2.1.

1 Group 1 (Alkali Metals)

How Alkali Metals React with Water

 The group 1 metals are known as the alkali metals

o They form alkaline solutions when they react with water
 The group 1 metals are lithium, sodium, potassium, rubidium, caesium and francium
and they are found in the first column of the periodic table
 The alkali metals share similar characteristic chemical properties because they each
have one electron in their outermost shell
 Some of these properties are:
o They are all soft metals which can easily be cut with a knife
o They have relatively low densities and low melting points
o They are very reactive (they only need to lose one electron to become highly
stable)

The alkali metals lie on the far left of the periodic table, in the very first group

Reaction with water

 The reaction of the group 1 metals with water provides evidence for categorising
these elements into the same chemical family
 The general pattern shown is:

group 1 metal + water ⟶ metal hydroxide + hydrogen

2M (s) + 2H O (l) ⟶ 2MOH (aq) + H2 (g)

2

where M is Li, Na, K, Rb or Cs

 The hydroxides formed all have the same general formula and are colourless, aqueous
solutions
 The metals are so named because they form alkalis in water

Exam Tip
Remember the group 1 metals all produce alkaline solutions (>pH 7) when they react with
water.Lithium will produce a solution of lithium hydroxide; sodium will produce a solution
of sodium hydroxide and so on.Make sure you can give the reaction equations with the
correct state symbols to show what is happening during the reactions!

Trends in Group 1

 The differences between the reactions of the group 1 metals with water and oxygen
provide evidence of trends within the group

Reactions with Water

 The reactions of the alkali metals with water get more vigorous as you descend the
group

Summary of the Reactions of the First Three Alkali Metals with Water

Reactions with Oxygen

 The alkali metals react with oxygen in the air forming metal oxides, which is why the
alkali metals tarnish when exposed to the air
 The metal oxide produced is a dull coating which covers the surface of the metal
 The metal tarnish more rapidly as you go down the group

Summary of the Reactions of the First Three Alkali Metals with Oxygen
Physical Trends

 Apart from the chemical trends there are also patterns to be seen in the physical
properties
 The alkali metals are soft and easy to cut, getting softer as you move down the group
 The first three alkali metals are less dense than water
 They all have relatively low melting points which decrease as you move down the
group, due to decreasing attractive forces between outer electrons and positive ions

The melting point of the group 1 metals decreases as you descend the group

Exam Tip
Trends are patterns of behaviour that change as you go down a group or across a period.
Trends are not the same as rules, so sometimes there are odd properties that seem
inconsistent, but the overall patterns remain the same.

Predicting Properties in Group 1

 Following these trends, we can say that:

 o Rubidium, caesium and francium will react even more vigorously with air and
water than the first three alkali metals
 Of the alkali metals, lithium is the least reactive (as it is at the top of group 1) and
francium would be the most reactive (as it’s at the bottom of group 1)
 Using the information given in the trends we would predict that rubidium:
o would be a soft grey solid
o appears shiny when freshly cut
o is more dense than potassium (> 0.86 g cm-3)
o has a lower melting point than potassium (< 63.5 oC)

Exam Tip
You could be asked to make predictions about how rubidium would be expected to react with
water, knowing that it lies below potassium in group 1. Words like 'explosively' and
'violently' would be good ones to choose when describing the reaction.

2.1.2 Group 1: Reactivity & Electronic

Configurations

Group 1: Reactivity & Electronic Configurations

 The reactivity of the group 1 metals increases as you go down the group
 When a group 1 element reacts its atoms only need to lose electron, as there is only 1
electron in the outer shell
o When this happens, 1+ ions are formed
 The next shell down automatically becomes the outermost shell and since it
is already full, a group 1 ion obtains noble gas configuration
 As you go down group 1, the number of shells of electrons increases by 1
o This means that the outermost electron gets further away from the nucleus, so
there are weaker forces of attraction between the outermost electron and the
nucleus
o Less energy is required to overcome the force of attraction as it gets weaker,
so the outer electron is lost more easily
o So, the alkali metals get more reactive as you descend the group
These electron shell diagrams of the first 3 alkali metals show that the group 1 metals have
1 electron in their outer shell

Exam Tip
In your exams, you could be asked to explain the trend in reactivity of the alkali metals -
make sure you answer this question using their electronic configuration to support your
answer.

2.2.1 Group 7 (Halogens)

Physical Properties

 The elements in group 7 are known as the halogens

o These are fluorine, chlorine, bromine, iodine and astatine
 These elements are non-metals that are poisonous
 All halogens have similar reactions as they each have seven electrons in their
outermost shell
 Halogens are diatomic, meaning they form molecules made of pairs of atoms
sharing electrons (forming a single covalent bond between the two halogen atoms)

Trends in Physical Properties

 At room temperature, the halogens exist in different states and colours, with
different characteristics

The Appearance, Characteristics and Colour in Solution of the Halogens

 The melting and boiling points of the halogens increase as you go down the group
 This is due to increasing intermolecular forces as the atoms become larger, so
more energy is required to overcome these forces

This graph shows the melting and boiling points of the group 7 halogens

 At room temperature (20 °C), the physical state of the halogens changes as you go
down the group
o Fluorine and chlorine are gases, bromine is a liquid and iodine is
crumbly solid
 The colours of the halogens also change as you descend the group - they become
darker
 The physical states and colours of chlorine, bromine and iodine at room
temperature

Exam Tip
Exam questions on this topic occur often so make sure you know and can explain the
trends of the group 7 elements in detail, using their electron configurations.

