………….

(b) Complete the passage below by using the words given below. (solid, gas, liquid, exactly balanced, decreased, increased)

Freezing point is the temperature at which a …………………… changes into

………………………

During the freezing process, the temperature remains unchanged because the heat lost

to the environment is …………………….….. by the heat released when the liquid

particles rearrange themselves to become solid.

B THE ATOMIC STRUCTURE

Activity 3

1. Complete the table and draw the structure of each atomic model.

Model Structure Characteristic

Dalton’s atomic model proposed by …………………… in 1805

The atom was imagined as a small indivisible ball similar to a very tiny ball.

Thomson’s atomic model proposed by …………………….. in 1897

J.J Thomson discovered ……………….., a negatively-charged particle. The atom was described as a sphere of positive charge embedded with electrons.


describe the development of atomic model,

state the main subatomic particles of an atom,

compare and contrast the relative mass and the relative charge of the protons, electrons and neutrons,

define proton number,

define nucleon number,

determine the proton number,

determine the nucleon number,

relate the proton number to the nucleon number,

relate the proton number to the type of element,

write the symbol of elements,

determine the number of neutrons, protons and electrons from the proton number and the nucleon number and vice versa,

construct the atomic structure.


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Model Structure Characteristic

Rutherford’s atomic model proposed by …………………… in 1911

Ernest Rutherford discovered ……………., a positively-charged particle in an atom. The central region of atom has a very small positively-charged …………………..…, which contains almost all the mass of the atom.

Bohr’s atomic model proposed by ……………………in 1913

The electrons in an atom move in ………..……… around the nucleus which contains protons.

Chadwick’s atomic model proposed by …………..…………. in 1932

Chadwick proved the existence of ……………….., the neutral particle in the nucleus. The nucleus of the atom contains protons & neutrons, and the nucleus is surrounded by electrons.

Activity 4 Fill in the blanks and complete the table.

1. Atoms are made up of subatomic particles namely protons, …………… and ………….. 2. ………………………and …………….. are found in the nucleus of an atom while electrons

surround the nucleus.

3.

Subatomic particle Symbol Relative mass Relative electric charge

Proton

Neutron

Electron

4. The proton number of an element is the number of ………………… in its atom.

Proton Number, Z = Number Of Proton

5. Protons and neutrons are collectively called nucleons.

The nucleon number of an element is the total number of ………………….. and

……………….. in its atom.


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Nucleon Number, A = Number Of Proton + Number Of Neutron

6. The nucleon number is also known as the ……………………………..

……………………… = Nucleon Number -- Proton Number = A -- Z

7. The standard representation for an atom of any element shows the proton number and the nucleon number of the element. It can be written as follows:

A X Z

# A – Nucleon number Z – proton number X – symbol of element

1 H 1

21 Sc 45

Proton number Nucleon number

2 He 4

3 Li 7

4 Be 9

5 B 11

6 C 12

7 N 14

8 O 16

9 F 19

10 Ne 20

11 Na 23

12 Mg 24

13 Al 27

14 Si 28

15 P 31

16 S 32

17 Cl 35

18 Ar 40

19 K 39

20 Ca 40

21 Sc 45

By referring to part of the Periodic Table of Element above, complete the table below.

Element Symbol Proton number

No. of neutrons

Nucleon number

No. of electrons

Standard representation

Scandium

Aluminium

Argon

Beryllium

Boron

Calcium

Carbon

Chlorine

Fluorine


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Helium

Hydrogen

Lithium

Magnesium

Neon

Nitrogen

Oxygen

Phosphorus

Potassium

C ISOTOPES AND THEIR IMPORTANCE Activity 5 Fill in the blanks. 1. Isotopes are atoms of the same element with the ………………………. of proton but ……………………………….. of neutron. 2. Complete the table below:

Element Number of isotopes

Symbol of isotopes

Number of protons

Number of electrons

Number of neutrons

Name of isotope

Hydrogen

3

H1

1 1

1 Hydrogen-2

1 2

Oxygen

3

O16

8 8 8 Oxygen-16

8 9

O18

8

Carbon

3

6 Carbon-12

6 7

C14

6 6 8


state the meaning of isotope,

list examples of elements with isotopes,

determine the number of subatomic particles of isotopes,

justify the uses of isotope in daily life.


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Chlorine 2 Cl35

17 17

17 20

Bromine 2 35 Bromine-80

35 35 Bromine-81

3. For each of the isotope list below, state one of its uses. a) Gamma rays of Cobalt-60: …………………………………………………………………. b). Carbon-14: ………………………………………………………………………………….. c). Phosphorus-32: …………………………………………………………………………….. d). Sodium- 24: ………………………………………………………………………………… e). Iodine -131: ……………………………………………………………………………………

D THE ELECTRONIC STRUCTURE OF AN ATOM

Activity 6 1. Electrons are filled in specific shells, starting with the shell nearest to the nucleus of the

atom. Every shell can be filled only with a certain number of electrons.

The first shell can be filled with a maximum of ……………. electrons

The second shell can be filled with a maximum of ……………. electrons

The third shell can be filled with a maximum of …………….electrons

Use ‘x’ as symbol for electrons. Draw the maximum number of electrons in each shell.


describe electron arrangements of elements with proton numbers 1 to 20,

draw electron arrangement of an atom in an element,

state the meaning of valence electrons,

determine the number of valence electrons from the electron arrangement of an atom.


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2(a) Draw the electron arrangement of a sodium atom,23

11Na in the box and complete the

table given.

(b) Draw the electron arrangement of a chlorine atom, Cl in the box and complete the table given.

3 Valence electrons are electrons in the ………………..………… shell of a neutral atom. 4 Identify the number of valence electrons in these atoms according to its electron arrangement.

Atom of Element Electron Arrangement Number of valence electrons

Oxygen 2.6

Aluminium 2.8.3

Chlorine 2.8.7

Neon 2.8

Potassium 2.8.8.1

Magnesium 2.8.2

Number of protons

Number of electrons

Number of neutrons

Proton number

Nucleon number

Electron arrangement

Number of protons

Number of electrons

Number of neutrons

Proton number

Nucleon number


x

x


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Carbon 2.4

Phosphorus 2.8.5

Helium 2


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CHAPTER 3 : CHEMICAL FORMULAE AND EQUATIONS

A RELATIVE ATOMIC MASS (RAM) AND RELATIVE MOLECULAR MASS (RMM) Learning Outcomes

You should be able to:

state the meaning of relative atomic mass based on carbon-12 scale,

state the meaning of relative molecular mass based on carbon-12 scale,

state why carbon-12 is used as a standard for determining relative atomic mass and relative molecular mass,

calculate the relative molecular mass of substances.

Activity 1 (refer text book pg 28 )

Relative atomic mass of an element , Ar = The average mass of an atom of the element 1/12 x the mass of an atom of carbon-12

Example: Ar of C=12 Ar of O=16 Ar of Mg=24

1. The Relative atomic mass of an element is

……………………………………………………………...…………………………………. when compare with 1/12 of the mass of an atom of carbon – 12.

2. Carbon-12 is chosen because it is a ………………………. and can be easily handled. 3. Find the relative atomic masses of these elements.

Element Relative Atomic Mass

Element Relative Atomic Mass

Calcium, Ca Argon, Ar Sodium, Na Silver, Ag

Iron, Fe Caesium, Cs Copper, Cu Lead, Pb Carbon, C Chlorine, Cl

Hydrogen, H Flourine, F Potassium, K Aluminium, Al

Lithium, Li Zinc, Zn Bromine, Br Helium, He


Relative molecular mass of a substance, Mr = The Average mass of a molecule of the substance 1/12 x the mass of an atom of carbon-12


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Calculating Relative molecular mass,Mr

Mr= The sum of Ar of all atoms present in one molecule Example:

Mr of Water, H2O = 2(1) + 16 = 18

Mr of Carbon dioxide, CO2 = 12 + 2(16) = 44 For ionic substance , Relative formula mass , Fr = The sum of Ar of all atoms present in the formula Example: Fr of Magnesium oxide, MgO = 24 + 16 = 40 Fr of Sodium chloride, NaCl = 23 + 35.5 = 58.5

1. The relative molecular mass of a molecule is ……………………………………………………………………………………………… when compared with 1/12 of the mass of one atom of ……………………………………………

2. Calculate the relative molecular masses of the substances in the table below.

Substance

Molecular formula Relative molecular mass, Mr

Hydrogen gas H2 2(1) = 2 Propane C3H8

Ethanol

C2H5OH

Bromine gas Br2 Methane CH4 Glucose C6H12O6

Ammonia NH3

[Relative atomic mass : H,1; C,12; O,16; Br,80 ; N,14 ] 3. Calculate the relative formula masses of the following ionic compounds in the table.

Substance

Compound formula Relative formula mass, Fr

Potassium oxide K2O

2(39) + 16 = 94

Aluminium sulphate

Al2(SO4)3

2(27)+3[32+4(16)]=342

Zinc nitrate

Zn(NO3)2

2 Hydrogen

atoms Molecular

formula

Relative atomic mass

for Oxygen Relative atomic mass

for Hydrogen

All Ar, Mr and

Fr have no unit


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Aluminium nitrate

Al(NO3)3

Calcium carbonate

CaCO3

Calcium hydroxide

Ca(OH)2

Hydrated copper(II) sulphate

CuSO4.5H2O

64 + 32 + 4(16) + 5[2(1) + 16]=250

Hydrated sodium carbonate

Na2CO3.10H2O

Sodium hydrogen sulphate

NaHSO4

Aluminium chloride

AlCl3

Copper(II) sulphate

CuSO4

Zinc carbonate

ZnCO3

Potassium carbonate

K2CO3

[Relative atomic mass: O,16; C,12; H,1; K,39 ; Cu,64 ; Zn, 65; Cl, 35.5 ; Al, 27 S,32 ; Ca, 40; Na,23; N, 14]


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B THE MOLE AND THE NUMBER OF PARTICLES


define a mole as the amount of matter that contains as many particles as the number of atoms in 12 g of

12C,

state the meaning of Avogadro constant,

relate the number of particles in one mole of a substance with the Avogadro constant,

solve numerical problems to convert the number of moles to the number of particles of a given substance and vice versa.


1. To describe the amount of atoms, ions or molecules , mole is used. 2. A mole is an amount of substance that contains as many particles as the

……………….. …………………………………………………………….. in exactly 12g of carbon-12.

3. A mole is an amount of substance which contains a constant number of particles atoms, ions, molecules which is 6.02 x 1023 4. The number 6.02 x 1023 is called …………………………………… (NA) 5. In other words:

1 mol of atomic substance contains ……………………………. atoms

1 mol of molecular substance contains ……………………………. molecules

1 mol of ionic substance contains ……

…………………………….. formula units 6. Relationship between number of moles and number of particles

(atom/ion/molecules):

x Avogadro Constant

∻ A vogadro Constant

Number of moles Number of particles 0.5 mol of carbon atoms

…………………………………… atoms of carbon

number of moles number of particles


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0.2 moles of hydrogen gas ( H2)

(i) …………………………..molecules of hydrogen gas (ii) …………………………….Atoms of hydrogen

2 mol of carbon dioxide molecules

………………x 10 23 molecules of carbon dioxide gas contains : ………………. atoms of C and …………………. atoms of O

0.007 mol of calcium ions

……………………… calcium ions

…………………………. mol of water

6.02 x 10 25 molecules of water

0.4 mol of ozone gas ( O3)

………………….x 10 23 molecules of ozone, contains : ……………………… atoms of Oxygen.

7. Complete these sentences .

a) 1 mol of calcium contains ………………………………………….. atoms b) 2 mol of iron contains ……………………………………………….. atoms c) 2 mol of magnesium oxide, (MgO) contains …………………………………………ions d) 2 mol of sodium carbonate, (Na2CO3) contains ………………………………………. e) 3 mol of carbon dioxide, (CO2) contains …………………………………….. molecules f) 0.5 mol Copper (II) nitrate, Cu(NO3)2 contains ………………………………. Cu2+ ions and …………………………………………………. NO3

- ions


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C NUMBER OF MOLES AND MASS OF SUBSTANCES


state the meaning of molar mass,

relate molar mass to the Avogadro constant,

relate molar mass of a substance to its relative atomic mass or relative molecular mass,

solve numerical problems to convert the number of moles of a given substance to its mass and vice versa.


1. The molar mass of a substance

= The molar mass of _________________ mole of the substance. = The mass of (NA) number of particles = The mass of ____________________ particles

x Molar mass

2. Calculating the Mass from a number of Moles

Number of moles = . mass of the substance . Mass of 1 mole of the substance Therefore : Mass of substance = Number of moles x Mass of 1 mole Example 1 : What is the mass of 2 moles of carbon ? Mass = 2 x 12 = 24g Example 2 : What is the mass of 2 moles of H2O ? Mass = 2 x [ 2(1) + 16 ] = 36g

Number of moles Mass in g

∻ Molar mass


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3. Calculate the masses of these substances

a) 2 moles of aluminium atoms Mass =

b) 10 moles of iodine atoms Mass =

c) 3 moles of lithium atoms Mass =

d) 0.5 moles of oxygen gas (O2) Mass =

e) 0.1 moles of sodium Mass =

f) 2 moles of chlorine molecules (Cl2) Mass =

g) 1 mole of carbon dioxide ( CO2) Mass =

h) 3 moles of nitric acid, ( HNO3 ) Mass =

i) 2 moles of calcium carbonate (CaCO3 ) Mass =

j) 0.25 moles of calcium chloride (CaCl2 ) Mass =

k) 0.25 moles of sodium hydroxide (NaOH) Mass =

l) 0.25 moles of sodium carbonate (Na2CO3) Mass =

m) 0.5 moles of potassium manganate (VII) (KMnO4) Mass =

n) 0.25 moles of hydrated magnesium sulphate (MgSO4.7H2O) Mass =

Activity 5 4. Calculate the Number of Moles from a given Mass Example : How many moles are there in 88g of CO2 Number of moles = 88 = 2 moles 44

a) 2g of helium atoms Number of moles =

b) 6g of carbon atoms Number of moles =


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c) 16g of helium atoms Number of moles =

d) 4g of sulphur atoms Number of moles =

e) 4g of oxygen molecules (O2) Number of moles =

f) 213g of chlorine molecules (Cl2) Number of moles =

g) 0.56g of nitrogen molecules (N2) Number of moles =

h) 254g of iodine molecules (I2) Number of moles =

i) 88g of carbon dioxide (CO2) Number of moles =

j) 3.1g of sulphur dioxide (SO2) Number of moles =

k) 560g of potassium hydroxide (KOH)

Number of moles =

l) 392g of sulphuric acid (H2SO4) Number of moles =

m) 170g of ammonia (NH3) Number of moles =

n) 120g of magnesium oxide (MgO) Number of moles =

o) 4g of sodium hydroxide (NaOH) Number of moles =

p) 73g of hydrogen choride (HCl) Number of moles =

q) 15.8g of potassium manganate (VII) KMnO4

Number of moles =

r) 8g of ammonium nitrate (NH4NO3) Number of moles =

s) 0.78g of aluminium hydroxide Al(OH)3

Number of moles =

t) 0.92g of ethanol (C2H5OH) Number of moles =


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Activity 6 5. Complete the following table.

Element/compound

Chemical formulae

Molar mass

Calculate

Copper

Cu

RAM= 64

(a)Mass of 1 mol = ……………g (b) Mass of 2 mol = …………. g (c)Mass of ½ mol = ………….g (d)Mass of 3.01x1023 Cu atoms =

Sodium hydroxide

NaOH

RFM= 40

(a) Mass of 3 mol of sodium hydroxide = (b) Number of moles of sodium hydroxide in 20 g =

Zinc nitrate

Zn(NO3)2

RFM =

a) Number of moles in 37.8 g of zinc nitrate :

D NUMBER OF MOLES AND VOLUME OF GAS Learning Outcomes


state the meaning of molar volume of a gas,

relate molar volume of a gas to the Avogadro constant,

make generalization on the molar volume of a gas at a given temperature and pressure,

calculate the volume of gases at STP or room conditions from the number of moles and vice versa,

solve numerical problems involving number of particles, number of moles, mass of substances and volume of gases at STP or room conditions.

