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Unit 1
Section (a) – Reaction Rates
1. The rate of a reaction can be followed by measuring changes in concentration, mass
or volume.
Change in …………………………, ……………………….. or ………………………….
Rate =
………………………………… taken
Units are mol l-1 s-1 (for concentration change), ………………. (for mass change), …………………… (for volume change)
2. The rate of a reaction is proportional to 1/........................................ (units s-1).
3. Reaction rate increases as the concentration of a reactant.........................................
As the concentration of a reactant increases the number of collisions between reactant
particles ............................................ .
4. Reaction rate increases as the particle size of solid reactants .........................................
The smaller the particle size the ....................................... the surface area of the solid.
The bigger the surface area the ................................ collisions there are between reactant
particles.
5. Reaction rate increases as the temperature at which the reaction is carried out ...............................
Temperature is a measure of the average ............................................. energy of the particles of a
substance. The higher the kinetic energy of the particles the ............................................ the number
of collisions between reactant particles.
6. The .......................................... energy (EA) is the minimum kinetic energy required
by colliding particles before reaction will occur. Increasing the temperature
......................................... the number of particles which possess the
.................................................. energy (E A) and so increases the rate of the reaction.
A ............................... increase in temperature can give a large increase in reaction rate.
7. With some chemical reactions light energy can be used to increase the number of particles with
the necessary activation energy. One such reaction is ........................................................... in which
green plants use light energy to make starch.
8. From a balanced chemical equation the amount of a reactant in excess can be calculated. (See
separate booklet on Higher Chemistry calculations).
9. A .................................................... catalyst is in the same physical state as the reactants.
A ..................................................... catalyst is in a different physical state to at least one of the
reactants.
10. Heterogeneous catalysts work by adsorbing a reactant onto the surface of the catalyst.
The reactant that is absorbed forms ........................ bonds with the active site on the catalysts
surface. The reactant adsorbed and the .......................... site are complementary in size and shape.
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11. The surface activity of a catalyst can be ............................................. by a catalyst poison. For example,
cars fitted with catalytic converters have to use .......................................... petrol as any .......................
would poison the catalyst. The catalytic converter converts harmful carbon monoxide and oxides of
nitrogen into ..................................................................... and .........................................................
12. The aluminium oxide catalyst used in cracking long chain hydrocarbons often gets covered in soot
and so the soot has to be .......................... off and the catalyst regenerated.
13. ...................................... catalyse the reactions taking place inside the cells of living plants and animals.
14. ..................................... are also used to catalyse many industrial processes. For example, zymase
is used to catalyse the fermentation of sugar into ..................................................
Section (b) – Enthalpy
1. In an ............................................... reaction heat energy is released to the surroundings.
In an ............................................... reaction heat energy is absorbed from the surroundings.
2. In an exothermic reaction the reactants have ........................... stored energy than the products.
In an endothermic reaction the reactants have ........................... stored energy than the products.
3. A potential energy diagram can be used to show the energy stored in both
reactants and products. The enthalpy change (H) for the reaction is given
by Hproducts - .............................................. The enthalpy change has a negative
value in an ............................................... reaction and a positive value in an
.................................................... reaction.
exothermic reaction
4. The activated ................................... is an unstable arrangement of atoms formed at the maximum of
the potential energy barrier during a reaction .
5. The ........................................... ................................... (E A) is the minimum energy required by the reactant
particles in order to form the activated complex.
6. A catalyst speeds up a reaction by providing an alternative reaction pathway with a ...............................
activation energy.
7. The enthalpy of ...................................................... is the amount of heat energy give out when one ............
of a substance is completely burned in oxygen.
8. The enthalpy of ...................................................... is the enthalpy change when one ............of a substance
is dissolved in water.
9. The enthalpy of ........................................................ is the mount of heat energy give out when one ............
of water is formed in the neutralisation of an acid.
10. Enthalpy change can be calculated using the relationship H= - cmT where
c is the ..................................................................................... of the substance (usually water with a value
of 4.18 kJ kg-1 oC-1)
m is the mass in ................................................ (kg) of the solution being heated.
T is the temperature change (in ............).
2
Section (c) – Patterns in the Periodic Table.
1. The Modern day Periodic Table is based on the work of the Russian chemist .............................................
who arranged the elements in order of increasing .................................... mass and left spaces for yet
to be .................................................... elements.
The present day Periodic Table actually arranges the elements in order of increasing atomic number
(the number of .............................. in the nucleus of an atom).
2. The covalent radius of the elements ................................................ from left to right across a period
and .................................................. from top to bottom within a group.
3. The .................................................. in covalent radius from left to right across a group is due to an
................................................. in the number of protons (positive charges) in the nucleus attracting the
electron energy levels ......................................... to the nucleus.
4. The .................................................. in covalent radius from top to bottom in a group is due to an
.............................................. in the number of occupied electron energy levels as we go down the group.
Each new energy level is ..................................... away from the nucleus making the atom .............................
in size.
5. The first ........................................................ energy is the amount of energy that has to be supplied to
remove one mole of electrons from one mole of neutral atoms in the gas state.
ie for sodium Na(g)
................ + ........... H= .................... kJ mol-1 (data booklet page 10)
6. The second ...................................... energy is the amount of energy to remove one mole of electrons
from one mole of +1 ions in the gas sate.
+
ie for sodium Na (g)
................... + ...................... H= .................. kJ mol-1 (data booklet page 10)
7. The second ionisation energy is higher than the first due to the electron having to be removed
from a ......................... ion rather from a .......................... atom.
8. The first ionisation energy ..................................... from left to right across a group due to the
increasing number of ................................. in the nucleus attracting the electrons more ...........................
This increased attraction means that more energy has to be supplied to remove an electron from
the atom.
9. The first ionisation energy ............................... from top to bottom within a group due to the
increasing number of occupied electron ................................... .......................................... Each new
energy level that is occupied is ........................... away from the nucleus. Each new energy level
provides additional screening and so it requires ............................. energy to remove an electron.
10. Atoms of different atoms have different attractions for bonding electrons. The
.................................................................. value is a measure of the attraction that an atom has for
electrons it shares with other atoms.
11. The higher the electronegativity value of an atom the more ...................................... it attracts
bonding electrons towards itself.
3
12. Electronegativity values ..................................... from left to right across a period. This is due
to an increase in the number of ................................ in the nucleus of the atoms attracting the
shared electrons in the bond more strongly.
13. Electronegativity value ...................................... from top to bottom within a group. This is due to
the greater ..................................... effect on more occupied electron energy levels.
