Williamwood High School
Chemistry
National 4 and 5
Learning Outcomes
Name
________________________
Class________
Unit 1 Chemical Changes and Structure
Key area: Atomic theory
Elements
1.
Everything in the world is made from over 100 elements. (Nat 4)
2.
An element is a substance containing only 1 type of atom. (Nat 4)
3.
Chemists have classified elements by arranging them in the Periodic Table. (Nat 4)
4.
Each element has a name and a symbol. (Nat 4)
5.
Elements are classified in different ways: metal / non-metal, naturally occurring / manmade, solid / liquid / gases. (Nat 4)
6.
A group is a column of elements in the Periodic Table. (Nat 4)
7.
A period is a row of elements in the Periodic Table. (Nat 4)
8.
The groups include
Group 1 the Alkali metals (Nat 4)
Group 7 the Halogens (Nat 4)
Group 8 (0) the Noble gases (Nat 4)
9.
The alkali metals are a family of very reactive metals and are stored under oil. (Nat 4)
10.
The halogens are a family of very reactive non-metals. (Nat 4)
11.
The noble gases are a family of very unreactive non-metals. (Nat 4)
12.
The transition metals are found between Groups 2 and 3 in the Periodic table. (Nat 4)
13.
Elements in any one group of the Periodic Table show similar chemical properties.
(Nat 4)
Compounds and mixtures
14.
Mixtures occur when two or more substances come together without reacting. (Nat 4)
15.
Air is a mixture of gases. (Nat 4)
16.
The gases in air include 79% nitrogen, 20% oxygen, carbon dioxide and Noble gases.
There is no hydrogen in air. (Nat 4)
17.
The test for oxygen is that it relights a glowing splint. (Nat 4)
18.
There is not enough oxygen in the air for the test to be positive. (Nat 4)
19.
A solution is a mixture formed when a solute dissolves in a solvent. (Nat 4)
20.
A compound is a substance that is made up of two or more elements that are
chemically joined. (Nat 4)
1
Solubility
21.
A substance which dissolves in a liquid is soluble; a substance which does not dissolve is
insoluble. (Nat 4)
22.
A saturated solution is one in which no more substance can be dissolved at that
temperature. (Nat 4)
23.
A dilute solution has a lower concentration of dissolved substance than a concentrated
solution. (Nat 4)
24.
A solution is diluted by adding more solvent. (Nat 4)
25.
State symbols (s) (l) (g) and (aq) can be used to show the state of a substance. (Nat 4)
Chemical reactions
26.
All chemical reactions involve the formation of one or more new substances. (Nat 4)
27.
Compounds are formed when elements react together. (Nat 4)
28.
Chemical reactions can be identified by changes in appearance of
substance, including colour change, gas evolved, a solid formed. (Nat 4)
29.
Chemical reactions can be identified by energy changes. (Nat 4)
30.
Exothermic reactions release energy to the surroundings and the
products have less chemical energy than the reactants. (Nat 4)
31.
Endothermic reactions take in energy from the surroundings and the products have
more chemical energy than the reactants. (Nat 4)
Naming compounds
32.
Most compounds with a name ending in ‘-ide’ contain the two elements indicated. (Nat 4)
33.
The ending ‘ite’ or ‘-ate’ indicates the additional element oxygen. (Nat 4)
Sub-atomic particles
34.
Every element is made up of very small particles called atoms. (Nat 4)
35.
The atom has a nucleus, which contains protons and neutrons, with electrons moving
around outside the nucleus in specific energy levels. (Nat 4)
36.
Atoms of different elements have a different number of protons, called the atomic
number. (Nat 4)
37.
The elements of the Periodic Table are arranged in terms of their atomic number and
chemical properties. (Nat 4)
38.
Elements with the same number of outer electrons are in the same group and have similar
chemical properties. (Nat 4)
2
Mass & charge of particles
39.
Protons have a charge of one-positive, neutrons are neutral and electrons have a
charge of one-negative. (Nat 4)
40.
An atom is neutral because the numbers of protons and electrons are equal. (Nat 4)
41.
Protons and neutrons have an approximate mass of one atomic mass unit and
electrons, in comparison, have virtually no mass. (Nat 4)
42.
An atom has a mass number which equals the number of protons plus neutrons. (Nat 4)
43.
The number of protons, neutrons and electrons can be found from the atomic number and
mass number, and vice versa. (Nat 4)
44.
Candidates should be able to draw and label a diagram of an atom showing the location,
charge and mass of the sub-atomic particles. (Nat 5)
Nuclide notation
45.
Atoms can be represented by nuclide notation, e.g.
35
17
Cl. (Nat 5)
Isotopes
46.
Isotopes are atoms with the same atomic number but different mass numbers. (Nat 5)
47.
Most elements exist as a mixture of Isotopes. (Nat 5)
RAM
48.
The relative atomic mass of an element is an average of all the isotopes of that element.
(Nat 5)
49.
The relative atomic mass of an element is rarely a whole number. (Nat 5)
50.
Candidates should be able to work out which is the most abundant isotope given the mass
numbers of the isotopes and the relative atomic mass of the element. (Nat 5)
Key area: Rates of reaction
Factors affecting rate
51.
The rates of reactions are affected by changes in concentration, particle size and
temperature. (Nat 4)
52.
The collision theory can be used to explain the effects of concentration and surface area
on reaction rates. (Nat 5)
53.
Catalysts are substances which speed up some reactions and are not used up by the
reactions. They also allow the reaction to take place at a lower temperature. (Nat 4)
3
54.
Heterogeneous catalysts are catalysts which are in a different state from the reactants.
(Nat 5)
55.
Homogeneous catalysts are catalysts which are in the same state as the reactants. (Nat
5)
56.
Catalysts are used in many industrial processes. (Nat 4)
57.
There are many everyday examples of uses of catalysts. (Nat 4)
58.
Enzymes are biological catalysts which catalyse the chemical reactions which take place
in the living cells of plants and animals. (Nat 4)
Measuring reaction rates
59.
Reactions can be followed by measuring changes in concentration, mass and volume
of reactants and products. (Nat 4)
60.