Activity 7 (refer text book pg 36, 37 )

1. The molar volume of a gas is defined as the ………………………………………………….……………………………………………………


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2. One mole of any gas always has the …………………………………………… under the same temperature and pressure. 3. The molar volume of any gas is

24 dm3 at ……………………………………………… or

22.4 dm3 at ……………………………………………. Example : 1 mol of oxygen gas, 1 mol of ammonia gas, 1 mol helium gas and 1 mol sulphur dioxide gas occupies the same volume of 24 dm3 at room condition x 22.4 / 24 dm3 x 22.4/24 dm3

∻22.4/24 dm3

4. Calculate the volume of gas in the following numbers of moles at STP Example : Find the volume of 1 mole of CO2 gas

Volume = number of moles x 22.4 dm3 = 1 x 22.4 dm3 = 22.4 dm3

a) 3 moles of oxygen Volume =

b) 2 moles of CH4 Volume =

c) 0.3 moles of Argon Volume =

d) 0.2 moles of SO3 Volume =

e) 0.1 moles of N2 Volume =

f) 1.5 mol of N2 Volume =

Number of moles of gas Volume of gas


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5. Complete the diagram below . (Refer to Page 33,34 & 38-Chemistry textbook)

Activity 8 Solve these numerical problems

1. What is the volume of 0.3 mole of sulphur dioxide gas at STP? [Molar volume: 22.4 dm3 mol-1 at STP]

(Ans: 6.72 dm3)

2. Find the number of moles of oxygen gas contained in a sample of 120 cm3 of the gas at room conditions.

[Molar volume: 24 dm3 mol-1 at room conditions]

(Ans: 0.005 mol)

3. Calculate the number of water molecules in 90 g of water, H2O. [Relative atomic mass: H, 1; O, 16. Avogadro constant, NA: 6.02 x 1023 mol-1] (Ans; 3.01x 10

24 molecules)

Volume of gas (dm3)

Number of moles Mass in gram No of particles


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4. What is the volume of 24 g methane ,CH4 at STP? [Relative atomic mass: H, 1; C, 12. Molar volume: 22.4 dm3 mol-1 at STP] (Ans: 33.6 dm

3)

5. How many aluminium ions are there in 20.4 g of aluminium oxide, Al2O3? [Relative atomic mass: O, 16; Al, 27. Avogadro constant, NA: 6.02 x 1023 mol-

(2 x 0.2 x 6.02 x1023

)

6. Calculate the number of hydrogen molecules contained in 6 dm3 of hydrogen gas at

room conditions. [Molar volume: 24 dm3 mol-1 at room conditions Avogadro constant, NA: 6.02 x 1023 mol-1]

(Ans: 1.505x1023

molecules)

7. Find the volume of nitrogen in cm3 at STP that consists of 2.408 x 1023 nitrogen molecules. [Molar volume: 22.4 dm3 mol-1 at STP. Avogadro constant, NA: 6.02 x 1023 mol-1]

(Ans: 8.96 dm

3 )


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E CHEMICAL FORMULAE

Learning Outcomes You should be able to

state the meaning of chemical formula

state the meaning of empirical formula

state the meaning of molecular formula

determine empirical and molecular formula of substances

compare and contrast empirical formula with molecular formula

solve numerical problems involving empirical and molecular formula.

write ionic formula of ions

construct chemical formulaf ionic compounds

state names of chemical compounds using IUPAC nomenclature.

use symbols and chemical formula for easy and systematic communication in the field of chemistry.

ACTIVITY 9 (Refer text book pg 40)

1) A Chemical formula - A representation of a chemical substance using letters for

……………………………………… and subscripts to show the numbers of each type of

…………………….. that are present in the substance.

2) Complete this table

Chemical subtance Chemical

formulae

Notes

Water

……………..

2 atoms of H combine with 1 atom of O

………..

NH3

……. atoms of H combine with 1 atom of N

Propane

C3H8

…….. atoms of C combine with ……. atoms of

H

Magnesium oxide

……………..

…………………………………………….

………………..

H2SO4

……………………………………………

H2 Subscript shows 2 hidrogen atoms in a molecule

The letter H shows ……………. …………….


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3). There are two types of chemical formulae. Complete the following:

** Empirical Formula The simplest ………… ……….. ratio of atoms of each

in the compound.

** Molecular Formula The actual …………… of atoms of each …………… that are

present in a molecule of the compound

Remember:

Example: (i) Compound – Ethene (ii) Compound – Glucose

Molecular formula - 42HC Molecular formula - 6126 OHC

Empirical formula - ................... Empirical formula - ....................

Activity 10

1 Find the empirical formula of a compound

Example of calculation:

a) When 11.95 g of metal X oxide is reduced by hydrogen, 10.35 g of metal X is

produced. Find the empirical formula of metal X oxide [ RAM; X,207; O,16 ]

Element X O

Mass of element(g) 10.35 11.95-10.35

Number of moles of

atoms

10.35÷207 (11.95-10.35)÷16

Ratio of moles

Simplest ratio of moles

Empirical formula : …………

b) A certain compound contains the following composition:

Na 15.23%, Br 52.98% , O 31.79%, [ RAM : O, 16; Na, 23; Br,80]

(Assume that 100g of substance is used)

Element Na Br O

Mass of element(g) 15.23 52.98 31.79

Number of moles atoms 15.23 ÷23 52.98÷80 31.79÷16

Ratio of moles

Simplest ratio of moles

Molecular formula = (Empirical formula)n


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Empirical formula:: ……………………………………………….

c) Complete the table below.

Compound Molecular Formula Empirical formula Value of n

Water H2O

Carbon Dioxide CO2 CO2

Sulphuric Acid H2SO4

Ethene C2H4 CH2

Benzene C6H6

Glucose C6H12O6

d) 2.52g of a hydrocarbon contains 2.16 g of carbon. The relative molecular mass of the

hydrocarbon is 84. [RAM H,1; C,12]

i. Find the empirical formula of the hydrocarbon

ii. Find the molecular formula of the carbon.

Activity 11 :Chemical Formula for ionic compounds:

Complete the table below :

Cation Formula Anion Formula

Hydrogen ion H Flouride ion F

Lithium ion Chloride ion

Sodium ion Bromide ion

Potassium ion Iodide ion

Magnesium ion Hydroxide ion

Calcium ion 2Ca Nitrate ion


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Barium ion 2Ba Manganate(VII) ion

Copper(II) ion Ethanoate ion COOCH3

Iron(II) ion 2O

Iron (III) ion Sulphate ion

Lead (II) ion Sulphide ion 2S

Zinc ion Carbonate ion

Chromium (III) ion Dichromate (VI) ion 2

72OCr

Aluminium ion 3Al 3

4PO

Ammonium ion Chromate (VI) ion

Avtivity 12

a) Chemical formula of an ionic compound comprising of the ions Xm+ and Yn- is constructed

by exchanging the charges of each element. The formula obtained will XnYm

Example : Sodium oxide Copper (II) nitrate

Na+ O2- Cu2+ NO3-

+1 -2 +2 -1

2 1 1 2

= Na2O = ....................

b) Construct a chemical formula for each of the following ionic compounds:

(i) Magnesium chloride

(ii) Potassium carbonate

(iii) Calcium sulphate

(iv) Copper (II) oxide

(v) Silver nitrate

(vi) Zinc nitrate


28

(vii) Aluminium oxide

(viii) Iron(II) hydroxide

(ix) Lead(II) sulphide

(x) Chromium(III) sulphate

CHEMICAL EQUATIONS

Learning Outcomes

You should be able to

1. state the meaning of chemical equation identify the reactants and products of a chemical equation

2. write and balance chemical equations 3. interpret chemical equations quantitatively and qualitatively 4. solve numerical problems using chemical equations 5. identify positive scientific attitudes and values practiced by scientist in doing research 6. justify the need to practice positive scientific attitudes and good values in doing researsh 7. use chemical equations for easy and systematic communication in the field of chemistry.

Activity 13 (refer text book pg 48)

Example: C (s) + 2O (g)

2CO (g)

Reactant product

1) Qualitative aspect of chemical equation:

a) Arrow in the equation the way the reaction is occurring

b) Substances on the left-hand side ……………………..

c) Substances on the right-hand side ………………………

d) State of each substance ………: (s), ………………(l), gas ……….and aqueous

solution ……………….

2) Quantitative aspect of chemical equations

Coefficients in a balanced equation the exact proportions of reactants and products in

equation.

Example: 22H (g) +

2O (g) 2 OH 2(l)

(Interpreting): 2 molecules (2 mol) of 2H react with 1 molecule (1 mol) of

2O to produced 2

molecules(2 mol) of water


29

Complete the following word equations and write in chemical equation

a) Sodium + chlorine …………………………..

………… + …………… NaCl

b) Carbon + ……….. Carbon dioxide

………. + ………… ……………………..

c) Sulphur + oxygen ……………………………

……….. + ……….. …………………………..

d) Zinc + oxygen ………………………………..

………… + O2 ………………………………..

3) Write a balanced equation for each of the following reactions and interpret the equations

quantitatively.

(a). Carbon monoxide gas + oxygen gas carbon dioxide gas

………………………………………………………………………………………………………

Interpreting:

……………………………………………………………………………………………………………

(b). Hydrogen gas + nitrogen gas ammonia gas

……………………………………………………………………………………………………….

Interpreting:

…………………………………………………………………………………………………………..

(c). Aluminium + Iron (III) oxide aluminium oxide + Iron

……………………………………………………………………………………………………….

Interpreting:

……………………………………………………………………………………………………………

.

Activity 14

** Numerical Problems Involving Chemical Equations

Hydrogen peroxide decomposes according to the following equation:

222OH (l) 2 OH 2

(l) + 2O (g)

1). Calculate the volume of oxygen gas, 2O measured at STP that can be obtained from

the decomposition of 34 g of hydrogen peroxide, 22OH .

[Relative atomic mass : H, 1 ; O, 16. Molar volume : 22.4 3dm 1mol at STP]


30

(Ans: 11.2 dm3)

2).Silver carbonate Ag2CO3 breaks down easily when heated to produce silver metal

2 Ag2CO3(l) 4 Ag (s) + 2 2CO (g) +

2O

Find the mass of silver carbonate that is required to produce 10 g of silver

[Relative atomic mass: C, 12 ; O, 16 ; Ag, 108]

(Ans : 12.77g)

3). 16 g of copper (II) oxide, CuO is reacted with excess methane, 4CH . Using the equation

below, find the mass of copper that is produced.

[Relative atomic mass : Cu, 64 ; O, 16]

4 CuO (s) + 4CH (g) 4 Cu (s) +

2CO (g) + 2 OH 2 (l)

(Ans : 12.8 g)


31

4). A student heats 20 g of calcium carbonate 3CaCO strongly. It decomposes according to

the equation below:

3CaCO (s) CaO (s) + 2CO (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.

[Relative atomic mass: C, 12 ; O, 16; Ca, 40. Molar volume :

24 dm3 1mol at room conditions]

(Ans : (a). 4.8 dm 3

(b) 11.2 g)


32

CHAPTER 4 :PERIODIC TABLE OF ELEMENTS

A ANALYSING THE PERIODIC TABLE OF ELEMENTS Learning Outcomes You should be able to:

describe the contributions of scientists in the historical development of the Periodic Table,

identify groups and periods in the Periodic Table,

state the basic principle of arranging the elements in the Periodic Table from their proton numbers,

relate the electron arrangement of an element to its group and period,

explain the advantages of grouping elements in the Periodic Table,

predict the group and the period of an element based on its electron arrangement.

Activity 1 Draw lines to match name of scientist with their contribution

Name of Scientist Contribution

Lothar Meyer

First scientist to classify substances

Henry J. G. Moseley

Classified the elements into group of three with similar chemical properties

Antoine Lavoisier

Arranged the known elements in order of increasing atomic mass . Elements with similar properties recurred at every eighth element.

John Newlands

Plotted a graph of the atomic volume against the atomic mass of elements

Johann W. Dobereiner

Left gaps in the table to be filled by undiscovered elements

Dimitri Mendeleev

Rearranging the elements in order of increasing proton number

Activity 2 1 Arrangement of elements in the Periodic Table

a) Elements are arranged in an increasing order of ………………………… b) Each vertical column is called a ………………..

c) Each horizontal rows is called ……………………………. d) Elements with similar chemical properties are placed in the same

………………………


33


Element Electron arrangement

Number of valence electrons

Group Number of shells occupied with electrons

Period

H1

1

He4

2

1

Li7

3

2.1

B11

5

2.3

C12

6

2

2

N14

7

5

15

O16

8

F19

9

17

Ne20

10

2.8

Mg24

12

Al27

13

2.8.3

Si28

14

S32

16

Cl.35

17

K39

19

Ca40

20


34

3. An atom of element E has 10 neutrons. The nucleon number of element E is 19. In which group and period is element E located in the Periodic Table? Answer :Group ………………………… Period ………………………… 4. An atom of element G has 3 shells occupied with electrons. It is placed in group 17 of the Periodic Table. What is the electron arrangement of atom G? Answer : Electron arrangement of atom G …………………

B ANALYSING GROUP 18 ELEMENTS Learning Outcomes: You should be able to:

list all Group 18 elements,

state in general the physical properties of Group 18 elements,

describe the changes in the physical properties of Group 18 elements,

describe the inert nature of elements of Group 18,

relate the inert nature of Group 18 elements to their electron arrangements,

relate the duplet and octet electron arrangements of Group 18 elements to their stability,

describe uses of Group 18 elements in daily life..

Activity 3 : Analysing Group 18 Elements 1 Physical Properties of Group 18 Elements [Circle the correct answer].

Down the Group :

a) Atomic radius /Atomic size increase/ decrease b) Melting point and Boiling point increase/ decrease c) Density increase/ decrease

2 Complete the figure below by giving the uses of Group 18 .

3 All noble gases are ……………………. which means chemically unreactive. This is

because they have ………………………. electron arrangement.

Uses of Group

18

Krypton

*

Helium

*

Neon

*

Argon

*

Radon

*

Xenon

*


35

Example : (i) Helium atom has …………………. valence electrons which is called a

…………………………… electron arrangement

(ii) Other noble gases has …………………… valence electrons which is

called an ………………………. electron arrangement.

4.The noble gases exists as ……………………. gases because their electron arrangement

are very …………………………..

C ANALYSING GROUP 1 ELEMENTS Learning Outcomes You should be able to:

list all Group 1 elements.

state the general physical properties of lithium, sodium and potassium,

describe changes in the physical properties from lithium to potassium,

list the chemical properties of lithium, sodium and potassium,

describe the similarities in chemical properties of lithium, sodium and potassium,

relate the chemical properties of Group 1 elements to their electron arrangements,

describe changes in reactivity of Group 1 elements down the group,

predict physical and chemical properties of other elements in Group 1

Activity 4 A State 3 physical properties of group 1 elements.

a) ………………………………………………………………………………………………

b) ………………………………………………………………………………………………

c) ………………………………………………………………………………………………

B Chemical Properties of Group 1 Elements (Alkali metals )

1. Group 1 elements react vigorously with water to produce alkaline metal hydroxide solutions and hydrogen gas

Example: Li2 + OH 22 LiOH2 + 2H

Write down the balanced equation when potassium reacts with water …………………………………………………………………………………………………..

2. Group 1 elements react with oxygen gas rapidly to produce white solid metal oxides.

Example: Li4 + 2O OLi22

Write down the balanced equation when rubidium reacts with oxygen …………………………………………………………………………………………………

3. Group 1 elements react t with chlorine gas 2Cl , to form white solid metal chlorides.

Example: Na2 + 2Cl NaCl2

Write down the balanced equation when potassium reacts with chlorine gas ………………………………………………………………………………………………..


36

4 The reactivity of Group 1 elements increases when going down the group. Explain why. a) When going down Group 1 the single valence electron in the outermost

occupied shell become ………………….. from the nucleus. b) The attraction between the nucleus and the valence electron become

…………. . c) Therefore it is ……………. for the atom to donate the single valence electron to

achieve the stable electron arrangement. 5 Potassium reacts more vigorously with water as compared to sodium. Explain. (Proton number: Na, 11 ; K, 19) ………………………………………………………………………………………………… ………………………………………………………………………………………………… Activity 5 Complete the table below and answer the following questions :

Alkali Metal Proton number

Number of electrons


Number of valence

electrons Lithium 3 Sodium 11

Potassium 19 Rubidium 37 37 2.8.18.8.1 Caesium 55 55 2.8.18.18.8.1

1) Lithium , sodium and potassium have similar chemical properties because their atoms have ………….…………electron in their outermost occupied shell. 2) How an atom of alkali metal achieve a stable electron arrangement of inert gas ? ……………………………………………………………………………………………………… 3) What is the charge of an alkali metal ion ? . …………............................................................................................……………………… 4) Reactivity of alkali metals increases from Lithium to Caesium . Explain why . ………………………………………………………………………………………………………

………………………………………………………………………………………………………


37

Activity 6 To Investigate The Chemical Properties of Lithium, Sodium & Potassium

1) The Reaction of alkali metals with water, OH 2

Problem Statement: How does the reactivity of Group 1 elements change when they react with water? Hypothesis: When going down Group 1, alkali metals become more reactive in their reactions with water. Variables: Manipulated variable – Different types of alkali metals Responding variable – Reactivity of metals with water Fixed variables – water, size of metals

a) Write the procedure to carry out this experiment. (refer to practical book pg 39)

b) Complete this table (Data & Observation )

Alkali metal Observation Lithium

Sodium

Potassium

2) The Reaction of alkali metals With Oxygen, O2

(This procedure can also be used to test the reaction of alkali metals with chlorine gas!)

a) Problem Statement:

…………………………………………………………………………………………

b) Hypothesis:

…………………………………………………………………………………………

c) Variables: Manipulated variable ……………………………………………………………………

alkali metal

water


38

Responding variable – ……………………………………………………………………. Fixed variables – ………………………………………………………….......................

d) Write the procedure of this experiment. (refer to practical book pg 36)

e) Data & Observation

Alkali metal Observation Lithium

Sodium

Potassium

f) Based on your results, arrange the alkali metals in ascending order of

reactivity. ……………………………………………………………………………………

Alkali metal

oxygen


39

D ANALYSING GROUP 17 ELEMENTS Learning Outcomes You should be able to:

list all Group 17 elements,

state the general physical properties of chlorine, bromine and iodine,

describe changes in the physical properties from chlorine to iodine,

list the chemical properties of chlorine, bromine and iodine,

describe the similarities in chemical properties of chlorine, bromine and iodine,

describe changes in reactivity of Group 17 elements down the group,

predict physical and chemical properties of other elements in Group 17,.

