545 chemistry - college of education and external …cees.mak.ac.ug/sites/default/files/545...

CHEMISTRY

O-LEVEL 2006-2010

(May not be taken with 500 General science)

Introduction

This syllabus is designed to meet the needs of those carrying on to “A” level Chemistry

and those leaving school after senior 4 (or UCE). Although a good deal of the content

will be familiar to teachers, the treatment must be based on practical situations whenever

possible. A large portion of the examination questions will demand knowledge and

understanding of practical situations. Candidates will be at a considerable disadvantage if

the teacher has not adopted an approach based on practical work.

The course has been broadly divided into two parts. The first part, Elementary

Chemistry, can be viewed as a complete two-year introductory course in Chemistry. It

should form the basis for UCE Chemistry. The second part for years 3 and 4, constitutes a

deeper, more theoretical treatment of the topics introduced in Elementary Chemistry.

SI units and chemical formulae or names as specified in the ‘UNEB SI’ booklet

and “Chemical Nomenclature’ booklet respectively will be used. Alternative names of

substances will be given wherever ambiguity might otherwise arise. Two areas of subject

matter are worth noting:

1. Applied Chemistry: (section 19)

The aim is to give a clear understanding of the applications of Chemistry in

society with particular emphasis on aspects relevant to East Africa. The section

should be based on student practical work whenever possible and reference has

been made to some suitable experimental work. The application of chemical

principles and the use of particular compounds can be integrated naturally into

other sections of the course.

As an aid to integration the course has been cross-referenced.

2. Organic Chemistry:

Emphasis is placed upon simple organic molecules as the basic units of macro

molecules and the importance of carbon compounds in the natural and

technological environment. (See Section 11- Carbon Chemistry and Section 19.3-

Natural and synthetic materials).

Aims:

i. To help students to appreciate the importance and application of Chemistry in

every day life.

ii. To assist a student to learn how to think critically in any given learning situation.

iii. To guide a student to discover knowledge, from known to unknown, using a

practical approach.

iv. To relate and to be able to apply the discovered knowledge in every day life.

Testable objectives:

It should be noted that each section of subject matter is preceded by a statement of the

specific objectives associated with the section. The section objectives should be regarded

as being examinable. They can be conveniently expressed in the more general form

outlined below. The detailed breakdown given under the three main headings is given as

amplification.

There is no suggestion that such clear distinctions can always be made in constructing

examination questions.

1. Knowledge.

a) Knowledge of chemical terminology and conventions.

b) Knowledge of a variety of experimental methods from work in the laboratory,

from demonstration on films, to provide experience of techniques, of equipment

and of observation.

c) Knowledge of the main facts that have been established.

d) Knowledge of the main generalizations that have been made and of the theories

that are widely held.

Note: abilities-‘recall’, ‘state’, ‘recognise’, ‘name’, ‘define’, fall in the above categories.

2. Comprehension, Application and Evaluation:

a) The ability to understand and interpret scientific information presented in verbal,

numerical or graphical form and to translate such information from one form to

another.

b) The ability to explain familiar phenomena in terms of the relevant models, laws

and principles.

c) The ability to select and apply known laws and principles to given situations.

d) The ability to analyse given data.

e) The ability to recognize mistakes and misconceptions.

Note: Abilities- ‘explain’, ‘interpret’, ‘predict’, ‘judge’, ‘classify’, fall in the above

categories.

3. Practical abilities

(a) The ability to apply the above knowledge and abilities to practical situations.

(b) The ability to handle apparatus and perform experiments in given situations.

(c) The ability to make accurate observations.

(d) The ability to devise simple experiments.

(e) The ability to record results accurately and present data in an appropriate

form.

(f) The ability to make correct deductions from observations.

Examination format:

There will be three papers.

Paper 1 (1 ½ hours)

This will consist of 50 compulsory objective type questions covering the whole syllabus.

(50 marks)

Paper 2 (2 hours)

This will consist of two sections, A and B. section A will consist of ten compulsory

structured questions requiring short answers.

Section B will consist of four structured and semi-structured essay questions. Candidates

will be required to answer two. In both sections questions will be set on any part of the

syllabus. (80 marks)

Paper 3 (2 hours)

This will be a practical test designed to test the abilities specified above.