The rate of a reaction, or stage in a reaction, is proportional to the reciprocal of the time
taken,
i.e.
61.
Rate =
1
t
(Nat 5)
The average rate of a reaction, or stage in a reaction, can be calculated from a graph by
Average rate of reaction =
∆𝒒𝒖𝒂𝒏𝒕𝒊𝒕𝒚
∆𝒕𝒊𝒎𝒆
(Nat 5)
Key area: Bonding, structure and properties
Covalent bonding
62.
Atoms can be held together by bonds. (Nat 5)
63.
In forming bonds, atoms can achieve a stable electron arrangement the same as a
noble gas. (Nat 5)
64.
In covalent bonds atoms share pairs of electrons. (Nat 5)
65.
A covalent bond is a result of two positive nuclei being held together by their
common attraction for the shared pair of electrons. (Nat 5)
66.
Covalent bonds are strong forces of attraction. (Nat 5)
67.
More than one bond can be formed between atoms leading to double and triple bonds.
(Nat 5)
4
Covalent molecular structure
68.
A molecule is a group of atoms held together by covalent bonds. (Nat 5)
69.
A diatomic molecule is made up of two atoms chemically joined. (Nat 5)
70.
Candidates should be familiar with the seven diatomic elements. (Nat 5)
71.
Hydrogen, nitrogen, oxygen, the halogens and carbon monoxide are examples of diatomic
molecules. (Nat 5)
72.
Bonding diagrams can be drawn to show how the outer electrons are shared to form the
covalent bond(s) in a molecule. (Nat 5)
73.
Diagrams can be drawn to show the shape of simple two-element molecules. (Nat 5)
74.
Candidates should be able to draw the shapes of linear, bent, pyramidal and
tetrahedral two-element molecules. (Nat 5)
75.
Covalent substances can consist of discrete molecules. (Nat 5)
76.
The molecular formula for a discrete covalent substance gives the actual number of
atoms in each molecule. (Nat 4)
77.
Formulae can be written from models or given molecular pictures. (Nat 4)
78.
The bonds between the molecules are weaker than the covalent bonds within
molecules. (Nat 4)
Covalent network structure
79.
A covalent network structure consists of a giant lattice of covalently bonded atoms. (Nat
5)
80.
The formula for a covalent network substance gives the simplest ratio of atoms of each
element. (Nat 5)
Ions
81.
An ion is a charged particle and they are formed when an atom loses or gains
electrons. (Nat 5)
82.
Positive ions are formed by metal atoms losing electrons and negative ions are
formed by non-metal atoms gaining electrons. (Nat 5)
83.
Candidates should be able to use nuclide notation for ions, e.g.
16
8
O2- (Nat 5)
Ionic bonding
84.
Ionic compounds are usually formed when metals combine with non-metals. (Nat 5)
85.
Ionic bonding is the electrostatic force of attraction between oppositely charged ions.
(Nat 5)
5
Ionic lattice
86.
An ionic structure consists of a giant lattice of oppositely charged ions. (Nat 5)
87.
The formula for an ionic compound gives the simplest ratio of positive ions to negative
ions. (Nat 5)
Metallic bonding
88.
Metallic bonding is the electrostatic force of attraction between positively charged ions
and delocalised electrons. (Nat 5)
Physical properties of substances
Electrical conductivity of elements and compounds
89.
Metal elements (solids, liquids) and carbon (graphite) are conductors of electricity
because they contain free, delocalised electrons. (Nat 4)
90.
Most non-metal elements are non-conductors of electricity. (Nat 4)
91.
Covalent substances (solids, liquids, solutions) do not conduct electricity since they are
made up of molecules which are uncharged. (Nat 4)
92.
Ionic compounds only conduct electricity when dissolved in water or when molten.
(Nat 4)
93.
Ionic compounds do not conduct electricity in the solid state since the ions are not free
to move, but these compounds do conduct electricity when dissolved in water or when
molten as the ions are now free to move. (Nat 5)
Electrolysis
94.
An electric current is a flow of charged particles. (Nat 4)
95.
Electrolysis is the breakup of an ionic compound into its elements using electricity. (Nat
4)
96.
Electrolytes are ionic solutions and molten ionic compounds which complete the circuit.
(Nat 4)
97.
A d.c. supply must be used if the products of electrolysis are to be identified at each
electrode. (Nat 5)
98.
Positive metal ions gain electrons at the negative electrode and negative non-metal ions
lose electrons at the positive electrode. (Nat 5)
99.
Certain ions are able to be identified by their colour e.g. copper ions blue. (Nat 4)
100.
The results of electrolysis experiments can be illustrated by the migration of coloured ions
e.g. copper chromate experiment. (Nat 4)
6
Melting and boiling points and solubility
101.
Ionic compounds and covalent network substances have high melting and boiling
points due to the strong forces of attraction which need to be overcome. (Nat 5)
102.
Discrete covalent substances have low melting and boiling points due to the weak
forces of attraction between the molecules that need to be overcome. (Nat 5)
103.
Ionic compounds are usually soluble in water. (Nat 5)
104.
Covalent substances which are insoluble in water may dissolve in other solvents. (Nat
5)
Key area: Chemical formulae
Write formula from named compounds
105.
The ratio in which elements combine to form two element compounds can be determined
by the valency rule. (Nat 4)
Chemical & ionic formulae (including Group Ions)
106.
Formulae can be written for names using prefixes, including mono-, di-, tri-, tetra-. (Nat
4)
107.
Formulae for two-element compounds are written using valencies from the Data booklet.
(Nat 4)
108.
Formulae can be written for compounds which include Roman numerals in their names,
e.g. iron (III) chloride. (Nat 5)
109.
Formulae can be written for compounds involving group ions but not requiring brackets,
e.g. Na2SO4. (Nat 5)
110.
Formulae requiring brackets can be written for compounds, e.g. Mg(OH)2. (Nat 5)
111.
Candidates should be able to extend the use of brackets to writing formulae for
compounds which include Roman numerals in their names, e.g. iron (III) hydroxide,
Fe(OH)3. (Nat 5)
112.
Candidates should be able to use nuclide notation for ions, e.g.
23
+
11 Na .
(Nat 5)
Chemical equations
113.
Word equations can be written to describe the progress of a chemical reaction from