Activity 7 1) State the uses of Chlorine and iodine

a) Chlorine – …………………………………………………………………………….

b) Iodine – ………………………………………………………………………………..

2) (a) Give the physical state of halogens below at room temperature :

i) Fluorine: ……………………………… ii) Chlorine:………………………………………… iii) Bromine: ……………………………… iv) Iodine : ………………………………………….

(b) Fill in the blanks below.

2Cl melting and colour of Density

2Br boiling points halogens ……….

2I ……….. becomes

…………. 3) When going down the Group 17, the melting and boiling points increase. Explain ……………………………………………………………………………………………………… ……………………………………………………………………………………………………… 4) Chemical Properties of Group 17 Elements

a) Group 17 elements react with water to form two acids

Example: 2Cl + OH 2 HCl + HOCl

Hydrochloric acid hypochlorus acid Write a balanced equation when bromine reacts with water.

…………………………………………………………………………………………


40

b). In gaseous state they react with hot iron to form a brown solid, iron (III) halides.

Example: Fe2 + 23Br 32FeBr

Write a balanced equation when iodine vapour reacts with iron

…………………………………………………………………………………………

c). Group 17 elements react t with sodium hydroxide solution, NaOH , to form

sodium halide, sodium halate (I) and water

Example: 2I + NaOH2 NaI + NaOI + OH 2

Write a balanced equation when chlorine reacts with sodium hydroxide solution

…………………………………………………………………………………………

5)The reactivity of Group 17 elements decreases when going down the group. Explain why.

a) When going down the Group 17 atomic size …………………………….. b) The outermost occupied shell becomes …………………. from the nucleus. c) Therefore the strength to attract one electron into the outermost occupied

shell by the nucleus becomes ………………………

6) Chlorine gas reacts more vigorously with hot iron as compared to bromine gas. Explain (Proton number: Cl, 17 ; Br, 35)

………………………………………………………………………………………………… …………………………………………………………………………………………………

Activity 8

To investigate the Chemical properties of Group 17 elements.

The Reaction of halogens with iron (refer practical book pg 44)

1) Data and Observation (Complete the following table)

Halogen Reactant

Observation Chlorine Bromine Iodine

Water

Iron wool

Sodium hydroxide , NaOH solution


41

2) Based on your results, arrange the halogens, 2Cl , 2Br , 2I in ascending order of

reactivity. ……………………………………………………………………………………………………… 3) Element E is placed below element D in Group 17 of the Periodic Table. (a) Compare the melting and boiling points of element D with element E. Explain your answer. ……………………………………………………………………………………………………… ……………………………………………………………………………………………………… (b) Write a chemical equation for the reaction between element D and hot iron. ………………………………………………………………………………………………………

E ANALYSING ELEMENTS IN A PERIOD Learning Outcomes You should be able to:

list all elements in Period 3,

write electron arrangements of all elements in Period 3,

describe changes in the properties of elements across Period 3,

state changes in the properties of the oxides of elements across Period 3,

predict changes in the properties of elements across Period 2,

Activity 9 Period 3 in the Periodic Table – Properties of Elements 1) Complete the table and answer the question given below

Element

Na Mg Al Si P S Cl Ar

Proton number

11 12 13 14 15 16 17 18


Number of valence electrons

Atomic radius (pm)

186 160 143 118 110 104 100 94

Physical state at room temperature

Solid Solid Solid Solid Solid Solid Gas Gas

Electronegativity

0.9 1.2 1.5 1.8 2.1 2.5 3.0 -


42

2) Fill in the blanks with the correct answer.

a). The proton number ……………………….. by one unit from one element to the next

element

b) All the atoms of elements have ……………………….. shells occupied with electrons.

c) The number of valence electrons in each atom …………………………. from 1 to 8.

d) The physical state at room temperature changes from …………………….. to

…………………………

e) The atomic radius (atomic size) of elements ………………… from left to right across the

period 3

f) The electronegativity of elements. …………………….. from left to right across the period

3 .

Activity 10

1) Below are some oxides of elements of Period 3.

(a) Which of these oxides can react with

(i) dilute nitric acid? ………………………… (ii) sodium hydroxide solution? ……………………… (b) Based on your answers in (a), what inferences can you make about the properties of each of the oxides? (Acidic/Basic/Amphoteric)

Oxides Sodium oxide

Magnesium oxide

Aluminium oxide

Silicon (IV) oxide

Phosphorus (V) oxide

Sulphur dioxide

Oxides properties

2)

The diagram above shows the symbols of lithium, carbon and fluorine. (a) Which period in the Periodic Table can you find the three elements? Explain. ………………………………………………………………………………………………………

*Sodium oxide, ONa2 *Silicon (IV) oxide, 2SiO

*Aluminium oxide, 32OAl *Sulphur dioxide, 2SO

Li7

3, C12

6, F19

9


43

(b) Arrange the three elements in order of increasing atomic size. ………………………………………………………………………………………………… (c) Compare the electronegativity of the three elements. Explain your answer.

The electronegativity of the elements (i)………………………… from Li , C , F This is due to the (ii)………………….. nuclei attraction on the valence electrons and the (iii)……………………………… in atomic size.

F TRANSITION ELEMENTS Learning Outcomes You should be able to:

identify the positions of transition elements in the Periodic Table,

give examples of transition elements,

describe properties of transition elements,

state uses of transition elements in industries. Activity 11 1 (a) Transtition elements are elements from Group ………………. ……..to Group ……………………… (b) State 3 examples of transtition elements found in Period 4 …………………………………………... 2 Complete the diagram below.

3 Transition elements and their compounds are useful catalysts. Complete the table below

Chemical Process Product Catalysts Haber Ostwald Contact

Special characteristics of

Transition elements


44

4 Transition elements form coloured ions or compound. Complete the table below

Ion of transition element Formula of the ion Colour of aqueous solution

Cooper (II) ion

Iron (II) ion

Fe 2+

Iron (III) ion

Yellowish Brown

Chromium (III) ion

Chromate (VI) ion

Dichromate (VI) ion

Manganese (II) ion

Manganate (VI) ion

4 Transition elements form ions with different Oxidation Numbers.

Elements Compound Chemical Formula Oxidation Number

Manganese Manganese (II) chloride

Manganese (IV) oxide

Potassium manganate

(VI)

Iron Iron (II) chloride

Iron (III) chloride

Copper Copper (I) chloride

Copper (II) oxide

**(Precious stones such as emerald, rubies, sapphire and jade are beautiful due to the

colours of the transition element compounds present in them )


45

Activity 12 1 Diagram 1 shows part of the Periodic Table of the Elements. D, E, G, L, M, and J, that do not

represent the actual symbol of the element

Diagram 1

Using the letters in the Periodic Table of the Elements in Diagram 1, answer the following questions.

(a) (i) State the position of element E in the Periodic Table.

…………………………………………………………………………………………………….

(ii) Choose the element which exhibit different oxidation numbers in its compounds.

…………………………………………………………...……………………….. ……………..

(b) Element D combines with element L to form a compound.

Write the chemical formula of this compound.

..............................………………………………………………………………….................

(c) D and E have the same chemical properties

(i) Which element is more reactive? .........................................................................................................................................

(ii) Explain your answer in (c) (i).

……………………………………………………………………………………………………. …………………………………………………………………………………………………….

(d) Which element exists as diatomic molecules?

……..….………………………………………………………………………… ……………….

L

D

M

G

J

E


46

2 The information shows the chemical symbols which represent elements W, X, Y and Z.

W X Y Z (a) State three subatomic particles in an atom. .

……….............……………………………………………………...............………..

(b) (i) What is the meaning of the “period” in the Periodic Table of element?

............................................................................................................................

(ii) State the period of element W in the Periodic Table of element.

Explain.

....….…………………………………………………..............….......………………

............................................................................................................................

(c) (i) Compare the atomic size of element W and X.

............................................................................................................................

(ii) Explain your answer in (c) (i).

............................................................................................................................ ............................................................................................................................

27

13

35

17

12

6

23

11


47

CHAPTER 5 : CHEMICAL BONDS A Formation of Compounds Learning Outcomes: You should be able:

explain the stability of noble gases

explain the conditions for the formation of chemical bonds

state the types of chemical bonds

Activity 1: Formation of chemical bonds Choose the correct answer from the table 1 Noble gases are ………………… gases. They exist as……………………….. gases and

are chemically unreactive . They have ………….………….. octet or ……………..……… electron

…………………………..

2 Other atoms besides noble gases tend to achieve the stable electron arrangement through the

formation of …………………………

3 Two types of chemical bonds :

(i) …………………….. bond

- formed when atoms join together by transferring of electrons

(ii) …………………….. bond - formed when atoms join together by ………………………………..of electrons

B IONIC BONDS Learning outcomes: You should be able to:

explain the formation of ions

write the electron arrangements for the ions formed

explain the formation of ionic bonds

illustrate electron arrangement of an ionic bond

illustrate the formation of ionic bonds

Activity 2 : Formation of ions 1 Underline the correct answer.

To achieve a stable electron arrangement : (i) A metal atom (donates / accepts) electrons , forming a (positive / negative) ion

called cation .

(ii) A (non-metal / metal) atom accepts electrons , forming a (positive / negative) ion

Sharing ionic stable chemical bonds monoatomic arrangement inert covalent duplet


48

called anion .

2

Complete the diagram below.

(a)

(b)

(Refer to page 84 – 85 - F4 Chemistry text book)

Activity 3 : Formation of ionic bonds Fill in the blanks with the correct words. 1 Formation of ionic compound, sodium chloride ( NaCl )


49

Electron arrangement of sodium atom is .............................................................

A sodium atom ………………….one electron to achieve the ………………electron

arrangement which is 2.8.

Sodium ion, ……..…........ is formed

Electron arrangement of chlorine atom is…………………………………………..

Electron from sodium atom is transferred to a …………………………….…atom

A chlorine atom …………………………electron from sodium atom to

………………………….. the stable electron………………………which is 2.8.8

Chloride ion,………………….. is formed

The sodium ion, Na and chloride ion,

Cl formed are ……………………..to one another

to form an ionic compound …………………….., NaCl .

The strong ……………………………….forces between the opposite-charged ions is

called ………………………….bond.

(Refer to page 86 - F4 Chemistry text book)

2 Complete the diagram below.

2.8.1 2.8.7 ............. ...............

Sodium atom,Na Chlorine atom,Cl Sodium ion, …… Chloride ion, …..

3 (a)

Fill in the blanks with the correct answers. Formation of ionic compound magnesium chloride, MgCl2 .

Electron arrangement of magnesium atom is ..................................................

A magnesium atom ………………….two electrons to achieve the ………………electron

arrangement which is, 2.8.

Magnesium ion, ……..…........ is formed

Electron arrangement of chlorine atom is…………………………………………

Electrons from magnesium atom is transferred to two ………………….…atoms

A chlorine atom …………………………electron from magnesium atom to ……………..

the stable electron …………………………which is 2.8.8.

Chloride ion,………………….. is formed


50

The magnesium ion,Mg2+ and two chloride ions,Cl formed are ………………to one

another to form an ionic compound …………………………., MgCl2

The strong ……………………………….forces between the opposite-charged ions is

called …………………………..bond


(b) Complete the diagram below.


C Covalent Bonds

Learning Outcomes: You should be able :

state the meaning of covalent bonds

explain the formation of covalent bonds

illustrate the formation of covalent bonds

compare and contrast the formation of ionic and covalent bonds Activity 4 : Formation of covalent bonds Fill in the blanks with the correct words. 1 Covalent bonds are formed when ..………………… atoms …………………….. electrons to

achieve ……………………. electron arrangements .

2 Types of covalent bonds:-

(i) …………………………………………………………………………………………………

(ii) …………………………………………………………………………………………………

(iii) …………………………………………………………………………………………………

3 A single bond is formed when …………………of electrons is shared between two atoms.

A double bond is formed when ………………..of electrons is shared between two atoms.

A triple bond is formed when ………………….of electrons is shared between two atoms.

4 Formation of hydrogen molecules, H2 :-


51

A hydrogen atom has ……valence electron, with an electron arrangement of..…….

It needs ……….. more electron to achieve the …………….. electron arrangement

………..hydrogen atoms ………………… one electron each for ………………

Shared-paired electrons forms a …………….. bond in the hydrogen molecule, H2

Single bond holds the two hydrogen atoms together because the shared-pair of electrons

is attracted to the………………….. of both atoms

5 Complete the diagram below. (a)

(b) A covalent bond can be illustrated by using……………………………….

Activity 5 : Formation of covalent bonds

1 Formation of oxygen molecules, O2 :

An oxygen atom has …… valence electron, with an electron arrangement of…..…

It needs ……….more electrons to achieve the …………….. electron arrangement

…....oxygen atoms share………pairs of electrons forming a…..……………bond

2 (a)

Draw the electron arrangement for the formation of oxygen molecule. [Proton number : O, 8 ;]

(b) Illustrate the formation of oxygen molecule using the Lewis structure.


52

3 Formation of a nitrogen molecule, N2 :

A nitrogen atom has …… valence electron, with an electron arrangement ..……..

It needs ……… more electrons to achieve the …………….. electron arrangement

………..nitrogen atoms share………………pairs of electrons forming a ………………

covalent bond

4 Draw the electron arrangement for the formation of nitrogen molecule. [Proton number : N, 7]

(b) Illustrate the formation for nitrogen molecule using the Lewis structure.

Activity 6 1 Draw the electron arrangements for the formation of the following ionic compounds:

[Proton number : Li, 3 ; Ca, 20 ; O, 8 ; Cl, 17 ] (a) Lithium oxide, OLi2

(b) Calcium chloride,

2CaCl


53

2 Draw the electron arrangements for the formation of the following covalent compounds [Proton number : C, 6 ; Cl, 17, : O, 8]

(a) Tetrachloromethane, CCl4

(b) Carbon dioxide molecule, CO2

3 Complete the table below to compare the characteristics for the formation of ionic and covalent

bonds.

Ionic bond Characteristic Covalent bond

Valence electrons

Electrons involved

Elements Non-metals atom and non metal atoms

Electron transfer to achieve stable electron arrangement

Bond formation

Particles

D PROPERTIES OF IONIC AND COVALENT BONDING Learning outcomes You should be able to:

list the properties of ionic compounds.

list the properties of covalent compounds

explain the differences in the electrical conductivity of ionic and covalent compounds

describe the differences in melting and boiling points of ionic and covalent compounds

compare and contrast the solubility of ionic and covalent compounds

state the uses of covalent compounds as solvents.

Activity 7 : Physical properties of ionic and covalent compounds 1. Complete the table of the properties of ionic and covalent compounds, using the words given in the box :

conduct electricity in aqueous solution or molten state do not conduct electricity low solid high insoluble soluble solid, liquid, gas


54

Ionic compound Properties Covalent compound

Physical states at room temperature

Melting points

Boiling points

Electrical Conductivity

Solubility in water

Solubility in organic solvent

2 Explain why ionic compounds are able to conduct electricity in aqueous solution or in

molten state but not in solid state.

……………………………………………………………………………………………………….

……………………………………………………………………………………………………….

……………………………………………………………………………………………………….

3 Explain why covalent compound do not conduct electricity in all states.

……………………………………………………………………………………………………….

……………………………………………………………………………………………………….

4 Explain why ionic compounds have higher melting points than covalent compounds.

……………………………………………………………………………………………………….

……………………………………………………………………………………………………….

………………………………………………………………………………………………………

5 Example of covalent compounds ;-

............................................................................................................................................

Uses of covalent compounds as solvent in our daily lives:

………..................................................................................................................................

………………………………………………………………………………………………………

(Refer to page 90 - 91 - F4 Chemistry text book) Activity 8

1

Atom A Atom B

(a) Write the electron arrangement for atom A. ………………………………………………………………………………………………..