Predicting Properties in Group 7

 Chlorine, bromine and iodine react with metals and non-metals to form compounds

Metal Halides

 The halogens react with some metals to form ionic compounds which
are metal halide salts
 The halide ion carries a -1 charge so the ionic compound formed will have different
numbers of halogen atoms, depending on the valency of the metal
 E.g., sodium is a group 1 metal:
o 2 Na + Cl2 → 2 NaCl
 Calcium is a group 2 metal:
o Ca + Br2 → CaBr2
 The halogens decrease in reactivity moving down the group, but they still form
halide salts with some metals including iron
 The rate of reaction is slower for halogens which are further down the group such
as bromine and iodine
Sodium donates its single outer electron to a chlorine atom and an ionic bond is
formed between the positive sodium ion and the negative chloride ion

Non-metal Halides

 The halogens react with non-metals to form simple molecular covalent structures
 For example, the halogens react with hydrogen to form hydrogen halides (e.g.,
hydrogen chloride)
 Reactivity decreases down the group, so iodine reacts less vigorously with
hydrogen than chlorine (which requires light or a high temperature to react with
hydrogen)
  Fluorine is the most reactive (reacting with hydrogen at low temperatures in the
absence of light)

Displacement Reactions

 A halogen displacement reaction occurs when a more reactive halogen displaces

a less reactive halogen from an aqueous solution of its halide
 The reactivity of group 7 elements decreases as you move down the group
 You only need to learn the displacement reactions with chlorine, bromine and iodine
o Chlorine is the most reactive and iodine is the least reactive

Chlorine with Bromides & Iodides

 If you add chlorine solution to colourless potassium bromide or potassium iodide

solution a displacement reaction occurs:
o The solution becomes orange as bromine is formed or
o The solution becomes brown as iodine is formed
 Chlorine is above bromine and iodine in group 7 so it is more reactive
 Chlorine will displace bromine or iodine from an aqueous solution of the metal
halide:

Cl2 + 2KBr → 2KCl + Br2

chlorine + potassium bromide → potassium chloride + bromine

Cl2 + 2KI → 2KCl + I2

chlorine + potassium iodide → potassium chloride + iodine

Bromine with Iodides

 Bromine is above iodine in group 7 so it is more reactive

 Bromine will displace iodine from an aqueous solution of the metal iodide

bromine + potassium iodide → potassium bromide + iodine

Br2 + 2KI → 2KBr + I2

 This table shows a summary of the displacement reactions of the halogens:

chlorine, bromine and iodine
Exam Tip
Displacement reactions are sometimes known as single replacement reactions.

2.2.2 Group 7: Reactivity & Electronic

Configurations
Group 7: Reactivity & Electronic Configurations

 When halogen atoms gain an electron during reactions, they form -1 ions called halide
ions
 We can use electronic configuration to explain the trends in chemical reactivity down
group 7

The atoms of the elements of group 7 all have 7 electrons in their outer shell

 Reactivity of group 7 non-metals decreases as you go down the group

 o As you go down group 7, the number of shells of electrons increases, the
same as with all other groups
 However, halogen atoms form negative ions when they gain an electron to obtain a
full outer shell
o This means that the increased distance from the outer shell to the nucleus as
you go down a group makes the halogens become less reactive
 Fluorine is the smallest halogen, which means its outermost shell is the closest to the
positive nucleus of all the halogen
o Therefore, the ability to attract an electron is strongest in fluorine making it
the most reactive
o As you move down the group, the forces of attraction between the nucleus
and the outermost shell decreases
o This makes it harder for the atoms to gain electrons as you descend the group
o Therefore, the halogens are less reactive the further down the group you go

Exam Tip
Exam questions on this topic occur often so make sure you know and can explain the
reactivity trends of the group 7 elements in detail, using their electron configurations.

2.3.1 Composition of Air

Composition of Air

 The proportion of gases in the air has not changed much in 200 million years
 About four fifths of the air is nitrogen and one fifth is oxygen
 The remaining gases include carbon dioxide, water vapour and trace quantities of the
noble gases
 Pie chart showing the composition of clean air

 Scientists know the historic composition of the air by analysing the tiny air bubbles
trapped in ice cores taken at the poles
 The air bubbles were trapped as the snow and ice was laid down tens of thousands of
years ago and provide a snapshot of what our atmosphere was like back then

Exam Tip
Although the proportion of carbon dioxide is very small, it plays a substantial role in global
warming as a greenhouse gas.

Finding the Percentage of Oxygen

 The percentage of oxygen in air can be found by reacting a metal or non-metal with
the oxygen in a fixed volume of air
 One way to carry this out is to burn a small amount of phosphorus in a bell jar that is
sitting in a trough of water
 Initially the water levels are the same inside and outside the jar
 The percentage of oxygen in air can be determined by burning phosphorus in air and
measuring the volume change

 As the phosphorus burns it uses up the oxygen inside the bell jar and the water level
rises
 By making careful measurements of water levels before and after the experiment you
can determine the percentage of oxygen in the air
 Phosphorus is very suitable for this experiment as it burns readily until all the
available oxygen is used up
 A disadvantage of this experiment is that phosphorus is toxic, so it is hazardous and
great care must be taken to handle it safely

2.3.2 Practical: Determine the % of Oxygen

in Air
Practical: Determine the Percentage of Oxygen in Air
Aim:

To determine the percentage of oxygen in air using the oxidation of iron

Diagram:
 Apparatus to determine the percentage of oxygen in the air

Method:

1. Firstly, you will need to measure the volume between the final mark on the scale and
the tap (stopcock)
2. Fill the burette with water up to lowest mark, 50.0 mL, and then let it drain into a
small measuring cylinder
3. Measure the volume of water
4. Add a little water to moisten the inside of the burette
5. Make sure the tap is closed and sprinkle some iron filings or push a piece of iron wool
into the bottom of the burette
6. Invert the burette into a trough of water and clamp the burette vertically
7. Note and record the position of the water level
8. After 3-4 days note the new position of the water level

Results:

Volume occupied between 50 mL & the tap = 3.8 mL

Initial water level = 2.6 mL

Final water level = 12.7 mL

Data Processing:

Initial volume of air = (50.0 + 3.8) - 2.6 = 51.2 mL

Final volume of air = 53.8 - 12.7 = 41.1 mL

Volume of oxygen = 51.2 - 41.1 = 10.1 mL

Percentage of oxygen = (10.1 ÷ 51.2) x 100 = 19.7%

Conclusion:
The oxygen takes up approximately 20% of the air

2.3.3 Combustion
Combustion

 Combustion is the scientific word for burning

 All combustion reactions involve a chemical change in which oxygen reacts with
elements or compounds to produce oxides
 Combustion reactions give out heat, so they will always be exothermic reactions
 You need to be able to describe the combustion reactions of magnesium, hydrogen
and sulfur:

The Combustion Reactions of Mg, H2 & S

Exam Tip
Combustion reactions can also be classified as oxidation reactions.

2.3.4 Carbon Dioxide from Thermal

Decomposition
Carbon Dioxide from Thermal Decomposition

 Thermal decomposition is the term used to describe reactions where a substance

breaks down due to the action of heat
 One such reaction is the thermal decomposition of metal carbonates
 Carbonates of metals from the lower half of the reactivity series tend to decompose on
heating to produce the metal oxide and carbon dioxide gas:

metal carbonate → metal oxide + carbon dioxide

 The thermal decomposition of copper(II)carbonate

 The thermal decomposition of copper(II)carbonate occurs readily on heating

 Copper(II) carbonate is a green powder and slowly darkens as black copper(II)
oxide is produced
 The carbon dioxide given off can be tested by passing the gas through limewater and
looking for it to turn milky
 The equation for the reaction is

CuCO3 (s) → CuO (s)+ CO2 (g)

copper(II) carbonate → copper(II) oxide + carbon dioxide

Exam Tip
The release of carbon dioxide from calcium carbonate in the production of cement is a
contributing source of rising atmospheric CO2 levels that contributes to the enhanced
greenhouse effect.

2.3.5 The Greenhouse Effect

The Greenhouse Effect

 When shortwave radiation from the sun strikes the Earth’s surface it is absorbed
and re-emitted from the surface of the Earth as infrared radiation
 Much of the radiation, however, is trapped inside the Earth’s atmosphere
by greenhouse gases which can absorb and store the energy
 Carbon dioxide, methane and water vapour are gases that have this effect
 Increasing levels of carbon dioxide, although present in only a small amount, is
causing significant upset to the Earth’s natural conditions by trapping extra heat
energy
 This process is called the enhanced greenhouse effect
 Greenhouse gases trap some of the Sun's radiation causing the Earth to warm up

Carbon dioxide

 Sources: Combustion of wood and fossil fuels, respiration of plants and animals,
thermal decomposition of carbonate rocks and the effect of acids on carbonates

Exam Tip
It is important to understand the difference between the greenhouse effect and the enhanced
greenhouse effect. The greenhouse effect ensures the mean global temperature is around 15oC
and without greenhouse gases the surface of the Earth would swing between extreme heat and
extreme cold. The enhanced greenhouse effect, due an increase in greenhouse gas
concentrations, most scientists believe, is leading to global warming.

2.4.1 Metals Reacting with Water & Acids

Metals Reacting with Water & Acids

 The chemistry of the metals is studied by analysing their reactions with water and
acids
 Based on these reactions a reactivity series of metals can be produced
 The series can be used to place a group of metals in order of reactivity based on the
observations of their reactions with water and acids
Reaction with water

 The reactions of potassium and sodium have already been seen previously in the
alkali metals, but the reaction with calcium and water is given here for reference:

Ca (s) + 2H2O (l) ⟶ Ca(OH)2 (aq) + H2(g)

calcium + water ⟶ calcium hydroxide + hydrogen

 The reactions with magnesium, iron and zinc and cold water are very slow

Reaction with dilute sulfuric or hydrochloric acids

 Only metals above hydrogen in the reactivity series will react with dilute acids
 The more reactive the metal then the more vigorous the reaction will be
 Metals that are placed high on the reactivity series such as potassium and sodium are
very dangerous and react explosively with acids
 When acids react with metals they form a salt and hydrogen gas:
 The general equation is:

metal + acid ⟶ salt + hydrogen

 Some examples of metal-acid reactions and their equations are given below:

Acid-Metal Reactions Table

2.4.2 Metal Displacement Reactions
Metal Displacement Reactions

 The reactivity of metals decreases going down the reactivity series.