reactants to products. (Nat 4)
7
114.
Formulae equations can be written to describe the progress of a chemical reaction from
reactants to products. (Nat 4)
115.
Formulae equations can be balanced to show the relative number of moles of
reactant(s) and product(s). (Nat 5)
Formula mass
116.
The relative formula mass of a substance can be calculated from the relative atomic
masses. (Nat 4)
Gram Formula Mass
117.
The gram formula mass of any substance is known as one mole. (Nat 5)
Mole calculations and calculations based on equations
118.
Candidates should be able to calculate the mass and the number of moles using the
triangle below. (Nat 5)
m
n
119.
GFM
Where:
m= Mass (g)
n= number of moles (moles)
GFM= Gram Formula Mass (GFM)
The mass of a reactant or product can be calculated using a balanced equation. (Nat 5)
Key area: Acids and bases
Acids, alkalis and the pH scale
120.
The pH scale is a continuous range from below 0 to above 14. (Nat 4)
121.
Acids have a pH of less than 7; alkalis have a pH of more than 7; pure water and
neutral solutions have a pH equal to 7. (Nat 4)
122.
Acids and alkalis are in common use in both the laboratory and the home. (Nat 4)
123.
Household acids include vinegar, lemonade, soda water and Coke. (Nat 4)
124.
Household alkalis include baking soda, dishwashing powder and bleach. (Nat 4)
125.
Universal indicator and pH paper can be used to determine whether solutions are acidic
or alkaline. (Nat 4)
126.
Red cabbage is an example of a natural indicator. (Nat 4)
8
Uses of acids in food and drink and their impact on health
127.
Acids play an important role in the food and drink industry. (Nat 4)
128.
These acids have an impact on human health. (Nat 4)
129.
Drinks with a low pH can cause teeth to erode. (Nat 4)
130.
Positive use of acids (as preservatives in foods, HCl used in the body for digestion). (Nat
4)
131.
Lightning provides nitrates to soil. (Nat 4)
Dilution effect on pH
132.
The effect of dilution on the pH of an acid or alkali is explained in terms of the
decreasing concentration of hydrogen and hydroxide ions. (Nat 5)
Formation of acids and alkalis
133.
Non-metal oxides which dissolve in water produce acidic solutions. (Nat 5)
134.
An acidic solution contains more hydrogen ions than hydroxide ions. (Nat 5)
135.
Metal oxides and hydroxides which dissolve in water produce alkaline solutions. (Nat 5)
136.
An alkaline solution contains more hydroxide ions than hydrogen ions. (Nat 5)
137.
Ammonia dissolves in water to produce an alkali. (Nat 5)
Sources of carbon dioxide
138.
Carbon dioxide is a by-product of burning fossil fuels. (Nat 4)
139.
Another source of CO2 is during cement manufacturing for use in new buildings. (Nat 4)
Other non-metal oxide pollutants
140.
Carbon dioxide, sulphur dioxide and oxides of nitrogen are produced as a result of
our continued use of fossil fuels. (Nat 4)
141.
These non-metal oxides are produced in nature and are linked to environmental
problems, including acid rain, global warning and ocean acidification. (Nat 4)
142.
Acid rain has damaging effects on buildings made from carbonate rock, structures
made of iron and steel, soils and plant and animal life. (Nat 4)
pH in relation to hydrogen ion concentration
143.
pH is a measure of the hydrogen ion (H+) concentration. (Nat 5)
Neutralisation reactions
144.
Neutralisation is the reaction of acids with bases. (Nat 4)
145.
Metal oxides, metal hydroxides and metal carbonates are examples of bases. (Nat 4)
9
146.
Neutralisation moves the pH of an acid up towards 7. (Nat 4)
147.
Neutralisation moves the pH of an alkali down towards 7. (Nat 4)
148.
In the reaction of an acid with an alkali the hydrogen ions and hydroxide ions form water.
(Nat 4)
149.
Bases which dissolve in water form alkalis. (Nat 4)
150.
Everyday examples of neutralisation include the treatment of acid indigestion and
using lime to reduce acidity in soil and lochs. (Nat 4)
151.
An acid reacts with some metals to give off hydrogen gas. (Nat 4)
152.
In the reaction between some metals and acids, hydrogen ions form hydrogen
molecules. (Nat 4)
153.
The test for hydrogen is that it burns with a ‘pop’. (Nat 4)
Preparation of salts
154.
A salt is a compound in which the hydrogen ions of an acid have been replaced by
metal ions (or ammonium ions). (Nat 4)
155.
Salts are formed in the reaction of acids with bases or metals. (Nat 4)
156.
Hydrochloric acid forms chloride salts, sulphuric acid forms sulphate salts and nitric
acid forms nitrate salts. (Nat4)
157.
Some nitrogen salts, including ammonium nitrate, ammonium sulphate and potassium
nitrate are made by neutralisation reactions for use as fertilisers; these salts are soluble
in water. (Nat 4)
158.
In the preparation of a soluble salt, it is often easier to use an insoluble metal
carbonate or metal oxide as the base. (Nat 4)
159.
In the reaction of an acid with a metal oxide the hydrogen ions and the oxide ions
form water. (Nat 5)
160.
In the reaction of an acid with a metal carbonate the hydrogen ions and carbonate
ions form water and carbon dioxide. (Nat 5)
Precipitation
161.
Precipitation is the reaction of two solutions to form an insoluble product, called a
precipitate. (Nat 5)
162.
Insoluble salts can be formed by precipitation (Nat 5)
Dissociation of water
163.
A very small proportion of water molecules will dissociate into an equal number of
hydrogen and hydroxide ions. (Nat 5)
10
164.
In water and neutral solutions, the concentration of hydrogen ions is equal to the
concentration of hydroxide ions. (Nat 5)
165.
An acidic solution contains more hydrogen ions than hydroxide ions. (Nat 5)
166.
An alkaline solution contains more hydroxide ions than hydrogen ions. (Nat 5)
Spectator ions
167.
Spectator ions are ions which do not take part in a chemical reaction. (Nat 5)
168.
Spectator ions can be identified in neutralisation and precipitation reactions and the
equations can be rewritten omitting these ions. (Nat 5)
Concentration and the mole
169.
The concentration of a solution is expressed in mol l-1. (Nat 5)
170.
The number of moles of solute, volume and concentration of a solution can be