55

(b) A and B can form a compound

(i) What type of bond holds atom A and B together ?

…………………………………………………………………………………………..

(ii) What will happen to atom A during the formation of the compound with atom B? …………………………………………………………………………………………..

(iii) Draw the electron arrangement of the compound formed in (b)(ii).

(iv) State one physical property of the compound formed. …………………………………………………………………………………………..

(c) Carbon atom, C, with an electron arrangement of 2.4 can combine with atom B to form a compound.

(i) What is the molecular formula of the compound formed? ………………………………………………………………………………………….

(ii) If the relative atomic masses of carbon is 12 and B is 32, what is the relative molecular mass of the compound in c(i). …………………………………………………………………………………………..

Activity 9 1 The diagram below shows the proton number and the nucleon number for three atoms

of elements, E , G and W. The letters used do not represent the actual symbols of the elements.

12

E 6

23

G 11

35

W 17


56

(a) Construct a table to show the information of the three elements in terms of

The number of protons

The number of neutrons

The electron arrangement

The number of valence electrons [4 marks] Answer:

(b) The reaction between atoms of element G and W forms an ionic compound whereas the

reaction between atoms of element E and W forms a covalent compound. Explain how these ionic and covalent compounds are formed.

Answer: (Ionic compound)

Answer: (Covalent compound)


57

( c) The ionic compound formed from the reaction between atoms of element G and W does conduct electricity when it is in solid state but can conduct electricity when dissolved in water.

Describe an experiment to investigate this property. [8 marks] Answer:


58

CHAPTER 6: ELECTROCHEMISTRY

A ELECTROLYTES AND NON-ELECTROLYTES

Activity 1 1. State the meaning of electrolyte: An electrolyte is a substance that can conduct (a) ……………………in (b)……………… state or (c) ……………………. (d)……………………. and undergo (e)…………………… (f)…......................…. 2. A non-electrolyte is a substance that cannot conduct (g)……………………. either in (h)……………………. state or (i)……………………….. solution. 3. Ionic compounds in molten state or in aqueous solution are electrolytes

because these substances contain freely (j) …………………. ……………………. 4. Covalent compounds are non-electrolytes and these substances contain neutral (k) ................... and no freely (l) ……………………… …………… However, certain covalent compounds such as hydrogen chloride, ammonia and ethanoic acid when dissolved in water are electrolytes. This is because these compounds react with water to produce freely (m) ………………… …………..

Learning Outcomes: You should be able to:

state the meaning of electrolyte,

classify substances into electrolytes and non-electrolytes.

relate the presence of freely moving ions to electrical conductivity.


59

5. Ionic compounds in molten state or in aqueous solution are electrolytes while covalent compounds and ionic compounds in solid state are non-electrolytes. Metals are non-electrolytes but are good conductors of electricity. Classify the substances in the text box below into electrolyte and non-electrolyte

Electrolyte Non-electrolyte

Solid lead(II) chloride, molten aluminium oxide, lead(II) nitrate solution, solid sodium chloride, sodium chloride solution, magnesium, molten lead(II) chloride, glucose solution, glacial ethanoic acid, dilute ethanoic acid, molten naphthalene, ethanol, silver, tetrachloromethane, sodium hydroxide solution, aqueous ammonia


60

B ELECTROLYSIS OF MOLTEN COMPOUNDS Activity 2 1. What do you understand by the term electrolysis? Electrolysis is a process whereby compounds in (a) .................... or (b) ..……………

states are broken down (or decomposed) into their constituent (c) …………………… by

passing (d) ……………………. through them.

2. Anode is the electrode which is connected to the (e) …………………. terminal of a battery.

3. Cathode is the electrode which is connected to the (f) ……………………… terminal of a battery.

4. Carbon or platinum is chosen as electrodes as they are chemically inert or unreactive. 5. The diagram below shows the set-up of apparatus of electrolysis of molten lead(II)

bromide. Name the main apparatus and materials in the diagram.

Activity 3 1. Diagram 3.1 shows the relationship between the presence of freely moving ions and

electrical conductivity. The box below shows a list of statements that explain the why ionic compound in solid state do not conduct electricity but will conduct electricity in aqueous solution. The statements are arranged in random order. Choose the correct statement from the box below and write it into the correct text box in Diagram 3.1.

During electrolysis cations are attracted to the cathode and anions are attracted to the anode.

Solid sodium chloride contains sodium ions and chloride ions which are in fixed position and not freely moving.

In solid state, sodium ions and chloride ions are strongly attracted by electrostatic forces in a lattice.

Electric circuit is complete due to the flow of electrons along the connecting wires and movement of ions in the solution.

If the electrodes are placed further apart, the ammeter reading will decrease because there will be an increase in internal resistance.

Aqueous sodium chloride contains freely moving ions to conduct electricity.

Learning Outcomes: You should be able to:

describe electrolysis,

describe electrolytic cell,

identify cations and anions in a molten compound,

describe evidence for the existence of ions held in a lattice in solid state but move freely in molten state,

describe electrolysis of a molten compound,

write half-equations for the discharge of ions at anode and cathode,

predict products of the electrolysis of molten compounds.


61

Diagram 3.1


62

2. Colour all the cations red and the anions blue in solid sodium chloride and in the electrolyte in Diagram 3.1 above.

Activity 4 1. Given below is a list of ionic compounds in molten state. Identify the cation and anion in

each electrolyte.

Electrolyte (Molten)

Cation Anion Name Formula Name Formula

Sodium chloride

Lead(II) oxide

Potassium bromide

2. Given below is a list of electrolytes and products discharged at both electrodes. Based

on the given substance discharged at the electrode, write a half equation to represent the reaction occurring at the electrode.

Electrolyte

(molten)

Substance discharged at the electrodes and the half equation

Anode Cathode

(i) Aluminium oxide

Oxygen gas Half equation: ……………………………………..

Aluminium Half equation: ……………………………………..

(ii) Potassium iodide

Iodine Half equation: ……………………………………

Potassium Half equation: ……………………………..………

(iii) Sodium chloride

Chlorine gas Half equation: …………………………………..…

Sodium Half equation: ………….…………………………

(iv) Zinc bromide

Bromine gas Half equation: ………………………………..……

Zinc Half equation: …………….………………………


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Activity 5 1 The flow chart below is used to predict the products formed at the electrodes during the

electrolysis of molten lead(II) bromide. 2. In the spaces below, draw a similar flow chart (as in question 1) to predict the products

formed at the electrodes from the electrolysis of molten zinc chloride, ZnCl2.

d

Molten lead(II) bromide

a

g f

e

c b

Consists of

(Ions that are present)

( Movement of ions)

To anode

To cathode

(Half equation)

At Anode

At Cathode

(Products formed)

At Anode

At Cathode


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C ELECTROLYSIS OF AQUEOUS SOLUTION Learning Outcomes: You should be able to: identify cations and anions in an aqueous solution, describe the electrolysis of an aqueous solution, explain using examples factors affecting electrolysis of an aqueous solution, write half equations for the discharge of ions at the anode and the cathode, predict the products of electrolysis of aqueous solutions.

Activity 6 1. State three factors that may influence the selective discharge of ions during the

electrolysis of an aqueous solution.

(a) ………………………………………………………………………………………………… (b) ………………………………………………………………………………………………….

(c) ………………………………………………………………………………………………….

2. In an aqueous solution of sodium chloride, apart from sodium ions, Na+ and chloride

ions, Cl-- , ………………………ions, ……… and ……………………………….ions, …… from the slight dissociation of water are also present.

3. List the electrochemistry series (cations and anions) in order of increasing ease of discharge.

Ease of discharge increases Cation: ……………………………………………………………………………………………. Anion: ……………………………………………………………………………………………..

4. The following statements refer to the factors that affect the electrolysis of an aqueous solution. Fill in the blanks.

(a) The ions that are placed …………………… in the electrochemical series will be

……………….. discharged. (b) If the concentration of a particular ion is …………………, the ion is ……………..

………………………

(c) In the electrolysis of copper(II) sulphate, CuSO4 ……………..using copper electrodes, no ions are discharged at the anode. Instead, the copper anode …………………… and ………………………… in the electrolyte.


65

Activity 7 The diagram below shows the set-up of apparatus of an electrolytic cell containing concentrated copper(II) sulphate solution. Two test tubes filled with copper(II) sulphate solution were placed over the electrodes J and K to collect any gas evolved. The switch is then turned on so that electrolysis of copper(II) sulphate solution can occur.

(a) Identify the cations and the anions present in the aqueous solution.

Cations: …………………………………….. Anions: …………………..………………………

(b) Identify which electrode ( J or K ) is the anode and the cathode: Anode ……………………………….. Cathode ……………………………………….

(c) (i) Which ion is selectively discharge at the anode? ……………………………………… (ii) Give a reason for your answer in (c) (i). ………………………………………………….. …..……………………………………………………………………………………………. (iii) What do you observe at the anode? ……………………………………………………… (iv) Give one test to confirm the gas released at K. …………………………………………

………………………………………………………………………………………………. (v) Write a half equation to represent the discharge of ions at anode.

…….……………………………………………………………………………………………

(d) (i) Which ion is selectively discharge at the cathode? …………………………………………

(ii) Give a reason for your answer in (d) (i) ………………………………………………….

…………………………………………………………………………………………………

(iii) Which do you observe at the cathode? …………………………………………………

(iv) Write a half equation to represent the discharge of ions at the cathode. …………………………………. …………………………………………………………. (e) What do you observe about the copper(II) sulphate solution? …………………………………………………………………………………………………..


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Activity 8 (a) The table below shows two electrolytic cells with electrolytes of different concentration. You are required to answer each section by writing your answer in the spaces provided.

1. In the diagrams, label the cathode

with the symbol “” and the anode with the symbol “+”.

2. Show the direction of the flow of the

electrons with arrowheads, “ > “

3. Write the formula of all ions in the

electrolyte.

4. a. Write the formula of ions which are

attracted to the cathode. b. Underline the formula of ion which

is selectively discharged. c. State the factor that affect the selective discharged of ion

5. Write the half equation to represent

the reaction at the cathode.

6. What will you observe at the

cathode?

7. a. Write the formula of ions which are

attracted to the anode. b. Underline the formula of ion which

is selectively discharged. c. State the factor that affect the selective discharged of ion


the reaction at the anode.

9. What will you observe at the anode?


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(b) The table below shows two electrolytic cells with the same electrolytes with different electrodes. You are required to answer each section by writing your answer in the spaces provided.

1. In the diagrams, label the cathode

with the symbol “” and the anode with the symbol “+”.

2. Show the direction of the flow of the

electrons with arrowheads, “ > “

3. Write the formula of all ions in the

electrolyte.

4. a. Write the formula of ions which

are attracted to the cathode. b. Underline the formula of ion which is selectively discharged.

c. State the factor that affect the selective discharged of ion


the reaction at the cathode.


cathode?

7. a. Write the formula of ions which

are attracted to the anode. b. Underline the formula of ion which is selectively discharged.

c. State the factor that affect the selective discharged of ion


the reaction at the anode.


anode?

10 What do you observe about the copper(II) sulphate solution?


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Explain

D ELECTROLYSIS IN INDUSTRIES Learning Outcomes: You should be able to: state uses of electrolysis in industries, explain the extraction, purification and electroplating of metals involving electrolysis in industries, write chemical equations to represent the electrolysis process in industries, justify uses of electrolysis in industries,

describe the problem of pollution from electrolysis in industry.

Activity 9

1. Fill in the blanks.

The application of electrolysis in industries are (a) ……………………………………….

(b) ……………………………………………… and (c) ……………………………………

In the extraction of aluminium from its ore, (d) ….……………… electrodes are used

and (e) ……..………………. is added to aluminium oxide to lower its melting point.

In purification of metals, the pure metal is made the (f) ….………………. and the

impure metal is made the (g) ……….……………. The electrolyte used is an aqueous

salt solution of the metal ions.

In electroplating of metals, the (h) ………..…………….is made the anode and the (i)

……………… to be (j) ..…..………………….. is made the cathode. The electrolyte

used is an aqueous salt solution of the electroplating metal.

The purposes of electroplating metals are to make the electroplated object more (k)

………………………………….. and (l) …………………..……………… to corrosion.

2. Below are shown the three uses of electrolysis in industries. Fill in the blanks.

Extraction of aluminium from bauxites

Purification of copper from impure mined copper

Electroplating of iron spoon with silver

1. Substance used as cathode and anode

Cathode: Anode:

Cathode: Anode:

Cathode: Anode

2. Electrolyte used


69

3. Half equation representing the process.

Cathode: Anode:

Cathode: Anode:

Cathode: Anode:

E VOLTAIC CELLS Activity 10

1. A simple voltaic cell can be constructed by immersing two ……………………….

metals in an ………………………. connected by ………………… 2. In a voltaic cell, ……………………… energy is converted to ……………………. energy.

3. THE ELECTROCHEMICAL SERIES is an arrangement of metals based on the tendency of each metal atom to donate electrons. Complete the table below.

Electrochemical series of metals Cation formed and number of

electron(s) released during the process

K K+ + e

Al Al3+ + 3e

* Note: Hydrogen is not a metal, but it is included in the Electrochemical Series

Learning outcomes: You should be able to:

describe the structure of a simple voltaic cell and Daniell cell,

explain the production of electricity from a simple voltaic cell,

explain the reactions in a simple voltaic cell and Daniell cell,

compare and contrast the advantages and disadvantages of various voltaic cells,

describe the differences between electrolytic and voltaic cells.

Tendency of metal atoms to donate electrons to form ions increases


70

4. The diagram below shows an example of a simple voltaic cell.

In the text box below are sentences explaining the production of electricity from a simple voltaic cell. The sentences are listed in random order. You are required to arrange these sentences in the best possible order so as to give a clear description of the reactions occurring in a simple voltaic cell.

[If you have any problem, you can refer to page 104 of the text book for guidance.] Answer: (a) ………………………………………………………………………………………………………

……………………………………………………………………………………………………… (b) ………………………………………………………………………………………………………

……………………………………………………………………………………………………… (c) ……………………………………………………………………………………………………..

………………………………………………………………………………………………………

An example of a simple voltaic cell is a magnesium strip and a copper strip immersed in dilute sodium chloride solution.

The electrons then flows from the magnesium ribbon to the copper plate through the wire and this results in the flow of electrical current.

Hence magnesium atom releases electrons more easily than a copper atom and the magnesium act as the negative terminal of the cell.

The overall equation for the reaction is given as follows. Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)

Magnesium is placed higher than copper in the electrochemical series.

At the negative terminal, each magnesium atom releases two electrons and the Mg2+ formed moved into the solution.

Mg(s) Mg2+ (aq) + 2e

At the positive terminal which is the copper plate, the electrons are accepted by the H+ ions in sodium chloride solution.

2H+ + 2e H2 (g)


71

(d) ………………………………………………………………………………………………………

………………………………………………………………………………………………………

(e) ………………………………………………………………………………………………………

……………………………………………………………………………………………………… (f) ……………………………………………………………………………………………………… …………………………………………………………………………………………………………… (g) ………………………………………………………………………………………………………. …………………………………………………………………………………………………….. Activity 11 1. (a) Draw and label the set-up of apparatus of a Daniell cell consisting of a salt bridge.

(b) (i) Which metal in the Daniell cell is the negative terminal?

………………………………………………………………………………………………… (ii) Give reason for your answer in (b)(i).

………………………………………………………………………………………………… (iii) Write a half equation to represent the reaction occurring at the negative terminal.

………………………………………………………………………………………………….

(c) Write a half equation to represent reaction at positive terminal.

…………………………………………………………………………………………………… (d) What do you observe at (i) negative terminal ………………………………………………………………………………………………….. (ii) positive terminal ………………………………………………………………………………………………….. (e) State two functions of the salt bridge.

1. ……………..……………………………………………………………………………… 2. ………………………………………………………………………………………………


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2. The table below listed are five types of voltaic cells commonly used in our daily lives. Each voltaic cell has its advantages and disadvantages. Complete the table by stating the advantages and disadvantages of each voltaic cell.

Voltaic cell Advantage

s Disadvantage

s 1. Lead-acid accumulator

2. Dry cell

3. Mercury cell

4. Alkaline cell

5. Nickel-cadmium cell

Activity 12 What are the differences between an electrolytic cell and a voltaic cell? Table 12.1 are statements showing differences between an electrolytic cell and a voltaic cell. Complete Table 12.2 by choosing the correct matching statements.

It does not require a source of electric current

It requires a source of electric current

The electrical energy causes chemical reactions to occur at the electrodes. Electrical energy chemical energy

The chemical reaction that occur at the electrodes produces electric current. Chemical energy electrical energy

The electrodes must be of two different metals

The electrodes may be of the same material such as carbon

Electrons flow from the positive electrode (anode) to the negative electrode (cathode).

Electrons flow from the more electropositive metal (negative terminal) to the less electropositive metal (positive terminal).