Questions may be set on any section of the syllabus. (30 marks)

Detailed syllabus:

A feature of the syllabus is that the specific objectives of each section are indicated.

These statements are designed to: (i) indicate the depth of treatment to be given by the

teacher and (ii) indicate the level of understanding to be attained by the pupils.

Since some topics may be treated in an elementary manner early in the course and then in

more depth later, teachers must read the section objectives carefully to ascertain the

treatment. For example terms such as ‘recall’, ‘state’, ‘name’, demand low-level abilities,

while ‘explain’, ‘predict’, ‘demonstrate’, ‘analyse’, etc. demand higher-level abilities.

PART I: ELEMENTARY CHEMISTRY: (Year 1 and 2)

The general aims of elementary Chemistry are that students should be able to:

i. Handle apparatus, make accurate observations and use their experience of

practical work in new situations.

ii. Identify patterns in the chemical behaviour of single substances.

Section 1: Some observations and experiments. Simple classification.

1. Handle simple laboratory apparatus.

2. Select and use the laboratory apparatus for simple investigation.

3. Recall simple laboratory apparatus.

4. Make accurate observations and draw simple conclusions during experimental

work.

5. State the criterion for a pure substance.

6. Explain the three states of matter in terms of the simple kinetic theory of particles.

7. Recall the effect of heat on a range of substances.

8. Recognise non-permanent and permanent change, energy changes, and the

formation of new substances in chemical reactions.

9. Define an element and compound in terms of atoms and ‘bonded atoms’

respectively.

Note: many of the objectives specified for section 1 will be common for other sections of

the Elementary Chemistry Course. Teachers should bear in mind when teaching other

sections’ objectives in Elementary Chemistry.

Content:

1.1 Mixtures and pure substances: separation of mixtures. Miscible and immiscible

liquids, distillation, crystallisation, paper chromatography. Methods of gas

collection, drying of solids, liquids and gases. Simple criteria of purity, melting-

point and boiling-point.

1.2 Heating substances:

States of matter, melting and boiling, simple kinetic (particle) theory (see 9.1);

atoms and molecules; definitions of element and compound; permanent and non-

permanent changes (using iodine, wax, copper (II) sulphate crystals, sand, copper

(II) nitrate, potassium manganate (VII), and zinc oxide). Test for water of

crystallisation.

(A qualitative appreciation of bond formation/breakage and heat is needed).

Section 2: The atmosphere and combustion:

Objectives

At the end of this topic students should be able to:

1. State the composition of air.

2. Recall the preparation, properties and uses of oxygen.

3. Recall the combustion of specified elements in air and oxygen.

4. Name the conditions for rusting, its prevention and apply the principles to new

situations.

Content:

2.1 Composition of air:

Approximate volumetric ratio of nitrogen/oxygen, quantitative determination

using copper. Burning in air (C, S, P, Na, Cu), oxygen as ‘active part’, mass

changes involved. Burning of a candle. Rusting (rusting as a product of iron,

oxygen and water).

2.2 Oxygen

Laboratory preparation (prepared by hydrogen peroxide with manganese (IV) oxide or

sodium peroxide with water); combustion of elements in oxygen; uses.

Competition for oxygen: reactivity series K, Na, Ca, Mg, Al, C, Zn illustrate with CO2

/Mg, PbO/Mg. Experiments carried out on a ‘bottle top’. Take care.

Section 3: Acids and Bases (see section 15)

Objectives


1. State the colours and corresponding pHs of universal indicator.

2. Recall the acidic/neutral/basic (alkaline) properties of some compounds.

3. State the reactions of dilute acids with metals and metal oxides, hydroxides,

carbonates and sulphates.

Content

Flower extracts as simple indicators. Universal indicator, pH scale. Acidic, neutral and

basic/alkaline solutions. Simple properties of mineral acids. Test solutions: NH4CI,

(NH4 ) 2 SO4, NH3 (aq), NaCI, NaOH, CO2 (aq), SO2 (aq), H2 SO4 (aq), HNO3 (aq),CH3

COONa, NaHCO3, Ca(OH)2, MgO.