calculated from the other two variables using the triangle below (Nat 5):
Where:
n= number of moles (moles)
C= Concentration (mol l-1)
V= Volume (l)
n
C
171.
V
Candidates should be able to carry out calculations involving mass from mol l-1 and
mass per volume for a requested concentration using the two triangles below (Nat 5):
n
C
m
V
n
GFM
Volumetric titrations
172.
The concentration of acids/alkalis can be calculated from the results of volumetric
titrations. (Nat 5)
173.
Candidates are expected to calculate the volumes of acids/alkalis required for
neutralisation from titration data. (Nat 5)
Indicators
174.
Universal indicator and pH paper can be used to determine whether solutions are acidic or
alkaline. (Nat 4)
175.
Some neutralisation reactions do not need indicators to determine their end-point (Nat 4)
176.
Natural indicators can be used to determine end-points (e.g. red cabbage). (Nat 4)
11
Unit 2: Nature's Chemistry
Key area: Homologous series
Fractional Distillation
1.
Distillation is the process of separating a mixture of liquids using the difference in boiling
points. (Nat 4)
2.
Crude oil is a mixture of chemical compounds, mainly hydrocarbons. (Nat 4)
3.
Fractional distillation is the process used to separate crude oil into fractions according to
the boiling points of the components of the fractions. (Nat 4)
4.
A fraction is a group of hydrocarbons with boiling points within a given range. (Nat 4)
5.
Ease of evaporation, viscosity, flammability and boiling point range of the fractions are
properties related to molecular sizes of the molecules within the fractions. (Nat 4)
6.
The uses of the fractions can be prepared are related to the ease of evaporation,
viscosity, flammability and boiling point range of the fractions. (Nat 4)
Hydrocarbons
7.
A hydrocarbon is a compound which contains only hydrogen and carbon. (Nat 4)
8.
A homologous series is a set of compounds with the same general formula and similar
chemical properties. (Nat 4)
Alkanes
9.
The alkanes are a subset of the set of hydrocarbons and are identified from the 'ane'
name ending. (Nat 4)
10.
The alkanes are saturated hydrocarbons. (Nat 5)
11.
Saturated hydrocarbons contain only carbon to carbon single covalent bonds.(Nat 5)
12.
Straight-chain alkanes can be named from molecular formulae, shortened and full
structural formulae (only C1 to C8). (Nat 4)
13.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the names of straight-chain alkanes (only C1 to C8). (Nat 5)
14.
The general formula for the alkanes is CnH2n+2. (Nat 5)
Alkenes
15.
The alkenes are a subset of the set of hydrocarbons and can be referred to as a
homologous series. (Nat 4)
16.
The alkenes are unsaturated hydrocarbons. (Nat 5)
12
17.
Unsaturated hydrocarbons contain at least one carbon to carbon double covalent bond.
(Nat 5)
18.
An alkene can be identified from the carbon to carbon double bond and the ‘-ene’ name
ending. (Nat 4)
19.
Straight-chain alkenes can be named, including the position of the double bond, from
shortened and full structural formulae (only C2 to C8). (Nat 4)
20.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the names of alkenes (only C2 to C8). (Nat 5)
21.
The general formula for the alkenes is CnH2n. (Nat 5)
Cycloalkanes
22.
The cycloalkanes are a subset of the set of hydrocarbons and can be referred to as a
homologous series. (Nat 5)
23.
The cycloalkanes are saturated hydrocarbons. (Nat 5)
24.
Saturated hydrocarbons contain only carbon to carbon single covalent bonds. (Nat 5)
25.
A cycloalkane can be identified from the name. (Nat 5)
26.
Cycloalkanes can be named from molecular formulae, shortened and full structural
formulae (only C3 to C8). (Nat 5)
27.
Molecular formulae can be written; and shortened and full structural formulae can be
drawn, given the names of the cycloalkanes (only C3 to C8). (Nat 5)
28.
The general formula for the cycloalkanes is Cn H2n. (Nat 5)
29.
The properties of the alkanes, alkenes and cycloalkanes should be known including the
trends in their boiling points. (Nat 5)
Systematic Naming of Hydrocarbons
30.
Branched-chain alkanes can be systematically named from shortened and full structural
formulae (up to C8). (Nat 5)
31.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the systematic names of branched-chain alkanes (only C4 to C8). (Nat 5)
32.
Branched-chain alkenes can be systematically named from shortened and full structural
formulae (up to C8). (Nat 5)
33.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the systematic names of branched-chain alkenes (only C4 to C8). (Nat 5)
34.
The properties of the branched-chain hydrocarbons should be known including the trends
in their boiling points. (Nat 5)
13
Reactions of Carbon Compounds
35.
It is possible to distinguish an unsaturated hydrocarbon from a saturated hydrocarbon
using bromine solution. (Nat 5)
36.
The reactions of an alkene with bromine, hydrogen and water are addition reactions. (Nat
5)
37.
Candidates should be able to name and draw the full structural formula of the alkane
product of the addition reaction. (Nat 5)
38.
Candidates should be able to draw full structural formulae for possible products of the
addition of bromine or water to an alkene. (Nat 5)
39.
An alkene reacts with hydrogen to form the corresponding alkane. (Nat 5)
40.
Fractional distillation of crude oil yields more long-chain hydrocarbons than are useful for
present-day industrial purposes. (Nat 4)
41.
Cracking is an industrial method for producing a mixture of smaller, more useful
molecules, some of which are unsaturated. (Nat 4)
42.
The catalyst allows the reaction to take place at a lower temperature. (Nat 4)
43.
Cracking can be carried out in the laboratory using an aluminium oxide or silicate catalyst.
(Nat 4)
Isomers
44.
Isomers are compounds with the same molecular formula but different structural formulae
and have different properties. (Nat 5)
45.
Isomers can be drawn for given molecular formulae, shortened and full structural formulae
including alkanes, branched alkanes, alkenes, branched alkenes and cycloalkanes. (Nat
5)
Key area: Everyday consumer products
Fermentation
46.
