mole concept subtopic: conversions week one: le

MINISTRY OF EDUCATION

SECONDARY ENGAGEMENT PROGRAMME

CHEMISTRY

GRADE 10

TOPIC: MOLE CONCEPT

SUBTOPIC: CONVERSIONS

WEEK ONE: LESSON 1

Definitions:

The mole is a unit of measurement. It is the amount of substance which contains the same number

of particles as there are in 12 g of carbon – 12.

The mole is the amount of substance which contains 6.0 X 1023 particles (Avogadro’s number).

Molar mass is the mass (in grams) of one mole of a substance.

The relative atomic mass, Ar, of an element is the ratio of the average mass of one atom of an

element compared to 1/12 the mass of carbon-12.

The relative molecular mass, Mr, is the average mass of one molecule or formula unit of the

compound compared with 1/12 the mass of one atom of Carbon-12.

Finding the molar mass:

• STEP 1: Write the formula for the compound

• STEP 2: Multiply the relative atomic mass (from the periodic table) of each element

present by the number of atoms of that element in the compound

• STEP 3: Add them all

Consider sulphuric acid, H2SO4

Atoms present # of atoms in the

molecule

Relative atomic mass

of one atom

Total mass of given

atoms

H 2 1.0 2.0

S 1 32.0 32.0

O 4 16.0 64.0

TOTAL = 98 gmol-1

Converting mass to moles

Example

A chemist has a jar containing 388.2 grams (g) of iron (Fe) filings. How many moles of iron

does the jar contain?

1. ANALYZE

● What is given in the problem? mass of iron in grams

● What are you asked to find? amount of iron in moles

2. PLAN

• What step is needed to convert from mass of Fe to number of moles of Fe?

The molar mass of iron can be used to convert mass of iron to amount of iron in moles.

3. COMPUTE

No of Moles = 𝒎𝒂𝒔𝒔

𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔

No of Moles = 𝟑𝟖𝟖.𝟐

𝟓𝟓.𝟖𝟓 = 6.95 mols

ACTIVITY 1

1. Find the relative molecular mass of:

i. Calcium Oxide, CaO

ii. Water, H2O

iii. Hydrogen Peroxide, H2O2

iv. Sulphuric acid, H2SO4

v. Ethanoic acid, CH3COOH

2. Calculate the number of moles in each of the following masses:

i. 64.1 g of aluminum, Al

ii. 28.1 g of silicon, Si

iii. 0.255 g of sulfur, S

iv. 850.5 g of zinc, Zn

TOPIC: MOLE CONCEPT


WEEK ONE: LESSON 2

Converting moles to mass

Example

A student needs 0.366 mol of zinc for a reaction. What mass of zinc in grams should the student

obtain?

1. ANALYZE

• What is given in the problem? amount of zinc needed in moles

• What are you asked to find? mass of zinc in grams

From: No of Moles = 𝒎𝒂𝒔𝒔

𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔

Mass = No. of Moles x Molar Mass

= 0.33 Moles x 65.39 g/mol

= 21.5 grams.

Converting no. of particles to moles

Example

How many moles of lithium are there in 1.204 x 1024 lithium atoms?

1. ANALYZE

• What is given in the problem? number of lithium atoms

• What are you asked to find? amount of lithium in moles

Items Data

Number of lithium of atoms 1.204 X 1024 atoms

Avogadro’s constant—the

number of atoms per mole

6.022 X 1023

atoms/mol

Amount of lithium ? mol

2. PLAN

What step is needed to convert from Avogadro’s constant is the number

number of atoms of Li to moles of Li?

3. COMPUTE

No. of Moles = 𝑁𝑜.𝑜𝑓 𝑎𝑡𝑜𝑚𝑠

𝐴𝑣𝑜𝑔𝑎𝑑𝑟𝑜′𝑠 𝑛𝑢𝑚𝑏𝑒𝑟

No. of Moles = 1.204 𝑥 10^24

6.022 𝑥 10^23 = 1.99 mols

ACTIVITY 2

1. Calculate the mass of each of the following amounts:

i. 1.22 mol sodium, Na

ii. 14.5 mol copper, Cu

iii. 0.275 mol mercury, Hg

iv. 9.37 x 103 mol magnesium, Mg

2. Calculate the amount in moles in each of the following quantities:

i. 3.01 x 1023 atoms of rubidium

ii. 8.08 x 1022 atoms of krypton

iii. 5.7 x 109 atoms of lead

iv. 2.997 x 1025 atoms of vanadium

TOPIC: MOLE CONCEPT


WEEK ONE: LESSON 3

Converting moles to no. Of particles

How many atoms are there in 2 mols of lithium?

Example

Number of atoms = No. of Moles x Avogadro’s constant

Number of atoms = 2 mols x 6.022 X 1023

= 1.204 x 1024

Converting mass to no. Of particles

Example

How many boron atoms are there in 2.00 g of boron?

1. ANALYZE

• What is given in the problem? mass of boron in grams

• What are you asked to find? number of boron atoms

Items Data

Mass of boron 2.00 g

Molar mass of boron 10.81 g/mol

Avogadro’s constant—the

number of boron atoms per

mole of boron

6.022 x 1023

atoms/mol

Number of boron atoms ? atoms

2. PLAN

• What steps are needed to convert from grams of B to the number of atoms of B?

STEP 1: Convert the mass of boron to moles of boron by using the molar mass of boron.

STEP 2: Use Avogadro’s constant to convert the amount in moles to the number of atoms of

boron.

3. COMPUTE

No. of Moles = 𝑚𝑎𝑠𝑠

𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

No. of Moles = 2𝑔

10.81 𝑔/𝑚𝑜𝑙 = 0.185 mol

No.of Atoms = No. of Moles x Avogardo’s constant

= 1.85 x 6.022 x 1023

= 1.1 x 1024 atoms f

Converting no. Of particles to mass

STEP 1: Use Avogadro’s constant to convert the number of particles to amount in moles.

STEP 2: Then convert the number of moles to mass.

ACTIVITY 3

1. Calculate the number of atoms in each of the following quantities:

i. 1.004 mol of bismuth

ii. 2.5 mol of manganese

iii. 2.0 x 10-7 mol of helium

iv. 32.6 mol of strontium

2. Calculate the number of atoms in each of the following masses:

i. 54.0 g of aluminum

ii. 69.45 g of lanthanum

iii. 0.697 g of gallium

iv. 0.000 000 020 g beryllium

3. Calculate the mass of the following:

i. 54.0 g of aluminum6.022 x 1024 atoms of tantalum

ii. 3.01 x 1021 atoms of cobalt

iii. 1.506 x 1024 atoms of argon

iv. 1.20 x 1025 atoms of helium

TOPIC: MOLE CONCEPT

SUBTOPIC: PERCENTAGE COMPOSITION BY MASS

WEEK TWO: LESSON 4

Definitions:

The percentage composition by mass of a compound is the percent by mass of each element in the

compound.

Finding the %composition by mass

% composition by mass of a particular element = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑒𝑙𝑒𝑚𝑒𝑛𝑡

𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 X 100

Consider Al2(SO4)3:

Find the mass of the compound

Atoms present # of atoms in the

molecule

Relative atomic mass

of one atom

Total mass of given

atoms

Al 2 27 54

S 3 32 96

O 12 16 192

TOTAL = 342

% composition by mass of Al = 54

342 X 100 = 15.78 %

% composition by mass of S = 96

342 X 100 = 28.08 %

% composition by mass of O = 192

342 X 100 = 56.14 %

ACTIVITY 4

1. Which substance has the greater percent by mass of hydrogen: C4H8 or C8H18?

2. Ammonium sulphate, (NH4)2SO4, is an important ingredient in many artificial fertilisers

supplying to plants the essential mineral elements of nitrogen and sulphur.

3. Calculate the percentage of nitrogen and the percentage of sulphur in ammonium sulphate.

Calculate the percentage of sulphate ion in ammonium sulphate.

4. What is the percentage of carbonate ion in sodium carbonate? (Na2CO3)

5. Calculate the percentage water of crystallisation in magnesium sulphate crystals,

MgSO4.7H2O, known as Epsom salt.

TOPIC: MOLE CONCEPT

SUBTOPIC: EMPIRICAL AND MOLECULAR FORMULA

WEEK TWO: LESSON 5

The empirical formula is the simplest whole no. ratio of atoms or ions in compound.

The molecular formula is the actual number of atoms or ions in a compound.

Finding the empirical formula

Example

Determine the empirical formula of a compound containing 41.0 % by mass potassium, 33.7 % by

mass sulphur and 25.3 % by mass oxygen.

K S O

% composition by

mass

41 % 33.7 % 25.3 %

Mass in 100 g of

compound

41 g 33.7 g 25.3 g

# of moles 41

39 = 1.05

33.7

32 = 1.05

25.3

16 = 1.58

Mole ratio (divide by

the smallest moles)

1.05

1.05 = 1

1.05

1.05 = 1

1.58

1.05 = 1.5

Simplest whole

number ratio

2 2 3

Empirical Formula = K2S2O3

NB: Since the mole ratio for O is not a whole number, multiply throughout by the smallest whole

number to produce a whole number, in this case the number 2.

Finding the molecular formula

The molecular formula can be calculated given the empirical formula and the molar mass.

Molecular formula = no. of empirical units X empirical formula

Consider a compound with empirical formula CH and a molar mass of 39 gmol-1.

# of empirical units = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑒𝑚𝑝𝑖𝑟𝑖𝑐𝑎𝑙 𝑢𝑛𝑖𝑡 =

39

13 = 3

NB: mass of the empirical unit, CH = 13 g

Molecular formula = no. of empirical units X empirical formula

= 3 X CH = C3H3

ACTIVITY 5

1. Find the empirical formula of the following compounds.

i. Pb = 92.8 %, O = 7.2 %

ii. Na = 43.4%, C = 11.3 %, O = 45.3%

iii. K = 42.4%, Fe = 15.2 %, C = 19.5 %, N = 22.8%.

2. Find the empirical and molecular formula of the following compounds.

i. A gaseous compound of molar mass 44 gmol-1 containing 27.3 % carbon and 72.7 %

oxygen.

ii. An oxide of phosphorous, with a relative molecular mass of 284 containing 43.7 %

phosphorous and 56.3 % oxygen.

iii. A compound of relative molecular mass 545, containing 20 % carbon, 2.2 % hydrogen

and the other element being chlorine.