Ions receive electrons at the positive terminal. (Reduction)

Ions donate electrons at the positive terminal. (Oxidation)

Ions receive electrons at the negative terminal. (Reduction)

Ions donate electrons at the negative terminal. (Oxidation)

Table 12.1


73

DIFFERENCES Electrolytic cell Aspect Chemical cell

Source of

electric current

Conversion of energy

Type of electrodes

Direction of flow of electrons

Type of reaction at positive terminal

Type of reaction at negative terminal

Chemistry Form 4 [email protected]

74

F THE ELECTROCHEMICAL SERIES Learning Outcomes: You should be able to: describe the principles used in constructing the electrochemical series, construct the electrochemical series, explain the importance of electrochemical series, predict the ability of a metal to displace another metal from its salt solution, write the chemical equations for metal displacement reactions.

Activity 13 Three experiments were conducted to investigate the potential differences between three pairs of metals in a voltaic cell. An electrochemical series for four metals P, Q, S and T is then constructed based on the potential difference obtained. Three pair of metals used as electrodes in different voltaic cells are: P and Q, Q and S and S and T. All the metals are cleaned with sandpaper before used. 50 cm3 of 1.0 mol dm-3 sodium nitrate solution is poured into a beaker as electrolyte. Experiment I The electrodes P and Q are immersed into the solution. The two electrodes are connected to a voltmeter using copper wires. Electrode Q is the positive terminal. The voltmeter reading is recorded. Experiment II The electrodes Q and S are immersed into the solution. The two electrodes are connected to a voltmeter using copper wire. Electrode Q is the positive terminal. The voltmeter reading is recorded. Experiment III The electrodes S and T are immersed into the solution. The two electrodes are connected to a voltmeter using copper wire. Electrode T is the positive terminal. The voltmeter reading is recorded. Based on Experiment I, II and III, answer the questions below. (a) Record the voltmeter reading of each experiment in the spaces provided. Experiment 1 Experiment II Experiment III Voltmeter reading: …………. Voltmeter reading: …………… Voltmeter reading: ….…….. (b) Construct a table to record the data from the above experiments.

0

1 2 3 4 5

6

V

0

1 2 3 4 5

6

V

0

1 2 3 4 5

6

V


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(c) List the apparatus and materials that you will need to carry out this experiment.

Apparatus: ……………………………………………………………………………………….

…………………………………………………………………………………………………….. Materials: ……………………………………………………………………………………….. ……………………………………………………………………………………………………

(d) State all the variables: 1. Manipulated variable: ………………………………………………………………….. 2. Responding variable: …………………………………………………………………….

3. Controlled variable: ………………………………………………………………………

(e) State the hypothesis: ……………………………………………………………………….

…………………………………………………………………………………………………..

(f) Based on the information obtained in Experiment I, what can you infer about metal P and Q?

…………………………………………………………………………………………………… (g) Write a half equation for the reaction occurring in negative of Experiment I, assuming the cation

has a +2 charge. ……………………………………………………………………………………………………….

(h) Arrange the metals P, Q, S and T in descending order of their tendency to donate electrons. ...……………………………………………………………………………………………………

(i) Another voltaic cell is set-up using metals T and Q as electrodes. Predict the potential difference produced in the cell.

……………………………………………………………………………………………………..

(j) Given that metal X is placed between metal S and metal Q in the electrochemical series, can metal X displace metal S from its salt solution? Give an explanation for your answer

………………………………………………………………………………………………………. ……………………………………………………………………………………………………….

(k) Given that copper is more electropositive than metal T, a displacement reaction will occur when copper is immersed into a salt solution of metal T, TNO3.

Write the chemical equation for this reaction. ……………………………………………………………………………………………………….

(l) State three important uses of the electrochemical series

…………………………………………………………………..………………………………… ……………………………………………………………………………………………………. ……………………………………………………………………………………………………...

Activity 14 : Displacement Reactions 1. Metals placed ……… up in the Electrochemical series ar able to displace metals placed …….

them from their salt solutions.


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magnesium zinc Lead(II) copper(II) nitrate nitrate nitrate nitrate solution solution solution solution magnesium strips

K > Na > Ca > Mg > Al > Zn > Fe Sn > Pb > H > Cu > Ag

2. Example : Ewacrion between xinc and copper(II) sulphate solution Chemical equation : Zn + CuSO4 ……………… + ………….

…….. is place higher position than ……… in electrochemical series, , …………….. is more

electropositive than …………….., hence ………….. can displace ……………from

………………………. solutions.

Copper(II) sulphate solution Zinc atom releases 2 electrons to form zinc ion, Zn2+ : Half equation : …………………. Copper(II) ion Cu2+ receives 2 electrons to form copper atom : Half equation :…………….. Ionic equation : ………………………………………………………. 3. Experiment : To construct the electrochemical series using the principle of displacement of metals (Displacement Reaction) Problem Statement : ……………………………………………………………………………… ……………………………………………………………………………… Hypothesis: The greater the number of metals that can be displaced by a metal from their solutions, the higher is its position in the electrochemical series. Variables: Manipulated variable : ……………………………………

Responding variable : ………………………………

Fixed variables : ……………………………………

List of Materials and Apparatus: Test tubes, sand paper, 1 mol dm-3 zinc nitrate solution, 1 mol dm-3 lead (II) nitrate solution, 1 mol dm-3 copper (II) nitrate solution, magnesium strips, zinc strips, lead strips and copper strips.

zinc

Observation:

1. Zinc ………………….

2. ……………. Solid formed

3. Intensity of blue solution …………….


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Procedure:

1. Pour 5 cm3 of magnesium nitrate solution, zinc nitrate solution, lead (II) nitrate solution, and copper (II) nitrate solution into four separate test tubes.

2. For each test tube, place a strip of magnesium into each solution. 3. Record all the observations. 4. Repeat steps 1 to 3 using strips of zinc, lead and copper to replace the magnesium strip. For

each repetition, use a fresh salt solution. Data and Observation

Magnesium nitrate solution

Zinc nitrate solution

Lead (II) nitrate solution

Copper (II) nitrate solution

Magnesium

Zinc

Lead

Copper

a. Which metal can displace the most number of other metals from their solutions? ……………………………………………………………………………………………. b. Write the half equation of the reaction occurred in magnesium strip for this experiment. …………………………………………………………………………………………………… c. Which metal can be displaced by all other metals in the experiment? ………………………………………………………………………………………………….. d. Arrange the metals in descending order based on the electrochemical series and the number of metals displaced by it. …………………………………………………………………………………………………… e. Write the ionic equations to show all displacement reactions by zinc ………………………………………………………………………………………………….. …………………………………………………………………………………………………..

Metal

strip

Salt

solution

Complete this table please!


78

CHAPTER 7: ACIDS AND BASES A ACIDS AND BASES

Learning Outcomes


State the meaning of acid, base and alkali

State uses of acids, bases and alkalis in daily life

Explain the role of water of water in the formation of hydrogen ions to show the properties of acids

Explain the role of water in the formation of hydroxide ions to show the properties of alkalis

Describe chemical properties of acids and alkalis

Activity 1 Meaning of acid ,base and alkali Fill in the blanks with the correct words: 1

An acid is a chemical substance which ionises in………………. to produce ………….ions.

The hydrogen ion combines with a water molecule, H2O to form a ……………………….. ion,

H3O+.

2 Acid can be classified as a……………………….acid or a ………………….acid based on its

basicity.

3 Basicity is the number of ionisable …………………. atoms per molecule of an acid.

4 A base is a substance that reacts with an acid to form a …………... and ……………. only.

Bases include metal hydroxides and metal oxides.

5 Give the names of acids, their formulae and its basicity:

Name of acid Formula of acid Basicity

(i) Hydrochloric acid

(ii) H2SO4

(iii) HNO3

(iv) Ethanoic acid

6 Complete the ionization of acids below :

a)

b)

c)

d)

HCl (aq) → . . .…….(aq) + Cl- (aq)

……… (aq) → H+ (aq) + NO3 - (aq)

H2SO4 (aq) → ……. …… + ………..

CH3COOH (aq) …………. + CH3COO- (aq)

(Refer to page 117 -118 - F4 Chemistry textbook)


79

Activity 2 : Fill in the blanks with correct words: 1 A base is a substance that reacts with an acid to form ...................................................only.

An alkali is a water-soluble base which ionises in water to produce………..………..........,OH-

2 Complete the table below

Name of alkali Formula of alkali

(i) Sodium hydroxide

(ii) KOH

(iii) Ammonia aqueous

3 Complete the ionization of following alkalis :

NaOH (aq) → …….(aq) + OH- (aq)

……… (aq) + H2O (l) NH4+ (aq) + …………. (aq)

4 Uses of acids, bases and alkalis

(a) To use as …………………. … reagent

Example: sodium hydroxide solution, sulphuric acid, hydrochloric acid

(b) To ..................................................................................................... ……………

Example: Ethanoic acid (vinegar), benzoic acid

(c)To make various ………………..

Example: Magnesium oxide antacid medicine, Ascorbic acid vitamin C

(d)To produce …………., detergent and ………………

Example: sodium hydroxide to make soap and detergent

Magnesium hydroxide added to tooth-paste

(e)To manufacture dyes,……………………. and drugs

Example: methylamine

(f)Used in rocket fuel

Example:…………………………………………………………………………………………

(Refer to page 118 - F4 Chemistry textbook)

Activity 3: Role of water and its properties of acids and alkalis Fill in the blanks with correct words. 1 An acid only shows its …………..…… properties when dissolve in ……………………….

2 In the presence of water, the acid ionises to form ……………………………………..ions .


80

3 Without water, an acid still exists as…………………and there are no ................ ions present.

Glacial Ethanoic acid, CH3COOH is an example of acidic substance.

Complete the following table to show the role of water in acidic properties.

Condition of ethanoic acid Effect on the blue litmus paper

Inference

Glacial ethanoic acid,

Ethanoic acid in water

Ethanoic acid in dry propanone ( organic solvent )

5 An alkali only shows its ……………….. properties when dissolve in ………………………..

6 In the presence of water, the base dissociates to produce …………………………... ions that

are responsible for the ………………… properties

7 Ammonia,NH3 is an example of alkali.

Complete the table below to show the role of water in alkaline properties.

Condition of ammonia Effect on the red litmus paper Inference Dry

Aqueous (dissolved in water)

Dissolved in propanone, (organic solvent)

(Refer to page 118 -119 F4 Chemistry textbook and page 84 – 86 F4 Chemistry Practical Book )

Activity 4: Properties of acids and alkalis Tick (√ )

Statement True or False

1. All acids are dangerous

2. All alkalis are dangerous

3. Acids taste sweet

4. Alkalis taste bitter

5. Acids taste sour

6. Most acids can burn skin

7. Alkalis feel soapy

8. Acids produce H+ ions in solution

9. Acids produce OH- ions in solution

10. Acids can corrode

11. Acids have a pH above 7


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12. Acids have a pH below 7

13. Alkalis turn moist red litmus paper blue

14. Acids turn moist red litmus paper blue

(Refer to page 119 - F4 Chemistry textbook and page 87 – 90 - F4 Chemistry Practical Book )

Activity 5 : Chemical properties of acids

1

Acids react with bases to form salts and water only. Bases are metal oxides or metal hydroxide.

Write the chemical equation for the reaction between sulphuric acid and copper(II) oxide.

……………………………………………………………………………………………………...

2

Acids react with reactive metals to produce salts and hydrogen gas.

Write the chemical equation for the reaction between hydrochloric acid and zinc .

……………………………………………………………………………………………………...

3

Acids react with metal carbonates to produce salts, water and carbon dioxide gas.


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Write the chemical equation for the reaction between nitric acid and calcium carbonate.

……………………………………………………………………………………………………...

Activity 6 : Chemical properties of alkalis 1

Alkalis react with acids to form salts and water only

Write the chemical equation for the reaction between sodium hydroxide and benzoic acid.

……………………………………………………………………………………………………...

2

When a mixture of an alkali and an ammonium salt is heated, ammonia gas is liberated.

Write the chemical equation for the reaction between sodium hydroxide and ammonium chloride .

……………………………………………………………………………………………………...

3

Alkalis react with most metal ion solutions ( cations ) to produce the insoluble metal hydroxides or precipitate - (precipitation reaction)


83

Write the chemical equation for the reaction between sodium hydroxide and iron(II) sulphate.

……………………………………………………………………………………………………...

(Refer to page 120 - F4 Chemistry textbook and page 91- F4 Chemistry Practical Book )

B: The strength of acids and alkalis Learning Outcomes: You should be able to:

State the use of a pH scale

Relate pH value with acidic or alkaline properties of a substance

Relate concentration of hydrogen ions with pH value

Relate concentration of hydroxide ions with pH value

Relate strong or weak acid with degree of dissociation

Conceptualise qualitatively strong and weak acids

Conceptualise qualitatively strong and weak alkalis

Activity 7: The pH scale Fill in the blanks with correct words 1 The pH scale ( 0 to 14 ) , is used to indicate the degree of ……………… or ……………. of a

solution. 2 pH value less than 7 , indicates an…………………………solution

pH value equal to 7 , indicates a…………………………solution

pH value more than 7 , indicates an…………………………solution

3 pH value can be determined by using………….................,pH paper or…………………indicator.

4

Acids Neutral Alkalis 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

Acidity ......................................(increase or decrease ) Alkalinity.........................(increase or decrease )

(Refer to page 121 – Chemistry text book)


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Activity 8 : Strong and weak acid 1. Complete the flowchart below to understand the concept of strong acid and weak acid Strong acid Weak acid

Degree of dissociation Ionization in water Concentration of ions

pH value

Examples Activity 9: Strong and weak alkali 1. Complete the flowchart below to understand the concept of strong alkali and weak alkali. Strong alkali Weak alkali

Degree of dissociation Ionization in water Concentration of ions

pH value

Examples


85

C: CONCENTRATIONS OF ACIDS AND ALKALIS Learning Outcomes You should be able to:

State the meaning of concentration

State the meaning of molarity

State the relationship between the number of moles with molarity and volume of a solution

Describe the methods for preparing standard solutions

Describe the preparation of a solution with a specified concentration using dilution method

Relate pH values with the molarity of acids and alkalis

Solve numerical involving molarity of acids and alkalis

Activity 10 : Concentration of acids and alkalis Fill in the blanks with the correct answers. 1

2

The concentration of a solution refers to the quantity of solute in 1dm3 of solution Concentration can be defined in two ways :-

(a)

Concentration in g dm-3 =

(b) Concentration in mol dm-3 =

(Concentration in mol dm-3 is also known as molarity or molar concentration (M ) )

2

The two units of concentration can be inter-converted:

Work this out. 3 5.0 g of copper(II) sulphate is dissolved in water to form 500 cm3 solution. Calculate the

concentration of copper(II) sulphate solution in g dm-3 ?

[Answer: 10.0 g dm-3]

4

What is the mass of sodium carbonate required to dissolve in water to prepare 200 cm3 solution that contains 50 g dm-3 ?

[Answer: 10 g]

(Refer to page 123 -124 - F4 Chemistry textbook )

Concentration in ...................

Concentration in............................

X Molar mass

÷ Molar mass


86

5 4.0 g of sodium carbonate powder, Na2CO3 , is dissolved in water and made up to 250 cm3. What is the molarity of the sodium carbonate solution? [Relative atomic mass: C, 12; O, 16; Na, 23]

[Answer: 0.15 mol dm

-3]

Activity 11

The number of moles of solute, n in a given volume of solution V and the molarity of M can be calculated by using the formula :

n = Number of moles of solute M = Molarity of solution (mol dm-3) V = Volume of solution (dm3)

If the volume is in cm3 – convert the volume of solution from cm3 to dm3

5 Calculate the number of moles of ammonia in 150 cm3 of 2 mol dm-3 aqueous ammonia.

[Answer: 0.3 mol] 6 A student pipetted 20.0 cm3 of potassium hydroxide , KOH solution into a conical flask. The

concentration of the alkali was 1.5 mol dm-3 . Calculate the number of moles of potassium , KOH in the flask.

[Answer: 0.03 mol dm-3

]

7 Calculate the number of moles of hydrogen ions present in 200 cm3 of 0.5 mol dm-3 sulphuric acid, H2SO4.

[Answer: 0.2 mol of H+ ions]

Activity 12 : Preparation of Standard solutions 1

What is a standard solution? ……………………………………………………………………………………………………

n = MV

n = M x V

1000 n = MV

1000 or


87

2

Preparation of standard solutions by Weighing method (mass of solute) :-

Step 1 : Calculate the mass of solute needed . mass = n X molar mass

= MV X molar mass 1000 Example: To prepare 100 cm3 of 2.0 mol dm-3 sodium hydroxide solution. Calculate the mass of NaOH needed. [Relative atomic mass: Na, 23 ; O, 16 ; H, 1] mass = n X molar mass

= MV X molar mass 1000 = 2.0 X 100 X 40 = 8 g 1000 Try this: (a) To prepare 250 cm3 of 1.0 mol dm-3 sodium carbonate solution. Calculate the mass of Na2CO3 needed. [Relative atomic mass: Na, 23 ; O, 16 ; C, 12]

[Answer : 26.50g]

(b) 0.25 mol dm-3 solution of sodium hydroxide was prepared by dissolving x g of sodium hydroxide in 750 cm3 of water. What is the value of x ? [Relative atomic mass: Na, 23 ; O, 16 ; H, 1]

[Answer : 7.5 g] (Refer to page 126 - F4 Chemistry textbook )

Step 2 : Match the descriptions / procedures with the correct diagram below.

