IGCSE Chemistry
IGCSE Chemistry
From the Edexcel IGCSE 2009 Syllabus including triple science
statements
CGPwned
When Chemistry and CGP books get pwned…
The Periodic Table!
For the PDF, go to
http://chemistry.about.com/library/PeriodicTableallcol
or.pdf
Unit 1: The Periodic Table
The periodic table contains about a hundred or so elements that have been currently discovered. The rows are
known as periods and elements of the same period have the same number of electron shells. The columns are
known as groups and elements of the same group have the same number of electrons on their outer shell. Group one
has one outer electron,
An element is a substance that cannot be broken down into anything simpler. KCl for example (potassium chloride) is
NOT an element because it can be broken down into K (potassium) and Cl (chlorine). The potassium and chlorine are
the elements.
A compound is two or more elements chemically bonded together. An example would be KCl (potassium chloride),
which consists of the elements potassium and chloride chemically bonded together.
Atoms are the building blocks of substances.
Molecules are two or more atoms bonded together. It doesn’t have to be a compound. Elements such as O 2 and Br2
are diatomic molecules – they exist in pairs.
Atomic Structure
Atoms are made up of protons, neutrons and electrons.
Protons are positively charged. Electrons are negatively
charged. Neutrons don’t have a charge.
An atom consists of a nucleus, which contains protons and
neutrons; and some electron shells which surround the nucleus
and contain electrons. The neutrons however, are different.
The number of protons and the number of neutrons add up to
make the mass number of an element.
Understanding the Lack of Reactivity in Noble Gases (Group 0)
Noble gases have eight electrons on their outer shell, therefore, there is no need for them to gain or lose electrons.
Basically they have a full outer shell so they don’t need to react. This is what makes them so unreactive.
How to Read Each Square on The Periodic Table
You probably already know that the periodic table is made up of
lots and lots of squares, each containing an element and
information about it.
Anyways we already know what the atomic mass number is (the
number of protons + neutrons). It says 12.011 here but this is
probably because this picture came from some super complicated
periodic table. In IGCSE level however, the atomic mass should read 12. Anyways, the atomic number is the number
of protons (and electrons), so to find the number of neutrons, if asked to, simply subtract the atomic mass by the
atomic number.
Example: Calculate the number of neutrons Carbon has.
The answer: 12 – 6 = 6 neutrons
The Arrangement of Electrons
Atoms are surrounded by electron shells which contain electrons.
But the arrangement is the same for ALL the elements, not matter
how different they are.
Each shell can only hold a certain number of electrons. The very first shell can hold only two electrons. The second
shell can hold eight. The third sometimes appear full with eight but can expand to a total of eighteen. However, this is
beyond GCSE level, and for now, the shells only hold eight.
So how do you find the electron configuration? Well let’s use potassium (K) as an example.
Look up the atomic number of potassium. It should say 19. This tells you the number of protons, which is equal to the
number of electrons so we can use that.
Arrange the electrons in shells, always filling up the inner shell before you go to the outer one. Remember the first,
innermost shell can only take 2 electrons, the second one can take 8, and the third one, 8. You will find that you have
one electron left. That goes on the fourth shell.
Your electron configuration should look like this: 2, 8, 8, 1.
Example: Work out the electron configuration of chlorine.
Chlorine has an atomic number of 17 – so 17 electrons.
17 – 2 (as the innermost shell only holds two electrons) = 15
15 – 8 (as the second shell only holds eight electrons) = 7 (This number is the number of electrons Chlorine
has on its outer shell).
7 electrons does not fill up the third shell so we are left with the configuration: 2, 8, 7.
Isotopes
The number of neutrons in an atom can vary slightly. For example, there are three kinds of carbon atom, called
carbon-12, carbon-13 and carbon-14. They all have the same number of protons, but the number of neutrons vary.