TOPIC: MOLE CONCEPT

SUBTOPIC: AVOGADRO’S LAW AND BALANCING EQUATIONS

WEEK TWO: LESSON 6

Definitions:

Avogadro’s Law states that equal volumes of gases under the same conditions of

temperature and pressure contain the same number of particles.

The volume which contains 1 mole of a gas is known as the molar volume and is

the same for ALL GASES measured under the same conditions of temperature and

pressure.

The molar volume of a gas is usually quoted under two sets of conditions:

RTP: Room temperature and pressure

Temp: 298 K / 25 oC Pressure: 101 Kpa / 1 atm

Molar Volume of all gases: 24 dm3 / 24,000 cm3

STP: Standard temperature and pressure

Temp: 273 K / 0 oC Pressure: 101 Kpa / 1 atm

Molar Volume of all gases: 22.4 dm3 / 22,400 cm3

Example: What volume is occupied by 8g of hydrogen at STP?

# of mol = 𝑚𝑎𝑠𝑠

𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 =

8𝑔

2𝑔/𝑚𝑜𝑙

# of mol = 4

# of mol = 𝑣𝑜𝑙𝑢𝑚𝑒

𝑚𝑜𝑙𝑎𝑟 𝑣𝑜𝑙𝑢𝑚𝑒

Volume = # of mol x molar volume

Volume = 4 mol X 22.4 dm3mol-1 = 89.6 dm3

The law of conservation of matter: Matter can change form through physical and chemical

changes, but through any of these changes, matter is conserved. The same amount of matter exists

before and after the change—none is created or destroyed.

Based on the law of conservation of matter: the no. of moles of reactants must equal the no. of

moles of products. Hence, the need for balancing chemical equations.

Balancing Equations

• Atoms are NOT changed in a chemical reaction – ONLY the way they are joined changes.

• Eg. Chlorine is no longer joined to hydrogen it is now joined to sodium

HCl(aq) + Na2CO3(s) NaCl(aq) + H2O(l) + CO2(g)

• NOW because of the law of conservation of matter – YOU MUST HAVE the same no. of

the same types of atoms on both sides of an equation ie. It must be balanced.

2HCl(aq) + Na2CO3(s) 2NaCl(aq) + H2O(l) + CO2(g)

ACTIVITY 6

1. What is the volume of 3.5 g of hydrogen at STP?

2. What is the volume of 11.6 g of butane gas at RTP?

3. What mass of chlorine (Cl2) does 6.0 dm3 of the gas contain at RTP?

4. What mass of methane (CH4) is contained in a 10 dm3 cylinder at three times normal

atmospheric pressure at room temperature?

5. Write balanced equations for the following reactions:

i. Zinc metal reacts with copper (II) sulphate solution to form zinc sulphate solution

and copper metal

ii. Lithium metal reacts with fluorine gas to form solid lithium fluoride

iii. Ammonium chloride solution and silver nitrate solution react to form ammonium

nitrate solution and solid silver chloride

TOPIC: MOLE CONCEPT

SUBTOPIC: BALANCING IONIC EQUATIONS

WEEK THREE: LESSON 7

Ionic equations show ONLY the species taking part in a reaction between IONIC SUBSTANCES.

The substances that do not undergo a change are known as spectator ions and must be removed.

Ensure that the net charges are equal on both sides.

Eg. Aqueous barium nitrate reacts with sulphuric acid to form solid barium sulphate and nitric

acid.

Balanced Chemical equation: Ba(NO3)2(aq) + H2SO4(aq) BaSO4(s) + 2HNO3(aq)

Full balanced ionic equation:

Ba2+(aq) 2NO3(aq)

- + 2H+(aq) SO4

2-(aq) BaSO4(s) + 2H+

(aq) + 2NO3(aq)-

Remove spectator ions: Ba2+(aq) 2NO3(aq)

- + 2H+(aq) SO4

2-(aq) BaSO4(s) + 2H+

(aq) + 2NO3(aq)-

Balanced Ionic equation: Ba2+(aq) + SO4

2-(aq) BaSO4(s)

The mole and Solutions

Standard solution: a solution of known concentration. It can be used to find the concentration of

another solution by reacting the standard with the other solution of unknown concentration. This

analytical method of analysis is known as titration.

Mass concentration: the mass of the solute (in grams) dissolved in 1000 cm3 (1 dm3) of solution.

Molar concentration: the number of moles of solute dissolved in 1000 cm3 (1 dm3) of solution.

NB: 1000 cm3 = 1 dm3 = 1 L

Finding mass concentration

Example:

Find the mass concentration of the following solution:

25 cm3 of Na2CO3 solution containing 0.5 g of Na2CO3.

Since the mass concentration is the mass dissolved in 1000 cm3 of solution

Mass concentration = 𝑚𝑎𝑠𝑠

𝑣𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑐𝑚3 x 1000 = 20 g dm-3

Mass concentration = 0.5 𝑔

25 𝑐𝑚3 x 1000 = 20 g dm-3

Finding molar concentration

Example:

Find the molar concentration of 750 cm3 of solution that contains 1.5 mol H2SO4.

Molar concentration = # 𝑜𝑓 𝑚𝑜𝑙𝑠

𝑣𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑐𝑚3 x 1000

Molar concentration = 1.5 𝑚𝑜𝑙

750 𝑐𝑚3 x 1000 = 2 mol dm-3

ACTIVITY 7

1. Write balanced ionic equations for the following reactions:

i. Aqueous potassium iodide reacts with aqueous lead nitrate to produce solid lead

iodide and aqueous potassium nitrate.

ii. Aluminium metal reacts with hydrochloric acid to form aqueous aluminium

chloride and hydrogen gas.

iii. Aqueous calcium hydroxide reacts with nitric acid to form aqueous calcium nitrate

and water.

iv. Aqueous sodium hydrogen carbonate and sulphuric acid reacts to form aqueous

sodium sulphate, carbon dioxide and water.

2. What is the concentration in g dm-3 of the following solutions?

i. 21.25 cm3 of solution containing 6.25 of HNO3.

ii. 10 dm3 of solution containing 1 Kg of CuSO4.

3. What is the concentration in mol dm-3 of the following solutions?

iii. 7.2 dm3 of solution containing 10 mols of NaOH.

iv. 25 cm3 of solution containing 2.4 x 10-3 mols of NaCl.

TOPIC: MOLE CONCEPT

SUBTOPIC: MISCELLAENOUS QUESTIONS


ACTIVITY 8

1. If 10 g of calcium carbonate (limestone) is thermally decomposed what volume of carbon

dioxide is formed at room temperature and pressure?

MgCO3(s) + H2SO4(aq) MgSO4(aq) + H2O(l) +CO2(g)

i. What mass of magnesium carbonate is needed to make 6 dm3 of carbon dioxide?

[RAM: Mg = 24, C = 12, O = 16, H =1 and S = 32]

2. Six grams of a hydrocarbon gas has a volume of 4.8 dm3. Calculate its molecular mass.

3. Given the equation: Ca(s) + 2HCl(aq) CaCl2(aq) + H2(g)

What volume of hydrogen is formed when:

i. 3g of calcium is dissolved in excess hydrochloric acid?

ii. 0.25 moles of hydrochloric acid reacts with calcium?

4. Given the equation ... (RAM Mg = 24, H = 1, Cl = 35.5)

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

i. How much magnesium is needed to make 300 cm3 of hydrogen gas?

5. What volume of carbon dioxide is formed at RTP when 5g of carbon is completely burned?

C(s) + O2(g) CO2(g)

6. What volume of carbon dioxide gas is formed at RTP if 1Kg of propane gas is burned?

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

7. 3.27g of a metal ‘M’ was dissolved in dilute sulphuric acid to form 1.2 dm3 of hydrogen

gas at RTP. Calculate the atomic mass of the metal.

M + H2SO4 MSO4 + H2

TOPIC: ACIDS, BASES AND SALTS


Definitions:

Acids: proton donors; produces H+ ions in solution

Acid anhydride: non-metal oxide that dissolves in H2O to form an acid; form salts when reacted

with bases

Bases: proton acceptors, produces OH- when in soln

Alkali: soluble base

Salt: compound formed when the H+ in acid is replaced by metal or ammonium ions

Acidic oxides: oxides of non-metals; react with H2O to form acids eg. SO2, CO2

Basic oxides: oxides of metals; react with H2O to form alkalis eg. Na2O, MgO, CaO

Amphoteric oxides: metal oxides that can react with both acids and bases eg. ZnO, Al2O3

Neutral oxides: have no acidic or basic properties; do not form salts when reacted with acids or

bases eg. CO, NO, N2O

Strengths of acids and bases

Strong acids and bases: Completely ionized when in dilute solution

Strength of acids increase with a decrease in pH

Strength of bases increase with an increase in pH

Weak acids and bases: only partially ionized when in dilute solutions

Reaction of acids with metals:

Acid + metal salt + hydrogen e.g. H2SO4(aq) + Zn(s) ZnSO4(aq) + H2(g)

Reaction of acids with bases:

Acid + base salt + water e.g. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Reaction of acids with hydrogen carbonates:

Acid + hydrogen carbonate salt + water + carbon dioxide

e.g HNO3(aq) + NaHCO3(aq) NaNO3(aq) + H2O(l) + CO2(g)

Reaction of acids with carbonates:

Acid + carbonate salt + water + carbon dioxide

2HCl(aq) + Na2CO3(s) 2NaCl(aq) + H2O(l) + CO2(g)

Reaction of bases with ammonium salts:

Base + ammonium salts salt + water + ammonia

Ca(OH)2(s) + (NH4)2SO4(s) CaSO4(s) + 2H2O(l) + 2NH3(g)

Solubility Rules

Salts Solubility characteristics

Nitrates All nitrates are soluble

Halides All halides are soluble except Ag and Pb; lead

chloride and lead bromide are soluble in hot

water

Sulphates All sulphates are soluble except Ba and Pb;

Ca and Ag are slightly soluble

Carbonates All carbonates are insoluble except Na, K and

NH4

+

Hydrogencarbonates Most are soluble

ACTIVITY 9

MULTIPLE CHOICE QUESTIONS

1. Metals react with acids to give a salt and ________.

A. Water.

B. Hydrogen.

C. Carbon dioxide.

D. Alkali.

2. Which of the following is not a property of alkalis?

A. They displace ammonia from ammonium salts

B. They are soapy to touch

C. They have a bitter taste

D. They do not react with acids

3. Which of the following statements about all bases are true?

A. Metallic oxides and hydroxides are alkalis

B. They react with acids to form a salt and water only

C. They are conductors of electricity in aqueous solutions

D. They produce hydroxide ions in aqueous solutions

4. Which of the following is the formula of a dibasic acid?

A. HNO3

B. CH3COOH

C. H2CO3

D. H3PO3

5. Which of the following statement/s is/are true?

I. Metal oxides are bases

II. Metal hydroxides are bases

III. Soluble bases are alkalis

IV. Alkalis have a pH of less than 7

A. I only

B. I, II and III

C. IV only

D. III and IV

TOPIC: ACIDS, BASES AND SALTS

SUBTOPIC: NEUTRALISATION REACTIONS

WEEK FOUR: LESSON 10

Definitions:

Acids salts: form acidic solutions when dissolved in a solvent. They still contain replaceable H+

ions from the parent acid. Eg. NaHSO4

Normal salts: all the hydrogen ions from the acid were replaced. They form neutral solutions. Eg.