Fermentation is the breakdown of glucose to form ethanol and carbon dioxide. (Nat 4)
47.
An enzyme in yeast acts as a catalyst for the reaction. (Nat 4)
48.
Ethanol, for alcoholic drinks, can be made by fermentation of glucose derived from any
fruit or vegetable. (Nat 4)
49.
Enzymes work best under optimum conditions (Nat 4)
50.
The alcohol content of drinks is measured in units. (Nat 4)
14
51.
Alcoholic drinks, if taken in excess, can have damaging affects to health and mind. (Nat
4)
52.
There is a limit to the ethanol concentration of fermentation products because yeast
becomes inactive (the enzyme is denatured) when ethanol concentrations get too high.
(Nat 5)
53.
Distillation is a method of increasing the ethanol concentration of fermentation products in
the manufacture of ‘spirit’ drinks. (Nat 4)
Alcohols
54.
An alcohol can be identified from the hydroxyl functional group -OH and the ‘-ol’ name
ending. (Nat 5)
55.
The general formula for the alcohols is CnH2n+1OH. (Nat 5)
56.
Alcohols can be formed from the addition of water to an alkene to form the corresponding
alcohol. This is called hydration. (Nat 5)
57.
The removal of water from an alcohol produces the corresponding alkene. This is called
dehydration. (Nat 5)
58.
Straight-chain and branched-chain alcohols can be named, including the position of the
hydroxyl group, from shortened and full structural formulae (only C1 to C8). (Nat 5)
59.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the names of straight-chain alcohols (only C1 to C8). (Nat 5)
60.
The properties of the alcohols should be known including the trends in their boiling points
and solubility. (Nat 5)
61.
Alcohols can be used as solvents. (Nat 5)
Alkanoic Acids
62.
An alkanoic acid can be identified from the carboxyl functional group -COOH and the ‘-oic’
name ending. (Nat 5)
63.
Vinegar is a solution of ethanoic acid. (Nat 5)
64.
Vinegar is used in household cleaning products designed to remove limescale (a build-up
of insoluble carbonates on plumbing fixtures) and as a preservative in the food industry.
(Nat 5)
65.
Straight-chain and branched-chain alkanoic acids can be named from shortened and full
structural formulae (only C1 to C8). (Nat 5)
66.
Molecular formulae can be written and shortened and full structural formulae can be
drawn, given the name of straight-chain alkanoic acids (only C1 to C8). (Nat 5)
67.
The properties of the alkanoic acids should be known including the trends in their boiling
points and solubility. (Nat 5)
15
Esters
68.
An ester can be identified from the ester functional group and the ‘-oate’ name ending.
(Nat 5)
69.
Some uses of esters are in food flavouring, industrial solvents, fragrances and materials.
(Nat 5)
70.
The ester link is formed by the reaction of a hydroxyl functional group with a carboxyl
functional group. (Nat 5)
71.
An ester is made from an alcohol and an alkanoic acid reacting together with the loss of a
water molecule. (Nat 5)
72.
An ester can be named given the names of the parent alcohol and alkanoic acid, or from
shortened and full structural formulae. (Nat 5)
73.
Esters are formed by the condensation reaction between a carboxylic acid and an alcohol.
(Nat 5).The reaction can also be referred to as esterification.
74.
In a condensation reaction, the molecules join together by the reaction of the functional
groups to make water. (Nat 5)
75.
Shortened and full structural formulae for esters can be drawn given the names of the
parent alcohol and alkanoic acid or the names of esters. (Nat 5)
76.
Esters can be broken down to the alcohol and carboxylic acid. This reaction is called
hydrolysis.
77.
The making and breaking of esters can be referred to as a reversible reaction.
78.
The products of the breakdown of an ester can be named, or shortened and full structural
formulae can be drawn, given the name of the ester or from the shortened or full structural
formulae of the ester. (Nat 5)
Carbohydrates
79.
Carbohydrates are compounds which contain carbon, hydrogen and oxygen with the
hydrogen and oxygen in the ratio of two to one. (Nat 4)
80.
The production of carbon dioxide and water, on burning, indicates the presence of carbon
and hydrogen in a carbohydrate. (Nat 5)
81.
Plants are a source of carbohydrates and oil which can be used for food or fuel. (Nat 5)
82.
Carbohydrates supply the body with energy. (Nat 5)
83.
Photosynthesis is the process by which plants make glucose from carbon dioxide and
water using light energy in the presence of chlorophyll; oxygen is released in the process.
(Nat 4)
84.
Plants convert the glucose into starch for storing energy. (Nat 4)
85.
Respiration is the process by which the body obtains a supply of energy by breaking down
starch (using oxygen) to give carbon dioxide and water. (Nat 4)
16
86.
Carbohydrates release energy, producing carbon dioxide and water when burned. (Nat 4)
87.
Respiration and photosynthesis are important in maintaining the balance of carbon dioxide
and oxygen in the air. (Nat 4)
88.
Glucose is a simple carbohydrate.
89.
The molecular formula of glucose is C6H12O6 (Nat 4)
90.
Carbohydrates can be divided into sugars and starches.
91.
Examples of sugars include glucose, fructose, maltose and sucrose (table sugar). (Nat 4)
92.
Glucose and fructose are monosaccharides and they have the formula C6H12O6. (Nat 5)
93.
Maltose and sucrose are disaccharides and they have the formula C12H22O11. (Nat 5)
94.
Starch is a complex carbohydrate formed by joining many glucose molecules together.
95.
Starch is a polysaccharide with the formula (C6H10O5)n. (Nat 5)
96.
Starch can be distinguished from other carbohydrates by the iodine test. (Nat 4)
97.
Glucose is sweet and dissolves in water whereas starch is not sweet and does not
dissolve well in water. (Nat 4)
98.
Most sugars can be detected by the Benedict’s test; sucrose is an exception. (Nat 4)
99.
Starch is formed from glucose by a process known as polymerisation (condensation
polymerisation).
100.
Glucose is a small molecule and is known as a monomer.
101.
The large molecule (starch) is known as a polymer.
102.
Starch is broken down (hydrolysed) in the body, during digestion.
103.
Glucose, due to its small size, can pass through the gut wall into the bloodstream to be