Section 4: Water and Hydrogen

Objective


1. Explain the experiments showing that water contains hydrogen.

2. State the products of the reaction of water and steam with different metals.

3. List the reactivity series obtained from metal/water reactions.

4. Recall the preparation and test for hydrogen.

5. Define oxidation and reduction as gain/removal of oxygen.

6. Explain metal oxide/hydrogen reaction in terms of this definition.

Content

Burning of organic matter (energy source). Water as oxide of hydrogen. Reactions of

metals with water/steam. Burning hydrogen in air. (Burn candle in air, test for CO2 and

H2O vapour, Na, Ca, Mg with water and Mg, Zn, Fe with steam).

Metals and dilute acids:

Preparation and test for hydrogen, oxidation and reduction (oxygen gain/removal

only). Metal oxide (use copper (II) oxide, hydrogen reduction (see 16.3)

Section 5: The effect of Electricity on substances (see section 16)

Objectives


1. Define conductor/non conductor, electrolyte/non-electrolyte.

2. Recognise solutions and melts as electrolytes/non-electrolytes

3. State the relationship between electrolytes/non-electrolytes and the particles they

contain (ions, molecules).

4. Name the products of electrolysis of simple binary electrolytes.

5. Recognise that electrolysis is a means of obtaining elements from chemical

compounds.

Contents

Conductors and non-conductors; electrolytes and non-electrolytes; cathode and anode.

Ions as the particles in electrolytes solutions. Electrolysis, conduction. Test solids for

conductivity (metals, non-metals, plastics, wood)

Solutions: sugar, urea, sodium chloride and copper (II) chloride, dilute mineral acids.

Melts: PbBr2.

Electrolytes contain ions, non-electrolytes contain molecules. Avoid solutions where

water could interfere: use CuCl2, HCl, PbBr2.

Note: no explanation of the products of electrolysis is required in the elementary section.

Section 6: Ionic compounds (salts).

Objectives


1. Recall which compounds are soluble and insoluble.

2. Select the appropriate method for preparation of a particular salt.

3. Recall the action of heat on several salts.

4. Explain the terms ‘saturated solution’ and ‘crystallisation’.

Content

Preparation of soluble salts, acid with metals, metal hydroxide, metal oxides, metal

carbonates, soluble salts by precipitation. (One example of each will suffice) (See 16.1).

Solubility of well-known salts, sulphate, chlorides, nitrates, carbonates, hydroxide and

oxides. Relationship between method of preparation and solubility.

Action of heat on carbonates, nitrates, sulphates and hydrates.

Section 7: Structure of the Atom and the Periodic Table.

Objectives

At the end of the topic students should be able to:

1. Recall he structure of the atom and recognise the electron

2. Know the arrangements of electrons in the atoms of the first twenty elements.

3. Explain the structure of the atom from a simple energy model.

4. Define atomic number, relative atomic number, relative atomic mass and

isotope.

5. Explain the Periodic Table from he electronic arrangement of an element.

6. Predict ion formed from the electronic arrangement of an element.

Content.

7.1 Simple model of the atom, balance of the charges of protons and electrons;

neutrons

` Positive and negative charges should be introduced through simple electrostatics

experiments with charged rods and spheres.

Definitions of atomic number, relative atomic mass, significance of isotopes.

7.2 Ion formation; qualitative treatment of the energetics of electron loss from the atom

(use of ‘shells’ discouraged); derivation of the electron arrangements from Li+ ,Na+,

(Li has 3 electrons forming Li+ suggests 2.1 arrangements, similarly Na with 11

electrons forming Na+ suggests 10.1, becoming 2.8.1).

Build-up of the Periodic Table for the first twenty elements on this basis. (See 10.1).

Section 8: Chemical Families: Patterns in Properties:

Objectives


1. Recall the specified reactions of the alkali and alkaline-earth metals, the

halogens.

2. Recognise qualitatively the difference in ‘reactivity’ within the chemical

families.

3. State the noble gases family is comparatively unreactive.

4. Predict the reactions and ‘reactivity’ of elements within each family on

qualitative basis.