Na2SO4.

Neutralisation / equivalence point: where the acid is completely neutralized by the base

Endpoint: where a visible change marks the completion of neutralization. This can be either a

colour change using an indicator or a temperature change that is monitored using a thermometer.

Neutralisation Reactions

The reaction between an acid and a base is called a neutralisation reaction. The products are salt

and water. For example:

KOH(aq) + HCI(aq) → KCI(aq) + H2O(l)

One of the ways to monitor the formation of the product (in this case KCl and H2O) is by using an

indicator. Indicators are substances that change colour in acids and bases. One of the most popular

indicators for acid base titrations is phenolphthalein. Phenolphthalein is an indicator that is

colourless in acid and pink in base.

As the base is added dropwise to the reaction mixture, it reacts with the acid; when the reaction

has passed the equivalence point there is a pink colour owing to the excess base that is now left

unreacted in the mixture. This signals the endpoint of the reaction.

The volume of base used is read off the burette and this data is used to calculate the concentration

of the solution.

Example

Finding the concentration of a reactant using titration

To determine the concentration of a solution of sodium hydroxide, Paul titrates the base against

25.0 cm3 portions of hydrochloric acid solution of concentration 8.0 g dm-3 using phenolphthalein

indicator to identify the endpoint. The titration was repeated three times.

Below are the burette readings:

Titration no.

1 2 3

Final burette

reading

15.3 16.6 18.9

Initial burette

reading

0.5 1.9 4.2

Titre (volume of

NaOH)

14.8 14.7 14.7

STEP 1: Determine the volume of sodium hydroxide needed to neutralise 25 cm3 of 8.0 g dm-3

hydrochloric acid.

This is the average volume used.

volume of sodium hydroxide = 14.8+14.7+14.7

3 = 14.7 cm3

STEP 2: Calculate the number of moles of hydrochloric acid used in the titration.

Volume of HCl used = 25 cm3 = 0.025 dm3

Concentration = 8 g dm-3

# of moles = volume X concentration

Molar concentration = 𝑚𝑎𝑠𝑠 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛

𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

Molar concentration = 8𝑔/𝑑𝑚3

36.5 𝑔 /𝑚𝑜𝑙 = 0.2 mol dm-3

# of moles = 0.025 dm3 X 0.2 mol dm-3 = 0.005 mol

STEP 3: Write a balanced equation for the reaction.

HCl + NaOH NaCl + H2O

STEP 4: Determine the number of moles of sodium hydroxide used in the titration.

From the equation: 1 mole of hydrochloric acid reacts with 1 moles of sodium hydroxide

Mole ratio of HCl : NaOH is 1:1

Since the # of moles of HCl was found to be 0.005 moles then the # of moles is also 0.005 moles.

STEP 5: Determine the mass concentration of the sodium hydroxide.

The # of moles of NaOH reacted = 0.005 moles

The volume reacted = 14.7 cm3

Mass concentration = 𝑚𝑎𝑠𝑠

𝑣𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑐𝑚3 x 1000

Convert # of moles NaOH to mass

Mass = # of moles X molar mass

Mass = 0.005 mols X 40 gmol-1 = 0.2 g

=> Mass concentration = 0.2𝑔

14.7 𝑐𝑚3 x 1000 = 13.6 g dm-3

ACTIVITY 10

1. In a titration of sulphuric acid against sodium hydroxide, 32.2 mL of 0.250 moldm-3 NaOH is

required to neutralize 26.6 mL of H2SO4. Calculate the molarity of the sulphuric acid.

2. It takes 38 mL of 0.75 moldm-3 NaOH solution to completely neutralize 155 mL of a sulphuric

acid solution (H2SO4). What is the concentration of the H2SO4 solution?

3. To determine the mole ratio in which alkali X and acid Y react, Susan placed 25 cm3 of alkali

X of concentration 1.0 mol dm-3 in a polystyrene cup and added acid Y of concentration 1.0

mol dm-3 from the burette. She stirred the solution and quickly recorded its temperature after

each 2 cm3 of acid. The thermometer readings are shown below.

Volume of acid added (cm3) Temperature of solution

(oC)

0 29.2

2 30.2

4 31.3

6 32.3

8 33.3

10 34.4

12 35.4

14 35.0

16 34.2

18 33.6

20 32.8

i. Plot temperature against volume of acid Y added and draw TWO straight lines of best

fit.

ii. Use your graph to determine the volume of acid Y needed to neutralise 25 cm3 alkali

X.

iii. Determine the nearest whole number mole ratio in which alkali X and acid Y reacts.

TOPIC: OXIDATION REDUCTION REACTIONS


Oxidation and reduction are opposite processes that occur together in certain reactions. These are

known as redox reactions.

The terms are defined as follows:

OXIDATION REDUCTION

Loss of electrons Gain of electrons

Gain in oxygen Loss in oxygen

Loss of hydrogen Gain of hydrogen

Increase in oxidation number Decrease in oxidation number

Oxidising and reducing agents

During any redox reaction:

• The oxidising agent causes another reactant to be oxidised, this oxidising agent is reduced

in the process.

• The reducing agent causes another reactant to be reduced, this reducing agent is oxidised

in the process.

In the following reaction, X has been oxidised and Y has been reduced:

The equation below shows a simple redox reaction:

Cu(NO3)2(aq) + Mg(s) → Cu(s) + MgO(aq)

The ionic form of the reaction is written as: Cu2+(aq) + Mg(s) → Cu(s) + Mg2+

(aq)

Identifying which species is oxidised and which is reduced by looking at changes in oxidation

state.

In the above reaction, magnesium reduces the copper (II) ion by transferring electrons to the ion

and neutralizing its charge; copper is moving from a +2 oxidation state to zero ie. reduction,

therefore, magnesium is a reducing agent. However, by losing electrons , magnesium is now

oxidised.

Rules for assigning oxidation number

The oxidation state of…

Examples

An atom in an element is zero Na(s) = 0

O2(g) = 0

A monoatomic ion (only one type of atoms) is

the same as its charge

Na+ = +1

Ca2+ = +2

Al3+ = +3

Fluorine is -1 in its compounds F in HF, PF3 = -1

Oxygen is usually -2 in its compounds

EXCEPT in peroxides (containing O22-) in

which oxygen is -1

O in H2O, CO2 = -2

O in H2O2 = -1

Hydrogen is +1 in its covalent compound H in H2O, HCl, NH3

The sum of the oxidation numbers of all the

atoms in a compound is equal to zero.

H2O: [(2 x +1) +(1 x -2)] = 0

The sum of the oxidation numbers in a

polyatomic ion is equal to the charge of the

ion

CrO42-: [(1 x +6) +(4 x -2)] = -2

ACTIVITY 11

1. Classify EACH of the following reactions as either oxidation or reduction. By reference to

electrons, give a reason for your classification in EACH case.

i. Fe3+(aq) + e- → Fe2+

(aq)

ii. Al(s) →Al3+(aq) + 3e-

iii. The formation of the oxide ion (O2-) when oxygen reacts with calcium.

2. Determine the oxidation number in the following:

i. N in NO2 -

ii. S in SO3

iii. Cl in ClO4-

iv. Cr in CrO42-

v. Mn in Mn2O7

vi. Cl in ClO34

vii. S in Na2SO3

viii. V in VO2+

ix. N in N2H4

x. P in H3PO4

3. Determine the change in oxidation number of nitrogen in the following reaction and use

this to decide if the ammonia has been oxidised or reduced.

2NH3(g) + 3CuO(s) →N2(g) + 3Cu(s) + 3H2O(l)

4. State, with reasons based on oxidation number, which reactant has been oxidised and which

has been reduced in EACH of the following reactions.

i. Mg(s) + CuSO4(aq) →MgSO4(aq) + Cu(s)

ii. Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)

iii. Cr+ + Sn4+ → Cr3+ + Sn2+

TOPIC: OXIDATION-REDUCTION REACTIONS


Common oxidising and reducing agents

Some substances always behave as oxidising agents and others always behave as reducing agents.

Others act as oxidising or reducing agents based on what they are reacted with. A visible change

may occur when some of these react.

• A colour change may occur.

• A precipitate may form.

• A particular gas may be produced.

ACTIVITY 12

1. State, with a reason based on oxidation number, if sulphur is acting as an oxidising or

reducing agent in each of the following reactions.

i. S(s) + O2(g) → SO2(g)

ii. Mg(s) + S(s) →MgS(s)

2. Using oxidation number to support your answer, state which reactant is acting as an

oxidising agent and which is acting as a reducing agent in the following reaction.

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

3. Three chemistry students find a bottle of colourless liquid in the laboratory and each makes

a different suggestion about the identity of its contents.

i. Complete the following table to summarise TWO different tests the students

could use to find out whose suggestion is correct.

Test reagent Results of test if:

Josh’s suggestion

is correct

Richard’s

Suggestion is

correct

Matthieu’s

suggestion is

correct

ii. Assuming Matthieu's suggestion is correct, explain the reason for the colour

change he observed and suggest a second reagent he could use to confirm

that he is correct.