used in cells, throughout the body, during respiration. (Nat 4)
104.
Starch is broken down (hydrolysed) by acid and by enzymes. (Nat 4)
Key area: Energy from fuels
Combustion of Fuels
105.
A fuel is a chemical which is burned to produce energy. (Nat 4)
106.
Combustion is another word for burning. (Nat 4)
107.
When a substance burns it reacts with oxygen. (Nat 4)
108.
Combustion is a reaction where energy is released and so is referred to as an exothermic
reaction. (Nat 4)
109.
The most common form of oxidation is the direct reaction of a fuel with oxygen through
combustion. (Nat 4)
110.
Energy is captured through processes such as photosynthesis and respiration. (Nat 4)
17
111.
Fossil fuels are so named because they originate from the decayed and fossilised remains
of plants and animals that lived millions of years ago. (Nat 4)
112.
Wood, petrol, coal, peat and any number of other fuels are energy rich, formed from the
energy of the sun, which is released when the fuel is burned. (Nat 4)
113.
Exothermic chemical reactions produce energy. (Nat 4)
114.
Endothermic chemical reactions take in energy.
115.
Fossil fuels are coal, oil, natural gas and peat. They are a finite resource. (Nat 4)
116.
Examples of non-fossil fuels are wood, biomass, hydrogen and ethanol.
117.
Finite sources of energy are sources of energy that are running out and cannot be
replaced. (Nat 4)
118.
Chemical fuels or the fossil fuels are useful reserves of fuels and are therefore used
extensively to satisfy the demands of an energy dependent world. (Nat 4)
119.
The fire triangle indicates that the three requirements for a fire are heat, fuel and oxygen.
If any one of these requirements is removed, the fire will go out.
120.
The chemical compounds which are found in oil and natural gas are mainly hydrocarbons.
(Nat 4)
121.
A hydrocarbon is a compound which contains only hydrogen and carbon. (Nat 4)
122.
The test for carbon dioxide is that it turns lime water milky/ cloudy (Nat 4)
123.
The test for the presence of water is blue cobalt chloride paper turns pink. (Nat 4)
124.
Hydrocarbons burn in a plentiful supply of air to produce carbon dioxide and water, this is
called complete combustion. (Nat 4)
Environmental Impact of Burning Fuels
125.
Carbon and carbon monoxide, a poisonous gas, are produced when hydrocarbons burn in
a supply of oxygen which is insufficient for complete combustion. This is called incomplete
combustion. (Nat 4)
126.
Nitrogen and oxygen from the air react inside a petrol engine to form nitrogen oxides
which are poisonous gases. (Nat 4)
127.
The burning of some fuels releases sulphur dioxide, a poisonous gas, into the
atmosphere. Sulphur dioxide is a major cause of acid rain. (Nat 4)
128.
Soot particles produced by the incomplete combustion of diesel are harmful. (Nat 4)
129.
Air pollution from the combustion of hydrocarbons can be reduced by the use of catalytic
converters which speed up the conversion of harmful pollutant gases to harmless gases.
(Nat 4)
130.
Catalytic converters reduce emissions of nitrogen oxides, carbon monoxide and unburnt
hydrocarbons. (Nat 4)
131.
Removing sulphur compounds before burning reduces air pollution. (Nat 4)
18
132.
Increasing the air to fuel ratio improves the efficiency of combustion thus decreasing
pollution. (Nat 4)
133.
Carbon dioxide in the atmosphere causes the greenhouse effect. (Nat 4)
134.
Extensive clearing of forests reduces the amount of carbon dioxide removed from the
atmosphere by photosynthesis. (Nat 4)
135.
Increased levels of carbon dioxide in the air may also be due to increased combustion of
fuels. (Nat 4)
136.
Many of our everyday actions involve the production of carbon dioxide and this is called
our carbon footprint. (Nat 4)
137.
An increase in the level of carbon dioxide in the atmosphere could cause the atmosphere
to retain more of the sun’s energy as heat, a process known as global warming. (Nat 4)
138.
Combustion of fossil fuels impacts on the environment and contributes to the carbon
cycle. (Nat 4)
Alternative Sources of Energy
139.
Carbon capture is a way to reduce carbon emissions by separating carbon dioxide from
other gases produced by the burning of fossil fuels. (Nat 4)
140.
Biofuels are used as alternative sources of energy to support society's energy needs. (Nat
4)
141.
Biomass, a source of biofuels, is a plant based material which can be burned to release
energy. (Nat 4)
142.
Biomass can also be converted to other useable forms of fuel. These include methane
gas or fuels used for transportation such as ethanol and biodiesel. (Nat 4)
143.
The benefits and risks of different energy sources and the impact on the carbon cycle can
be researched. (Nat 4)
144.
Understand the concept of conservation of mass through equations related to combustion
of hydrocarbons. (Nat 4) See balanced chemical equations from Unit 1.
Energy from Fuels
145.
Alkanes and alcohols can be used as fuels. (Nat 5)
146.
Ethanol, mixed with petrol, can be used as a fuel for cars. (Nat 5)
147.
The ethanol is obtained from sugar cane, a renewable source of energy. (Nat 5)
148.
A fuel releases energy on reaction with oxygen. (Nat 4)
149.
In an exothermic reaction, more energy is released in bond making than is required for
bond breaking. (Nat 5)
150.
When a substance is combusted the reaction can be represented using a balanced
equation. (Nat 5)
19
151.
The quantities of reactants and products in these reactions can be calculated from
balanced equations. (Nat 5)
152.
Different fuels provide different quantities of energy and this can be measured
experimentally and calculated using E = c m ∆T (Nat 5)
20
Unit 3 Chemistry in Society
(All Nat 5 level)
Key area: Metals
1.
Metals are found on the left side of the zig-zag line on the Periodic Table.
2.
Metallic bonding holds metal atoms together.
3.
Metallic bonds are the electrostatic attraction between negatively charged delocalised
electrons and the positively charged metal ion lattice.
4.
Delocalised electrons are electrons that are not ‘attached’ to a particular atom. They are
free to move.
5.
Physical properties of metals are as follows:

they are good conductors of heat and electricity

they are malleable (can be shaped)

they are ductile (can be drawn into wires)

they have high melting and boiling points

they are shiny

they are mostly solid, except mercury.
6.
Unreactive metals such as gold and silver are found uncombined in the Earth’s crust.
7.
Other metals are found in the ground in the form of metal ores. Metal ores are naturally
occurring metal compounds.
8.
Metals can be extracted from their ores by: heating, heating with carbon or carbon
monoxide, or by electrolysis.
9.
The extraction of a metal from an ore is an example of a reduction reaction.
10.
Metals can be placed in a reactivity series. The most reactive metals are placed at the
top and the least reactive are at the bottom.
11.
12.
When a metal reacts with oxygen, a metal oxide is formed.

metal + oxygen  metal oxide

e.g. copper + oxygen  copper oxide
The Alkali Metals (Group 1) and the Alkaline Earth Metals (Group 2) react with water to
form a metal hydroxide (an alkali) and hydrogen gas.
13.

metal + water  metal hydroxide + hydrogen

e.g. sodium + water  sodium hydroxide + hydrogen
All metals above copper react with acid to form a metal salt and hydrogen gas.

metal + acid  metal salt + hydrogen
21

e.g. zinc + nitric acid  zinc nitrate + hydrogen
14.
Copper does not react with dilute acid.
15.
The reactions of metals can be explained by the loss or gain of electrons in the reaction.
16.
A reaction in which electrons are gained is called REDUCTION.
17.

e.g. Mg2+(aq) + 2e-  Mg(s)

Reduction reactions will always have the electrons on the reactant side.
A reaction in which electrons are lost is called OXIDATION.

e.g. Fe(s)  Fe2+(aq) + 2e-

Oxidation reactions will always have the electrons on the right product side.
18.
Reactions in which both reduction and oxidation occur are called Redox reactions.
19.
The mnemonic: OIL RIG can be used to remember that Oxidation Is Loss and Reduction
Is Gain of electrons.
Electrochemical Cells
20.
A list of reduction reactions can be found on page 10 of the data booklet. These make up
the Electrochemical series.
21.
22.
23.
24.
In a battery, the electricity comes from a chemical reaction.

Electricity passing along metal wires is a flow of electrons

Batteries run out when the chemicals in the battery are used up.
Some batteries are rechargeable.

Nickel-Cadmium batteries can be recharged.

Lead-acid batteries in cars can be recharged.
Advantages of batteries vs advantages of mains electricity

Batteries are easy to transport and have a lower voltage, so are safer.

Mains electricity costs less than batteries.
The voltage produced in a cell depends on:

The bigger the difference between the metals on the electrochemical series, the
bigger the voltage produced.

Electrons always flow from the most reactive metal to the least reactive metal.

Therefore electrons flow from where oxidation occurs to where reduction takes
place.
25.
The purpose of the electrolyte is to complete the circuit.

An electrolyte is an ionic solution usually stored on an ion bridge.

Ion bridges are usually pieces of filter paper soaked in an ionic solution.

Positive ions move towards negative charges due to electrostatic attractions.

Negative ions move towards positive charges due to electrostatic attractions.
22
26.
Fuel cells use electrochemistry to generate electricity.
27.
Most fuel cells use hydrogen to generate electricity. The hydrogen reacts with water to
form water and oxygen.

The main advantage of using hydrogen fuel cells is that the products are nonpolluting.

The main disadvantage of using hydrogen fuels cells is that hydrogen gas is
difficult to store and is explosive.
28.
Displacement reactions are when a metal higher up the electrochemical series pushes
another metal lower down the electrochemical series, out of a solution of its own ions.
Key area: Properties of plastics
29.
Most plastics and synthetic fibres are made from crude oil.
30.
Synthetic means that the fibre has been made by scientists and is not naturally occurring.
31.
Both natural and synthetic fibres are examples of polymers.
32.
There are advantages and disadvantages of using natural or synthetic materials.
33.
Some plastics release toxic gases when they are burned e.g. carbon monoxide (CO).
34.
A biodegradable plastic is one that can be broken down by organisms such as bacteria.
35.
There are 2 types of plastic: thermoplastic and thermosetting.

Thermoplastics lose their shape when heated.

Thermosetting plastics keep their shapes when heated as the polymer chains are
cross-linked.
36.
There are 2 types of polymerisation:

Addition polymerisation

In addition polymerisation small unsaturated monomers containing a C=C double
bond undergo addition reactions to form a saturated polymer.

The polymer name starts with poly and then has the name of the monomer in
brackets, for example ethene monomers polymerise to make poly(ethene).

Condensation polymerisation

In condensation polymerisation monomers with functional groups at each end of
the molecule undergo condensation reactions to form a polymer and water.
37.
Polyesters are chains of polymers that contain an ester functional group. They are made
from a diacid monomer and a diol monomer.
38.
Kevlar is made by condensation polymerisation and is used in bullet proof vests. It is a
very strong fibre but is lighter than any other material with the same strength.
23
39.
Poly(ethanol) is a soluble plastic that can be used to make laundry bags and surgical
stiches.
40.
Poly(acrylate) is a hydrogel that has special water absorbing properties which allows
hydrogels to be used in a variety of applications such as nappies, contact lenses and as
medical bandages.
41.
Colour changing plastics can be used in food packaging to let consumers know the
condition of food inside.
42.
Conductive plastics are currently being researched which allow the development of flexible
touchscreens and e-paper in the near future.
Key area: Fertilisers
43.
The increasing population of Earth has led to a need for more efficient food production to
grow enough food to feed the increasing number of people on Earth.
44.
Growing plants require nutrients including compounds of: nitrogen (N), phosphorus (P) and
potassium (K).
45.
Different types of crops need fertilisers containing different proportions of N, P and K.
46.
Decomposition of protein in plants and animal remains recycles nitrogen in the nitrogen
cycle.
47.
Fertilisers are substances that restore the essential element for plant growth.
48.
Fertilisers need to be soluble to be absorbed through plant roots e.g. soluble compounds
of ammonium salts, potassium salts, nitrates and phosphates.
49.
Overuse of fertilisers can result in unused fertiliser being washed into rivers and lochs
causing damage to wildlife.
50.
Nitrifying bacteria in plant rood nodules can convert (fix) nitrogen from the air into
compounds containing nitrogen.
51.

Plants with such root nodules include clover, peas and beans.

The nitrogen compounds formed are nitrates (NO3).