Section (d) - Bonding, structure and properties of the first 20 elements.
1. Metallic bonding is an electrostatic attraction between the ................................charged nucleus and
the delocalised outer ..................................................
2. Atoms of non-metal elements bond together by sharing electrons – this type of bond is known as
a ......................................... bond. In a ..................................... bond the atoms are held together by an
electrostatic attraction between the ........................................ charged shared electrons and the
............................................ charged nuclei.
3, Ionic bonding usually occurs between metal and ........................................ atoms.
Ionic bonding is an electrostatic attraction between the .......................................... charged metal ions
and the ........................................... charged non-metal ions.
4. ............... der ........................ forces are weak forces of attraction that exists between covalent
molecules and also between the atoms of ............................. gases (group 8). .................... der ...................
forces are much .................................. than all other types of bonding.
5. Van der Waals forces are caused by small electrostatic attractions between ......................................
dipoles created in atoms by an uneven distribution of electrons in the electron cloud round an atom.
6. Van der Waals forces ................................... in strength as we go down groups 5, 6 7 and 8.
Bigger atoms have more electrons and so the size of the temporary dipoles .....................................
as we go down the group.
7. A molecule is described as being polar if the it has a ......................................................... dipole.
eg hydrogen chloride +H-Cl-.
8. The electrostatic attraction between molecules with permanent dipoles are .................................... than
the weak van der Waals forces between molecules with .................................................. dipoles.
9. Symmetrical molecules with polar bonds are themselves ................................................ – the symmetry
cancels out the polarity of the bonds.
For example, +O=C+=O- is a non-polar molecule as it has a ........................................ shape.
10. Bonds in which a hydrogen atom are bonded to a much more electronegative element are
very strongly polar. The elements which are much more electronegative are nitrogen, ........................,
....................................... and .......................................... Hydrogen bonds are .................................... forces of
attraction between these highly polar molecules eg +H-F+H-F hydrogen bonds).
11. Hydrogen bonds are ............................ than other forms of permanent dipole attractions but they are
weaker than covalent bonds.
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12. A metallic structure consists of a giant .................................... of positive ions surrounded by
...................................................... outer electrons.
-
+
delocalised outer
...............................
-
-
+
+
+
-
-
+
+
+
+
+
+
-
-
-
-
+
+
+
+
+
-
-
-
+
-
-
lattice of .........................
charged nuclei
+
+
-
-
13. A covalent molecular structure consists of discrete molecules with ............................... bonds
between the atoms inside the molecule and weak ...................... ..................... ............................. for ces
between the molecules.
H
eg
covalent bond
H
H
van der Waals force
H
14. A covalent network structure consists of a giant ........................................ of covalently bonded atoms.
covalent bond
eg diamond
15. An ionic structure consists of a giant .................................... of oppositely ...................................... ions.
eg sodium chloride
16. A monatomic structure consists of individual atoms with weak .................... ................. .....................
forces between the atoms.
He
van der Waals force
He
17. The first 20 elements in the Periodic Table can be categorised according to bonding and structure.

Metallic (Li, Be, .........., ............, .............., ..............., ................, )

Covalent molecular gases (H2, N2, ............, ................, .................)

Covalent molecular solids (P4, ................ and fullerines C60)

Covalent networks (B, ........... and .............)

Monatomic (He, Ne, .........., ............., ................, .................)
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Section (e) – Bonding and structure of compounds
1. Compounds adopt one of three different structures.
 Covalent molecular – discrete molecules with ................................ bonds between the atoms
inside the molecule and weak ................ ................. ..................... forces between the molecules.
eg methane (CH4)
.................................................... bond
between atoms within the molecule
weak van der Waals forces
....................... the molecules

Covalent network – a giant ............................ of covalently bonded atoms.
eg

silicon dioxide
Ionic - a giant ............................... of oppositely charged ions
eg sodium chloride
Section (f) – Properties associated with bonding
1. The melting and boiling points of polar substances is ................................... than the melting and boiling
points of non-polar substances with similar molecular mass.
2. Ionic and polar substances tend to be ........................................ in polar solvents such as water.
Non-polar substances tend to be .......................................... in polar solvents such as water
3. Water (H2O), ammonia (............) and hydrogen fluoride (...........) have ................................... than
expected boiling points due to ............................................ bonding between the polar molecules.
4. Hydrogen bonding between water molecules results in ice having a very open structure.
this results in ice having a ....................................... density than liquid water and so ice .....................
on water.
5. Diamond is a very .................... and strong substance due to its giant ....................................... structure.
In diamond every carbon atoms is bonded to .................. other atoms. This means that all 4 electrons
in a carbon atoms ............................ shell are bonded and so carbon does not ............................................
electricity.
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6. In graphite the carbon atoms are arranged in ........................................ with weak van der Waals forces
between the layers. This makes graphite ......................... and slippery. As only 3 of the 4 electrons in
the outer energy level are shared there are ........................................... electrons which makes it an
electrical ............................................
7. Silicon carbide (SiC) has a giant covalent ....................................... structure. This makes silicon carbide
very ........................ and strong. It is used in cutting and grinding tools.
Section (g) – The Mole
1. One mole of a substance contains 6.02 x 1023 formula units of the substance.
For example.
a) 1 mole of carbon (C) contains ............................................................ carbon atoms.
b) 1 mole of sulphur dioxide (SO2) contains 6.02 x 1023 sulphur dioxide ...............................................
c) 1 mole of sodium chloride (Na+Cl-) contains ................................................. Na +Cl- formula units.
2. The Avogadro Constant (L) has a value of ........................................... particles per mole.
3. The number of moles of a solid can be calculated using the relationship
where m = .............................. of the substance
n = number of .................. of the substance
m
n
gfm
gfm = ................... ........................ mass of the substance
4. The number of moles of a substance in solution can be calculated using the relationship
where n = number of .................. of the substance
n
C
V(l)
C = the ................................ of the solution
V(l) = the ............................. of the solution in litres
5. One mole of a gas occupies the ....................... volume at the same temperature and pressure.
The volume of gas that one mole occupies is called the .............................. volume of the gas.
6. The volumes of reactants and product gases can be calculated using the number of moles of
each reactant and product as shown in balanced equation.
There are many different types of calculation using the mole, Avogadro Constant and molar
volumers of a gas (See separate booklet on Higher Chemistry calculations).
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Unit 2
Section (a) – Fuels
1. Petrol can be produced by the ........................................... of the naphtha fraction from crude oil.
2. ........................................... alters the arrangement of atoms in a molecule without necessarily changing
the ..................................... of carbon atoms in the molecule.