Wash and rinse the weighing bottle or small beaker and filter funnel to ensure no solute remains in any of the apparatus used.

Transfer the dissolved solute into a suitable volumetric flask.

(a)

n = MV

1000

n = mass

molar mass


88

Add water slowly by using a dropper to bring the level of the solution to the calibration mark.

The volumetric flask is closed tightly and inverted several times to get a uniform or homogenous solution.

(b)

Calculate the mass of solute needed.

Weigh out the exact mass of solute needed in a weighing bottle.

Dissolved the solute in a small amount of distilled water.

(c)

Add more water carefully to the volumetric flask and swirl gently.

Shake well to ensure thorough mixing.

(d)

Activity 13 : Preparation of Standard solutions by Dilution method

1 Dilution method

Step 1 : Calculate the volume of stock solution required by using the equation:-

M1 = molarity of solution before dilution V1 = volume of solution before dilution M2 = molarity of solution after dilution V2 = volume of solution after dilution

Example: 50 cm3 of 0.1 mol dm-3 sodium hydroxide, NaOH solution from 2.0 mol dm-3

sodium hydroxide,NaOH solution

Before dilution After dilution M1 V1 M2 V2

2.0 mol dm-3 ? 0.1 mol dm-3 50 cm3

2.0 x V1 = 0.1 x 50 V1 = 0.1 x 50 = 2.5 cm3 2.0

Try this: 100 cm3 of 0.5 mol dm-3 potassium manganate(VII) ,KMnO4 solution is prepared from 1.0 mol dm-3 potassium manganate(VII) ,KMnO4 solution. Calculate the volume of the solution

M1V1 = M2V2

When using the equation M1V1 = M2V2 , make sure that both V1 and V2 are of the same unit.


89

[Answer : 50 cm3]

Step 2 Match the diagram with the correct descriptions below.

(a)

(b)

(c)

(d)

Add water slowly by using a dropper to bring the level of the solution to the calibration mark.

The volumetric flask is closed tightly and inverted several times to get a uniform or homogenous solution.

Transfer the stock solution to a suitable volumetric flask.

Calculate the volume of stock solution required.

Use a pipette to draw up the required volume of stock solution.

Activity 14 : The pH values and molarity of acids and alkalis Fill in the blanks with correct words . Use words given in the box.

Increases decreases concentration hydrogen dissociation higher hydroxide alkali

1 The pH value of an acid or an alkali depends on three factors :

(a) degree of…………………………………………………………………………………….

(b) molarity or ………………………………………………………………………………..

(c) ………………….. of the acid or …………………………………………………………..

2 The lower the pH value, the ……………….. the concentration of ……………………ions.

3 The higher the pH value, the …………….. the concentration of …………………... ions.

4 As the molarity of an acid increases , the pH value of the acid ………………………….

The pH value of an alkali increases when the molarity of the alkali ………………..…….

(Refer to page 128 - F4 Chemistry textbook )


90

D : NEUTRALISATION

Learning Outcomes

You should be able:-

Explain the meaning of neutralisation.

Write equations for neutralisation

Explain the applications of neutralization in daily life

Describe the titration process of acid-base

Determine the end-point of titration during neutralization

Solve numerical problems involving neutralisation

Activity 15 : Neutralisation

1 What is the meaning of neutralisation? ………………………………………………………………………………………………..

2 What are the only products of neutralisation? ………………………………………………………………………………………………….

3 Write a balanced chemical equation for the neutralisation of the following reactions:-

(a) nitric acid and barium hydroxide ……………………………………………………………………………………………

(b) sulphuric acid and sodium hydroxide ……………………………………………………………………………………………

(c) phosphoric acid and calcium hydroxide ……………………………………………………………………………………………

(d) ethanoic acid and potassium hydroxide ……………………………………………………………………………………………

4 Complete the flow chart below:-


91

(Refer to page 128 – 129 - F4 Chemistry textbook )

Activity 16 : Acid – base Titration

1 What is a titration?

………………………………………………………………………………………………….

2 What is the function of an indicator?

…………………………………………………………………………………………………..

3 Complete the table below.

Indicator

Colour in solution Acid Neutral Alkali

Red litmus paper Blue litmus paper Phenolphthalein Methyl orange

4 Write out the procedure for carrying out an acid-base titration to determine the volume of nitric acid 0.5 mol dm-3 needed to neutralise 25 cm3 potassium hydroxide 0.5 mol dm-3 . Label the diagram.


92

(Refer to page 130 – F4 Chemistry textbook and Page 103 – F4 - Chemistry Practical Book)

Activity 17 : Numerical problems involving neutralisation

Useful equations in solving numerical problems involving neutralisation.:

n = no of moles

M = Molarity of solution

V = Volume of solution in dm3

Ma = molarity of acid

Mb = molarity of alkali

Va = volume of acid

Vb = volume of alkali

a and b = mole ratio of acid to alkali in a balanced chemical equation

Example:

In an experiment, 25 cm3 of sodium hydroxide solution, NaOH of unknown concentration required 26.50 cm3 of 1.0 mol dm-3 sulphuric acid, H2SO4 for complete reaction in titration. Calculate the molarity of sodium hydroxide.

Write out a balanced chemical equation:

H2SO4 + 2NaOH Na2SO4 + 2H2O

a = 1 mol b = 2 mol

MaVa 1 MbVb = 2 ,

1.0 X 26.50 = 1 Mb X 25.00 2

Mb = 2 X 1.0 X 26.50 = 2.12 mol dm-3 (Molarity of sodium hydroxide)

25.00

1 What is the volume of 1.5 mol dm-3 aqueous ammmonia required to completely neutralise 30.00 cm3 of 0.5 mol dm-3 sulphuric acid ?

2NH3 + H2SO4 (NH4 ) 2SO4

[Answer: 20 cm3]

MaVa = a MbVb b n = MV n = mass

molar mass

mass


93

2 Calculate the volume in cm3 2.0 mol dm-3 hydrochloric acid that is required to react

completely with 2.65 g of sodium carbonate. [Relative atomic mass: Na, 23 ; O, 16 ; C, 12]

[Answer: 25 cm3]

3 25 cm3 of sulphuric acid was neutralised with 18.0 cm3 of sodium hydroxide 1.0 mol

dm-3. Calculate (a) the number of moles of sulphuric acid that is used in this reaction. (b) the molarity of sulphuric acid

[Answer (a) 0.009 mol (b) 0.36 mol dm-3

]

4 24 cm3 of 0.1 mol dm-3 NaOH is required to completely neutralise 20.0 cm3 of sulphuric acid. Calculate the concentration of sulphuric acid in (a) mol dm-3 (b) g dm-3

[Answer (a) 0.06 mol dm-3

(b) 5.88 g dm-3

]


94

1

3

5 What is the molarity of phosphoric acid if 15 cm3 of the acid is neutralized by 38.5 cm3 of 0.15 mol dm-3 NaOH ?

[Answer (a) 0.218 mol dm-3

] Activity 18 1 A student carried out an experiment to determine the end-point for the titration of

25.0 cm3 of 1.0 mol dm-3 sodium hydroxide solution with hydrochloric acid. Phenolphthalein is used as the acid-base indicator. Table 1 shows the three titrations that were conducted and the magnification of the burette readings.

Titration No. I II III

Initial burette

reading

1

2

…………………..

13

14

……………….

26

27

………………..

Final burette

reading

13

14

……………………

26

27

……………………

38

39

…………………….

Table 1

hydrochloric acid

hydrochloric acid hydrochloric

acid hydrochloric acid

hydrochloric acid

hydrochloric acid


95

(a) Record the burette readings for the three titrations in the space provided in Table 1.

(b) Construct a table and record the initial burette reading, final burette reading and the

volume of acid used for each titration.

(c) Calculate the average volume of hydrochloric acid used in the experiment.

(d) Calculate the concentration of hydrochloric acid used in the experiment.

(e) If the experiment is repeated by replacing 1.0 mol dm-3 of hydrochloric acid with 1.0 mol dm-3 of sulphuric acid, predict the end-point of the titration.

………………………………………………………………………………………………….

(f) Acids can be classified into strong acid and weak acid. Classify the following acids into strong acids and weak acids.

Ethanoic acid, hydrochloric acid, sulphuric acid,

carbonic acid, nitric acid,

Strong acids Weak acids


96

(g) State the colour change of the phenolphthalein indicator at the end point of titration. …………………………………………………………………………………………………

(h) If phenolphthalein is replaced with methyl orange as the acid-base indicator, state the colour change. …………………………………………………………………………………………………

(i) Write a chemical equation for the neutralisation reaction between hydrochloric acid and sodium hydroxide. ………………………………………………………………………………………………….


97

CHAPTER 8 : SALTS

A. SYNTHESISING SALTS


state examples of salts used in daily life, explain the meaning of salt identify soluble and insoluble salts, describe the preparation of soluble salts, describe the purification of soluble salts by recrystallisation, list physical characteristics of crystals, describe the preparation of insoluble salts, write chemical and ionic equations for reactions used in the preparation

of salts, design an activity to prepare a specified salt, construct ionic equations through the continuous variation method solve problems involving calculation of quantities of reactants or products in stoichiometric reactions

Activity 1 : Meaning and uses of Salts

1. A salt is an ……………………..……… formed when the ……………….. ion, from an ……………

is replaced by a ……………. ion or an …………………..ion.

2. Example of salts : Complete the table below

Acid Formula of acid

Salt Formula Cation Anion

Sodium chloride NaCl Na+ Cl-

Potassium carbonate K2CO3

Copper(II) sulphate CuSO4

Ammonium nitrate NH4NO3

Magnesium nitrate Mg(NO3)2

Sodium ethanoate CH3COONa 3. Match the following salts with their uses. Salts Uses

Barium sulphate BaSO4 Fungicide Calsium sulphate CaSO4 Bleaching agent Iron sulphate FeSO4 Paint for yellow line on road Ammonium nitrate NH4NO3 X-ray ‘meals’ in hospital Copper(II) sulphate CuSO4 Nitrogenous fertilizer Sodium chloride NaCl Toothpaste Sodium hydrogen carbonate Iron tablets, for anaemia

patient

Sodium nitrite NaNO2 Baking powder Sodium hypochlorite NaOCl

Preserve food

Tin(II) fluoride SnF2 A flavouring Lead(II) chromate PbCrO4 Plaster of Paris for broken

bone


98

Activity 2 : Identify soluble and insoluble salt. Fill in the blanks with the correct word(s) 1. All ………………. , ………………., and …………………. salts are soluble in water. 2. All ………………. salts are soluble in water.

3. All sulphate salts are soluble in water except …………….. sulphate, …………. sulphate and

………… sulphate.

4. All chloride salts are soluble in water except …………….. chloride , …………. chloride and ………… chloride.

5. All carbonate salts are insoluble in water except …………….. carbonate, …………. carbonate

and ………… carbonate 6. State whether each of the following salt is soluble or insoluble in water

No Formula of Salt Solubility ( , X ) No Formula of Salt Solubility ( , X )

1 PbCO3 21 MgCO3 2 NaCl 22 KCl 3 CaSO4 23 (NH4)2SO4 4 AgNO3 24 Cu(NO3)2 5 K2CO3 25 SnCO3 6 FeCl3 26 CaCl2 7 Na2SO4 27 BaSO4 8 NH4NO3 28 KNO3 9 CuSO4 29 Ag2CO3 10 PbCl2 30 MgCl2 11 ZnCO3 31 ZnSO4 12 Ca(NO3)2 32 Ba(NO3)2 13 Na2CO3 33 FeCO3 14 AgCl 34 NH4Cl 15 PbSO4 35 Fe(NO3)3 16 Pb(NO3)2 36 MgSO4 17 (NH4)2CO3 37 BaCO3 18 HgCl2 38 ZnCl2 19 Na2SO4 39 FeSO4 20 NaNO3 40 Mg(NO3)2

Water

Air

SO4 2-

Cl- CO3

2-

Na+

K+

NH4+

NO3 - Ba

2+

Ca2+

Pb2+

Pb2+

Ag+

Hg+


99

Activity 3 : Write chemical and ionic equations for reactions used in the preparation of soluble salts 1. Complete these general equation for preparing soluble salts.

a. metal + acid …………… + …………………….

b. metal oxide (or metal hydroxide) + acid …………… + ……………………

c. alkali + acid …………… + …………………….

d. metal carbonate + acid …………… + ……………… + ………………………. 2. Using the general equations in question 1, complete the following chemical equation. It may also

be necessary to balance the equation.

a. Mg + H2SO4 …………… + …………………….

b. (i) CuO + HCl …………… + …………………… (ii) Zn(OH)2 + HNO3 ................................... + .........................

c. NaOH + HCl ………………….. + ………………………. d. MgCO3 + H2SO4 …………… + ……………… + ……………………….

3. Deduce the identity of the acid, metal, salt, or other product by filling in the missing details in this

table of preparation of soluble salt

Method of Preparation

Reactants Salt Formed Other Product

a) metal + acid

Magnesium + ………………

Magnesium chloride

Hydrogen

b) metal oxide + acid

Copper(II) oxide + sulphuric acid

………………………

…………….

c) metal carbonate +

acid

……………… + ………………

Zinc sulphate

Water + ……………. ……………..

d) metal hydroxide +

acid

……………… + ………………

Potassium nitrate

……………

f) alkali + acid

……………… + ………………

Sodium chloride

……………


100

4. Name the reactants which are needed to prepare the following soluble salts:

(a) Copper(II) sulphate : …………………………………………………………………………….. (b) Zinc chloride : …………………………………………………………………………….

(c) Potassium nitrate : ……………………………………………………………………………..

(d) Ammonium sulphate : ……………………………………………………………………………..

(e) Magnesium nitrate : ……………………………………………………………………………..

5. Rewrite each of the following chemical equation as ionic equation. Shown below is an example

where a chemical equation can be simplified into an ionic equation.

Example : Chemical equation : Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g) Zn + 2H+ + SO4

2- Zn2+ + SO42- + H2

(s) (aq) (aq) (aq) (aq) (g)

Ionic equation : Zn(s) + 2H+ (aq) Zn2+ (aq) + H2(g)

a. Chemical equation : Mg(s) + 2HCl(aq) MgCl2(aq) + H2O Ionic equation : ......................................................................................................................... b. Chemical equation : MgO(s) + 2HCl(aq) MgCl2(aq) + H2O Ionic equation : ........................................................................................................................ c. Chemical equation : NaOH (aq) + HNO3 (aq) NaNO3 (aq) + H2O (l)

Ionic equation : .......................................................................................................................

d. Chemical equation : CuCO3 (s) + H2SO4 (aq) CuSO4 (aq) + CO2 (g) + H2O (l)

Ionic equation : .......................................................................................................................


101

Activity 4 : Write out the procedure for the preparation of soluble salts of sodium, potassium and ammonium

Soluble salt Sodium Chloride, NaCl

Name two chemical substances to prepare the salt

1. …………..…………………………..

2. ………………………………………

Chemical equation

Procedure: (Diagram) Description

Describe the physical characteristics of the crystals that you obtained


102

Activity 5 : Write out the procedure for the preparation of soluble salts (not sodium, potassium or ammonium salt)

Soluble salt Copper(II) sulphate, CuSO4

Name two chemical substances to prepare the salt

1. …………..…………………………..

2. ………………………………………

Chemical equation


Describe the purification process of the crystals


103

Activity 6 : Write chemical and ionic equations for reactions used in the preparation of insoluble salts 1. Insoluble salts can be prepared by ………………….. method through ……………………………..

reaction. In this reaction, two different aqueous solution mutually exchange their …….. to form

………………………….

Soluble salt solution + Soluble salt solution Insoluble salt MX containing cation M+ containing anion X-

Chemical equation : AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3(aq)

Ionic equation : Ag+ (aq) + Cl- (aq) AgCl (s) 2. Preparation of insoluble salts

Example 1: Barium sulphate, 4BaSO

Solution 1: ……………………………………... Solution 2: ……….………………………………….

Chemical equation : ………………………………………………………………..……………………..

Ionic Equation : …………………………………………………………………………………………….

Observation : White precipitate formed Example 2: Copper(II) carbonate, CuCO3

Solution 1: …………………………………….. Solution 2: ……………………………………….