4. Name one substance that behaves as both an oxidising agent and a reducing agent.

5. By referring to oxidation and reduction, explain EACH of the following statements.

i. The cut surface of an apple turns brown if the apple is left uneaten on a plate.

ii. Sodium chlorate(I) is a good bleaching agent.

iii. Sodium sulfite is a good preservative of some food items.

iv. Iron nails rust very easily when exposed to moist air.

MINISTRY OF EDUCATION

SECONDARY ENGAGEMENT PROGRAMME

CHEMISTRY

GRADE 10

TOPIC: OXIDATION-REDUCTION REACTIONS

SUB-TOPIC: MISCELLANEOUS QUESTIONS

WEEK FIVE: LESSON 13

ACTIVITY 13

1. Nitrogen has 5 valence electrons (Group V). It can gain up to 3 electrons, or lose up to 5

electrons. Fill in the missing names or formulae and assign an oxidation state to each of

the following nitrogen containing compounds:

Name Formula Oxidation state of N

NH3

Nitrogen

Nitrite ion

NO3-

Dinitrogen monoxide

NO2

Hydroxylamine NH2OH

Nitrogen monoxide

Hydrazine N2H4

2. Do you need an oxidising agent or reducing agent in order for the following reactions to

occur?

i. ClO3- → ClO2

ii. SO42- → S2-

iii. Mn2+ → MnO2

iv. Zn → ZnCl2

3. Given the following redox reaction: Zn2+(aq) + 2Al(s) 3Zn(s) + 2Al3+

(aq)

i. Identify, and write the half reaction for oxidation and the half reaction for reduction.

ii. Name the reducing agent and the oxidizing agent for the reaction

4. Write the half-cell reactions showing oxidation.

i. 2Na(s) + ½ O2(g) → Na2O(s)

ii. Fe(s) + Cu2+(aq) → Fe2+

(aq) + Cu(s)

iii. Sn4+(aq) + Fe2+

(aq) + e- → Sn2+(aq) + Fe3+

(aq)

5. Write the half-cell reactions showing reduction.

i. 2Mg(s) + O2(g) →2MgO(s)

ii. Ca(s) + Cl2(g) → CaCl2(s)

iii. 2H2(g) + O2(g) → 2H2O(l)


1. In a particular redox reaction, the oxidation number of phosphorus changed from -3 to 0.

From this it may be concluded that phosphorus ________________.

A. lost 3 electrons and was reduced.

B. lost 3 electrons and was oxidized.

C. gained 3 electrons and was reduced.

D. gained 3 electrons and was oxidized.

2. Which species is most readily oxidised to Mn2+?

A. Mn

B. MnO2

C. MnO4 –

D. Mn(OH)2

3. Which species is losing electrons in the following redox reaction?

SnO2 + 4Cl- + 4H+ SnCl2 + Cl2 + 2H2O

A. H+

B. Cl-

C. O

D. Sn

4. What is the oxidation number of N in NH2OH?

A. -2

B. -1

C. 0

D. +1

5. Which of the following represents a redox reaction?

A. 2HCl + Na2SO3 → 2NaCl + H2O + SO2

B. CuS+ H2→H2S + Cu

C. AgNO3+NaCl →AgCl + NaNO3

D. H2CO3 → H2O + CO2

6. Consider the following reaction: 2MnO4 + 5H2SO3 → 2Mn2+ + 3H2O + 5SO42- + 4H+

The species that undergoes reduction is:

A. S

B. H

C. O

D. Mn

7. Copper has an oxidation number of +1 in:

A. Cu (CH3COO)2

B. CuBr

C. CuC2O4

D. CuO

8. When ClO3- is oxidized, a possible product is:

A. Cl-

B. ClO-

C. ClO2-

D. ClO4-

9. Which of the following is an equation representing a redox reaction?

A. 2NO2(g) → N2O4(g)

B. Mg(s) + Cl2(g) → MgCl2(s)

C. Ag+(aq)+ C-

(aq) → AgCl(s)

D. NH2 (aq)+ H+ (aq)→NH4+(aq)

10. Consider the following:

O3(g) + H2O(l) + SO2(g) →SO42-

(aq) + O2(g) + 2H+(aq)

In the redox equation above, the chemical species which is oxidised is:

A. H+

B. O3

C. SO2

D. SO42-

11. A product of the oxidation of NO2 is:

A. NO

B. N2O

C. NO2-

D. NO3-

12. Electrons are lost by the __________.

A. reducing agent as it undergoes oxidation

B. reducing agent as it undergoes reduction

C. oxidizing agent as it undergoes oxidation

D. oxidizing agent as it undergoes reduction

TOPIC: ELECTROCHEMISTRY


Definitions:

The conductor is defined as the material which allows the electric current or heat to pass through

it. The electrons in a conductor freely move from atom to atom when the potential difference is

applied across them.

Metallic conductance involves the movement of electrons throughout a metal.

Electrolytic conduction involves the movement of ions throughout a pure liquid or solution. The

major difference between them is that one involves the movement of electrons and the other

involves the movement of ions.

Strong electrolytes ionize completely, while weak electrolytes ionize only partially.

Electrolytes are compounds that conduct an electric current and are decomposed by it.

An Electrochemical cell is a device that can generate electrical energy from chemical reactions

occurring in it or use the electrical energy supplied to facilitate chemical reactions in it.

ACTIVITY 14

1. Classify EACH of the following substances as a conductor or a non-conductor.

i. Graphite

ii. Solid sodium chloride

iii. Copper (II) sulphate solution

iv. Aqueous nitric acid

v. Carbon tetrachloride

vi. Sulphur


1. What is the difference between a conductor and an insulator?

B. An insulator allows electricity to flow through it easily and a conductor does not

C. A conductor allows electricity to flow through it easily and an insulator does not.

D. An insulator is magnetic and a conductor is not

E. A conductor is magnetic and an insulator is not

2. Conductors are materials that _____________.

A. Allow heat to pass through

B. Stop heat from passing through

C. Allow cold to pass through

D. Stop cold from passing through

3. What type of materials allow electricity to flow freely?

A. Conductors

B. Currents

C. Insulators

D. Inductors

5. Conductance is the inverse of what measurement?

A. Inductance

B. Resistance

C. Voltage

D. Capacitance

E. Current

6. Which of the following best defines an electrolyte?

A. A compound that remains neutral in solution

B. A compound that is only positively charged in a given solution

C. A substance that dissociates into ions in solution

D. A substance that doesn't conduct electricity

7. Which of the following describes a weak electrolyte?

A. A substance that partially breaks apart into ions in solution

B. A substance under which weak bases and acids can be classified

C. A compound that dissociates roughly 1-10% in solution

D. All of the answers are correct.

13. In an electrolytic cell the electrode at which the electrons enter the solution is called the

______ and the chemical change that occurs at this electrode is called _______.

A. anode, oxidation

B. anode, reduction

C. cathode, oxidation

D. cathode, reduction


SUBTOPIC: ELECTROLYSIS


Definitions:

Electrolysis is the decomposition of an electrolyte by the passage of an electric current through it.

Electrodes are the points where current enters and leaves an electrolyte. The anode is the positive

electrode and the cathode is the negative electrode.

When electrolytes conduct electricity, they are broken down (decomposed) by the

current. This is the process known as electrolysis and carried out in an

electrochemical cell as shown above.

The electrodes are basically the conducting rods and are typically made from

graphite or platinum. These electrodes are typically inert meaning they do not

participate in the chemical reaction taking place. However, electrodes can also be

active, i.e., where they participate in the chemical reaction.

During electrolysis:

The negative ions (anions) move towards the anode and the positive ions (cations)

move towards the cathode (as can be seen in the diagram above).

The negative ions lose electrons to the anode and become neutral atoms. In other

words, oxidation occurs at the anode.

The positive ions gain electrons from the cathode and become neutral atoms. In

other words, reduction occurs at the anode.

These neutral atoms are said to be discharged at their respective electrodes.

ACTIVITY 15

1. During the electrolysis of dilute sulphuric acid using inert electrodes, oxygen and hydrogen

gas are produced.

i. Identify the ions that are being produced.

ii. What specific compounds contribute to these ions?

iii. Which ions are attracted to the anode?

iv. Which ions are attracted to the cathode?

v. Write the equation for the reaction that occurs at the anode.

vi. Write the equation for the reaction that occurs at the cathode.

vii. Calculate the volume of the gases that will be produced.

MULITPLE CHOICE QUESTIONS

1. The half-reaction that occurs at the anode during the electrolysis of molten sodium bromide

is:

A. 2Br- Br2 + 2 e-

B. Br2 + 2 e- 2 Br-

C. Na+ + e- Na

D. Na Na+ + e-

1. Identify the ions present in molten sodium bromide.

A. Na+ and Br-

B. Na2+ and Br2-

C. Pb2+, Br2- and H+

D. Na- and Br

2. Identify the ions present in aqueous sodium bromide.

A. Na+ and Br-

B. Na2+ and Br2-

C. Pb2+, Br2- and H+

D. Na+ and Br-, OH-, H+

4. In an electrochemical cell, the cathode __________.

A. is reduced.

B. loses mass.

C. is the reducing agent.

D. is the site of reduction.

5. In an operating electrochemical cell the function of a salt bridge is to __________.

A. allow hydrolysis to occur.

B. allow a non-spontaneous reaction to occur.

C. permit the migration of ions within the cell.

D. transfer electrons from the cathode to the anode.

6. In an operating zinc-copper electrochemical cell, the oxidizing agent _________.

A. loses electrons at the anode.

B. loses electrons to the cations.

C. gains electrons at the cathode.

D. gains electrons from the anions.

7. Which of the following occurs as the cell operates?

A. Zinc electrode is reduced and increases in mass.

B. Zinc electrode is reduced and decreases in mass.

C. Zinc electrode is oxidized and increases in mass.

D. Zinc electrode is oxidized and decreases in mass.


SUB-TOPIC: ELECTROCHEMICAL SERIES

WEEK SIX: LESSON 16

The order of reactivity, with the most reactive at the top is called the electrochemical series

Hydrogen is also included in this series to show which metals displace hydrogen from acids.

As we GO UP the electrochemical series the metals:

• increase in reactivity.

• lose electrons more readily, so form positive ions more readily.

• become stronger reducing agents.