 This means that a more reactive metal will displace a less reactive metal from its
compounds
 Two examples are:
o Reacting a metal with a metal oxide (by heating)
o Reacting a metal with an aqueous solution of a metal compound
 For example it is possible to reduce copper(II) oxide by heating it with zinc.
 The reducing agent in the reaction is zinc:

Zn + CuO → ZnO + Cu

zinc + copper(II) oxide → zinc oxide + copper

Metal Oxide Displacement Table

Displacement reactions between metals & aqueous solutions of metal
salts

 The reactivity between two metals can be compared using displacement

reactions in salt solutions of one of the metals
 This is easily seen as the more reactive metal slowly disappears from the
solution, displacing the less reactive metal
 For example, magnesium is a reactive metal and can displace copper from
copper(II)sulfate solution:

Mg + CuSO4→ MgSO4 + Cu

 The blue colour of the CuSO4 solution fades as colourless magnesium sulfate
solution is formed
 Copper coats the surface of the magnesium and also forms solid metal which falls
to the bottom of the beaker

Diagram showing the colour change when magnesium displaces copper from
copper(II) sulfate

Other displacement reactions

Metal Solutions Displacement Table

Exam Tip
Displacement reactions occur when the solid metal is more reactive than the metal that
is in the compound.

2.4.3 Order of Reactivity

Order of Reactivity
Carbon and the reactivity series mnemonic

 Carbon is an important element and has its own place on the reactivity series
 Its use in the extraction of metals from their oxides and a more complete reactivity
series with an accompanying mnemonic to help you memorise it is below

The reactivity series mnemonic

 “Please send lions, cats, monkeys and cute zebras into hot countries signed
Gordon”

Reactivity Series Mnemonic Table

2.4.4 Rusting of Iron
Rusting of Iron

 Rusting is a chemical reaction between iron, water and oxygen to form hydrated
iron(III)oxide
 Oxygen and water must be present for rusting to occur
 Rusting is a redox process and it occurs faster in salty water since the presence of
sodium chloride speeds up the reaction

Iron + Water + Oxygen → Hydrated Iron(III) Oxide

4Fe (s) + 3O2 (g) + xH2O (l) → 2Fe2O3.xH2O (s)

Investigating Rusting
 Diagram showing the
requirements of oxygen and water for rust to occur: only the nail on the left rusts

Method:

 Set up the apparatus as shown in the diagram

 The water in the second test tube is boiled to remove any dissolved oxygen
 The oil provides a barrier to prevent oxygen diffusing into the boiled water
 Calcium chloride is a drying agent in the third test tube
 Leave the apparatus for a few days to give it time to react

Results:

 The nail on the left rusts as it is in contact with both air (which contains oxygen) and
water
 The nail in the middle does not rust as it is not in contact with air
 The nail on the right does not rust as it is not in contact with water (calcium chloride
absorbs any water molecules present due to moisture)
 The results show that both air and water must be present for rusting to occur

Damage to Iron Structures

 Rust is a soft solid substance that flakes off the surface of iron easily, exposing
fresh iron below which then undergoes rusting
 This means that over time all of the iron rusts and its structure becomes weakened

Preventing Iron Rusting

Barrier Methods

 Rust can be prevented by coating iron with barriers that prevent the iron from
coming into contact with water and oxygen
 However, if the coatings are washed away or scratched, the iron is once again
exposed to water and oxygen and will rust
 Galvanising / Sacrificial protection

 Iron can be prevented from rusting making use of metals higher in reactivity than
iron
 Galvanising is a process where the iron to be protected is coated with a layer of
zinc
 ZnCO3 is formed when zinc reacts with oxygen and carbon dioxide in the air and
protects the iron by the barrier method
 If the coating is damaged or scratched, the iron is still protected from rusting
because zinc preferentially corrodes as it is higher up the reactivity series than iron
 Compared to iron it loses its electrons more readily:

Zn → Zn2+ + 2e–

 The iron stays protected as it accepts the electrons released by zinc, remaining in
the reduced state and thus it does not undergo oxidation
 The electrons donated by the zinc react with hydrogen ions in the water producing
hydrogen gas:

2H+ + 2e– → H2

 Zinc therefore reacts with oxygen and water and corrodes instead of the iron

Sacrificial Corrosion

 Sacrificial corrosion occurs when a more reactive metal is intentionally allowed to

corrode
 An example of this occurs with ships’ hulls which sometimes have large blocks of
magnesium or magnesium alloys attached
 The blocks slowly corrode and provide protection to the hull in the same way the
zinc does by pushing electrons onto the iron which prevents it from being reduced
to iron(III) ions

Exam Tip
Corrosion and rusting are not the same process. Corrosion is the general term used to
describe the degradation of metal surfaces whereas rusting is the specific type of
corrosion that happens to iron.
2.4.5 Oxidation & Reduction
Oxidation & Reduction
Oxidation & reduction in terms of oxygen

 The reactions of metals with oxygen, such as in iron rusting can be classified
as oxidation
 Oxidation is any reaction in which a substance gains oxygen
 The opposite of oxidation is reduction
 Reduction is a reaction in which a substance loses oxygen
 For example, the displacement reaction between zinc and copper(II)oxide can be
classified as a redox reaction

Zn + CuO → ZnO + Cu

zinc + copper(II) oxide → zinc oxide + copper

 Oxidation cannot occur without reduction happening simultaneously, hence these are
called redox reactions
 The copper(II)oxide supplies the oxygen, so it is the oxidising agent
 The zinc is the reducing agent because it removes the oxygen

Oxidation & Reduction in terms of electrons

 Displacement reactions can be analysed in terms redox reactions by studying

the transfer of electrons
 For the example of magnesium and copper sulfate, a balanced equation can be written
in terms of the ions involved:

Mg (s) + Cu2+ (aq) + SO42- (aq) → Mg2+ (aq) + SO42- (aq) + Cu (s)