5. Explain the reactions of the chemical families in terms of their similar

electronic arrangements.

Content

8.1 Alkali metals (Li, Na, K only).

Reaction with air, water, chlorine. Similarity of chemical formulae of their

compounds ions, oxides, chloride.

8.2 Alkaline earth-metals (Ca, Mg)

Reaction with air, water, chlorine and dilute acids.

Similarity of ions and formulae of their compounds.

8.3 Halogens (chlorine, bromine and iodine)

Reaction with sodium, water (+bleaching action), zinc powder, sodium hydroxide

solution. Similarity of ions and compounds. (See 10.5.14).

8.4 Noble gases

Reaction of their low reactivity; appreciation in terms of the stable electron

arrangements. (See 10.2).

Note: It is not intended that full formulae equations be introduced in the Elementary

Section. Word equations should be given whenever possible. Where they arise naturally

the formulae of some simple compounds and elements may be given. O2, N2, CO2, H2O,

I2, CI2, Br2, NaCl, Cu, Fe, Zn, Mg, HCl, H2SO4, HNO3, CuO, NaOH, Na, K, Li, He, Ne,

Ar.

Part II: (For year 3 and 4)

Section 9: The Mole Concept: Formulae and Chemical Equations.

Objectives


1. State more experimental evidence for the existence of atoms, molecules, ions

and electrons.

2. Define the mole, molar solution, molar gas volume and other derived

quantities, and recognise their relationship to relative atomic and molecular

masses.

3. Use the mole, molar solution and molar gas volume in deriving chemical

formulae and equations from both experimental results and given data.

4. Represent a chemical reaction by either a full formula or ionic equation.

Content

Evidence of existence of atoms, molecules, ions and electrons; effect of heat and

electricity on elements and compounds; Brownian movement; diffusion experiments;

qualitative relationship between diffusion rate and density.

Use bromine, copper (II) sulphate crystals/water as illustrations. Effect of density on

ammonia and hydrogen chloride diffusion.

Use of kinetic theory to explain nature of solids, liquids and gases. Diffusion of gases,

evaporation of liquids, dissolving solute in solvent. Qualitative explanation of Boyle’s

law and Charles’ law for gases (no calculations to be set).

The mole as a basic unit (see revised UNEB ‘SI Units” booklet). Relative atomic mass.

The mole as a number of particles can be illustrated by ‘counting by weighing’

experiments.

Determination of formulae; ionic compounds, empirical and molecular formulae.

Quantitative determination of magnesium oxide (Mg/air) and copper (II) oxide (reduction

of copper (II) oxide with butane) should be carried out.

Molar gas volume (22.4 dm3 at S.T.P.); atomicity of gases; mass volume relationship for

gases. Molar solutions.

Stoichiometry of chemical reactions; quantitative work must be emphasised.

Reactions to be considered:

i. Ba2+ (aq) + CO

ii. Pb2+ (aq) + 2I‾ (aq) → PbI2 (s)

iii. Cu2+ (aq) + Fe (s) → Cu (s) + Fe 2+ (aq)

iv. CO2 evolution from Na2CO3 + HCl.

v. Titration of NaOH with HCl and H2SO4 recommended.

Note: Reactions in this section show quantitatively that mass is conserved.2+ (aq) → BaCO3 (s).3

Use of ionic equations and full formulae equations (latter recommended) in the

calculation of reacting masses. (See 16.1).

Note: The use of full formulae equations and ionic equations should be emphasised

henceforth.

Section 10: Atomic and Molecular Structure: Chemical Bonding.

Objectives


1. Recognise the importance of the outer electrons in chemical bonding.

2. Explain in a qualitative way the energetics of covalent and ionic formation.

3. Represent the covalent and ionic bonds in simple compounds by electron

sharing and transfer respectively.

4. Differentiate between different bond types on the basis of their chemical and

physical properties.

5. Demonstrate the changes in bond type within the groups and periods of the

Periodic Table

Content

10.1 The significance of the outer electrons in chemical bonding. Qualitative treatment

of the energetics of chemical bonding. Consider the molecules in terms of a

position of balance between p-p, e-e repulsions, and p-e attraction (the ionic bond

as an extreme example).