3. As a result of reforming many straight chain molecules in the naphtha fraction are changed into
........................................... chain , aromatic and .......................................... molecules.
4. ‘Knocking’ is caused by petrol auto-igniting. Straight chain hydrocarbon molecules are ...................
likely to auto-ignite than branched chain, ....................................... or ................................... hydrocarbon
compounds.
5. Lead compounds used to be added to petrol to prevent .............................................
6. Unleaded petrol contains has a ..................................... proportion of branched chain, aromatic and
cyclic compounds than leaded petrol.
7. In petrol engines the mixture of petrol and air is ignited by a ............................ from a ................... plug.
8. Petrol is a blend of hydrocarbons of different volatilities to suit different conditions.
In winter more of the ............................ volatile smaller molecules are used to prevent the petrol
from ....................................... in winter.
Section (b) – Alternative fuels
1. Ethanol can be produced by the ........................................... of sugar from sugar cane. This makes
ethanol a .......................................... source of energy.
2. The use of methanol as an alternative to petrol has both advantages and disadvantages.
Some advantages are that it is a ................................. burning fuel and it does not contain the
very harmful and carcinogenic compounds such as .................................. found in unleaded petrol.
Some disdvantages are it is a very .......................... compound and it absorbs water so the risk
of .................................. of the metal components in a car engine is increased.
3. ............................ is a renewable fuel produced when biological materials undergo anaerobic
respiration. ............................... contains a mixture of methane and carbon dioxide.
4. Hydrogen is a possible fuel for the future. Hydrogen can be produced by the ........................................
of water – hydrogen forming at the ...................................... electrode. If the electricity for the
.................................................. can be produced fron renewable and non-polluting sources such as
wind power, ................................... or ...................................................... the hydrogen will be a much more
environmentally fuel than fossil fuels.
5. The burning of hydrogen produces only ......................................... Replacing petrol with hydrogen will,
therefore, reduce the amount of .................................... ........................................... in the atmosphere and
reduce the effects of Global Warming.
8
Section (c) – Nomenclature and structural formulae.
1. Alkanes are saturated hydrocarbons having the general formula ...................................
Saturated hydrocarbons contain only .................................... covalent bonds.
2. Cycloalkanes are saturated hydrocarbons having the general formula ...................................
Cycloalkanes have a ........................... ring of carbon atoms
3. Alkenes are unsaturated hydrocarbons having the general formula ...................................
Unsaturated hydrocarbons contain a carbon to carbon .......................... covalent bond.
4. Alkynes are unsaturated hydrocarbons having the general formula ...................................
Alkynes contain a carbon to carbon .......................... covalent bond.
5. Alcohols contain the ......................................... (-OH) functional group and have names that end in –ol.
The general formula for straight chain alcohols (alkanols) is ...........................................
6. Aldehydes and ketones contain the .......................................... (-C=O) functional group.
In aldehydes the ................................. (-C=O) group is on the ................ of the carbon chain.
In ketones the .................................... (-C=O) group is not on the ..................... of the carbon chain.
eg
H
H
H
C
C
H
H
O
H
H
and
C
H
C
C
H
O
C
H
H
H
is the .............................. called ............................. and this is the ............................. called ...........................
7. Aldehydes have a name ending in ..................... while ketones have a name ending in ........................
8. Carboxylic acids contain the ......................................... ( -COOH) functional group. Carboxylic acids have
a name ending of -......................... acid.
O
H
H
H
eg
H
C
C
C
H
H
H
C
O
H
is called ..................................................... acid.
9. You must be able to give the systematic names, molecular formulae, condensed and structural
formula of the alkanes, cycloalkanes, alkenes, alkynes, alcohols, aldehydes, ketones and carboxylic
acids with up to 8 carbon atoms per molecule.
10. Esters contain the ester functional group ( .....................) and have a name ending of –oate.
The first part of an esters name comes from the parent alchol eg ethanol gives an .......................
ester.
The second part of the ester name comes from the parent carboxylic acid eg propanoic acid gives
a ......................................... ester.
eg
H
H
H
C
C
H
H
O
O
H
.......................................
+
H
O
H
C
H
+ ............................... acid -->
9
H
H
C
C
H
H
O
O
+
C
H
O
H
.............................................
+
water
H
11. Aromatic compounds are a homologous series of compounds containing a .............................. type
ring structure.
.............................................. (C 6H6) is the simplest aromatic compound.
12. Benzene does not decolourise ..................................... solution. This shows that there are no
carbon to carbon ................................ bonds in the molecule. Benzene does not react by addition.
The benzene molecule has the following structure
the ring shows the ............................................. electrons
which make the structure very stable.
13. A benzene ring in which one hydrogen atom has been replaced by another atom, or group of atoms,
is known as a ....................................... group. A ................................. group has the formula C 6H5.
Section (d) – Reactions of organic compounds.
(i) Addition reactions
1. Alkenes react quickly by .......................................... across the carbon to carbon ................................. bond.
2. Alkenes can undergo addition of ................................................. (hydrogenation) to form alkanes.
3. Alkenes can undergo addition of .............................................. (halogenation) to form dihalogenalkanes.
eg
H
H
C
H
H
C
+
C
H
Br
H
Br
H
H
H
C
C
C
C
Br
Br
H
H
C
H
H
H
H
H
.................................................. +
bromine
-->
...............................................................
4. Alkenes can undergo addition of hydrogen halides to form halogenoalkanes.
H
eg
H
H
H
C
H
C
H
H
C
+
H
Br
H
..................................... +
H
C
C
H
H
Br
.....................................
H
H
C
H
hydrogen bromide
H
H
H
C
C
C
H
H
H
Br
.....................................
5. Alkenes undergo addition of steam (hydration) to form alcohols.
H
H
H
C
C
H
H
C
C
H
H
H
C
H
+
H
O
C
H
H
O
H
H
.....................................
H
.....................................
H
+
water
10
H
H
H
H
C
C
C
H
H
H
O
H
.....................................
6. Alkynes also react by addition across the ............................. covalent bond.
e.g
H
H
H
C
C
H
+
H
H
C
Br
......................... + hydrogen bromide
H
C
C
H
C
H
H
+
H
Br
Br
Br
.....................................
Br
bromoethene + hydrogen bromide
Br
H
H
C
C
H
H
Br
.....................................
(ii) Dehydration of alcohols
Alcohols can be dehydrated (loss of .............................) by passing the alcohol vapour over heated
aluminium oxide (catalyst) to form alkenes.
eg.