In electrolysis, only one type of cation or anion is discharged. This is called preferential discharge

of ions. There are three factors determining this:

1. The position of the ion in the electrochemical series. Ions lower in the electrochemical

series are discharged in preference to the ones above them. So, if Cu2+ and H+ ions are

present, Cu2+ ions are discharged at the cathode. We can also arrange anions in a discharge

series. In concentrated aqueous sodium chloride, Cl- ions are discharged in preference to

OH- ions.

2. The concentration of the solution. For anions, the most concentrated ion tends to get

discharged in preference to the less concentrated ion. So, Cl - ions are discharged in

preference to OH- ions when a concentrated aqueous solution of sodium chloride is

electrolysed. But if the sodium chloride solution is dilute, the OH- ion is discharged in

preference.

3. Inert or inactive electrodes. Graphite or platinum electrodes do not take part in the

chemical reaction ie. they are inert. Active electrodes, eg. copper, takes part in the reaction.

ACTIVITY 16

1. The diagram below shows the electrolysis of molten lead (II) bromide using inert

graphite electrodes.

i. Label W, X, Y and Z.

ii. Write ionic equations to show the formation of X and Z.

2. The electrochemical series of metals can be used to predict various chemical reactions.

i. Metal F is above zinc in the electrochemical series. Would you expect F to react with

zinc sulphate solution? Give a reason for your answer.

ii. Metal G is below hydrogen in the electrochemical series. Would you expect G to react

with hydrochloric acid? Give a reason for your answer.

iii. Write an equation for the reaction between magnesium and copper (II) sulphate solution.


SUB-TOPIC: QUANTITATIVE ELECTROLYSIS

WEEK SIX: LESSON 17

The mass of a substance produced at the electrodes (or consumed at a reactive anode) during

electrolysis is derived from the following equations:

Mass is directly proportional to electric charge ie. m α Q

Electric charge (Q) = Current (I) x time (t)

Q = It

Charge is measured in coulombs, C; current is measured in amperes, A and time is measured in

seconds, s.

The Faraday constant, F, is the quantity of electric charge carried by one mole of electrons or one

mole of singly charged ions.

F = 96,500 coulombs per mole (Cmol-1)

ACTIVITY 17

1. A steady current of 2.5 Amperes flows for 2 hours, 8 minutes and 40 seconds through dilute

sodium chloride solution.

i. Calculate the quantity of electricity flowing.

ii. Write an equation for the reaction occurring at the anode.

iii. Determine the number of moles of oxygen produced at the anode.

2. How long must a steady current of 5·0 A flow through dilute sulphuric acid in order to

produce 3.0 g of hydrogen at the cathode?

3. Calculate the increase in mass of a spoon if a current of 2.0 Amperes is allowed to flow for

32 minutes and 10 seconds through the electrode.


SUB-TOPIC: APPLICATIONS OF ELECTROLYSIS

WEEK SIX: LESSON 18

1. Purification of metals

Many metals can be purified by electrolysis. The impure metal serves as the anode and a thin sheet

of pure metal is the cathode. The electrolyte is a soluble salt of the pure metal.

E.g., In the purification of copper, the copper atoms at the anode lose electrons and form Cu2+ ions

according to the following equation:

Cu(s) →Cu2+(aq) + 2e-

The anode becomes thinner and the impurities fall to the bottom of the cell as an anode sludge

whereas the copper ions at the cathode gain electrons and form copper atoms according to the

following equation:

Cu2+(aq) + 2e- →Cu(s)

The cathode becomes thicker because the pure metal is deposited on it.

2. Electroplating

Electroplating involves coating of the surface of one metal with a layer of another, usually less

reactive, metal. Metals are electroplated because it makes them more resistant to corrosion, e.g.

chromium plating, nickel plating as well as improves their appearance, e.g. plating with silver.

In electroplating, the anode is the pure metal and the cathode is the metal to be electroplated

whereas the electrolyte is a soluble salt of the pure metal at the anode.

In silver plating, silver ions (Ag+) are formed at the anode from silver atoms (Ag). The Ag+ ions

accept electrons from the cathode and become silver atoms. These form the layer (silver plating)

on the cathode.

3. Anodising

Anodising is the process of increasing the thickness of an unreactive oxide layer on the surface of

a metal. It is used to reduce the reactivity of metals, such as nickel or aluminium, so that they can

be used under a variety of conditions to increase corrosion resistance and reduce wear.

The anode is the metal. When anodising aluminium, the thin oxide layer normally present on the

surface of the metal is first removed by reaction with sodium hydroxide. The cathode is usually

unreactive, e.g. carbon and the electrolyte is sulphuric acid. During the reaction, the sulphuric acid

is electrolysed to form oxygen and hydrogen.

Oxygen gas is produced at the anode according to the following equation:

4OH-(aq) → O2(g) + 2H2O(l) + 4e-

The oxygen gas reacts with the anode and forms a thick oxide layer according to the following

equation:

4Al(s) + 3O2(g) → 2AI2O3(s)

ACTIVITY 18

1. Suggest TWO reasons for anodising a saucepan.

2. A student wishes to demonstrate the principle of electroplating a spoon with silver to his

fellow students. What should he use as the anode, cathode and electrolyte?

3. Explain why electrolysis is NOT suitable for purifying metals above hydrogen in

the electrochemical series.


SUB-TOPIC: MISCELLANEOUS QUESTIONS

WEEK SEVEN: LESSON 19

ACTIVITY 19

1. Find the quantity of electricity that results from a current of 3.50 amperes flowing for 6

minutes.

2. Calculate the mass of zinc plated onto the cathode of an electrolytic cell by a current of

750 milliamperes in 3.25 hours.

3. Calculate the mass of silver deposited at the cathode during the electrolysis of silver nitrate

solution if you use a current of 0.1 amperes for 10 minutes.

4. Calculate the volume of hydrogen produced (measured at room temperature and pressure,

RTP) during the electrolysis of dilute sulphuric acid if you use a current of 1.0 ampere for

15 minutes.


1. In the anodizing of an aluminum pot, which of the following is true?

A. H2 is given off at the anode

B. A layer of aluminum hydroxide forms on the pot

C. The aluminum pot is the anode in the cell

D. The electrolyte is a solution of sodium chloride

2. Which of the following metals can be extracted by electrolysis?

A. Potassium

B. Calcium

C. Aluminum

D. Iron

TOPIC: RATES OF REACTION


The rate of a reaction is the change in concentration of reactants or products in unit time at a given

temperature.

Collision theory explains how various factors affect rates of reaction. According to this theory,

chemical reactions can occur only when reacting particles collide with each other and with

sufficient energy as well as the correct orientation. The minimum amount of energy that particles

must have to react is called the activation energy.

Factors which affect the rate of chemical reactions include:

1. Concentration of reactants: An increase in concentration means more molecules are present

in a given volume, therefore the probability of more frequent effective collisions occurring

increases thus leading to an increased rate of reaction. The reverse occurs when there is a

decrease in concentration.

2. Pressure of reacting gases: increasing pressure has the same effect as increasing concentration

since the same amount of molecules are now present in a smaller volume.

3. Surface area of solid reactants: increasing the surface area causes an increase in rate of

reaction since more surfaces are now exposed for reacting particles to collide with.

4. Temperature: increasing the temperature of the reaction causes an increase in the energy of

the reactant molecules. The molecules can now move faster which increases the frequency of

collisions leading to increased rate. Additionally, more molecules will have energies that are

greater than or equal to the activation energy so more effective collisions will occur. The

reverse occurs when there is a decrease in temperature.

5. Presence of catalysts: catalysts increases the reaction rate by providing a different path (with

a lower activation energy) for the reaction to occur.

ACTIVITY 20


1. As the temperature of a reaction is increased, the rate of the reaction increases because:

A. reactant molecules collide less frequently

B. reactant molecules collide more frequently and with greater energy per collision

C. activation energy is lowered

D. reactant molecules collide less frequently and with greater energy per collision

2. The rate of a reaction depends on __________.

A. collision frequency

B. collision energy

C. collision orientation

D. all of the above

3. Of the following, __________ will lower the activation energy for a reaction.

A. increasing the concentration of reactants

B. raising the temperature of the reaction

C. adding a catalyst to the reaction

D. removing products as the reaction proceeds

4. The minimum amount of energy needed to start a reaction is called the

A. activation energy.

B. energy of reaction.

C. entropy of reaction.

D. reaction mechanism energy

5. A catalyst increases the rate of a reaction by __________.

A. increasing the concentration of reactant/s.

B. decreasing the concentration of the reactant/s.

C. increasing the activation energy of the overall reaction.

D. decreasing the activation energy of the overall reaction.

6. Which of the following would NOT increase the rate of reaction?

A. raising the temperature

B. adding a catalyst

C. increasing the concentration of the reactants

D. increasing the volume of the container

7. The rate of a chemical reaction can be expressed in ______.

A. grams per mole.

B. energy consumed per mole.

C. volume of gas per unit time.

D. molarity per second.

8. When the concentration increases, the rate of reaction increases because the ______ of

collisions increases.

9. When the temperature increases, the frequency of collisions increases and so does the

______ of the collisions.

TOPIC: RATES OF REACTION

SUBTOPIC: MISCELLANAEOUS QUESTIONS


ACTIVITY 21

1. During any reaction, reactants are used up and the rate of reaction decreases.

Explain, in terms of particles, why the rate of reaction decreases.

2. To determine how the rate of a reaction varies as the reaction proceeds, Keenan reacted

calcium carbonate crystals with excess hydrochloric acid and measured the volume of

carbon dioxide produced every 30 seconds. His results are in the table below.

i. Draw a labelled diagram to show how Keenan’s experimental setup.

ii. Plot the results on a graph.

TOPIC: ENERGETICS

WEEK EIGHT: LESSON 22

Exothermic Reactions are chemical changes that result in the increase of the temperature of the

surroundings.

Endothermic reactions are chemical changes that result in a fall of the temperature of the

surroundings.

Energy changes occur during the course of chemical reactions, as reactants form new products.

This energy change is represented by ΔH (read as delta H). Energy change measured at constant

pressure is referred to as enthalpy change and is also represented as ΔH.

ΔH is the difference between the energy of the products (Hp) and the energy of the reactants (Hr).

ΔH = Hp – Hr

If Hp is less than Hr, then ΔH is negative and the reaction is exothermic.

If Hp is greater than Hr, then ΔH is positive and the reaction is endothermic.

ACTIVITY 22

1. There are two types of chemical reactions based on energy changes occurring. Explain how

you can differentiate them.