10.2 Significance of the noble gas configuration; covalent bond as electrons sharing;

ionic bond as electron-transfer. Consideration of C-C and C=C (see 11.5)

10.3 Influence of bond type on physical and chemical properties (melting point,

solubility and electrical conductivity).

Molecular, giant atomic and giant ionic structures (iodine, carbon (diamond), and

sodium chloride respectively.

10.4 Metallic bond related to electrical conductivity only.

10.5 Periodic of bond type.

Elements Na, Mg, Al, Si,S,Cl, Ar: their electronic structures, their ions (valency),

mode of combination in oxide and chloride; inertness of noble gases; chemical

and physical properties of metals and non-metals (across a period).

Elements: fluorine, chlorine, bromine and iodine (down group). Electronic

configuration, graduation in size of atom and ion, reactions, (see 8.3).

Elements: Li, Na, K (as above including ease of oxidation, reaction with water,

chlorine). (See 8.1)

Section 11: Carbon Chemistry

Objectives


1. Explain the physical properties of the carbon allotropes in terms of the

bonding.

2. Recall the preparation and properties of carbon dioxide.

3. State the chemical reactions of carbonates and hydrogen carbonates.

4. Recognise the importance of carbon compounds in the natural

environment and industry.

Content

11.1 Diamond and graphite (note charcoals)

Structure and physical properties.

11.2 Carbon monoxide

Combustion, reducing action, poisonous fumes (car exhausts, coke fire “sigiri”).

Laboratory preparation not required (see 19.1).

Reducing action illustrated with copper (II) oxide + blast furnace. (See extraction

of ion 19.4).

11.3 Carbon dioxide

Preparation, reaction with water, limewater.

Weathering, boiler ‘scale’; no detailed treatment of hard water.

11.4 Carbonate and hydrocarbonates

Action of heat, dilute acids.

Production of soda ash (Lake Magadi). (See 19.4d)

Carbon and carbon dioxide cycles. (See equilibration of atmosphere 19.1)

11.5 Alkanes (methane to butane)

Formulae and combustion only.

Natural gas.

Paraffin wax.

Ethane: formulae and combustion. (See combustion of fuels, 19.1).

Preparation of ethane by dehydration of ethanol: reaction with bromine (see 10.2)

Note: a detailed study of the organic chemistry of alkanes, alkenes, alkanols etc. is not

required.

Section 12: Nitrogen

Objectives


1. Recognise the unreactive nature of nitrogen in comparison with oxygen.

2. Recall the laboratory preparation of ammonia and list its properties.

3. Recognise the difference in chemical reactions of ammonia gas and its aqueous

solution.

4. Recall the reactions of both dilute and concentrated nitric acid and recognise the

difference between their chemical action.

5. Name the products when different metal nitrates are heated.

Content

12.1 Nitrogen

Inert character. Compare combustion of Na, Ca, P, S, in N2 andO2.

(See 19.1 Industrial solution of component of air).

12.2 Ammonia

Laboratory preparation, solubility in water, reaction with air/oxygen (catalysed

and uncatalysed), copper (II) oxide and chlorine. Reactions of aqueous solution:

metal ions, dilute acids.

Manufacture of ammonia (Haber process). (See 19.4 (a))

Ammonium salts as fertilizers. (See 19.2).

12.3 Nitric Acid

Dilute: reaction with metals, carbonates, hydroxides, oxides.

Concentrated: oxidizing action; Fe (II) solution, sulphur, copper metal.

Nitrates: action of heat (Na, K, Cu, Pb, Ag).

Acidic nature of nitrogen (IV) oxide. (See 19.4)

Manufacture of nitric acid (see 19.4(b))

Section 13: Sulphur

Objectives


1. Recall the preparation of sulphur dioxide and state its properties and uses.

2. Recall the preparation and properties of dilute sulphuric acid.

3. Illustrate the difference in chemical action between dilute and concentrated

sulphuric acid.

4. Recognise the atmospheric pollution caused by sulphur containing compounds.

Content

13.1 Sulphur dioxide

Preparation (Sulphite +Acid), acid character, bleaching action, test with potassium

dichromate (IV)

Reducing action is not required.