H
H
H
H
H
C
C
C
H
H
H
O
H
--H2O
H
-H2O
H
C
C
H
C
H
H
propan-1-ol
..................................
(iii) Oxidation reactions
1. Alcohols burn in oxygen, or air, forming ..................................... ................................... and water.
eg.
Br
.................. + ............................
C2H5OH + 3O2
2. Alcohols can be classified as primary, ........................................... or .................................................
eg
H
H
H
H
Br
O
O
H
H
C
C
C
H
H
H
H
H
C
H
H
C
H
C
C
H
H
H
H
H
H
H
C
C
C
H
H
H
O
H
H
propan-2-ol is a
........................... alcohol
2-methylpropan-2-ol is a
............................... alcohol
propan-1-ol is a
............................ alc ohol
3. Primary and secondary alcohols can be oxidised using oxidising agents such as
a) acidified potassium dichromate solution. In this reaction the orange .............................. ion
is reduced to the ............................ chromium(III) ion.
b) hot copper(II) oxide. In this reaction the black copper(II) oxide is reduced to .......................
copper metal.
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4. Oxidation of a primary alcohol first produces an ............................................... which then is oxidised
further to a ............................................... acid.
eg
H
H
H
C
C
H
H
H
O
-- 2H
H
H
C
- 2H
H
+ O
C
H
O
H
H
C
H
+O
O
C
O
H
ethanol
..................................
............................. acid
5. Oxidation of a secondary alcohol produces a ...........................................
eg
O
OH
CH3
-- 2H
CH
CH3
CH3
C
- 2H
CH3
propan-2-ol
...................................
6. Aldehydes but not ketones can be oxidised to ................................................. acids.
Suitable oxidising agents include Fehlings solution.
The blue copper(II) ion in the Fehling’s solution is reduced to ............................................ copper(I) ions.
eg
O
O
CH3
CH2
CH2
+ O
CH
CH3
CH2
CH2
C
+O
OH
butanal
.......................................................
Section (e) - Making and breaking esters.
1. Esters are made in a ................................................... reaction between and alcohol and a
...................................................... acid.
eg
O
H
H
H
C
C
+
H
H
O
O
C
+
H
H 2O
....................................
H
H
.................................... acid
The
...........................
methylethanoate + water
shows that the reaction is ....................................................
2. The breaking of an ester into an alcohol and a carboxylic acid is an example of a .................................
O
reaction.
H
C
eg
H
H
C
C
H
O
H
C
H
H
+ H2O
C
H
H
........................
H
............................................................... water
ethanol
12
.............................
+
propanoic acid
Section (f) - Percentage yields.
Percentage yields can be calculated using the masses of reactant(s) and product(s) from balanced
equations (See separate booklet on Higher Chemistry calculations).
Section (g) - Uses of organic compounds.
1. Esters are used as solvents, ....................................... bases. and food ..............................................
2. Carboxylic acids are used to make esters and polyesters. A dilute solution of ethanoic acid is
used a .....................................( a food preservative). Benzoic acid is used as an anti-.................................
food preservative.
3. Halogenoalkanes such as chlorofluorocarbons (CFCs) were widely used as ............................................ in
fridges and freezers, .................................... propellants and for degreasing circuit boards.
CFC’s are however also known to destroy ..................................... (O 3) in the upper atmosphere.
As ozone in the upper atmosphere protects us from ..................................... radiation from the Sun,
the depletion of the ozone layer posed an increased threat of skin ............................. and cataracts.
Manufacture and use of ................... has now been banned.
4. Benzene and its related compounds are important feedstocks – chemicals from which other
chemicals can be synthesised.
Section (h) - Polymers.
1. Ethene is a starting material of major importance in the manufacture of addition polymers.
Ethene is obtained by the ............................... of ethane found in natural gas.
C2H6
ethane
............................. +
ethene
+
H2
hydrogen
2. Ethene can also be obtained by cracking the .......................................... fraction obtained from the
distillation of crude oil.
3. Propene can be obtained by the cracking of .................................... or the cracking of naphtha.
4. Poly(ethene) and poly(propene) are .......................................... polymers made by addition polymerisation.
Show below how 3 propene monomers can undergo addition polymerisation
5. Condensation polymers are made from monomers with ...................... functional groups per molecule.
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6. Polyesters are condensation polymers made by a reacting a ................................... (two –OH groups)
with a ........................................... ( two –COOH) groups
eg Shown below a two molecules of ethane-1,2-diol and benzene-1,4-dicarboxlic acid
Complete the diagram to show part of the polyester formed when these 4 molecules react.
Once drawn, circle the ester links in the polymer.
O
H
O
H
H
C
C
H
H
O
C
O
H
O
C
O
O
H
H
O
H
H
C
C
C
O
H
H
H
O
C
O
O
H
H
H
7. Polyesters are manufactured for use as textile fibres and as resins.
Polyesters used as textile fibres have a ....................................... structure. Polyesters used as resins
have a ........................................ - ............................................. 3 dimensional structure.
8. An amine is a molecule which contains the ............................... (-NH2) functional group.
9. Polyamides are condensation polymers made by reacting a diamine (a molecule with two ...............
groups) with a diacid ( a molecule with two ............................... groups).
Nylon is an example of a polyamide. One version of nylon can be made by reacting
1,6-diaminohexane with hexane-1,6-dioic acid.
Complete the diagram to show part of the nylon polymer formed when these 4 molecules react.
Once drawn, circle the amide links in the polymer
O
H
O
H
H
H
H
C
C
C
C
N
C
H
H
H
H
O
O
CH
H
H
C
C
H
H
H
C
N
O
H
H
H
H
C
C
C
C
H
H
H
H
H
O
O
H
H
H
C
N
H
O
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
H
C
H
H
H
H
C
C
C
C
H
H
H
H
N
O
H
H
H
O
H
H
10. Nylon is so strong that it can be used as an engineering plastic. The strength of nylon is due to the
many .................................. bonds between polymer chains..
O
H
C
H
H
H
H
H
H
O
H
H
O
H
C
C
C
C
H
H
H
H
C
O
14
O
C
O
H
11. Methanal (H2C=O) is an important feedstock in the manufacture of ...................................................
polymers such as Bakelite and urea-formaldehyde resin.
12. Methanal is formed by the oxidation of .....................................................
13. Methanol (CH3OH) is made from .......................................... gas ( a mixture of hydrogen and carbon
monoxide)
CO(g) + 2H2(g)
.......................