 The sulfate ions, SO42-, appear on both sides of the equation unchanged hence they
are spectator ions and do not participate in the chemistry of the reaction so can be
omitted:

Mg (s) + Cu2+ (aq) → Mg2+ (aq) + Cu (s)

 This equation is an example of a balanced ionic equation which can be further split
into two half equations illustrating oxidation and reduction individually:

Mg → Mg2+ + 2e–

Cu2+ + 2e–→ Cu

 The magnesium atoms are thus oxidised as they lose electrons

 The copper ions are thus reduced as they gain electrons
  These equations illustrate the broader definition of oxidation and reduction in terms of
electron transfer:
o Oxidation is the loss of electrons
o Reduction is the gain of electrons

'OIL RIG' is a useful mnemonic to help remember the definitions of oxidation and
reduction

2.4.6 Practical: Investigate Metals Reacting

with Acids
Practical: Investigate Metals Reacting with Acids
Aim:

To investigate the reactions between dilute hydrochloric and sulfuric acids with the metals
magnesium, iron and zinc

Diagram:

Investigating the reactions of dilute acids with metals

Method:

 Wear some safety glasses before handling acids

 Using a small measuring cylinder, add 5 cm3 of dilute hydrochloric acid to each of
three test tubes
 Add about 1 cm length of magnesium ribbon to the first tube, observe and note down
what you see
 Use a lighted splint to test for any gases given off
 To the second test tube add a few pieces of iron filings and to the third some zinc
turnings
 Observe what happens, test for any gases and note down your observations
 Repeat the experiment with dilute sulfuric acid

Results:

Metals with Acids Observations Table

Equations for the Reactions

Conclusions:

 The metals can be ranked in reactivity order Mg > Zn > Fe

 The three metals react in the same with both acids
 Hydrogen and a metal salt solution is produced
2.5.1 Sources of Metals
Sources of Metals

 The Earth’s crust contains metals and metal compounds such as gold, copper, iron
oxide and aluminium oxide
 Useful metals are often chemically combined with other substances forming ores
 A metal ore is a rock that contains enough of the metal to make it worthwhile
extracting
 They have to be extracted from their ores through processes such as electrolysis,
using a blast furnace or by reacting with more reactive material
 In many cases the ore is an oxide of the metal, therefore the extraction of these metals
is a reduction process since oxygen is being removed
 Common examples of oxide ores are iron and aluminium ores which are
called haematite and bauxite respectively
 Unreactive metals do not have to be extracted chemically as they are often found as
the uncombined element
 This occurs as they do not easily react with other substances due to
their chemical stability
 Examples include gold and platinum which can both be mined directly from the
Earth’s crust

Exam Tip
A metal can reduce another metal (remove oxygen) only if it is more reactive than the metal
that is bonded to the oxygen.

2.5.2 Extracting Metals

Extracting Metals
Extraction of metals and the reactivity series

 The most reactive metals are at the top of the series

 The tendency to become oxidised is thus linked to how reactive a metal is and
therefore its position on the reactivity series
 Metals higher up are therefore less resistant to oxidation than the metals placed lower
down which are more resistant to oxidation
 The position of the metal on the reactivity series determines the method of extraction
 Higher placed metals (above carbon) have to be extracted using electrolysis as they
are too reactive and cannot be reduced by carbon
 Lower placed metals can be extracted by heating with carbon which reduces them
 Metals Extraction Method Table

The extraction method depends on the position of a metal in the reactivity series

Exam Tip
Make sure you can explain why aluminium is extracted by electrolysis while iron is extracted
by reduction as it is a question that often comes up.
2.5.3 Using Metals
Using Metals

 The uses of aluminium, copper and steel are summarised in these tables:

Uses of Aluminium

Uses of Copper

Uses of Steel
2.5.4 Alloys
Alloys

 Alloys are mixtures of metals, where the metals are mixed together physically but are
not chemically combined
 They can also be made from metals mixed with nonmetals such as carbon
 Alloys often have properties that can be very different to the metals they contain, for
example they can have greater strength, hardness or resistance to corrosion or
extreme temperatures
 Alloys contain atoms of different sizes, which distorts the regular arrangements of
atoms
 This makes it more difficult for the layers to slide over each other, so they are usually
much harder than the pure metal
 Brass is a common example of an alloy which contains 70% copper and 30% zinc

Particle diagram showing a mixture of elements in an alloy. The different sizes of the two
types of atoms prevent the layers of atoms from sliding over each other, so the alloy
becomes less malleable than the pure metal

Exam Tip
Questions on this topic often give you a selection of particle diagrams and ask you to choose
the one which represents an alloy. It will be the diagram with uneven sized particles and
distorted layers or rows of particles.
2.6.1 Indicators
Two Colour Indicators

 Two colours indicators are used to distinguish between acids and alkalis
 Many plants contain substances that can act as indicators and the most common one
is litmus which is extracted from lichens
 Synthetic indicators are organic compounds that are sensitive to changes in acidity
and appear different colours in acids and alkalis
 Phenolphthalein and methyl orange are synthetic indicators frequently used in acid-
alkali titrations

Two Colour Indicators Table

 Synthetic indicators are used to show the endpoint in titrations as they have a very
sharp change of colour when an acid has been neutralised by an alkali and vice-versa
 Litmus is not suitable for titrations as the colour change is not sharp and it goes
through a purple transition colour in neutral solutions making it difficult to determine
an endpoint
 Litmus is very useful as an an indicator paper and comes in red and blue versions, for
dipping into solutions or testing gases