Combination with oxygen (laboratory demonstration, Pt. Catalyst).

13.2 Sulphuric acid

Dilute: reaction with metals carbonates and bases.

Concentrated: dilution with water, copper (II) sulphate crystals, ethanol, sucrose.

Test for sulphate: barium nitrate solution.

Manufacture of sulphuric acid (Contact process). (See 19.4(c)).

13.3 Hydrogen sulphide (iron (II) sulphide and acid).

‘Bad eggs’ smell.

Pollution by S, SO2, H2S. (See 19.1 The atmosphere.)

Section 14: Chlorine

Objectives

At the end of this topic students should be able:

1. Recall the preparation and properties of hydrogen chloride and chlorine.

2. Explain how the composition of hydrogen chloride can be deduced from a series

of chemical reactions.

3. Recognise that an aqueous solution of hydrogen chloride shows typical acid

properties.

4. Name some uses of chlorine-containing compounds.

Content

14.2 Hydrogen chloride (common salt + conc. H2SO4).

Deduction of composition of salt gas reactions.

(i) KMno4, (ii) iron metal, (iii) direct H2 and Cl2 combination

Tests for gas with ammonia.

Aqueous solution of gas; acid properties: mental, carbonates, etc.

14.2 Chlorine

Preparation: conc. HCl + Potassium magnate (VII), electrolysis of chloride

solutions.

Reactions:

(i) Hydrogen and hydrocarbons (turpentine)

(ii) Metals (Na, Zn, Mg)

(iii) Non-metals (P, S).

(iv) Water and dilute alkali (see 8.3)

(v) Bromides and iodides.

(vi) Bleaching action.

14.3 Test for chlorides: (a) dry solid, (b) aqueous solution. Uses of compounds:

disinfectant, bleach, chloroform. (See 19.2 water treatment).

Section 15: Acids and Bases (Acidity and Alkalinity) (See section 3)

Objectives


1. Define acids as proton donors and bases as proton acceptors.

2. Recognise the difference between weak and strong acids and bases.

3. Explain the role of the solvent in the acidity of hydrogen chloride.

4. Write ionic and formula equations for specified acid-base reactions.

Content

15.1 Weak and strong acids.

PH, electrical conductivity, rate of reaction with marble chips, magnesium, for

acids and bases. (Use hydrochloric/ethanoic and sodium hydroxide, aqueous

ammonia as illustration.).

Note: other examples are tartaric and citric acids instead of HCI.

15.2 Role of the solvent

Hydrogen chloride or tartaric acid in methyl benzene compare with aqueous

solutions. React with dry litmus, magnesium, marble chips.

Reactions of dry and aqueous ammonia. Importance of H+ (aq) and OH–(aq).

15.3 Acids (proton donors), bases (proton acceptors). Use of ionic equations.

15.4 Amphoteric oxides (Al2O3, PbO, ZnO): react with acid and alkali (no equations

for reaction with alkali).

Section 16: Ion Chemistry

Objectives


1. Recognise the precipitates and complex ions produced by specified cation-anion

reactions.

2. Identify ions from a series of specified reactions.

3. Differentiate between ions using a series of ionic reactions.

4. Explain a redox reaction in terms of electron transfer.

5. Compare the oxidizing and reducing power of ions from displacement reactions.

6. Recognise the role of water in the products of electrolysis.

7. Explain an electrochemical cell in terms of electron transfer processes.

Content

16.1 Precipitation reactions involving the following ions:

Mg2+ (aq), Ca2+ (aq), Fe3+ (aq), Al3+ (aq), Zn2+ (aq), Cu2+ (aq), Fe2+(aq),

with Cl–(aq), OH–(aq), CO

(See 9.7)

16.2 Complex ions: limited to dissolving of specific metal hydroxides in excess

ammonia solution or sodium hydroxides. Formula of the following required:

Cu (NH3)2+ Pb(OH)2–, Al(OH).

Note: No instructions on equations required.

16.3 Redox reactions.

Electron transfer.