These different atoms of carbon are called isotopes. Isotopes are atoms that have the same atomic number, but
different mass numbers. They have the same number of protons, but different numbers of neutrons. The fact that
they have varying numbers of neutrons makes no difference whatsoever to their chemical reactions. The chemical
properties are governed by the number and arrangement of the electrons.
Calculating Relative Atomic Mass (R.A.M.)
Lets start this off with an example!
Example: Naturally occurring silver is 51.84% silver-107 and 48.16% silver-109. Calculate the relative atomic mass of
silver.
r.a.m. (Ag) = (51.84/100 x 107) + (48.16/100 x 109)
= 55.469 + 52.494
=107.96
Now what did we do there? Well I simply calculated 51.84% of 107 (of silver) and 48.16% of 109 (of silver), and
added the two answers! What we end up with is 107.96. Round that up to a whole number and the average relative
atomic mass of silver is 108.
Calculating the Abundance (percentage) of an Isotope
Example: Copper consists of two isotopes, copper-63 and copper-65. Its relative atomic mass is 63.62. Find the
abundance of each isotope.
Let y/100 = abundance of copper-63
Let (100-y)/100 = abundance of copper-65
63.62 = (y/100 x 63) + [(100-y)/100 x 65]
63.62 = 63y +6500 – 65y
-2y = -135
y = 69
Abundance of copper-63 = 69%
Abundance of copper-65 = 100 – 69 = 31%
About Metals and Non-Metals
The IGCSE spec. states you have to recall the positions of metals and nonmetals on the periodic table. That’s easy! Its on page two. Have a look. Its
colour-coded.
Anyways, this section covers 2.2, 2.3 and 2.5.
Metals
Metals tend to be shiny. They tend to have high melting and boiling points
because of powerful attractions. Metals conduct heat and electricity because delocalized electrons are free to move
throughout the structure. Metals are usually easy to shape due to their regular packed molecules. Metals react with
water to form bases, and their oxides are also bases. They are good reducing agents because they lose electron.
Non-Metals
Non-metals tend to be brittle. They are poor conductors of heat and electricity. They form acidic oxides and are good
oxydising agents because they gain electrons.
Aluminium Oxide
Aluminium oxide is amphoteric. It can neutralize both an acid and a base.
Reaction with acids
Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. That
means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride
solution.
In this (and similar reactions with other acids), aluminium oxide is showing the basic side of its amphoteric nature.
Reaction with bases
Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodium
hydroxide solution. Various aluminates are formed - compounds where the aluminium is found in the negative ion.
This is possible because aluminium has the ability to form covalent bonds with oxygen.
Group 1: The Alkali Metals
Alkali metals are metals that are part of group one. They are extremely reactive metals, and reactivity
increases DOWNWARDS – in other words, lithium is the least reactive and francium.
Some Basic Physical Properties
Metal
Lithium
Sodium
Potassium
Rubidium
Francium
Melting Point (0C)
181
98
63
39
29
Boiling Point (0C)
1342
883
760
686
669
Density (g/cm3)
0.53
0.97
0.86
1.53
1.88
You can see that as reactivity increases, the melting and boiling points decreases; however, density increases. These
points are very low for metals. Remember that potassium, sodium and lithium would float on water due to their
densities. But why are they so reactive? Well they only have one electron to lose!
The metals are also very soft and easy to cut, becoming softer as you go down the group. They are shiny and silver
when cut, but tarnish within seconds on exposure to air.
Storage and Handling
All these metals are extremely reactive. Anyways the metals will quickly react with air to form oxides, and react
between rapidly and violently with water to form strongly alkaline solutions of metal hydroxides.
To stop them reacting with oxygen or water vapour in the air, lithium, sodium and potassium are stored under oil.
Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of
oxygen getting at them.
Great care must be taken not to touch any of these metals with bare fingers. There could be enough sweat on your
skin to give a reaction producing lots of heat and a very corrosive metal hydroxide.
Reactions with Water
All these metals react with water to produce a metal hydroxide and hydrogen.