The pH Scale

 The pH scale goes from 0 – 14 (extremely acidic substances can have values of below
0)
 All acids have pH values of below 7, all alkalis have pH values of above 7
 The lower the pH then the more acidic the solution is
 The higher the pH then the more alkaline the solution is
 A solution of pH 7 is described as being neutral
 The pH scale showing acidity, neutrality and alkalinity

Universal Indicator

 Universal indicator is a wide range indicator and can give only an approximate value
for pH
 It is made of a mixture of different plant indicators which operate across a broad pH
range and is useful for estimating the pH of an unknown solution
 A few drops are added to the solution and the colour is matched with a colour chart
which indicates the pH which matches with specific colours
 Universal indicator colours vary slightly between manufacturer so colour charts are
usually provided for a specific indicator formulation

pH scale with the Universal Indicator colours used to determine the pH of a solution

Exam Tip
A common error is to suggest using universal indicator as a suitable indicator for an acid-base
titration. This is incorrect as a sharp colour change is required to identify the end-point,
which cannot be achieved with universal indicator.
2.6.2 Acids, Alkalis & Neutralisation

Acids & Alkalis

 When acids are added to water, they form positively charged hydrogen ions (H+)
 The presence of H+ ions is what makes a solution acidic
 When alkalis are added to water, they form negative hydroxide ions (OH–)
 The presence of the OH– ions is what makes the aqueous solution an alkali
 The pH scale is a numerical scale which is used to show how acidic or alkaline a
solution is, in other words it is a measure of the amount of the ions present in solution

Neutralisation

 A neutralisation reaction occurs when an acid reacts with an alkali

 When these substances react together in a neutralisation reaction, the H+ ions react
with the OH– ions to produce water
 For example, when hydrochloric acid is neutralised a sodium chloride and water are
produced:

 The net ionic equation of all acid-base neutralisations and is what leads to a neutral
solution, since water has a pH of 7:

H+ + OH– ⟶ H2O

 Neutralisation is very important in the treatment of soils to raise the pH as some crops
cannot tolerate pH levels below 7
 This is achieved by adding bases to the soil such as limestone and quicklime

Exam Tip
Not all reactions of acids are neutralisations. For example, when a metal reacts with an acid,
although a salt is produced there is no water formed so it does not fit the definition of
neutralisation.
2.6.3 Acid-Alkali Titrations
Acid-Alkali Titrations

 Titrations are a method of analysing the concentration of solutions

 Acid-base titrations are one of the most important kinds of titrations
 They can determine exactly how much alkali is needed to neutralise a quantity of acid
– and vice versa
 You may be asked to calculate the moles present in a given amount,
the concentration or volume required to neutralise an acid or a base
 Titrations can also be used to prepare salts

How to carry out a titration

Performing a titration
Method:

1. Use the pipette and pipette filler and place exactly 25 cm3 sodium hydroxide solution
into the conical flask
2. Place the conical flask on a white tile so the tip of the burette is inside the flask
3. Add a few drops of a suitable indicator to the solution in the conical flask
4. Perform a rough titration by taking the burette reading and running in the solution in 1
– 3 cm3 portions, while swirling the flask vigorously
5. Quickly close the tap when the end-point is reached (sharp colour change) and record
the volume, placing your eye level with the meniscus
6. Now repeat the titration with a fresh batch of sodium hydroxide
7. As the rough end-point volume is approached, add the solution from the burette one
drop at a time until the indicator just changes colour
8. Record the volume to the nearest 0.05 cm3
9. Repeat until you achieve two concordant results (two results that are within 0.1 cm3 of
each other) to increase accuracy

Results:
Record your results in a suitable table, e.g:

Exam Tip
Use a funnel to fill the burette but be sure to remove it before starting the practical as it can
drip liquid into the burette, making the initial reading false.

2.7.1 Solubility Rules

Solubility Rules

 Ionic compounds are generally soluble in water compared to covalent substances,

but there are exceptions
 A knowledge of the solubility of ionic compounds helps us to determine the most
appropriate method for the preparation of salts
 The solubility of common ionic compounds is shown below:
 Solubility of Ionic Compounds Table

 Calcium hydroxide is slightly soluble in water

Exam Tip
Calcium hydroxide solution is more commonly know as limewater and is used to test for
carbon dioxide.

2.7.2 Acids, Bases & Protons

Acids, Bases & Protons
Proton transfer

 The earlier definition of an acid and a base can be extended

 In terms of proton transfer, we can further define each substance in how they interact
with protons

Acids

 Acids are proton donors as they ionize in solution producing protons, H+ ions
 These H+ ions make the aqueous solution acidic

Bases (Alkalis)

 Bases (alkalis) are proton acceptors as they ionize in solution producing OH- ions
which can accept protons
 These OH- ions make the aqueous solution alkaline
 Dia
gram showing the role of acids and bases in the transfer of protons

2.7.3 Reactions of Acids

Reactions of Acids
Reactions of acids with metals

 Only metals above hydrogen in the reactivity series will react with dilute acids
 The more reactive the metal then the more vigorous the reaction will be
 Metals that are placed high on the reactivity series such as potassium and sodium are
very dangerous and react explosively with acids
 When acids react with metals they form a salt and hydrogen gas:
 The general equation is:

metal + acid ⟶ salt + hydrogen

 Some examples of metal-acid reactions and their equations are given below:

Acid-Metals Reactions Table

 In general, we can summarise the reaction of a metal that forms a +2 ion as follows:
 Acids-Metals Summary Table