Useful illustrations Fe2+, Fe3+ (with 2 – (aq) H2O2/ H+ ) 3

16.4 Displacement reactions as redox reactions.

(a) Reducing power: reaction of metal/cation

(b) Oxidising power of halogens: Cl2, Br2, I2 only. (See 8.3, 10.5)

16.5 The role of water in electrolysis products.

Preferential discharge of hydrogen and oxygen where appropriate from the

following solutions: sodium chloride, dilute sulphuric acid (acidified water),

magnesium sulphate.

16.6 Electrochemical cell

Qualitative treatment of the electron flow in

Zn/Zn2+ (aq)/Cu2+ (aq)/Cu cell (See 17.3).

(See electrolytic processes, 19.4 (a)).

Section 17: Energy Changes in Chemical, Physical Reactions.

Objectives


1. Define endo- and exo-thermic reactions using ∆H notation qualitatively.

2. Recognise that energy changes in chemical and physical changes are due to bond

formation and bond breaking.

3. Compute enthalpy values using data both given and obtained experimentally.

4. Explain the construction of the Zn―Cu cell.

Content

17.1 Molar heat of vaporization and boiling-point (latent heats). As evidence for

interparticle forces.

Enthalpy notation (ΔH) for exo and endo-thermic reactions.

17.2 Enthalpy chemical reactions

Students should carry out simple quantitative work (See 15.1), e.g. enthalpy of

combustion (methanol, ethanol), enthalpy of displacement (Cu2+ (aq) + Fe (s))

(see 16.3); enthalpy of solution (NH4NO3, NaOH, and conc. H2SO4) (see the

combustion of fuel and the internal combustion engine 19.1).

17.3 Electrochemical cell

Qualitative look at the electrochemical series.

Students can use copper reference, sodium chloride electrolyte and metal strip

electrodes.

Zn-Cu cell: teacher demonstration and measurement of potential difference.

The potential E is not required (see 19.1).

17.4 Simple treatment of solar energy as energy from atoms.

Section 18: Reversible Reactions and Reaction Rates:

Objectives


1. Recognise the effects of different factors on reaction rate.

2. Name some methods used to measure reaction rate.

3. Illustrate reaction rate graphically and explain the representation in a qualitative

way.

4. To recall some simple reversible reactions.

5. To recognise the sign and explain how reversible reactions reach a state of

‘balance’

Content

18.1 Reaction rate

The effect on rate of: concentration, pressure, temperature, surface area, light and

catalysts.

Only qualitative, descriptive, graphical representation required, quantitative data

given to illustrate a qualitative effect.

(i) Marble chips/dilute acids.

(ii) Decomposition of H2O2.

(iii) Manganese (IV) oxide to catalyse H2O2 decomposition.

(iv) Plantinised asbestos to catalyse SO2/O2 combination.

Note: Candidates will be expected to appreciate the applications of reaction rate to

laboratory and industrial processes. (See 19.4).

18.2 Reversible reaction

Elementary treatment incorporating the idea that two-way reactions can reach a

“balance”, “equilibrium” is avoided. Examples: acid alkali plus indicator,

chromate/dichromate/acid, hydrated and anhydrous copper (II) sulphate. (See

10.4)

Note: The effect of changing concentration, pressure, temperature on position of

equilibrium not required. The use in industrial processes should be regarded as

‘optimum’ only. (Section 19.4 (a), (b), (c) can be used as illustrations).

Section 19: Applied Chemistry

Objectives


1. Explain the industrial isolation of nitrogen, oxygen and the inert gases.

2. Recognise the advantages and disadvantages of different fuels.

3. Recall the industrial and domestic uses of water and explain the sources of

pollution and their treatment

4. State the specific uses of chemical compounds both locally and worldwide.

5. Name the natural resources available locally and illustrate their use in local

industry.

6. Illustrate the operation of chemical properties already encountered in both natural

environment and specific industrial processes.

7. Explain the synthesis and breakdown of natural and synthetic materials.

8. Recognise the relative advantages of synthetic materials over those of natural

origin in terms of both structure and properties.

Content

The depth of treatment envisaged should not introduce the students to any new concepts.