Chemical equation : …………………………………………………………………………………….

Ionic Equation : ………………………………………………………………………………………….

Observation : ……………….. precipitate formed

Example 3: Lead(II) chromate(VI), 4PbCrO

Solution 1: …………………………………….. Solution 2: …………………………………….

Chemical equation : ………………………………………………………………………………….

Ionic Equation : ………………………………………………………………………………………..

Observation : ……………….. precipitate formed


104

Activity 7 : Describe the preparation of insoluble salts

Insoluble salt Lead(II) iodide, PbI2

Name two chemical

substances to prepare the

salt

1. …………..…………………………..

2. ………………………………………

Chemical equation

Ionic equation



105

Activity 8 : Construct ionic equation for the formation of lead(II) chromate through the continuous variation method,

Figure 1

Figure 1 shows seven test tubes for the reaction between lead(II) nitrate Pb(NO3)2 0.5 mol dm-3 and potassium chromate(VI) K2CrO4 0.5 mol dm-3.

(a) Calculate the number of moles of lead(II) nitrate Pb(NO3)2 and potassium chromate(VI) K2CrO4

used in test tubes 1-7. Using a ruler, measure the height of lead(II) chromate(VI) precipitate formed. Record all these in Table 1as well as complete Table 1.

Test Tube 1 2 3 4 5 6 7 Volume of Pb(NO3)2 /cm

3

5.0 5.0 5.0 5.0 5.0 5.0 5.0

No of mole of Pb(NO3)2

Volume of K2CrO4 /cm

3

1.0 2.0 3.0 4.0 5.0 6.0 7.0

No of mole of K2CrO4

Height of precipitate / cm

Colour of solution above precipitate

Table 1

Test tube 2… 2.cm3

of potassium

chromate (VI), test

tube 3…3 cm3

varying the volumes

of potassium

chromate (VI)…

Fixed the volumes

of lead (II) nitrate

at 5.0 cm3.

1 2 3 4 5 6 7


106

(b) Based on Table 1, draw a graph of the height of the precipitate against volume of lead (II) nitrate solution on the graph paper.

(c) On the graph that you have drawn in (b), (i) mark and write the minimum volume of potassium chromate(VI) solution needed for complete

reaction with 5.0 cm3 of lead(II) nitrate solution 0.5 mol dm-3. (ii) Calculate the number of moles of chromate(VI) ions that has reacted with 1 mole of Pb2+. ions. (iii) Write the formula of lead(II) chromate.

……………………………………………………………………………………………………………… (iv) Write the ionic equation for the formation of lead(II) chromate(VI). ……………………………………………………………………………………………………………… (d) What can you observed about the height of the precipitate in Figure 1? ……………………………………………………………………………………………………………… …………………………………………………………………………………………………………….. (e) What is your inference based on your answer in (d)?

………………………………………………………………………………………………………………. ……………………………………………………………………………………………………………….


107

Activity 9 : Solve problems involving calculation of quantities of reactants or products in stoichiometric reactions Example 1 : A student prepares copper (II) nitrate by reacting copper (II) oxide with 100 cm3 1.5 mol dm-3 nitric acid. Calculate the mass of copper (II) oxide need to react completely with the acid. [Relative atomic mass: Cu, 64 ; O, 16] Solution :

Chemical equation : CuO + 2HNO3 Cu(NO3)2 + H2O Mole ratio : 1 mole 2 mole 1 mole 1 mole Number of moles of HNO3 = 1.5 x 100 = 0.15 mol 1000 Mole ratio of CuO : HNO3 = 1 : 2 Number of mole of CuO = 1 x 0.15 = 0.075 mole

2 Mass of CuO = 0.075 x (64 + 16) = 6 g Question :

1 Excess zinc powder is added to react completely with 503cm of 2.0

3 dmmol hydrochloric acid.

(a) Write an ionic equation for the reaction between zinc and hydrochloric acid. (b) Calculate the number of moles of hydrochloric acid used. (c) Calculate the volume of hydrogen gas liberated at room conditions.

[Molar volume: 2413 moldm ]

2 Excess of magnesium carbonate powder, MgCO3, is reacted with 100 cm3 of a 1 mol dm-3

sulphuric acid H2SO4 , What is the mass of magnesium sulphate formed?

[Relative atomic mass : Mg =24, O=16, S = 32 ] 3. 0.12 g of magnesium reacts with excess hydrochloric acid to produce hydrogen gas. Given that

the relative molecular mass of H=1, Mg = 24, CI =35.5 and 1 mol of gas occupies 24 dm3 at room temperature and pressure. Fnd the (a) mass of salt formed (b) volume of gas produced


108

Example 2 : A sample of insoluble lead (II) sulphate is prepared by mixing 50 cm3 of 1.0 mol dm-3

lead (II) nitrate solution and y 3cm of 1.5 mol dm-3 sulphuric acid.

[Relative atomic mass: O, 16 ; S, 32 ; Pb, 207] (a) Calculate the volume, y, of the sulphuric acid needed to react completely with the lead (II) nitrate solution.

Solution : Chemical equation : Pb(NO3)2 + H2SO PbSO4 + 2 HNO3

Mole ratio : 1 mole 1 mole 1 mole 2 mole Number of moles of Pb(NO3)2 = 1.0 x 50 = 0.05 mol 1000 Mole ratio of Pb(NO3)2 : H2SO4 = 1 : 1

Number of mole of H2SO4 reacted = 0.05 mol 1.5 x y = 0.05 mole

1000 y = 0.05 x 1000 = 33.33 cm3

1.5 (b) Calculate the mass of lead (II) sulphate obtained.

Solution : Number of mole of PbSO4 = Number of moles of Pb(NO3)2 = 0.05 mol Mass of PbSO4 = 0.05 x (207 + 32 + 4 x 16) g

= 15.15 g Question

4. A sample of insoluble silver chloride is prepared by mixing 503cm of 1.0 mol dm-3 silver nitrate

solution and z 3cm of 0.5

3 dmmol sodium chloride solution.

[Relative atomic mass: Ag 108; Cl 35.5] (a) Write the chemical equation for the reaction between silver nitrate and sodium chloride. (b) Calculate the volume, z, of the sodium chloride needed to react completely with the silver nitrate solution. (c) Calculate the mass of silver chloride obtained.


109

B. SYNTHESISING QUALITATIVE ANALYSIS OF SALTS

Learning Outcomes You should be able to: state examples of salts used in daily life, explain the meaning of salt identify soluble and insoluble salts, describe the preparation of soluble salts, describe the purification of soluble salts by recrystallisation, list physical characteristics of crystals, describe the preparation of insoluble salts, write chemical and ionic equations for reactions used in the preparation

of salts, design an activity to prepare a specified salt, construct ionic equations through the continuous variation method solve problems involving calculation of quantities of reactants or products in stoichiometric reactions

Activity 10 : Qualitative Analysis 1. Qualtitative analysis of a salt is a chemical technique used to identify the …….. that are present

in a salt by analysing its ………………. and ……………………. properties.

2. Make inferences on the following substances based on their colour: (use formula of substance when

writing your answer. Make it is correct!)

Colour (solid or solution) Substance or cation or anion Green powder Salt: Cation Blue powder Cation: Brown powder Cation: Black powder Two metal oxides: Yellow powder when hot and white when cold

Brown powder when hot and yellow when cold

Blue solution Cation: Pale green solution Cation: Brown solution Cation: Solid : White Solution : colourless

6 cations :

Solid : White Solution : colourless

4 anions :

3. Complete the following table

Salts Solubility in water Colour

Insoluble white

Copper(II) carbonate Iron(II) sulphate

Soluble Brown

Lead(II) sulphate Magnesium carbonate Zinc chloride Ammonium carbonate

Insoluble Yellow


110

Activity 11 : Confirmatory Tests for gases,

Complete the observation for the confirmatory test for gases

Gas Method Diagram Observation

Carbon dioxide

Bubble the gas produced into lime water Heating

Carbonate salts

Oxygen Insert a glowing splinter into the test tube

Nitrogen dioxide

Observe the colour of gas produced. Bring a piece of moist blue litmus paper to the mouth of the test tube

Chlorine Observe the colour of the gas. Bring a piece of moist blue litmus

paper to the mouth of the test tube

Ammonia Dip a glass rod into concentrated hydrochloric acid and bring a drop of acid to the mouth of the test tube /place moist red litmus paper at the mouth of the test tube

.

Hydrogen

Bring a lighted splinter to the mouth of the test tube. Mg + HCl release

hydrogen gas

Tests For Gases


111

Activity 12 : Action of Heat On Carbonate Salts Carbonate salts (except Na+ & K+ ) decompose on heating giving off carbon dioxide gas and residue metal oxide Activity : Complete the chemical equation and observation for the action of heat on carbonate salt

Carbonate salt Action of heat

Potassium carbonate K2CO3 , Sodium carbonate Na2CO3

Not decompose by heat

Metal Carbonate metal oxide + carbon dioxide

Calcium carbonate CaCO3 CaO + CO2

Observation : White solid formed. Gas liberated turn lime

water chalky

Magnesium carbonate MgCO3 ……….. .. + …. ……… Observation : ………………………………………………. ……………………………………………………………….

Aluminium carbonate Al2(CO3)3 ……….. .. + …. ……… Observation : ………………………………………………. ……………………………………………………………….

Zinc carbonate ZnCO3 ……….. .. + …. ……… Observation : ………………………………………………. ……………………………………………………………….

Lead(II) carbonate PbCO3 ……….. .. + …. ……… Observation : ………………………………………………. ……………………………………………………………….

Hydrogen chloride

Dip a glass rod into concentrated ammonia solution and bring a drop of ammonia to the mouth of test tube

Metal oxide Colour

Copper (II)

oxide

Black

Zinc oxide Hot: yellow ;

Cold: White

Lead (II) oxide Hot: brown ;

Cold: Yellow

Iron(III) oxide Brown

Lime water

turn chalky


112

Copper(II) carbonate CuCO3 ……….. .. + …. ……… Observation : ………………………………………………. ……………………………………………………………….

Activity 13 : Action of Heat On Nitrate Salts Nitrates Salts - Decompose on heating liberate nitrogen dioxide gas and oxygen gas except NaNO3 and KNO3 which liberate oxygen gas only Activity: Complete the chemical equation and observation for the action of heat on nitrate salt

Nitrate salts Action by Heat

Metal Nitrate metal nitrite + oxygen

Potassium nitrate

2KNO3 2 KNO2 + O2

Observation : white solid formed, gas released relighted glowing splinter

Sodium nitrate

2NaNO3 ……….. .. + …. ……… Observation : ……………………………………………………….. ……………………………………………………………………….

Metal Nitrate metal oxide + nitrogen dioxide + oxygen

Calcium nitrate 2Ca(NO3)2 2CaO + 4NO2 + O2

Observation : white solid formed, Brown gas which turns moist blue

litmus red released. Another gas released relighted glowing splinter

Magnesium nitrate Mg(NO3)2 ……….. .. + …. ……… + ……………

Observation : ……………………………………………………….. ……………………………………………………………………….

Zinc nitrate Zn(NO3)2 ……….. .. + …. ……… + ……………

Observation : ……………………………………………………….. ……………………………………………………………………….

Brown gas turn moist blue litmus to red (NO2)

Colourless gas relighted glowing splinter (O2 ) Heat Heat


113

Lead(II) nitrate Pb(NO3)2 ……….. .. + …. ……… + ……………

Observation : ……………………………………………………….. ……………………………………………………………………….

Copper(II) nitrate Cu(NO3)2 ……….. .. + …. ……… + ……………

Observation : ……………………………………………………….. ……………………………………………………………………….

Activity 14 : Confirmatory Tests for Anions, 1. Write the ionic equation for the following reactions. 2. Which anion produce the following observations? a) b)

Inference :

Salt K1

Add BaCl2

solution + HCl acid

Add AgNO3

solution + HNO3 acid

Inference :

White

precipitate

formed

White

precipitate

formed

Salt K2

Anions

CO3 2-

Cl-

SO4 2-

NO3 -

+ HNO3

+ AgNO3

+ Dilute

acid

+ HCl + BaCl2

+ H2SO4

+ FeSO4

+ concentrated H2SO4

Effervescence – CO2

Lime water turns milky

Ionic equation : …………………………

White precipitate


White precipitate


Brown ring


114

c) d) Activity 15: Reaction of Cations with alkali solution 1. Positive ions are identified by their reactions with a. sodium hydroxide NaOH solution b. Ammonia solution NH3 2. In these reactions, the cations (positive metal ions) produce different coloured precipitate which

may or may not be soluble in excess alkali

Look for precipitate

State whether each of the following precipitate is soluble or insoluble in excess alkali.

NaOH solution Ammonia Solution NH3

A little In excess A little In excess

Soluble ( , X ) Soluble ( , X ) Ca2+ White precipitate No change Zn2+ White precipitate White precipitate Al 3+ White precipitate White precipitate Pb 2+ White precipitate White precipitate Mg2+ White precipitate White precipitate Cu 2+ Blue precipitate Blue precipitate Fe 2+ Green precipitate Green precipitate Fe 3+ Brown precipitate Brown precipitate Which anion produces the following observations? a)

See if Precipitate dissolves

5 drops of alkali

(NaOH or NH3)

Solution of

cations

Salt K5

Add excess NaOH solution

White

precipitate

formed

Inference 2 :

White precipitate Dissolves in excess NaOH solution alkali

White precipitate does not dissolve in excess NaOH solution

Inference 3:

No white precipitate formed , when heated Alkali gas released (ammonia) released

Inference 1

Add 5 drops of NaOH solution

Q1

Inference : Effervescence,

Gas bubbles,

Gas turn lime

water chalky

Brown ring

formed

terhasil

Add FeSO4 solution + concentrated sulphuric acid

Add sulphuric acid

Inference :

Salt K3 Salt K4

Q2


115

b) Activity 16 : Confirmatory Tests for Fe2+, Fe3+, Pb2+ and NH4

+

(A) The table shows how confirmatory tests are conducted for ammonium ion, NH4

+ , Iron(II) ion, Fe2+ , Iron(III) ion, Fe3+ , and lead(II) ion, Pb2+ . Complete the confirmatory tests and observation.

Cation Name of Reagent Observation

2Pb

Add a few drops of …………………… to the test tube containing 2 cm3 of

lead(II) nitrate solution (2Pb ions)

Add 2 cm3 of distilled water and boil the mixture. Cool the contents using running water from the tap.

……………… precipitate is formed which ………………..in the hot water and is ……………… on cooling

2Fe


iron(II) sulphate solution (2Fe ions)

……………….. precipitate is formed

3Fe


iron(III) sulphate solution (3Fe ions)

……………….. solution is formed

4NH Add a few drops of …………………… to the test tube containing 2 cm3 of

ammonium chloride solution (

4NH

ions)

……………….. precipitate is formed

Salt K6

Add 5 drops of NH3 solution

Add NH3 solution in excess

Inference5 :

White precipitate Dissolve in excess NH3 solution

White precipitate does not dissolve in excess NH3 solution

White

precipitate

formed Inference 6:

No White precipitate formed

Inference 4


116

The diagram below shows the flow chart for the chemical test of Fe2+ ions and Fe3+ ions. Based on the flow chart, explain how to differentiate Fe2+ ions and Fe3+ ions. …………………………………………………………………………………………………………. …………………………………………………………………………………………………………. …………………………………………………………………………………………………………. …………………………………………………………………………………………………………. Activity 17 : Qualitative analysis to identify salts

(A). Identify the salt S1 The following tests were carried out to identify salt S1. Based on the observations given for

each test, state its inference. Finally, identify salt S1

Test Observation Inference 1. Heat S1 strongly in a test tube. Identify any gas liberated.

Brown gas and gas relight a glowing splinter liberated. Residue is brown when hot and yellow when cold

2. Dissolve a spatulaful of S1 in distilled water. Divide into four portions and carry out the following tests:

Residue dissolve in acid to produce colourless solution

(a) add NaOH solution until excess.

White precipitate, dissolve in excess NaOH solution

(b) add 3NH solution until

excess

White precipitate, insoluble in excess ammonia solution

(c) add potassium iodide solution

Yellow precipitate formed

(d) add dilute 42SOH ,

followed by 4FeSO

solution. Carefully add about

13cm of concentrated

42SOH

Brown ring formed

Solution contains Fe2+ ions or Fe3+ ions.

Light blue precipitate

Dark blue precipitate

Fe2+ ions

Fe3+ ions

K4Fe(CN)6

Potassium hexacyanoferrate(II)

Test II

Solution contains Fe2+ ions or Fe3+ ions.

No change

Blood red solution

Fe2+ ions

Fe3+ ions

Potassium thiocyanate

KSCN

Test I


117

(A). Conclusion for salt S1 : ……………………………………………………………………. (B). Identify the salt S2 The following tests were carried out on an aqueous solution of salt S2. Based on the

observations given for each test, state its inference. Finally, identify salt S2.