14. Synthesis gas, a mixture of ...................................... ................................... and .....................................
can be made by
a) steam reforming of coal
C(s) + H2O(g)
............. + ...............
b) steam reforming of methane CH4(g) + H2O(g)
............ +
................
15. Kevlar is an aromatic polyamide. Kevlar is very light and ....................................... Kevlar is used to
make ...................................................... vests.
16. Poly(ethenol) is a ............................. soluble polymer. It is used to make dissolving hospital laundry
bags.
17. Poly(ethyne) can be treated to make a polymer which ......................................... electricity.
18. Biopol is a ........................................................ polymers.
19. Poly(vinylcarbazole) is a photoconductive polymer used in the manufacture of ......................................
Section (i) – Natural products (Fats and Oils).
1. Fats and oils are naturally occurring esters. Fats and oils can be classified according to their origin
– animal, ......................... or ......................................
2. Oils have ........................................... melting points than fats. The ....................................... melting points
of oils is due to the oils being ............................................................. molecules. The fact that oils are
..................................................... can be shown by the fact that oils quickly decolourise .............................
solution.
3. Fat molecules are more linear in shape than oil molecules. This means that fat molecules can pack
together .................... closely than the more angular oil molecules. Since the fat molecules can pack
together more closely than oil molecules the van der Waals forces between fat molecules are
..................................... than between oil molecules.
4. The conversion of unsaturated oils into solid fats in margarine manufacture involves the addition
of ......................................... (hydrogenation) across some of the double bonds in the oil molecules.
5. Fats and oils are a ....................... concentrated source of energy in the body than carbohydrates.
6. The hydrolysis of a fat (or oil) produces a mixture of ...................................... (propane-1,2,3-triol) and
............................. acids (long chain carboxylic acids) in the ratio of one molecule of
........................................... to three molecules of fatty acid.
15
7. Glycerol (propane-1,2,3-triol) is a trihydric alcohol. This means it contains three ..........................
(-OH) functional groups per molecule.
8. Fats and oils are triglycrerides formed when one molecule of .......................................... combines with
............. fatty acid molecules. The fatty acid molecules may, or may not, be identical.
9. Soaps are produced by the alkaline hydrolysis of ................... and .....................
Section (j) – Natural products (Proteins and amino acids).
1. Proteins are condensation polymers made by combining together many ......................... ...........................
monomer molecules.
2. Proteins are required by the body for ........................ and tissue ...................................
3. Proteins and amino acids contain the elements carbon, hydrogen, oxygen and ......................................
4. The body cannot make all the amino acids required to make body proteins. The supply of these
.............................................. amino acids depends on the diet we eat.
5. Amino acids are molecules that contain a basic amino (...............) group and an acidic .............................
(-COOH) group.
The simplest amino acid is 2-aminoethanoic acid. The extended structural formula is shown below
6. Shown below is a small part of a protein polymer chain
H
H
C
H
O
H
N
H
N
H
H
C
H
C
C
H
C
O
O
H
N
H
C
C
H
N
H
H
Circle the peptide (amide) links shown in this part of the protein molecule.
Draw the extended structural formula of the two amino acids produced by the hydrolysis of
this part of the protein molecule.
16
7. Protein molecules can be classed as fibrous (..................... and thin) or globular (spiral chains folded
into ........................). Fibrous proteins are form many of the bodies major structural units such as
muscles and tendons. Globular proteins form many of the enzymes and hormones in the body. For
example haemoglobin in .................. blood cells and insulin are both ............................. proteins.
8. Enzymes are .................................... molecules. Enzymes are very specific ie they usually only catalyse
one particular reaction in the body. This is because the molecule which is to react must exactly
match the size and shape of the ................................. site on the surface of the enzyme molecule.
9. Changes in pH or an increase in temperature can .................................. the size and shape of the active
site. When this happens the enzyme loses its catalytic effect – we say the enzyme has been
.............................................................
17
Unit 3
Section (a) – The Chemical Industry
1. The chemical industry is a major contributor to the ...................................... of the UK.
2. Stages in the production of a new product involve research, pilot production, scaling up to
............................... scale production and review.
3. A ...................................... is a chemical from which other chemicals can be synthesised or extracted.
4. The major ............. materials used by the chemical industry are fossil ..................., metallic ..............
and minerals, air and ...............................
5. Chemicals may be manufactured in a batch or ................................................... process.
A batch process is used to make ........................ quantities of speciality chemicals where purity
is important eg in the manufacture of drugs and .......................................... The reaction vessels used
to make chemicals in a batch process can be used to make many different .................................... The
reaction vessel has to be carefully cleaned between the making of one product and another.
Continuous processes are more suited to situations where ......................... quantities of a single
product are being made eg fractional distillation of ..................... ......................, making cement and
making iron in a ............................ furnace.
6. The manufacture of chemicals involves both fixed costs and .................................... costs.
Fixed costs are the costs that are incurred no matter how much of the product is made eg
the cost of plant and ............................................, land ............................ and labour costs.
Variable costs depend on how much product is made eg energy costs associated with production.
the cost of ..................... materials and the costs of ......................................... finished product to the
consumer.
7. The chemical industry is capital intensive rather than ...................................... intensive.
Section (b) – Hess’s Law
1. Hess’s law states that the enthalpy change for a reaction is .................................................... of the
reaction pathway.
2. Enthalpy changes can be calculated by an application of Hess’s Law . (See separate booklet on
Higher Chemistry calculations).
Section (c) – Equilibrium
1. Reversible reactions attain a state of dynamic equilibrium when the ................... of the forward
reaction equals the .......................... of the ................................................. reaction.
2. At equilibrium the concentrations of the reactants and products remains ..........................................
It is important to realise that at equilibrium the concentrations of reactants and products will
not usually be .....................................
3. Increasing the concentration of one of the reactants will speed up the .................................... reaction
and shift the equilibrium from ....................... to right ie it will result in an increased yield of
...............................
18
4. ............................................ the temperature will favour the endothermic step in a reversible reaction.
............................................. the temperature will favour the exothermic step in a reversible reaction.
5. Increasing the pressure will favour the reaction which forms fewer ........................ phase molecules.
Decreasing the pressure will favour the reaction which forms more ....................... phase molecules.
6. In the Haber process
N2(g) + 3H2(g)
2NH3(g)
DH = --92 kJ
The yield of ammonia will increase by ........................................ the temperature, .................................. the
pressure and recycling unreacted nitrogen and hydrogen (as this has the same effect as increasing
the ........................................... of these reactants).