Metal + Water  Metal Hydroxide + Hydrogen
All the hydroxides are bases and turn pH paper purple.
With Sodium
The sodium floats because it is less dense than water. It melts because its melting point is low and a lot of heat is
produced by the reaction. Observations would be that the sodium would turn into a ball and whiz around the surface
of the water. It may form a white trail which is sodium hydroxide. This dissolves to make a strongly alkaline solution
with the water. When lit, it produces a yellow flame.
With Lithium
The reaction is very similar to sodium’s reaction, except it is slower. The lithium does not melt due to its higher
melting point. When lit, it produces a red flame.
With Potassium
Potassium’s reaction is faster than sodium’s. Enough heat is produced to ignite the hydrogen, which burns with a lilac
flame. The reaction often ends with the potassium spitting around.
With Rubidium and Caesium – The Two Baddies
The reaction is so violent it can be explosive. When lit, Rubidium forms a red flame and Caesium forms a blue flame.
Explaining the Increase in Reactivity
The differences between reactions depend in part on how easily the outer
electron of the metal is lost in each case. That depends on how strongly it
is attracted to the nucleus. The more electron shells an atom has, the less powerful the attraction forces are. For
example, Lithium is a lot less reactive than Potassium. This is because there are less shells which shield the full
attraction of the nucleus from the This makes the electron harder to lose. However, potassium has a lot more electron
shells which shield the outer electron from the nucleus. This weakens the attraction in compared to lithium, and
therefore, the electron is easier to lose.
Compounds of Alkali Metals
All group one metal ions are colourless. That means that their compounds will be colourless or white unless they are
combined with a coloured negative ion (remember metals would become positive ions because they lose electrons,
whereas, most non-metals gain electrons). Potassium dichromate is orange, for example, because the dichromate
ion is orange. Group one compounds are typical ionic solids and are mostly soluble in water.
Alkali Metals: Quick Notes






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




Group One so +1 charge
One electron on outer shell
Reactivity increases downwards
Density increases downwards
Melting and Boiling points both decrease downwards
Very soft and tarnish quickly in air
Li, Na and K are stored under oil, whilst Rb and Cs are stored in
sealed glass tubes
Reacts with air to form oxides
Reacts with water to form alkaline hydroxides, which turns pH paper
purple
Positive ions are formed and they are colourless
Flame Colours: lithium, red; sodium, yellow; potassium, lilac;
rubidium, red; caesium, blue.
Forget about Francium you don’t need to know much about it.
Group 2: Alkali Earth Metals
Alkali earth metals belong to Group two. They are beryllium, magnesium,
calcium, strontium, barium and radium. These metals are harder than
those in group one. They are silvery grey in colour. They tarnish quickly,
however they don’t just disappear into thin air because the oxides the
metals form when reacting with air would form an outer coat that protects
the metal from the air. They are good conductors of heat and electricity.
They burn in oxygen to form white oxides. They react with water to form
hydroxides and hydrogen, but the reaction is a lot less than that of group
one. Also, reactivity increases down the group.
Flame Colours
-Calcium
-Strontium
-Barium
–brick red
-crimson
-apple-green
Well that’s it for Group two!
Alkali Earth Metals: Quick Notes







Harder than group one metals
Two electrons on outer shell (2+ charge)
Form white oxides
Forms hydroxides and hydrogen when reacting with water. Reaction is less vigorous than that of group one
Reaction increases downwards
Silvery-Grey
Flame Colours: calcium, brick red; strontium, crimson; barium, apple-green.
Group 7: The Halogens
Element
State
Colours
Halogens are group seven elements. Their elements are
diatomic molecules. They exist in pairs, such as F2 and
Cl2. These two elements are gases, bromine is a liquid
and iodine is a solid. Astatine is radioactive.
Flourine
Gas
Yellow
Chlorine
Gas
Green
Bromine
Liquid
Orange – Brown vapour
Iodine
Solid
Dark Grey – Purple vapour
These vapours and gases are poisonous. All these
elements need to be handled in a fume cupboard.