Reaction of acids with bases

 When an acid reacts with a base, a neutralisation reaction occurs

 In all acid-base neutralisation reactions, a salt and water are produced:

acid + base ⟶ salt + water

 The identity of the salt produced depends on the acid used and the positive ions in the
base
 Hydrochloric acid produces chlorides, sulfuric acid produces sulfate salts and nitric
acid produces nitrates
 Metal oxides and metal hydroxides act as bases
 The following are some specific examples of reactions between acids and metal
oxides / hydroxides:

2HCl + CuO ⟶ CuCl2 + H2O

H2SO4 + 2NaOH ⟶ Na2SO4 + 2H2O

HNO3 + KOH ⟶ KNO3 + H2O

 In general, we can summarise the reaction of metals and bases as follows:

Acids and Metals Oxides or Hydroxides Summary Table

Reactions of Acids with Metal Carbonates

 Acids will react with metal carbonates to form the corresponding

metal salt, carbon dioxide and water
 These reactions are easily distinguishable from acid – metal oxide/hydroxide reactions
due to the presence of effervescence caused by the carbon dioxide gas

Acids & Metal Carbonates Reactions Table

 The following are some specific examples of reactions between acids and metal
carbonates:

2HCl + Na2CO3 ⟶ 2NaCl + H2O + CO2

H2SO4 + CaCO3⟶ CaSO4 + H2O + CO2

Exam Tip
If in an acid-base reaction there is effervescence produced then the base must be a metal
carbonate which produces carbon dioxide gas.
2.7.4 Bases
Bases
What makes a base act like a base?

 Bases are substances which can neutralise an acid, forming a salt and water
 The term base and alkali are not the same
 A base which is water-soluble is referred to as an alkali
o So, all alkalis are bases, but not all bases are alkalis
 Alkalis have pH values of above 7
 In basic (alkaline) conditions red litmus paper turns blue
 Bases are usually oxides, hydroxides or carbonates of metals
 The presence of the OH- ions is what makes the aqueous solution an alkali
 One unusual base is ammonia solution
o When ammonia reacts with water it produces hydroxide ions

Some Common Alkalis and the Ions They Contain

Exam Tip
Aqueous ammonia and ammonium hydroxide are the same thing. When ammonia gas
dissolves in water it forms ammonium hydroxide. Be careful to use the correct
terminology: ammonia is the gas, NH3, ammonium is the ion present in ammonium
compounds, NH4+

2.7.5 Prepare a Soluble Salt

Prepare a Soluble Salt

 A soluble salt can be made from the reaction of an acid with an insoluble base
 During the preparation of soluble salts, the insoluble reactant is added in excess to
ensure that all of the acid has reacted
  If this step is not completed, any unreacted acid would
become dangerously concentrated during evaporation and crystallisation
 The excess reactant is then removed by filtration to ensure that only the salt and
water remain
 Since all of the acid has reacted and the excess solid base has been removed then
the solution left can only be salt and water
 If a carbonate was used as the solid base instead of an oxide or hydroxide, then
any carbon dioxide gas produced would have been released into the atmosphere
 A common example is the preparation of copper(II) sulfate which can be made with
copper(II) oxide and dilute sulfuric acid:

CuO (s) + H2SO4 (aq) ⟶ CuSO4 (s) + H2O (l)

 The acid could also be reacted with a metal to produce the salt, as long as the
metal is above hydrogen in the reactivity series and not too reactive so that a
dangerous reaction does not take place

Exam Tip
Exam questions often ask why the solid oxide is added in excess. This is done to avoid
leaving any unreacted acid which would become dangerously concentrated during
evaporation and crystallisation.

2.7.6 Prepare a Soluble Salt II

Prepare a Soluble Salt II
Aim:

To prepare a sample of a dry salt starting from an acid and an alkali

Diagram:
 Diagram showing the apparatus needed to prepare a
salt by titration

Method:

 Use a pipette to measure the alkali into a conical flask and add a few drops of
indicator (phenolphthalein or methyl orange)
 Add the acid into the burette and note the starting volume
 Add the acid very slowly from the burette to the conical flask until the indicator
changes to appropriate colour
 Note and record the final volume of acid in burette and calculate the volume of acid
added (starting volume of acid - final volume of acid)
 Add this same volume of acid into the same volume of alkali without the indicator
 Heat to partially evaporate, leaving a saturated solution
 Leave to crystallise decant excess solution and allow crystals to dry

Results:
A dry sample of a salt is obtained

Exam Tip
When evaporating the solution some water is left behind to allow for water of crystallisation
in some salts and also to prevent the salt from overheating and decomposing.

2.7.7 Prepare an Insoluble Salt

Prepare an Insoluble Salt

 Insoluble salts can be prepared using a precipitation reaction

 The solid salt obtained is the precipitate, thus in order to successfully use this method
the solid salt being formed must be insoluble in water
 The preparation of a soluble salt follows this pattern:

soluble salt 1 + soluble salt 2 ⟶ insoluble salt + soluble salt 3

AB + CD ⟶ AD + CB

 The method involves measuring out a fixed volume of one solution and then adding
the second salt solution until it is in a slight excess
o This ensures the maximum amount of precipitate will be obtained
 The precipitate is recovered by filtration and then it must be washed with distilled
water remove reactants that are contaminating the residue (recovered solid)
o It is then left to dry
 This method is a good way to prepare silver and lead(II) salts which are often
insoluble; the starting material will usually be the nitrate of silver or lead(II) since all
nitrates are soluble

Exam Tip
This reaction is also known as a double decomposition reaction.