7. Low temperatures favour the formation of .......................................... in the Haber Process. However,
if the temperature is too low the ............................. of production of ammonia is too slow.
A compromise temperature of about ................. is used to give a reasonable yield at a reasonable rate.
8. A catalyst speeds up the rate at which a reaction reaches .................................................
Using a catalyst does not, however, alter the position of the equilibrium as it speeds up the forward
and reverse reactions by the ........................ amount.
Section (d) – Acids and Bases
1. The pH scale is a continuous number scale that runs from below .............. to above ...................
2. A solution with a pH of ......... has a hydrogen ion concentration of 1 x 10 -1 mol l-1.
A solution with a pH of ......... has a hydrogen ion concentration of 1 x 10 -2 mol l-1.
A solution with a pH of ......... has a hydrogen ion concentration of 1 x 10-3 mol l-1 etc.
3. In water and aqueous solutions with a pH of 7 the ......................................... ion concentration, [H +(aq)],
= the ................................... ion concentration, [OH -(aq)]. Both have a value of .......................... mol l-1.
4. The concentration of either hydrogen ions or hydroxide ions in an aqueous solution can be
calculated from the concentration of the other using the fact that
[H+(aq)] x [OH-(aq)] = ................................. mol2 l-2.
5. In water, and aqueous solutions, there is an .............................................. between hydrogen ions,
hydroxide ions and water molecules.
H2O(l)
H+(aq) + OH --(aq)
The equilibrium lies far to the ............................... ie most of the water exists as ...................................
6. In aqueous solution strong acids are .................................... ionised. Weak acids are only ........................
ionised.
7. When equimolar (same concentration) solutions of a strong acid and a weak acid are compared the
strong acid has a ................................. pH, a ...................................... conductivity and a ................................
reaction with magnesium than the weak acid.
However, the same volume of both the strong acid and the weak acid will neutralise the ...................
number of moles of a given alkali.
19
8. Carboxylic acids such as ethanoic acid are ........................ acids.
CH3COOH(l) + H2O(l)
......................................... + .......................
The equilibrium for the reaction lies far to the ............................ ie it exists mostly as
...................................... with very few ............................. present.
9. Solutions formed when sulphur dioxide and carbon dioxide dissolve in water are ...................... acids.
For example:SO2(g) + H2O(l)
....................... + ................................
sulphurous acid
As with all weak acids the equilibrium lies far to the ................................
10. In aqueous solutions strong bases are ................................ ionised. Weak bases are only ........................
ionised.
11. When equimolar (same concentration) solutions of a strong base and a weak base are compared the
strong base has a ................................. pH, and a ...................................... conductivity.
However, the same volume of both the strong base and the weak base will neutralise the ...................
number of moles of a given acid.
12. The solution formed when ammonia gas (NH3) dissolves in water is a ..................... base.
NH3(g) + H2O(l)
.......................... + ..............................
ammonium hydroxide solution
13. A soluble salt formed when a strong acid reacts with a strong base is ............................................
(pH = ..........)
eg sodium chloride is the salt formed when ............................................... acid reacts with
......................................... ............................................ Sodium chloride is a ................................ salt
14. A soluble salt formed when a strong acid reacts with a weak base is ............................................
(pH < ..........)
eg ammonium sulphate is the salt formed when ................................................. acid reacts with
...................................... ....................................... Ammonium sulphate is a ............................. salt.
15. A soluble salt formed when a weak acid reacts with a strong base is ............................................
(pH > ..........)
eg potassium carbonate is formed when ...................................... acid reacts with
..................................... ..................................... Potassium carbonate is a ........................... salt.
16. Soap solutions are ............................................ as soap is the salt of a weak acid and a strong base.
20
17. Ammonium chloride dissolves in water to form an ............................... solution.
Ammonium chloride is very ..................................... in water and produces lots of ammonium and
chloride ions in solution.
NH4Cl(s) + H2O(l)
NH4+(aq) + Cl --(aq)
In water there is an ........................................ between water molecules, hydrogen and hydroxide ions.
H2O(l)
H+(aq) + OH--(aq)
As ammonium hydroxide is a ......................... base there is a strong tendency for the ammonium and
hydroxide ions to react to form ammonia and water molecules.
NH4+(aq) + OH--(aq)
NH3(g) + H2O(l)
There is no such tendency for hydrogen and chloride ions to react as hydrochloric acid is a
................................. acid and remains fully ionised.
The removal of the ................................... ions leaves behind a solution with an excess of hydrogen
ions – ie an ............................... solution.
Section (e) – Oxidation and Reduction reactions
1. Oxidation involves a ........................... of electrons. Reduction involves a ...................... of electrons.
2. An oxidising agent is itself ......................................... An oxidising agent is an electron ...............................
A reducing agent is itself ........................................ A reducing agents is an electron ....................................
For example in the displacement reaction shown below
Zn(s) + 2Ag+(NO3--)2(aq)
Zn2+ (NO3--)2(aq) + 2Ag(s)
In this reaction the zinc is the .......................................... agent. The ion-electron equation for
the oxidation of the zinc atoms is:...........
................... + .................
In this reaction the silver(I) ions are the .......................................... agent. The ion-electron equation
for the reduction of the silver(I) ions is:............... + .............
..................
In this reaction the nitrate ions are called the ........................................................... ions.
3. To combine together the oxidation and reduction ion-electron equations into the balanced
redox equation we must first balance the number of .................................... lost and gained in the
oxidation and reduction reactions.
eg oxidation Fe2+(aq)
Fe3+(aq) + e -reduction MnO4--(aq) + 8H+(aq) + 5e -Mn2+(aq) + 4H2O(l)
To construct the redox equation we must first multiply the oxidation reaction by ...........
This gives
........... Fe2+(aq)
........ Fe3+(aq) + ........ e -Adding together the oxidation and reduction reactions gives the balanced redox equation
............................................................................................................................. .........................................................
From this equation it can be seen that ............ moles of Fe2+(aq) ions are oxidised by ............... mole of
MnO4-(aq) ions.
21
4. Given the reactant and product species an ion-electron equation reaction involving H+(aq) and H2O(l)
can be constructed.
eg Construct the ion-electron equation for the conversion of iodate (IO3-) ions to iodine.