Halogen
Flourine
Chlorine
Bromine
Iodine
Reaction with Hydrogen
Violent explosion, even in the cold and dark
Violent explosion if exposed to a flame or sunlight
Mild explosion if a bromine vapour/hydrogen
mixture is exposed to a flame
Partial reaction to from hydrogen iodide if vapour
is heated continuously with hydrogen
Reactions with Hydrogen
The halogens react with hydrogen to form
hydrogen halides such as hydrogen fluoride
and hydrogen chloride. These are all steamy,
acidic and poisonous gases. They are very
soluble in water, reacting with it to produce
solutions of acids. However as a gas, it is NOT
an acid.
Reaction Between Sodium and Chloride
Sodium burns in chlorine to produce the white solid sodium chloride – or salt!
2Na(s) + Cl2(g)  2NaCl(s)
In this reaction, sodium has been oxidized since it has lost electrons. Chlorine has been reduced.
Displacement Reactions with Halogens
Finding the reactivity of halogens are
done by reacting the elements with
potassium halides. Colour change
will indicate a reaction.
Observations
Potassium
Chloride
Chlorine
Bromine
Brown
Potassium
Bromide
Yellow to Brown
Potassium Iodide
Yellow to Brown
Brown to Dark Brown
Note: Colour changes are due to the
Iodine
Brown
Brown
element being displaced. For
example, the colour change from yellow to brown when chlorine reacted with potassium bromide was due
to the fact that the bromine was displace. It was the brown of the Bromine that turned the solution brown.
Potassium is only a spectator ion. It does not change.
But now we have a problem. To distinguish whether bromine or iodine has been displaced is difficult, as
both elements produce very similar shades of brown. What do we do? We add an organic solvent such
as Volasil. When Volasil is added, the iodine turns pink while the bromine stays brown. Pretty neat huh?
These reactions are known as redox reactions, where oxidation and reduction are occurring (not just one
of them).
Explaining the Trend in Reactivity of Halogens
As you go down the group, the oxidizing ability of the halogens falls due to the decreasing reactivity.
When a halogen oxidizes something, it does so by removing electrons from it. Chlorine is a strong
oxidizing agent because its atoms readily attract an extra electron to make chloride ions. Bromine is less
successful. Why? This relates to electron shells again. In Chlorine, there are three shells which shield the
nucleus’ attraction force from attracting another electron to gain a full outer electron shell. Bromine
however, has a lot more shells to shield the attraction, therefore, the force is much weaker.
Halogens: Quick Notes







Diatomic molecules
Seven electrons on outer shell
Highly reactive – only need one electron to fill outer shell
Form hydrogen halides when reacting with hydrogen
Reaction increases as you go up the group
Halogens can displace each other
Volasil turns iodine pink
The Difference Between Hydrogen Chloride and Hydrochloric Acid
Hydrochloric acid is basically a solution of hydrogen chloride gas in water.
The Bronsted-Lowry Theory
Bronsted and Lowry defined acids and bases as the following:
-An acid donates a proton.
-A base accepts a proton.
How is this related? Well, when hydrogen loses its only electron, it becomes a hydrogen ion (H+). In other words, it is
also a proton, because it has lost all of its electrons (it only has one remember?).
When hydrogen chloride dissolves in water, a proton (the hydrogen ion) is transferred to the water. This gives us the
equation:
H2O(l) + HCl(g)  H3O+(aq) + Cl-(aq)
The H3O+ ion is called a hydroxonium ion. We
normally write it as H+(aq). You can think of it as a
hydrogen ion riding on a water molecule.
So in this example, HCl is an acid because it
donates a proton (the hydrogen ion) to water.
So the real differences? Hydrogen chloride is NOT an acid and is a gas. Hydrochloric acid is an aqueous solution of
hydrogen chloride.