These bacterial methods for fixing nitrogen are cheaper than chemical methods.
Synthetic fertilisers can be made from nitrogen compounds such as ammonia (NH3) and
nitric acid (HNO3).
52.
Ammonia (NH3) is made by the Haber Process.

nitrogen + hydrogen ⇌ ammonia

N2(g) + 3H2(g) ⇌ 2NH3(g)
24
53.
The Haber Process is carried out at moderate temperature as high temperature leads to
the breakdown of ammonia intro nitrogen and hydrogen.
54.
Not all the reactants turn into ammonia as eventually the ammonia breaks down as quickly
as it is formed.
55.
The catalyst used in the Haber Process is iron.
56.
This reaction is reversible.
57.
Ammonia can be converted into ammonium compounds by reacting ammonia with a
strong alkali.
58.
Ammonia has the following properties:

colourless gas

pungent smell

soluble in water

dissolves to form an alkali.
59.
Ammonia has the chemical formula NH3.
60.
Ammonium has the chemical formula NH4.
61.
Nitric acid is made by the Ostwald Process.
62.
The Ostwald Process involves the catalytic oxidation of ammonia to form nitric acid:

Stage 1: ammonia + oxygen  nitrogen monoxide + water

Stage 2: nitrogen monoxide + oxygen  nitrogen dioxide

Stage 3: nitrogen dioxide + oxygen + water  nitric Acid
63.
The Ostwald Process is carried out at moderate temperature (900°C).
64.
The reaction is exothermic so once started, the reaction does not require further heating.
65.
A platinum catalyst is used in this process.
66.
When nitrogen dioxide is dissolved in water, nitric acid is formed.
67.
The percentage mass of elements in fertilisers can be calculated.

e.g. calculate the percentage mass of nitrogen in ammonium nitrate, NH4NO3.

Find the GFM of NH4NO3
2 X N = 2 X 14
= 28
4XH=4X1
=4
3 X O = 3 X 16
= 48
Total

= 80 g
Find the mass of N in the formula
2 X N = 2 X 14
= 28 g
25

Divide: mass in formula by GFM
% Mass =
28
80
x 100%
% Mass = 35%
Key area: Chemical analysis
68.
There are two types of chemical analysis:

Qualitative analysis

Quantitative analysis.
69.
Qualitative analysis allows the presence of a substance to be detected.
70.
Quantitative analysis allows the presence of a substance to be detected and allows us to
work out how much of the substance there is.
71.
When metal compounds are placed in a flame, characteristic colours are produced.
72.
Different metals give different colours when burned; therefore the presence of a metal in a
compound can be detected using flame colour.
73.
Flame colours can be found in the data booklet on page 6.
74.
Metal ions can also be detected using precipitation reactions.
75.
The colour of the precipitate formed (insoluble solid) allows us to determine which metal
ion was present.
76.
Non-metal ions can also be detected using precipitation.
77.
Chromatography can be used to separate mixtures of substances.
78.
Chromatography involves spotting small quantities of a substance on a piece of
chromatography paper, then placing this chromatography paper vertically in a solvent. The
solvent flows up the paper and separates the single spot for the substance into a spot for
each component of the mixture.
79.
Advanced forms of chromatography are available that allow better separation of mixtures.
High Performance Liquid Chromatography (HPLC) and Gas Phase Chromatography
(GPC) and Liquid Chromatography Mass Spectrometry (LCMS) are all example of such
techniques.
80.
A titration can be used to determine the concentration of acid or base used in a
neutralisation reaction.

In a titration a pipette is used to transfer a known volume of acid or base into a
conical flask. An indicator is then added to the conical flask. The indicator allows
the end point of the titration to be easily observed.
26

A burette is filled with acid or base of a known concentration.

The burette is then used to accurately add known volumes of acid or base into the
conical flask. When a colour change is observed, the reaction has reached its end
point.

The average titre can be worked out using concordant values.

Concordant values are values that are ±0.2 cm3 of each other.

Using the values obtained from the titration experiment, the formula below can be
used to calculate an unknown for the acid or the alkali:

nCV (acid) = nCV (alkali)

Where:

V (acid) = Volume of acid

V (alkali) = Volume of alkali

C (acid) = Concentration of acid

C (alkali) = Concentration of alkali

n (acid) = Number of moles of hydrogen ions (H+) in acid.

n (alkali) = Number of moles of hydroxide ions (OH-) in the alkali.
Key area: Nuclear chemistry
81.
There are many unstable isotopes of elements.
82.
Unstable isotopes can become more stable by emitting radiation.
83.
Isotopes which emit radiation are known as radioactive isotopes (radioisotopes.)
84.
Radioisotopes have many industrial and medicinal uses e.g. Co-60 is an artificially
produced radioisotope that is used to destroy cancer cells.
85.
The three types of radiation emitted from nuclei are alpha (α), beta (β) and gamma (γ).
86.
Alpha and beta radiations change an isotope of one element to an isotope of another
element.
87.
Gamma emissions are due to nuclei losing energy.
88.
Alpha particles are helium nuclei 42𝐻𝑒 i.e. they are heavy, positively charged particles.
89.
Alpha particles are slow moving and have low penetration – they will only travel a few
centimetres through air.
90.
Beta particles are electrons
0
−1𝑒
and are fast moving i.e. they are negatively charged.
27
91.
A beta particle is emitted when a neutron changes to a proton in the nucleus. Beta
particles are more penetrating than alpha particles but are stopped by a thin sheet of
aluminium.
92.
Nuclear equations are used to describe the transitions which produce radiations.
93.
The mass numbers and atomic numbers of isotopes are shown in nuclear equations.
94.
When an isotope emits an alpha particle, the atomic number of the new isotope decreases
by 2 and the mass number decreases by 4.
95.
When an isotope emits a beta particle, the atomic number of the new isotope increases by
1 and the mass number stays the same.
96.
When an isotope emits a gamma particle, the atomic number and mass number of the new
isotope stay the same.
97.
The time in which half of the nuclei of a radioisotope decay is known as the half-life.
98.
Half-lives for radioisotopes are unique and constant.
99.
The age of materials can be dated using the half-lives of radioisotopes.
100.
Nuclear fission is the splitting of nuclei by bombarding them with slow-moving neutrons.
101.
Nuclear fusion is when light nuclei combine to form heavier nuclei and produce a lot of
energy in the process.
28