2. During a chemical reaction, bonds are broken and bonds are formed, explain in terms of

energy what happens when bonds are broken and bonds are formed.


1. The diagram shows an energy profile diagram. What does the energy value of 1370 kJ

represent?

A. Activation energy

B. Product energy

C. Reactant energy

D. Released energy

3. Magnesium has to be heated before it reacts with oxygen in the air. What conclusion can

you draw from this?

A. The reaction is reversible

B. The reaction has a high activation energy

C. The reaction is exothermic

D. Magnesium has a high melting point

4. Which of the following statements correctly explains why a reaction is endothermic?

A. More energy is released when breaking reactant bonds than is absorbed when making

product bonds

B. More energy is absorbed when breaking reactant bonds than is released when making

product bonds

C. Less energy is absorbed when breaking reactant bonds than is released when making

product bonds

D. Less energy is released when breaking reactant bonds than is absorbed when making

product bonds

TOPIC: ENERGETICS


During chemical reactions bonds are broken and formed. Energy is used up to break

chemical bonds and released when new bonds are formed.

ENERGY PROFILE DIAGRAMS

NB: In all chemical reactions old bonds must be broken before new ones are formed.

For this reason, reactants must be supplied with energy.

ACTIVITY 23

1. When solid ammonium chloride is shaken with water, a colourless solution forms and the

temperature changes from 20°C to 16°C. Name the type of heat change occurring.

2. The reaction between nitrogen and hydrogen to form ammonia is exothermic and is

catalysed by finely divided iron. Draw energy profile diagrams for this reaction

TOPIC: ENERGETICS


Definitions:

The heat of neutralisation is the energy change per mole of water formed during

neutralisation of an acid by a base.

The heat of solution is the energy change when one mole of solute dissolves in a

particular volume of solvent to form a very dilute solution.

Specific heat capacity is the quantity of heat (in joules) required to raise the

temperature of a unit mass or a unit volume of substance by 1 degree Celsius (oC)

or 1 kelvin (K).

Calculating heat change

Heat change (ΔH) = mass (m) x specific heat capacity (c) x change in temperature

(ΔT)

ΔH = mcΔT

ACTIVITY 24

1. When 25 cm3 of hydrochloric acid of concentration 1.0 mol dm-3 is added to 25cm3 of

potassium hydroxide of concentration 1.0 mol dm-3, the temperature rises from 21. 1 °C to

27.3 °C. Calculate the heat of neutralisation for this reaction.

2. 50 cm3 of sodium hydroxide solution with a temperature of 29.4 oC and concentration of

2.0 mol dm3 is added to 50 cm3 of sulphuric acid of concentration 1.0 mol dm3 and

temperature 30 oC. The maximum temperature of the solution after mixing is 43.2 oC.

Determine the heat of neutralisation.

3. Dissolving 15.15 g of potassium nitrate in 100 cm3 of distilled water resulted in a

temperature decrease of 10.2 0C.

i. Determine the number of moles of KNO3 dissolved

ii. Determine the heat absorbed by dissolving 0.15 mol

iii. Calculate the heat of solution of potassium nitrate.

TOPIC: PRACTICE CSEC QUESTIONS

WEEK NINE: LESSON 25


1. Which of the following processes provide evidence of the particulate nature of matter?

I. Diffusion

II. Filtration

III. Osmosis

(A) 1 and 11 only

(B) 1 and 11I only

(C) 1I and 11I only

(D) I, II and III

Item 2 refers to the following table, which gives the melting points and boiling points of

the chlorides of four elements.

Element Melting Point Of

Chloride

(℃ )

Boiling Point Of

Chloride

(℃)

I 463 453

II 248 348

III 973 1693

IV 193 333

2. Which of the elements form a chloride that is ionic in nature?

(A) I

(B) II

(C) III

(D) IV

3. The arrangements of electrons in atoms of X and Y are 2,8,5 and 2,8,6 respectively. Which

of the following options represents X and Y?

X Y

(A) Metal Nonmetal

(B) Nonmetal Nonmetal

(C) Nonmetal Metal

(D) Metal Metal

4. Which of the following BEST describes the formation of a metallic bond?

A metallic bond is formed when

(A) Anions are held together by negative electrons.

(B) Metal atoms are held together by molecular forces.

(C) Cations are held together by a sea of mobile electrons.

(D) Positive metal ions are held together by a sea of anions.

5. Which of the following elements does NOT form simple ions by gaining or losing

electrons?

(A) Carbon

(B) Copper

(C) Calcium

(D) Chlorine

6. An element, X, has 8 electrons in its outer shell. Which of the following are possible

isotopes of X?

(A) (C)

8p

8n

8e-

8p

9n

8e-

8p

8n

8e-

8e-

8p

10n 9p

8n

(B) (D)

7. “Mass Number” is the number of

(A) Neutrons minus Protons

(B) Electrons plus Neutrons

(C) Neutrons plus Protons

(D) Electrons plus Protons

8. The mass of “1 mole of atoms of an element” refers to the quality of

(A) 1 atom of the element

(B) An element which contains 6.0 x 1023 atoms

(C) An element which occupies 24.0 dm3 at STP

(D) An element which combines completely with 12g of carbon -12

8e-

9p

11n

8e-

9p

12n

8e-

8e-

9p

8n

12p

8n

Item 9 refers to the following table

Particle Number of

Protons

Number of

electrons

Number of

neutrons

V 8 8 8

W 16 16 16

X 8 8 9

Y 16 18 16

9. Which of the following particles represents an anion?

(A) V

(B) W

(C) X

(D) Y

Items 10-11 refer to the following types of substances.

In answering items 11-12, each option may be used once, more than once or not at all.

10. Which of the above substances can be described as the oxide of a metal?

11. Which of the substances above supplies protons as the ONLY positive ions in aqueous

solutions?

12. Which of the following statements is true of an endothermic reaction?

(A) Heat is given up to the surroundings.

(B) Heat is absorbed from the surroundings

(C) The products have less energy than the reactants

(D) The products have the same energy as the reactants

13. The mass concentration of a potassium chloride solution is 60g dm-3. What is the mass of

potassium chloride in 25cm3 of this solution?

(A) 0.0015g

(B) 0.15g

(C) 1.5g

(D) 15.0g

14. When solid lead nitrate is heated, it decomposes giving off nitrogen (IV) oxide and oxygen.

The balanced equation for this reaction is

(A) 2Pb(NO3)2 (s)------> 2PbO(s) + 4NO(g) + O2(g)

(B) Pb(NO3)2 (s)------>PbO(s)+ NO2(g)+ O2(g)

(C) Pb(NO3)2 (s)------> PbO(s)+ 2NO2(g)+ O2(g)

(D) 2Pb(NO3)2 (s)------> 2PbO(s)+ 2NO2(g)+ O2(g)

(E)

Item 15 refers to the following information.

Element Atomic

Number

I 3

II 6

III 17

IV 18

15. Which two elements above, when combined, form MAINLY ionic compounds?

(A) 1 and II

(B) I and III

(C) II and IV

(D) III and IV




1. The equation for the reaction of silicon with chlorine is Si(s) + 2Cl2(g)

SiCl4(l)

In this reaction, silicon is the reducing agent and is

A. Oxidized with an increase in oxidation state

B. Oxidized with a decrease in oxidation state

C. Reduced with an increase in oxidation state

D. Reduced with a decrease in oxidation state

2. Item 2 refers to one mole of each of the following acids

I. H2SO4

II. CH3COOH

III. (COOH)2

Which of the acids above would require more than one mole of NaOH for complete

neutralisation?

A. II only

B. I and III only

C. II and III only

D. I, II and III

3. The atomic number of an element is defined as the number of

A. Electrons

B. Protons

C. Protons and neutrons

D. Electrons and neutrons

4. Which of the following statements is true for the equation below?

Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g)

A. Hydrogen is oxidised from 0 to +1

B. Hydrogen is reduced from +2 to 0

C. Mg is oxidised from 0 to +2

D. Mg is reduced from +2 to 0

5. Which of the following factors will usually increase the rate of catalytic

decomposition of hydrogen peroxide?

A. Increasing the temperature

B. Applying a greater external pressure

C. Using a larger volume of hydrogen peroxide

D. Using the catalyst in lumps rather than in powdered form

6. In which of the following substances does hydrogen have a negative

oxidation number?

A. CH4

B. H2O2

C. NH3

D. NaH

7. In the electrolysis of aqueous copper (II) sulphate solution using copper

electrodes, the anode

A. Decreases in mass

B. Increases in mass

C. Undergoes no change in mass

D. Is the site where reduction occurs

8. Which of the following energy changes occur during the breaking of a

chemical bond?

A. Energy is released

B. Energy is required

C. There is no energy change

D. The process is exothermic

9. Glucose is converted to starch or cellulose by

A. Dehydrogenation

B. Oxidation and reduction

C. Addition polymerization

D. Condensation polymerization

10. Which of the following halogens is a liquid at room temperature?

A. Chlorine

B. Bromine

C. Fluorine

D. Iodine

11. Which of the following substances conducts an electric current and remains

chemically unchanged?

A. Aqueous copper (II) sulphate

B. Aqueous sodium chloride

C. Sulphur

D. Copper

Items 12 – 14 refer to the following metals. Each option may be used once,

more than once or not at all.

A. Na

B. Fe

C. Cu

D. Al

12. Which metal does NOT react with an acid to produce hydrogen gas?

13. Which metal is covered with a passive layer of oxide?

14. Which metal reacts readily with water to produce a strongly alkaline

solution.

15. Which of the following mixtures can be referred to as a standard solution?

(A) Iodine in 50 cm3 of ethanol

(B) 30 g of sodium chloride in water

(C) Sodium chloride in 50 cm3 of water

(D) 30 g of iodine in 50 cm3 of ethanol




Items 1-2 refer to the following terms.

A. Ionic Crystals

B. Simple Molecular

C. Macromolecular

D. Metallic

In answering items 1-2, each term may be used once, more than once or not at all.