Hydrogen Chloride and Methylbenzene
Explaining Water Being a Polar Molecule
Water is a polar molecule. Electrons
in water are attracted towards the
oxygen end of the bond, which
leaves it slightly negative. This
leaves hydrogen slightly short of
electrons, and therefore, making it
slightly positive, just like the picture to the left.
Because of this electrical distortion, water is
described as a polar molecule.
When something such as sodium chloride is being dissolved in water, the slightly positive hydrogens cluster around
the chlorine, whereas, the slightly negative oxygen cluster around the sodium. The water molecules then literally pull
the sodium chloride crystal apart.
This pull doesn’t work on every molecule. Magnesium oxide isn’t soluble in water because the water molecules aren’t
strong enough to break the magnesium-oxygen attractions.
What’s so special about methylbenzene?
Well, methylbenzene is not a polar molecule. It is unable to pull the hydrogen and chlorine apart and therefore,
hydrochloric acid won’t be formed.
Oxygen and Oxides
(2.15) Composition of Air
This is the approximate composition of air. Memorize it.
There are also very small amounts of noble gases in the air.
Gas
Nitrogen
Oxygen
Argon
Carbon Dioxide
Amount in Air (%)
78.1
21.0
0.9
0.04
(2.16)Showing That Air Contains About 1/5 Oxygen
Using Copper
The apparatus originally contains
100cm3 of air. This is pushed
backwards and forwards of the heated
copper, which turns black as copper(II)
oxide is formed. This uses up the
oxygen. On cooling, around 79cm3 of
gas is left in the syringes – 21% has
been used up. Therefore, the air
contains 21% of oxygen.
Using the Rusting of Iron
Iron rusts in damp air, using oxygen up as it does so. The experiment shows some damp iron wool in a test tube
containing air. The tube is inverted in a beaker of water and the level of the water in the tube is marked by a rubber
band. The tube is left for a week or so for the iron to use up the oxygen to make…guess…iron oxide!
The water level rises in the tube as the oxygen is used up, and the new level can be marked using a second rubber
band. You can find the actual volumes of the gases at the end of the experiment by filling the tube with water to each
of the rubber bands in turn, and pouring it into a measuring cylinder. If the original volume was, say, 15cm3, and the
final volume was 12cm3, then the oxygen used up measures 3cm 3.
The percentage of oxygen in air was 3/15 x 100 = 20%.
Burning Phosphorus
This can be done by putting a bell jar into a beaker filled with water. Phosphorus on an
evaporating dish is placed onto the water (the jar has no bottom). It is then touched with a
hot metal rod, which starts the reaction between phosphorus and oxygen. Phosphorus uses
up the oxygen to form phosphorus oxide, lowering pressure in the jar and therefore, making
water levels rise in the jar. The water should rise up by 20%.
Equation:
(2.17)Making Oxygen in the Lab
Oxygen is most easily made in the lab from hydrogen peroxide solution using manganese(IV) oxide as a catalyst. The
reaction is known as the catalytic decomposition (splitting up using a catalyst) of hydrogen peroxide.
2H2O2(aq)  2H2O(l) + O2(g)
Reaction of Oxygen with Magnesium, Carbon and Sulfur
Magnesium
With Sulfur
With Carbon
Magnesium reacts with oxygen to produced white,
powdery magnesium oxide. It produces a bright white
flame during the reaction. It is a base.
Sulfur burns in oxygen with a tiny blue flame.
Poisonous, colourless sulfur dioxide is produced. It is
an acidic oxide.
Carbon burns in oxygen if heated strongly to give
colourless carbon dioxide. Depending on the purity of
the carbon, a small yellow-orange flame may be
produced.
Carbon Dioxide
Preparing It in the Lab
Carbon dioxide is made by the reaction between dilute hydrochloric
acid and calcium carbonate in the form of marble chips.