Test Observation Inference

1. Pour about 23cm of S2

into a test tube. Add

NaOH solution until excess

White precipitate, dissolve in excess NaOH solution

2. Pour about 23cm of S2

into a test tube. Add

3NH solution until excess

White precipitate, dissolve in excess ammonia solution

3. Pour about 23cm of S2 into a test

tube. Add dilute 3HNO , followed by

silver nitrate, 3AgNO solution

No change

4. Pour about 23cm of S2 into a test

tube. Add dilute HCl solution, then add BaCl2 solution

White precipitate

(B). Conclusion for salt S2 : ……………………………………………………………………..


118

Activity 18 : plan qualitative analysis to identify anions

Rajoo works in a laboratory. He noticed that there are two large bottles. However both the labels have fallen off. He found four labels beside the bottles. i.e ‘Sodium Chloride Solution’, ‘Sodium Carbonate Solution’, ‘Sodium Sulphate Solution’ and ‘Sodium Nitrate Solution’. So he has to carry out confirmatory test to identify the anion in both the solutions. Complete the graphic organizers describing four tests and their results. The charts can then be used by Rajoo to distinguish which bottle contains which solution. CHART A: SODIUM CARBONATE AND SODIUM NITRATE

Test 1

Add dilute HNO3

(or any dilute acid)

Test 2

Add dilute H2SO4

followed by …………

solution.

Carefully add 1 cm3 of

……………………

H2SO4

Test 3

Add dilute HNO3,

followed by

……………

…………., solution

Test 4

Add dilute HCl,

followed by

……………

……………. solution

SO

DIU

M C

AR

BO

NA

TE

SO

DIU

M N

ITR

AT

E

Result 1

Result 2

Result 3

Result 4

Result 1

Result 2

Result 3

Result 4


119

CHART B: SODIUM CHLORIDE AND SODIUM SULPHATE

Test 1

Add dilute HNO3

(or any dilute acid)

Test 2

Add dilute H2SO4

followed by …………

solution.

Carefully add 1 cm3 of

……………………

H2SO4

Test 3

Add dilute HNO3,

followed by

……………

…………., solution

Test 4

Add dilute HCl,

followed by

……………

……………. solution

SO

DIU

M C

HL

OR

IDE

SO

DIU

M S

UL

PH

AT

E

Result 1

Result 2

Result 3

Result 4

Result 1

Result 2

Result 3

Result 4


120

CHAPTER 9 : MANUFACTURED SUBSTANCES IN INDUSTRY A : SULPHURIC ACID Learning Outcomes You should be able to:

list uses of sulphuric acid

explain industrial process in the manufacture of sulphuric acid

explain that sulphur dioxide causes environmental pollution.

Activity 1 : SULPHURIC ACID 1. Sulphuric acid is manufactured through the …………………….Process

. 2. Contact Process consists of three stages:

(Complete the table below)

Stage 1

Preparation of (a)……………………………… gas

Chemical equation : S + O2 SO2

Stage 2

Conversion of sulphur dioxide to sulphur trioxide Chemical equation : (b)…………. + ………..… …….

Catalyst : (c)…………………………………….. Temperature : (d)……………….. oC

Pressure : (e)……….. atm

Stage 3

Production of sulphuric acid Chemical equation : (f)……… + H2SO4 ………….. (g)………. + H2O …………….

3. State six uses of sulphuric acid.

i)…………………………………………………………………………………………..

ii)…………………………………………………………………………………………..

iii)………………………………….. ……………………………………………………….. iv). …………………………………………………………………………………………… v)…………………………………… ………………………………………………………… vi). ……………………………………………………………………………………………..

4. ……………… ……………………….gas from the burning of product manufactured from sulphuric acid can cause ………………………disease and ……………… rain.

5. Figure 1 below shows the waste product from a factory which affect the quality of the

environment.


121

a) By referring to the Figure 1 above, state the following,

i) Types of waste products and their sources.

…………………………………………………………………………………

…………………………………………………………………………………

ii) How acid rain is formed and its effect.

Formation of acid rain:……………………………………………………….

…………………………………………………………………………………….

Effects on environment:……………………………………………………..

……………………………………………………………………………………… ……………………………………………………………………………………… iii) How does the toxic waste affect the environment and its effect

…………………………………………………………………………………….. ……………………………………………………………………………………… ………………………………………………………………………………………. B : AMMONIA Learning Outcomes


list uses of ammonia

state the properties of ammonia

explain the industrial process in the manufacture of ammonia

design an activity to prepare ammonium fertilizer.

Activity 2 : AMMONIA 1. Ammonia is manufactured through the Haber Process by combining

………………… gas and …………………………….gas. 2. (i). The reaction can be represented by the chemical equation

…………………………………………………………………………… (ii) State the condition necessary to produce ammonia.

Catalyst : (a) …………………………………………………………..

Figure 1


122

Temperature : (b) ……………………………………………………………. Pressure : (c)……………………………………………………………… Ratio N2 :H2 : (d) ……………………………………………………………… 3. The following statements refer to the uses of ammonia. Fill in the blank with the correct

words.

(a) To manufacture ……………………………….. such as ammonium sulphate and ammonium nitrate. The chemical equation for producing ammonium sulphate is given by

……………………………………………………………………………………………………

(b) Ammonia is used as raw material to produce …………………………. in the Ostwald Process.

(c) Ammonia is also used as an alkali to prevent the ………………………………of latex. 4. Listed below are three properties of ammonia. Fill in the blank according to the aspect

given.

(a) Colour: Ammonia is a……………………………….. gas. (b) Solubility: Ammonia is very……………………………in water.

(a) Smell : Ammonia has a …………………………… smell.

(b) Ammonia dissolves in water to produce an ………………………………. solution.

5. Ammonium fertiliser can be prepared in the laboratory by adding ammonia solution

and certain acids as shown in the table below.

Neutralisation reactions Alkali Acid

Name of ammonium salt (fertiliser)

Aqueous + Phosphoric

ammonia acid

Ammonium phosphate

(a) Formula:………………………………..

Aqueous + Nitric

ammonia acid

Ammonium nitrate

(b) Formula:………………………………..

Aqueous + Sulphuric

ammonia acid

Ammonium sulphate

(c ) Formula:……………………………….

(i) Calculate the percentage of nitrogen found in each of the ammonium

fertilisers. [Relative atomic mass: H = 1; N = 14; O = 16; P = 31; S = 32]

(ii) From the calculations in (b)(ii), deduce the type of ammonium compound that

is most

suitable for use as a nitrogenous fertiliser. Give reasons for your answer.

……………………………………………………………………………………


123

…………………………………………………………………………………….

C : ALLOY Learning Outcomes You should be able to:

relate the arrangement of atoms in metals to their ductile and malleable properties

state the meaning of alloy

state the aim of making alloys

list examples of alloys

list compositions and properties of alloys

relate the arrangement of atoms in alloys to their strength and hardness

relate properties of alloys to their uses.

Activity 4: ALLOY 1. What is alloy?

Alloy is a .......................of a pure metal with...................................in......................quantities

2. A pure metals contains atoms of the same size arranged in a regular and orderly

manner. Pure metal are ……………………………….. and ……………….…………… because the layers of atom……………………………………………when external force is applied on them.

3. In an alloy, the foreign metal atoms ………………………. …………...arrangement of metal

atoms and the layers of metal atoms are prevented from …………......over each other easily. 4. Complete the sequences by drawing the arrangement of atoms in the box below.

+

5. Three aims of alloying a pure metal are :

a. to increase the …………………………. and ……………………….. of metal. b. to prevent ……………………………… or rusting. c. to improve the …………………………. of metal surface.

Pure metal Another pure

metal

alloy

Figure 2


124

6. Examples of alloy.

(Complete the table below)

Alloy Composition Properties Uses

(i) ……………………

99% Iron 1% (ii)…………..

Hard

Bridges, vehicles, heavy machinery framework

(iii) ………………

97% (iv)………… 3% lead and antimony

Hard and shiny

Decorative ornaments, souvenirs

Bronze

90% (v)………… 10% tin

Hard and shiny

Decorative ornaments , art crafts

Brass

70% Copper 30% (vi)…………

Hard and shiny

Decorative ornaments, musical instrument

Magnalium

70% Aluminium 30% (vii)…………

Hard and light

(viii)……………… ……………………

D : POLYMERS


state the meaning of polymers

list naturally occurring polymers

list synthetic polymers and their uses

identify the monomers in the synthetic polymers

justify uses of synthetic polymers in daily life.

Activity 5: POLYMERS Fill in the blanks below.

1. Polymers are ……………….…………………... made up of many smaller and

identical separating unit called ……………………………………..

2. …………………………………… is the process by which the monomers are

joined together to form a big molecule known as the polymer.


125

3. Give at least two examples of:naturally occurring polymers and at least two examples of synthetic polymers.

4. Match the synthetic polymers with their respective monomer

Polyethylene Phenylethene

Polypropylene Chloroethene

Polyvinylchloride Ethene

Polystyrene Propene

Perspex Tetrafluoroethene

Teflon Methylmethacrylate

6. Complete the table.

Synthetic Polymer Uses

Polyethylene

Polypropylene

Polyvinylchloride

Polystyrene

Perspex

Teflon

Naturally occurring polymers Synthetic polymers

Synthetic Polymer Monomer


126

E : GLASS AND CERAMICS


list uses of glass

list uses of ceramics

list types of glass and their properties

state properties of ceramics.

Activity 6: GLASS AND CERAMIC Fill in the blanks below. 1. Main composition of glass is ………………………………, (SiO2).


Type of glass Chemical composition

Properties Examples of uses

(a) ……………………….

Silica 99% Boron oxide 1%

Very high softening point (1700oC). Transparent to ultra violet and infra red light. Difficult to be made into different shapes. Does not crack with sudden temperature change. Very resistant to chemical attack.

Mirrors, Lenses, Laboratory glass wares.

(b) ……………………….

Silica 70% Sodium oxide 15% Calcium oxide 10% Others 5%

Low softening point (700oC). Breaks easily. Cracks easily with sudden temperature changes. Less resistant to chemical attack. Easy to make into different shapes.

Bottles, Window, Light bulb, Bowl

(c) ……………………….

Silica 80% Boron oxide 15% Sodium oxide 3% Aluminium oxide 1%

High softening points(800oC). Does not crack easily with sudden temperature change. Transparent to ultra violet light. Very resistant to chemical attack.

Laboratory apparatus, Cooking utensils, Electrical tubes.

(d) ………………………..

Silica 55% Lead oxide 30% Potassium oxide 10% Sodium oxide 3% Alimunium oxide 2%

Low softening point (600oC). High density. High refractive index. Reflects light rays and appears shiny.

Decorative items, Crystal glass wares, Lens, Prisms Chandeliers


127

3. ……………………………are made from clay that is dried and then baked in a kiln at high temperatures. 4. The main constituent of clay is ……………………………………………….. ….

(aluminium oxide and silicon dioxide). 5 …………………... ……………………...consists of hydrated aluminosilicate crystals. (High

quality white clay)

6. Complete the table.

Composition Properties Examples of uses

Aluminosilicate (aluminium oxide and silicon dioxide)

(i) very hard and strong but brittle (ii)………………………………………. (iii)………………………………………. (iv)………………………………………. (v)………………………………………..

Construction materials, Tableware, Insulators in electric equipments, Refractories. Flowerpots

Activity 7

1. Compare and contrast between glass and ceramic.

Glass Ceramic

(a) Main components: ………………………………………..

(b) Main components: ……………………………………………... _______________________ (c) 4 types of glass:

………………………………………. ………………………………………. ……………………………………… ……………………………………… (d)Improved Glass: ………………………......................

(e) 4 examples of ceramics: …………………………………………….. ……………………………………………... …………………………………………….. …………………………………………….. (f) Improved Ceramics: ……………………………………………..

(g) 4 common Properties of glass and ceramic (i) very hard and strong but brittle (ii) ………………………………………………… (iii) ………………………………………………… (iv)…………………………………………………….


128

F : COMPOSITE MATERIALS Learning Outcomes You should be able to:

describe needs to produce new materials for specific purposes

state the meaning of composite materials

list examples of composite materials and their components

compare and contrast properties of composite materials with those of their original components.

Activity 8: COMPOSITE MATERIALS 1. Composite material is a structural material that is formed by

………….........…….or……………… different substances such as metal, alloys, glass,

ceramics and polymers.

2. Give three examples of composite materials.

(i)…………………………………………………………………………………………………….. (ii)…………………………………………………………………………………………………… (iii)…………………………………………………………………………………………………….

3. State the purpose of creating composite materials.

………………………………………………………………………………………………….. ………………………………………………………………………………………………….

4. Photochromic glass is an example of composite material.

a) Compare to a normal glass, what is added to a photochromic glass?

…………………………………………………………………………………………….. b) State the special feature of a photochromic glass.

…………………………………………………………………………………………… ……………………………………………………………………………………………

(h) 2 differences: (i) ………………………………………………… (ii) …………………………………………………


129


Example

Composition

Properties

Uses

Reinforced concrete

(i) …………………... …………………... ……………………

Strong but brittle, Weak in tension

(ii) ………………….. ………………….. …………………..

Superconductors

(iii) ………………….. ………………….. …………………..

Conducting electricity

(iv) …………………… …………………… ……………………

Fibre optic

(v) ………………….. ………………….. …………………..

Low material costs, High transmission capacity, chemical stability, Less susceptible to interference.

Transmit data, voice and image in a digital format.

Fibre glass

(vi) ………………….. ………………….. …………………..

High tensile strength, Easy to colour, Low in density, Very strong.

(vii) ………………….. ………………….. ……………………

Photochromic glass

(viii) ………………….. …………………... …………………..

When it is exposed to light, silver chloride is converted silver and darken the glass

(ix) ………………….. ………………….. …………………..

Activity 9

1. (a) Bronze is an alloy of copper. (i) Name the element that is added to copper to form bronze. …………………………………………………………………………………. .[1 mark]

(ii) Explain why bronze is harder than pure copper. ……………………………………………………………………………………………. ………………………………………………………………………………….. [2 marks]

(iii) Draw and label the arrangement of particles in pure copper and bronze. Pure copper Bronze [2 marks] (b) Synthetic polymers are widely used in our daily lives. (i) Complete the table with the correct monomers.

Polymer Monomer

Polyvinylchloride

Polyethene

[2 marks] (ii) State one example of polyvinylchloride commonly used n our daily lives. ……………………………………………………………………………… [1 mark]


130

(c) Glass and ceramic have similar characteristics. (i) State one similar characteristic of glass and ceramic.

……………………………………………………………………………… [1 mark]

(ii) What type of glass is used to make laboratory glassware? ……………………………………………………………………………… [1 mark]

2. Figure 3 shows the flow chart for the industrial manufacture of sulphuric acid and the production of fertilizer Z.

V2O5, 1 atm

Heat 450

oC-500

oC

Figure 3

Based on Figure 3, answer the following questions. (a) Name the process of manufacturing sulphuric acid. …………………………………………………………………………………………

[1 mark] (b) Name the substance X.

……………………………………………………………………………………………

[1 mark]

(b) Substance X could react directly with water to form sulphuric acid. Explain why this step is not carried out in the industrial process.

……………………………………………………………………………………………

…………………………………………………………………………………………… [1 mark]

(c) Write the chemical equation when oleum reacts with water to form sulphuric

acid.

…………………………………………………………………………………………….. [1 mark]

(d) Name the substance Y and the fertilizer Z.

Substance Y: ………………………………………………………………………

Fertilizer Z:……………………………………………………………………………...

Sulphur Sulphur

dioxide

X

Fertilizer Z Sulphuric

acid

Oleum

Concentrated H2SO4

Substance

Y


131

(e) Write a chemical equation when sulphur dioxide reacts with rain water.

……………………………………………………………………………………………….

[1 mark ] 3 A student conducts an experiment to study the hardness of two metallic plates,

R and S. He drops a steel ball on R a few times and each time, the diameter of the dent is measured. He repeats the same procedure on the S plate. The reading of the diameter of the dents made on each metallic plates are as follows,

Plate

Diameter of dent made (mm)

I II III Average size

R 2.4 2.3 2.3

S 3.1 3.2 3.2

(a) Write the average size of the dents made by R and S in the table above.

[ 2 marks] (b) What are the differences seen in the two types of metallic plates based on their

(i) properties ………………………………………………………………………… (ii) composition………………………………………………………………………

[ 4 marks] (c) From the observation made in the given table above, which plate would be made of

(i) iron? ………………………………………………………………………… (ii) steel? …………………………………………………………………………

[ 2 marks] (d) Give two uses of steel.

…………………………………………………………………………………… ……………………………………………………………………………………

[ 2 marks ]

The combustion of petrol in the engines of vehicles produce sulphur dioxide. This gas when dissolved

in rain water is corrosive.

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