1. Which of the terms above describes the structure of sodium chloride?

2. Which of the terms above describes the structure of copper?

3. Which of the following oxides react with both acids and bases?

A. MgO

B. CuO

C. Al2O3

D. CaO

4. Sulphur and oxygen are in the same group in the periodic table because

A. They can react with each other

B. The atomic number of sulphur is 16 and the relative atomic mass of oxygen is 16

C. They have the same number of electrons in their outer shell

D. They can form covalent compounds

5. Which of the following statements illustrates Brownian motion?

A. The random motion of pollen dust in water

B. Perfume scent throughout the air in a room

C. The swelling of red beans when soaked in water

D. Loss of heat from a hot body to a cold body

6. Radioactive isotopes are NOT normally used in the

A. Determination of the age of fossils

B. Treatment of cancer

C. Treatment of influenza

D. Powering of certain types of submarines

7. Sodium reacts with water according to the equation:

2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

The number of litres of hydrogen gas liberated when 0.1 mole of sodium reacts with

excess water is

A. 1.2

B. 2.4

C. 12

D. 24

8. The values of x and y respectively in the equation

x KOH + 3Br2 y KBr + KBrO3 + 3H2O

A. x = 5; y = 6

B. x = 6; y = 5

C. x = 5; y = 5

D. x = 6; y = 6

9. The number of shared electron pairs in a methane molecule is

A. 4

B. 6

C. 8

D. 10

10. How many covalent bonds are there in a nitrogen molecule?

A. 1

B. 2

C. 3

D. 4

11. A separating funnel can be used to separate a mixture of

A. Water and sodium chloride

B. Water and ethanol

C. Water and kerosene

D. Kerosene and sodium chloride

12. Which of the following atoms would NOT form a positive ion?

A. Magnesium

B. Aluminium

C. Sodium

D. Chlorine

13. Which of the following substances is the oxide of a metal?

A. Salt

B. Base

C. Alkali

D. Acid

14. Which of the following is an acid salt?

A. Na3PO4

B. Na2SO4

C. NaHSO4

D. Na2CO3

15. Which of the following statements about ionic compounds is true?

A. They contain molecules

B. They are solids and vaporize easily

C. They usually dissolve in organic solvents

D. They conduct electricity when melted or molten


WEEK TEN: LESSON 28

1.

a. Water exists in three states of matter while iodine exists in two.

i. List the THREE states of matter in which water exists.

ii. Describe the strength of the forces of attraction present between the particles in

EACH of the three states you have mentioned in (a) (i) above.

iii. When heated, iodine changes from one state into another. What is this process

called?

iv. Describe how the energy of the iodine particles changes as iodine undergoes

the process mentioned in (a) (iii) above.

b. Using appropriate diagrams, illustrate how the bonding in solid sodium chloride

differs from that of diamond.

c. Describe TWO tests that are performed in the laboratory to distinguish between an

‘ionic solid’ and a ‘molecular solid’. Suggest how the results of the tests described

can be used to distinguish between the two solids.

2.

a. Tums® and Epsom salts are items commonly found in most household medicine

cabinets. Calcium carbonate is the main active ingredient in Tums®, an antacid used

to relieve heartburn, acid indigestion and upset stomach.

i. Describe ONE method for the preparation of dry calcium carbonate in the

laboratory. In your answer, include an equation for the reaction as well as the

steps that are involved in its preparation.

ii. The main ingredient in Epsom salts is magnesium sulfate. List ONE use of

Epsom salts.

iii. In order to effectively use Epsom salts, it is usually made into a solution.

Explain why water molecules are able to dissolve Epsom salts.

b. When magnesium ions are present in natural water, it is referred to as ‘permanent

hard water’. Describe how permanent hard water is formed. Include balanced

chemical equations with state symbols to illustrate this process.


WEEK TEN: LESSON 29

1.

a. Sally and Ann live on opposite sides of an island. When Sally visits Ann she

observes that the soap takes longer to lather and produces more scum.

i. Explain to Sally why the soap may be producing more scum at Ann’s house

than at her house.

ii. Do you expect that Sally would get the same result if she uses soapless

detergent? State a reason for your answer.

b. Second generation detergents contained phosphates but their use was banned

because of the effect they had on the environment, particularly rivers and streams.

Outline the harmful effect that second generation detergents had on the environment.

c. In recent years, emphasis has been placed on preserving the environment and as such

a new area of chemistry, Green Chemistry, has evolved.

i. What is meant by the term ‘Green Chemistry’?

ii. Discuss TWO benefits of utilizing the principles involved in Green Chemistry.

2.

a. The figure below is a flow diagram of the industrial processing of sugar cane to

produce crystalline sucrose. Study the figure carefully and answer the questions

which follow.

i. Identify Process P and Process Q.

ii. State the importance of the centrifugation process.

iii. Identify Product X.

iv. The bagasse produced is used in the factory during the processing of sugar cane.

In which part of the factory is this bagasse used and what is it used for?

b. Many Caribbean islands are renowned for the quality of rum they produce. The

alcohol content of a typical rum averages between 40% and 55%.

i. During the fermentation process in the making of rum, yeast feeds on the

sucrose in molasses converting it into simpler sugars which are then converted

to ethanol. Outline, using balanced equations, the formation of ethanol from

sucrose.

ii. Jemina was presented with a flask that contains a mixture of diluted rum. Draw

a labelled diagram of the apparatus she should use in the laboratory to obtain a

concentrated sample of ethanol.

c. After accidentally leaving a bottle of wine open for several days, Jemina found that

the wine tasted slightly sour. She was given magnesium oxide to react with a sample

of the sour wine. Suggest the type of reaction that takes place when the magnesium

oxide is mixed with the sour wine. Write a balanced chemical equation for this

reaction.


WEEK TEN: LESSON 30

1. Compounds A and B are two colourless liquids. They both contain the elements

carbon, hydrogen and oxygen.

i. Given that 4.0 g of B contain 60% carbon, 13% hydrogen and 27% oxygen.

Calculate the empirical formula of B.

ii. The empirical formula of A is CH2O. Deduce whether or not A and B belong

to the same homologous series. Explain your answer.

iii. Calculate the molecular formula for A and B.

iv. Compounds A and B both react with sodium. Write the fully displayed

structural formulae for A and B.

2.

i. What will happen when the bulb is replaced with pure water?

ii. Explain your answer to part (i) above.

iii. It is observed that when the electrolyte in the cell above is dilute

sulphuric acid, the ratio of the gases collected in A and B is

approximately 2:1. Identify the gases in tubes A and B.

iv. By considering the reactions that occur at the electrode surfaces,

explain the occurrence of the 2: 1 ratio in the gases collected in

A and B. Use balanced ionic equations to support your answer.

v. If the bulb in the figure above is replaced with an ammeter, what

will happen to the reading on the ammeter if the electrolyte is

changed from dilute sulphuric acid to dilute methanoic acid.

Explain your answer.

vi. What adjustments have to be made to the figure above in order

to copper-plate a spoon?

CHEMISTRY GRADE 10

TOPIC: MISCELLANEOUS CSEC QUESTIONS

WEEK ELEVEN: LESSON 31

ACTIVITY 31

A student is required to investigate the rate of reaction in which a fixed mass of magnesium metal (0.12

g) is added to different volumes of 1.5 M hydrochloric acid. The acid is added from a burette and water

added to make the final volume of 50 cm3. The time taken for the magnesium ribbon to disappear is

recorded. Figure 1 below shows the burette readings for the volume of acid added and the time taken for

the magnesium to disappear for each reaction. The initial burette reading is always 0.0 cm3.

a. From the results shown in Figure 1 construct a table to show experiment number, volume of

acid added from the burette, volume of water added to the acid, and time taken for the

magnesium to disappear.

b. Using a graph paper, plot a graph of time taken for the magnesium ribbon to disappear against

volume of acid added from the burette.

c. Explain the shape of the graph.

d. Using the data from the graph determine the time it would take for 25 cm3 of the acid to react

with the magnesium ribbon.

e. Write a balanced equation for the reaction of magnesium with hydrochloric acid.

f. Calculate EACH of the following:

i. The number of moles of Mg in 0.12 g. (Relative Atomic Mass of Mg = 24)

ii. The volume of hydrogen gas produced at r.t.p. when all the magnesium ribbon reacts

with the acid. (1 mole of gas at RTP occupies 24 dm3)

g. Explain why it is necessary, for each experiment, to make up the volume of acid to 50 cm3 by

adding water.



ACTIVITY 32

The following instructions are given to students in the laboratory:

Write up the following parts of the report that the student would most likely present after

completing the above experiment.

a. Aim

b. Method

c. Discussion of results

i. Data to be collected

ii. Steps for doing calculations

iii. Discussion of results as it relates to the aim.

You are provided with 2M solutions of sodium hydroxide and hydrochloric acid,

measuring cylinders, a polystyrene cup and a thermometer.

1. Measure out 50 cm3 of the 2 M sodium hydroxide and 50 cm3 of the 2M solution of

hydrochloric acid.

2. Pour the sodium hydroxide solution into the polystyrene cup and measure its

temperature.

3. Add the acid to the sodium hydroxide. Stir the mixture and record the highest

temperature.

Note: • Assume 1 cm3 of water has a mass of 1 g.

• It takes 4.2 joules of energy to raise the temperature of 1 g of water by 1° C.



ACTIVITY 33

Figure 2 shows the labels from two bottles containing Suspensions P and Q respectively, showing

the formulae and percentage composition of the active ingredients.

a. Calculate the number of moles of AI(OH)3 in 250 g of Suspension P. (Relative Atomic

Mass: Al = 27, 0 = 16, H = 1)

b. A student attempts to prepare a sample of Suspension Q in the laboratory by reacting

calcium with dilute sulphuric acid. The reaction stops after a short while with only a small

amount of CaSO4 formed and most of the calcium unreacted.

i. Explain why the reaction stops after a while.

ii. Outline a suitable laboratory method for preparing a dry sample of calcium

sulphate. Include a relevant equation in your answer.

c. Which of the Suspensions, P or Q, could serve as an antacid? Explain your answer.

d. By mixing hot water and a concentrated aqueous solution of iron (Ill) chloride, a bright

yellow colloid is formed. State THREE ways in which a colloid differs from a suspension.