Step 1:- balance the iodine atoms on each side:
..........IO3--(aq)
I2(aq)
Step 2:- Add H2O(l) molecules to balance the oxygen atoms on each side
..........IO3--(aq)
I2(aq) + ........H2O(l)
Step 3 :- Add H+(aq) ions to balance the hydrogens on each side
..........IO3--(aq) + .........H+(aq)
I2(aq) + ........H2O(l)
Step 4:- Add electrons to balance the charge on each side
..........IO3--(aq) + .........H+(aq) + ....... e --
I2(aq) + ........H2O(l)
In this reaction the iodate is ................................. to iodine. The iodate is acting as an
............................................ agent.
5. You must be able to carry out calculations based on redox titrations. . (See separate booklet on
Higher Chemistry calculations).
Section (f) – Faradays Laws of Electrolysis
1. The production of one mole of an element always requires n x .............................. C where n is the
number of electrons in the relevant ion-electron equation.
For example:- the electrolysis of copper(II) chloride produces copper at the ..................................
electrode and chlorine at the ................................. electrode.
The ion-electron equation for the production of copper is:Cu2+(aq) + 2e -Cu(s)
The production of one mole of copper requires ........... x 96500C (or ........... Faradays)
The ion-electron equation for the production of chlorine is:2Cl --(aq)
Cl2(g) + 2e -The production of one mole of chlorine requires ........... x 96500C (or ........... Faradays)
2. You must be able to carry out calculations based on Faradays laws of electrolysis. (See separate
booklet on Higher Chemistry calculations).
3. The charge on one mole of electrons is one Faraday (................................ Coulombs)
Section (g) – Nuclear Chemistry
1. Radioactive decay involves changes inside the .................................. of atoms.
2. Unstable nuclei (radioisotopes) become more stable by the ........................ of energy and the emission
of radiation.
3. The stability of a nucleus depends on the ratio of the ............................ to ................................. inside
the nucleus.
22
4. The table below summarises the properties of alpha (), beta () and gamma () radiation.
Type of
radiation
Symbol
Charge
Mass
alpha
 or 42He 2+
.............
.............
towards the ........................... plate
 or
............
............
towards the ...................... plate
.............
.............
...................................................
beta
gamma
e --
0
-1

Deflection in
electric field
5. You must be able to write balanced nuclear equations involving protons, neutrons, alpha and
beta particles
eg Complete the nuclear equations for the following reactions:1) Uranium-238 decaying by alpha emission
U
238
92
.................... + ...............
2) Radium-228 decaying by beta emission
Ra
228
88
.............. + ..................
3) Nitrogen-14 capturing a neutron and the then breaking down into carbon-14 and a proton
N +
14
7
n
1
0
...................... + .................
6. The .................... -........................ is the time taken for the activity, or mass, of a radioisotope to
fall to half its original value.
7. ................ - ........................ of radioisotopes of independent of the chemical or physical state of the
radioisotope.
For example, the half life of lead-214 is ......................... minutes (see data book page 8). It could be
214
present as 10g of 214
Pb or as 0.5 g of 82Pb or as 5kg of 214
PbO the half life would be the ..................
82
82
in all three.
8. The mass of a radioisotope, half-life or time elapsed can be calculated given the value of the other
two variables.
For example:A 10g sample of Thorium-234 has a count rate of 440 counts per minute. Using the half life
given in the data booklet, calculate
i) the mass of Thorium-234 remaining after 96.4 days
ii) the count rate after 72.3 days.
23
9. Radioisotopes are used in
a) Medicine.
Cobalt-60 emits gamma rays which can be used to kill deep seated ................................. tumours.
Radioisotopes of iodine such as Iodine-131 can be used to detect anomalies in the ........................
gland.
b) In industry.
Gamma rays from a cobalt-60 source can be used to check for hairline ............................ in metals.
A beta emitter can be used to continuously monitor the ........................................... of paper in a
paper mill.
c) Scientific research
Radioactive carbon-14 can be used to date .......................................... remains. By comparing the
ratio of carbon-14 to carbon-.......... in the organic remains to a sample from the present day the
age of the remains can be estimated.
Radioactive isotopes of phosphorus can be used to determine the use of ..........................................
fertilisers by different plants.
d) Energy production by nuclear fission
Uranium-238 is a radioisotope of uranium that can undergo nuclear ................................... (splitting
of a ....................... nucleus into two, or more, lighter nuclei). In the process energy is
....................... ....................
The reaction is a chain reaction in which the fission is started when a uranium-238 atoms
absorbs a ............................ forming an unstable uranium-239 atom which then breaks down
(fissions) forming two or more stable atoms and releasing more ............................... which can go
on to react with more uranium-238 atoms.
e) Energy production in the Sun by nuclear fusion.
Inside the Sun light atoms are ............................. together to form heavier atoms and in the
process ......................... is given out.
4
eg 2H + 3H
He + ............
1
1
2
10. There are advantages and disadvantages of using fossil fuels and nuclear energy to
electricity .
The advantages of fossil fuels is that they are relatively .................... and readily .................................
The disadvantages of fossil fuels is that they give out harmful gases such as .......................................
.................................................... which is the main Greenhouse gas linked to Global warming.
The advantages of nuclear fuels is that they do not give out harmful ............................... which
cause damage to the environment.
The disadvantage of nuclear power stations is that they produce harmful ...........................................
waste which will remain highly .......................................... for .......................................... of years.
11. Elements are created inside ............................ by the process of nuclear ................................
24
Questions from past SQA Papers (Section B) for you to test your knowledge.
Unit 1
Section (a) – Reaction Rates (including calculations on excess)
2004 Question 6
2003 Question 7 (a)
2003 Question 11 (c)
2002 Question 8 (a)
2001 Question 3
2000 Question 2
Section (b) – Enthalpy
2004 Question 8 (a)
2001 Question 13(b)
2003 Question 11 (a) and (b)
2000 Question 5
Section (c) – Patterns in the Periodic Table.
2004 Question 9
2002 Question 1
2000 Question 7(a)
2000 Question 13 (a)
2003 Question 9
2002 Question 9
2002 Question 6
2000 Question 6
2001 Question 12
Section (d) - Bonding, structure and properties of the first 20 elements
2004 Question 1
Section (e) – Bonding and structure of compounds and Section (f) – Properties associated with
bonding
2002 Question 7
2003 Question 12
2001 Question 11(b)
Section (g) – The Mole
2004 Question 2 (b) (i)
2003 Question 10 (c)
2001 Question 4
2000 Question 13 (b)
2004 Question 11
2003 Question 11 (d)
2000 Question 7(b) (c)
2003 Question 3 (c)
2002 Question 12 (b)
2000 Question 10
Unit 2
Section (a) – Fuels
2004 Question 10 (b)
2001 Question 1
2003 Question 1 (b)
2001 Question 11 (a) (iii)
Section (b) – Alternative fuels
2004 Question 4 (c)
Section (c) – Nomenclature and structural formulae.