2.7.8 Practical: Prepare Copper(II)Sulfate

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Practical: Prepare Copper(II)Sulfate

Aim:

To prepare a pure, dry sample of hydrated copper(II) sulfate crystals

Materials:

 1.0 mol / dm3 dilute sulfuric acid

 Copper(II) oxide
 Spatula & glass rod
  Measuring cylinder & 100 cm3 beaker
 Bunsen burner
 Tripod, gauze & heatproof mat
 Filter funnel & paper, conical flask
 Evaporating basin and dish.

The preparation of copper(II) sulfate by the insoluble base method

Practical Tip:
 The base is added in excess to use up all of the acid, which would become dangerously
concentrated during the evaporation and crystallisation stages

Method:

1. Add 50 cm3 dilute acid into a beaker and warm gently using a Bunsen burner
2. Add the copper(II) oxide slowly to the hot dilute acid and stir until the base is in
excess (i.e. until the base stops dissolving and a suspension of the base forms in
the acid)
3. Filter the mixture into an evaporating basin to remove the excess base
4. Gently heat the solution in a water bath or with an electric heater to evaporate the
water and to make the solution saturated
5. Check the solution is saturated by dipping a cold glass rod into the solution and
seeing if crystals form on the end
6. Leave the filtrate in a warm place to dry and crystallise
7. Decant excess solution and allow the crystals to dry

Results:

Hydrated copper(II) sulfate crystals should be bright blue and regularly shaped
Exam Tip
Make sure you learn the names of all the laboratory apparatus used in the preparation of
salts.
2.7.9 Practical: Prepare Lead(II)Sulfate
Practical: Prepare Lead(II)Sulfate
Aim:

To prepare a dry sample of lead(II) sulfate

Diagram:
 The preparation of lead(II)sulfate by precipitation from two soluble salts

Method:

 Measure out 25 cm3 of 0.5 mol dm3 lead(II)nitrate solution and add it to a small beaker
 Measure out 25 cm3 of 0.5 mol dm3 of potassium sulfate add it to the beaker and mix
together using a stirring rod
 Filter to remove precipitate from mixture
 Wash filtrate with distilled water to remove traces of other solutions
 Leave in an oven to dry

Soluble salt 1 = lead(II) nitrate Soluble salt 2 = potassium sulfateEquation for the
reaction:
Pb(NO3)2 (aq) + K2SO4 (aq) → PbSO4 (s) + 2KNO3 (aq)

lead(II) nitrate + potassium sulfate → lead(II) sulfate + potassium nitrate

Exam Tip
Care should be taken with handling lead salts as they are toxic.
2.8.1 Tests for Gases
Tests for Gases

 Many reactions in the lab produce gases which then need to be tested
 The table below indicates the tests for the gases you should know:

Exam Tip
It is easy to confuse the tests for hydrogen and oxygen. Try to remember that a ligHted splint
has a H for Hydrogen, while a glOwing splint has an O for Oxygen.

2.8.2 Flame Tests

Flame Tests

 Metal ions produce a colour if heated strongly in a flame

 Ions from different metals produce different colours
 The flame test is thus used to identify metal ions by the colour of the flame they
produce
 Dip the loop of an unreactive metal wire such as nichrome or platinum in
concentrated acid, and then hold it in the blue flame of a Bunsen burner until there is
no colour change
 This cleans the wire loop and avoids contamination
o This is an important step as the test will only work if there is just one type of
ion present
o Two or more ions means the colours will mix, making identification erroneous
 Dip the loop into the solid sample and place it in the edge of the blue Bunsen flame
 Avoid letting the wire get so hot that it glows red otherwise this can be confused with
a flame colour

Diagram showing the technique for carrying out a flame test

 The colour of the flame is observed and used to identify the metal ion present
 Diagram showing the colours formed in the flame test for metal ions

Exam Tip
The sample needs to be heated strongly, so the Bunsen burner flame should be on a blue
flame.

2.8.3 Tests for Cations

Tests for Cations

 Metal cations in aqueous solution can be identified by the colour of the precipitate
they form on addition of sodium hydroxide and ammonia
 If only a small amount of NaOH is used then normally the metal
hydroxide precipitates

Analysing results

 The table below contains the results for each of the cations included in the syllabus
 If a precipitate is formed from NaOH then the hydroxide is insoluble in water

Testing for Cations Table

Exam Tip
Sometimes you may not see much of a precipitate because the cation you are testing is
present is very small amounts. However, every a slight cloudiness or colour change can
indicate a positive test result.

2.8.4 Tests for Anions

Tests for Anions
Testing for Anions Table

Exam Tip
When it comes to qualitative inorganic analysis, always remember that there will be a test for
the metal cation part of the molecule and another test for the anion part.

2.8.5 Tests for Water

Tests for Water

 Water can be identified using a chemical test and/or a physical test

Chemical test for water

 Anhydrous copper(II) sulfate turns from white to blue on the addition of water
 The equation is:
 CuSO4 (s) + 5H2O (l) → CuSO4.5H2O (s)

Copper sulfate turns a light blue colour in the presence of water

Physical test for water

 A physical test to see if a sample of water is pure is to check its boiling point
 A sample of the liquid is placed in a suitable container such as a boiling tube and
gently heated
 Using a thermometer, you can check if the boiling point is exactly 100 oC
 Any impurities present will usually tend to raise the boiling point and depress the
melting point of pure substance

Exam Tip
A lot of students are tempted to say you can identify water because it has no taste or smell.
While this may be true, it would be extremely hazardous to taste anything in the lab and
water is not the only colourless liquid to have no taste or smell!