The emphasis is on illustration and application of chemical principles already

encountered. Technical details of plant and process are not required. Many of the

compounds mentioned in 19.3 in particular have complex formulae. Students are not

required to memorise these formulae. The aim is to show the relationship between

structure, properties and uses. (P= practical work, the rest should be covered by the

teacher or by students’ private reading).

19.1 The atmosphere

The industrial isolation of nitrogen, oxygen and noble gases from air (see 12.1)

Their uses (to include: for oxygen: steel manufacture, life support, welding; for

argon; discharge tubes).

Principles and methods of extinguishing fires of different types (P).

Combustion of hydrocarbon fuels (practical work: charcoal, methylated spirit,

butane and ethyne (P))

Heat energy values of charcoal, fuel oil, menthylated spirit and natural gas. (See

17.2)

The internal combustion engine as a major source of atmospheric pollution.

(Refer to unburnt carbon, carbon monoxide, carbon dioxide, lead compounds, and

unburnt hydrocarbons)

Brief consideration of alternatives to the internal-combustion engine; steam

engines, fuel cells.

Sulphur dioxide as a pollutant from combustion of coal and heating oils.

Equilibrium of the atmosphere via the oxygen and carbon dioxide cycles

19.2 Water resources; pollution.

Industrial and domestic uses of water.

Water treatment: filtration; fluoridation and desalination.

Hard and soft water: causes and treatment; including ionization methods. Sewage:

methods of water treatment; production of methane and fertilizers.

The nitrogen cycle: elements necessary for plant growth: N, P, K, Ca, Mg, S.

Fertilizers as artificial replacements: ammonium salts, phosphates, nitrates,

sulphates.

Soaps and non-soapy detergents: manufacture, simple mode of action, and change

to biodegradable products.

(Teacher demonstration of laboratory preparation of detergents from castor oil

and concentrated H2SO4).

Pollution from fertilisers, insecticide, herbicides and other agricultural wastes

(only the pollution aspects).

Oil production of the sea and lakes; dispersal of oil slicks.

19.3 Natural and synthetic materials.

Note: Students should be introduced to the following topics and the underlying

chemical principles, but will not be expected to memorise complex formulae.

Natural cellulose materials e.g. cotton, wood, paper.

Natural protein fibres, wool and silk and natural dyes and colouring of fibres. (See

1.1 and3)

Constituent elements in food: Carbohydrates (See 19.4 (g)), sources of calorific

value.

Proteins; fats and oils (local sources).

Vitamins and minerals.

Fractional distillation of crude oil (five fractions) (D).

Uses of the five functions. Large scale isolation of hydrogen (Cracking process).

Synthesis of large molecule: polymerization.

(i) Look at polythene structure and C2H2 only (P).

(ii) Urea plus methanol (Structure not required) (D).

Natural rubber and its vulcanization.

Breakdown of large molecule: hydrolysis of starch, fermentation of sugar (aerobic

oxidation of alcohol to acid).

‘Cracking’ of Perspex. (P).

Simple classification of thermoplastics, thermosets, fibres, rubbers (p).

Simple relationship between physical properties, structure and uses.

Advantage and disadvantage of man-made polymers over those of natural origin.

19.4 Mineral resources/industrial processes.

The following process should be used to illustrate:

(i) The chemical principles already covered in the course.

(ii) The influence of the following factors (particular reference to East

Africa): availability of raw materials, choice of site, social and

economic factors, health and pollution problems, supply and demand.

(a) Herber process (Ammonia). (See 12.2)

(b) Manufacture of nitric acid. (See 12.3)

(c) Contact process (sulphuric acid) (P).

(d) Manufacture of soda and salt (Lake Magadi, Lake Katwe).

(See 11.4)

(e) Electrolytic processes:

- Sodium extraction

- Copper refining

- Electroplating

(f) Extraction of iron: manufacture of steel, examples of alloys,

brass, solder, duralumin and their composition.

(g) Large-scale extraction of sugar from sugar-cane

Note: Use of very simple floe-charts of the processes should be encouraged.

545 chemistry - college of education and external …cees.mak.ac.ug/sites/default/files/545...

Documents