WEEK TWELVE: LESSON 34

ACTIVITY 34

The figure below shows the arrangement of an apparatus that can be used for electrolysing a

number of electrolytes.

a. What will happen to the bulb when the electrolyte is replaced with pure water? Explain

your answer.

b. It is observed that when the electrolyte in the cell above is dilute sulphuric acid, the ratio

of the gases collected in A and B is approximately 2:1. Identify the gases in tubes A and

B.

c. By considering the reactions that occur at the electrode surfaces, explain the occurrence of

the 2: l ratio in the gases collected in A and B. Use balanced ionic equations to support

your answer.

d. If the bulb in the figure above is replaced with an ammeter, what will happen to the reading

on the ammeter if the electrolyte is changed from dilute sulphuric acid to dilute methanoic

acid? Explain your answer.

e. What adjustments have to be made to the figure above in order to copper-plate a spoon?



ACTIVITY 35

Atoms of an element, 19 9A, readily combine with those of another element, 27

13B

a. Draw a labelled diagram to illustrate the number of protons, neutrons and electrons present

in an atom of B.

b. To which group of the periodic table does element A belong? Give a reason for your

answer.

c. What type of chemical bonding will be formed when A combines with B? Give a reason

for your answer.

d. Write the chemical formula of the compounds expected to be formed when:

i. atoms of element A combine with those of element B.

ii. atoms of element A combine with each other.

e. Based on your answer in (d) (i) and the information given in the question, calculate the

mass of compound that will be formed when 54g of B reacts completely with A.

f. State THREE likely differences in the properties of the compound formed in (d) (i) when

compared to the compound formed in (d) (ii).



ACTIVITY 36

When chemical reactions occur, heat may be given off or taken in from the environment.

a. What changes occur during a reaction that can account for this fact?

b. Define the term 'heat of neutralization'.

c. It is observed that whenever a strong acid (such as HCl or HNO3) is completely neutralized

by a strong base (such as NaOH or KOH), the heat of neutralization (in kJ mol-1 ) is the

same. Account for this observation.

d. When 12.0 g potassium nitrate (KNO) is dissolved in 100 cm3 of water, the temperature

drops by 4.20 °C.

Relative atomic mass: K = 39, N = 14; O = 16

Specific heat capacity of water = 4.2 J g-1 K-1

Heat change= M x C x ΔT

1 cm3 of solution = 1 g

Using the above information calculate EACH of the following:

i. The number of moles of KNO3 used in the experiment.

ii. The heat change for the reaction.

iii. The enthalpy change in kJ mol-1 for the reaction.

e. State ONE assumption you made in your calculation.

f. Draw a labelled energy profile diagram to represent the enthalpy change for the reaction.


WEEK THIRTEEN: LESSON 37

ACTIVITY 37

The spacecraft Voyager detected the presence of a new element, Q, on the planet Mars. The

following data on this element were transmitted back to earth.

Relative Atomic Mass - 333

Melting Point - 1280 oC

Number of valence electrons - 2

a. Would Q to be electrically conducting or not? Give a reason for your answer.

b. State the expected reaction of Q with water. Include a chemical equation.

c. Further investigations reveal that on exposure to oxygen, the metal Q becomes chemically

inert. State a possible reason for this.

d. A compound of Q, suspected to be a nitrate, has also been detected on Mars. How would

the effect of heat on this metal nitrate of Q differ from that of sodium nitrate? Illustrate

your answer by means of suitable chemical equations.

e. Briefly outline how a dry sample of the sulphate of Q can be prepared from a sample of the

metal nitrate of Q. Include an ionic equation to illustrate your method of preparation.



ACTIVITY 38

Steve places a zinc nail in a beaker containing aqueous copper (II) sulphate. Figures 2a and 2b

indicate what happens to the contents of the beaker over a three-day period.

a. Based on the observed differences between Figures 2a and 2b, name TWO ions that would

most likely be present in the beaker after three days.

b. Write ionic equation(s) to show the reactions that take place in the beaker in Figure 2b.

Explain why these reactions occur.

c. Name the type of chemical reaction occurring in the experiment.

d. Comment on how the RATE OF REACTION might be affected if the zinc nail is replaced

by EACH of the following and explain your answer:

i. Magnesium metal

ii. Lead



ACTIVITY 39

Lauri conducts an experiment to investigate how the mass of a catalyst, manganese (IV) oxide,

affects the rate of production of oxygen during the decomposition of hydrogen peroxide (H2O2).

Water and heat are also produced during the decomposition.

a.

i. Name THREE factors OTHER THAN a catalyst, which can affect the rate of a

chemical reaction.

ii. Write a balanced equation to show the decomposition of hydrogen peroxide by

manganese (IV) oxide, MnO2.

b. The Figure below shows the rate of decomposition of hydrogen peroxide, as a plot of grams

of oxygen liberated per second against mass of catalyst for both 0.40 and 0.80M hydrogen

peroxide.

i. Explain why the plots are different.

ii. For 0.80 M H2O2 and 4.0 g of catalyst, determine EACH of the following:

1.The quantity of O2 produced in 16 seconds.

2.The number of moles of O2 produced in 16 seconds.

3.Volume of O2 produced at S.T.P. in 16 seconds.


WEEK FOURTEEN: LESSON 40

ACTIVITY 40

Chlorine has two isotopic forms with mass numbers 35 and 37 respectively. Chlorine has an atomic

number of 17.

a. What differences, if any, are expected between the chemical reactions of chlorine 35 and

chlorine 37? Explain your answer.

b. Determine the number of electrons, protons and neutrons in the anion formed from the

chlorine 37 atom.

c. Explain the term 'ionic crystal'.

d. The melting points of chlorine, sodium chloride and magnesium oxide are -101 °C, 800°C

and 2800°C respectively.

i. Explain why the melting point of sodium chloride is much higher than that of

chlorine.

ii. The crystal structures of magnesium oxide and sodium chloride are similar. Suggest

why the melting point of magnesium oxide is much higher than that of sodium

chloride.



ACTIVITY 41

The table below shows part of the periodic table.

Use only the elements indicated in the table to answer the questions which follow. Each element

may be used once, more than once or not at all.

a.

i. Which element reacts most readily with dilute hydrochloric acid?

ii. Write a balanced equation to illustrate the reaction occurring in (a) (i) above.

iii. Describe any difference(s) that may be observed in the reaction indicated in (a) (i)

above, if dilute sulphuric acid were to be used instead of dilute hydrochloric acid.

Give a reason for your answer.

b.

i. Identify TWO different elements from the table above which will combine by

covalent bonding.

ii. Draw dot-cross diagrams to illustrate the bonding in b (i) above.

c. An element X (not the actual symbol) has an atomic number of 19. Place element X in the

correct position in the Table above. Give a reason for your answer.



ACTIVITY 42

A student conducts an experiment to compare the effect of temperature on the solubility of two

salts, R and T. The data collected are represented in the table below.

Temperature (oC) Solubility (g per 100g water)

R T

10 25.0 40.0

30 50.0 43.0

50 90.0 45.0

70 140.0 48.0

90 200.0 55.0

100 250.0 58.0

a.

i. Using a graph paper, plot the data for the solubility of R and T given in the table

above.

Use the information from the graph to answer the following questions.

ii. Describe the effect of increasing temperature on the solubilities of R and T.

iii. Determine the temperature at which the solubilities of R and T are equal.

iv. Determine the solubility of T at 75°C.

v. Which of the two salts is more soluble at 5°C?

vi. Suggest a way for obtaining an essentially pure sample of R from a sample which

has been contaminated with a small amount of T.


WEEK FIFTEEN: LESSON 43

ACTIVITY 43

Electrolysis has many applications. The apparatus shown below can be used to purify copper by

electrolysis.

a. In order to adapt the electrolytic cell above for the purification of copper, what material

may be used as the:

i. Cathode?

ii. Anode?

iii. Electrolyte?

b. As electrolysis proceeds, describe the changes expected to be observed

i. At the cathode

ii. At the anode

iii. In the electrolyte

c. Write half-equations for the reactions occurring at the

i. Cathode

ii. Anode

d. A current of 5 amperes is passed for 2 hours during the period of the electrolysis. Calculate

EACH of the following:

i. The quantity of electricity passed in coulombs

ii. The mass of copper deposited (Relative Atomic Mass of Cu = 64; Faraday's

constant= 96500 coulombs)

e. In addition to extraction of metal from their compounds, electrolytic processes are also

widely used to protect metals from corrosion, as well as to make them attractive. Name

TWO such electrolytic processes.



ACTIVITY 44

Salts can be prepared by

- reacting acids with metals

- reacting acids with bases

- direct combination of the elements.

a. From the methods mentioned above, select ONE suitable method for preparing a dry

sample of potassium sulphate. In your answer, give full experimental details, including

1. a description of the method

2. ONE relevant equation.

b. Discuss why the two other methods not selected in (a) above would-be unsuitable choices

for the preparation of potassium sulphate.

c. The hydrochloric acid used in the laboratory has a lower pH than that of vinegar. Yet,

vinegar is recommended for removing limescale deposits in a kettle while hydrochloric

acid is not.

i. Explain why vinegar, and NOT hydrochloric acid, is recommended for removing

limescale deposits in kettles. Include TWO relevant ionic equations to support

your argument.

ii. Suggest what could possibly be done to a solution of hydrochloric acid to make

it suitable for removing the limescale deposits in a kettle.



ACTIVITY 45

1.

a. Ethanol and water, as well as black ink are examples of two mixtures.

For EACH of these two mixtures, outline a suitable technique which can be used to separate

them into their various components.

Include in your answer the

i. apparatus required

ii. principles involved in the separation process.

b. You are provided with a piece of plastic tubing and an iron rod.

i. Describe an experiment that could be used to determine the ability of each of these

two materials mentioned above to conduct electricity.

You should include in your answer a labelled diagram of the apparatus to be used.

ii. Distinguish between an electrolytic conductor and a metallic conductor.

iii. With reference to structure and bonding, explain which of these two pieces of

materials, plastic tubing or an iron rod, will conduct electricity.

iv. Explain how the experiment and the apparatus you have described in (b) (i) above

may be modified to test the ability of aqueous sodium chloride to conduct

electricity.

2. You are provided with a bottle of sea water which also contains some sand. Plan and design an

experiment to recover the solid "sea salt' from the mixture. Your answer should include the

following:

a. A suggested list of apparatus which you would use in recovering the solid sea salt from the

sandy sea water.

b. An outline of the steps for the procedure you could use.

c. List the main observations that would be expected at each stage of the experiment.

d. A student suggests that solid sea salt contains chloride ions. How would you go about

testing for chloride ions in the solid sea salt?

mole concept subtopic: conversions week one: le

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