2004 Question 10 (a) (iii)
2003 Question 1(a)
2001 Question 11 (a) (i)
2000 Question (b)
Section (d) – Reactions of organic compounds.
2004 Question 4 (a) and (b)
2004 Question 10 (a) (i) and (ii)
2003 Question 8
2002 Question 5 (a) (ii) to (iv)
2001 Question 2(a)
2001 Question 10 (a)
Section (e) - Making and breaking esters.
2004 Question 14
2003 Question 6 (a) and (b)
2000 Question 9 (a)
25
2001 Question 8 (a)(i)
2004 Question 13
2002 Question 13
2000 Question 11
2002 Question 5 (a)(i) (b)
Section (f) - Percentage yields.
2003 Question 6 (c)
2002 Question 10 (d)
2000 Question 14 (a) (iii)
Section (g) - Uses of organic compounds.
Section (h) - Polymers.
2003 Question 2
2001 Question 5
2003 Question 5
2001 Question 9
2002 Question 11
2000 Question 14(b)
Section (i) – Natural products (Fats and Oils).
2002 Question 3
2001 Question 17
Section (j) – Natural products (Proteins and amino acids).
2004 Question 7
2001 Question 8 (b)
2000 Question 3
Unit 3
Section (a) – The Chemical Industry
2003 Question 14
2000 Question 17
Section (b) – Hess’s Law
2004 Question 15 (a)
2001 Question 7
2003 Question 4
2001 Question 10(b)
2002 Question 4 (b)
2000 Question 12 (c)
Section (c) – Equilibrium
2004 Question 5
2001 Question 2(b)
2003 Question 10 (a) and (b)
2000 Question 8
2002 Question 10 (a) to (c)
Section (d) – Acids and Bases
2004 Question 8 (b)
2003 Question 13
2001 Question 13(a)
2004 Question 15 (b)
2002 Question 8 (b)
2000 Question 9 (b)
2003 Question 7 (c)
2001 Question 8(a)(iii)
Section (e) – Oxidation and Reduction reactions
2004 Question 12
2004 Question 16
2002 Question 15
2001 Question 14
2000 Question 15
Section (f) – Faradays Laws of Electrolysis
2004 Question 3
2002 Question 14 (a)
Section (g) – Nuclear Chemistry
2004 Question 2 (a) and 2 b(ii)
2001 Question 1
2003 Question 3 (a) and (b)
2000 Question 4
Miscellaneous problem solving type questions
2004 Question 17
2003 Question 7 (b)
2002 Question 4 (a)
2001 Question 6 (a) to (c)
2000 Question 14 (a)
2000 Question 16
26
2002 Question 14 (b)
2001 Question 15
2001 Question 6(d)
2002 Question 2
2003 Question 15
2001 Question 16)
Questions from past SQA Papers (Section A M/C) for you to test your knowledge.
Unit 1
Section (a) – Reaction Rates (including calculations on excess)
2004 Q 6
2003 Q 6
2002 Q 4
2001 Q 5
2000 Q 5,6,33
2002 Q 5, 6
2001 Q 6,7
2000 Q 7
Section (b) – Enthalpy
2004 Q 8
2003 Q 7
Section (c) – Patterns in the Periodic Table.
2004 Q 9, 11
2003 Q 9
2001 Q 8
2000 Q 8,9
Section (d) - Bonding, structure and properties of the first 20 elements
2004 Q 10
2003 Q 10
2001 Q 9
Section (e) – Bonding and structure of compounds and Section (f) – Properties associated with
bonding
2004 Q 12
2003 Q 8, 11, 12, 13
2002 Q 7, 8
2001 Q 10,11, 31
2000 Q 10
2001 Q 4, 12,13,14
2000 Q 11,32
Section (g) – The Mole
2004 Q 7, 13,14,15
2003 Q 5, 14, 15, 16
2002 Q 10, 11, 12
Unit 2
Section (a) – Fuels
2004 Q 16
2002 Q 13,14
2001 Q 16
Section (b) – Alternative fuels
2001 Q 15
Section (c) – Nomenclature and structural formulae.
2004 Q17, 18
2003 Q 18, 19
2002 Q15
2001 Q 17
Section (d) – Reactions of organic compounds.
2004 Q21
2003 Q 20, 21, 22,
2002 Q 17,18
2001 Q 18, 19, 21, 33
2000 Q 13
Section (e) - Making and breaking esters.
2002 Q 16
2001 Q 20
2000 Q 12
Section (f) - Percentage yields.
Section (g) - Uses of organic compounds.
2003 Q 23
2001 Q 22
2000 Q 14
Section (h) - Polymers.
2004 Q19, 23, 24,25
2003 Q 24
2002 Q 19,20,21 2001 Q 23
2000 Q 15, 16,17
Section (i) – Natural products (Fats and Oils).
2004 q 26, 27,
2003 Q26, 27
2000 Q 1934
Section (j) – Natural products (Proteins and amino acids).
2004 Q 28
2003 Q 25, 28
2002 Q 22
27
2001 Q 24, 32(a)
2000 Q 18
Unit 3
Section (a) – The Chemical Industry
2004 Q 29,
2003 Q 29
2001 Q 25, 26
2000 Q 20
Section (b) – Hess’s Law
2004 Q 30,
2003 Q 30
2002 Q 23
2000 Q 21
2002 Q 24
2001 Q 27, 34
2002 Q25,26
2001 Q 28, 29,35
Section (c) – Equilibrium
2004 Q 22, 31, 32,
2003 Q 31, 32
2000 Q 22,23
Section (d) – Acids and Bases
2004 Q 33, 34, 36,
2003 Q 33, 34,35
2000 Q 24,25,26, 31
Section (e) – Oxidation and Reduction reactions
2004 Q 37, 38,
2003 Q 36, 37
2002 Q 27
2000 Q 27
Section (f) – Faradays Laws of Electrolysis
2002 Q 28
2000 Q 28
Section (g) – Nuclear Chemistry
2004 Q 39, 40
2003 Q 38, 39, 40
2002 Q 29,30
2001 Q 30, 32(b)
2000 Q 29,30
Questions from S Chemistry course
2004 Q 1,2,3,4
2003 Q 1,2,3,4
2002 Q 1,2,3
Miscellaneous questions
2004 Q 5, 20, 35
2002 Q 9
28
2001 Q 1,2,3
2000 Q 1,2,3,4