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
Formation of Carbon Dioxide from Thermal Decomposition of
Metal Carbonates
Key thing here: When heating metal carbonate, you get:
Metal Carbonate  Metal Oxide + Carbon Dioxide
2Mg(s) + O2(g)  2MgO(s)
S(s) + O2(g)  SO2(g)
C(s) + O2(g)  CO2(g)
Here is the picture of the experiment setup:
Properties of Carbon Dioxide




Colourless gas, denser than air, slightly soluble in water
Used in carbonated (fizzy) drinks because it dissolves in water under pressure. When bottle is opened,
pressure falls and gas bubbles out of the solution.
Used in fire extinguishers to put out electrical fires, or those caused by burning liquids, where using water
could cause problems. The carbon dioxide sinks onto the flames and prevents any more oxygen from
reaching them.
Turns limewater cloudy white (limewater is calcium hydroxide – work out the equation yourself – water is
one of the products).
Carbon Dioxide and Sulfur Dioxide – Their Reactions With Water
Carbon Dioxide
Sulphur Dioxide
Carbonic acid is produced when carbon
dioxide reacts with water. It is a weak
acid. This reaction can be reversed by
simply heating or boiling the acid.
Sulfur dioxide reacts with water to form
a weak acid known as sulfurous acid,
CO2(aq) + H2O(l) ⇌ H+(aq) + HCO3–(aq)
H2O(l) + SO2(g)  H2SO3(aq)
Sulfur Dioxide, Nitrogen Oxide and the Environment
Acid rain is caused when oxygen and water in the atmosphere react with sulfur dioxide to produce sulfuric acid
(ouch), or with various oxides of nitrogen to give nitric acid. These mainly come from power stations, burning fossil
fuels, motor vehicles etc.
Acid rain can kill trees and make lakes so acidic it cannot support
life. Limestone and some metals such as iron are also attacked by
acid rain.
The solution to acid rain involves removing sulfur from fuels, using
catalytic converters in cars and scrubbing the gases from power
stations to remove the oxides. The catalyst helps convert nitrogen
oxides into harmless nitrogen gas but has no effect on sulfur dioxide.
Methods of Separation
Filtration: For separating an insoluble solid from a liquid, or
a soluble solid from an insoluble one.
Sand can be separated from water by pouring the mixture
down a funnel with filter paper. The sand will collect at the
filter paper.
It can also be used to separate sand from something like salt
by dissolving salt in water (which leaves you with sand mixed
with salt water). The mixture can then be poured down a
funnel. The sand that collects at the top should can be rinsed
and dried. The water can be evaporated from the salt by
heating with a Bunsen burner. Back
Chromatography: For separating liquids by
dissolving them in a solvent. The dyes that make up
the ink differ in two important ways:

How strongly they stick to the paper

How soluble they are with the solvent
An example would be separating ink colours or plant dyes. A dot of the ink/dye would be drawn onto a piece of paper.
It would then be left in water, which acts as a solvent. Because different colours have different solubility levels, some
colours would travel up further on the paper.
Crystallization: Mainly used for purifying substances by forming crystals from a precipitating solution.
Crystallization refers to the forming of solid crystals from a homogenous (solution) mixture.
An example would be forming pure salt crystals. This is done by dissolving the impure salt into a solvent such as
water. The salt solution is then allowed to cool. As it does, pure salt crystals would form at the bottom of the water,
whereas, the impure substances would be left in the water. The crystals can then be rinsed with a chilled solution and
dried.
Distillation: Distillation is good from separating
a liquid from a solution.
An example would be separating water from a salt
solution. The solution would be heated at the
liquid’s boiling point, in this case 1000C, so it will
leave the solution as a vapour. The vapour would
then condense into a liquid with the help of the
cooling water. The vapour, now as a liquid, would
fall into the beaker.
Fractional Distillation: Fractional Distillation
is used to separate two liquids based on their
boiling points.
An example would be separating ethanol from water.
Ethanol has a lower boiling water than water (at about
780C), therefore, the heating is monitored (using the
thermometer) to ensure that the temperature does not
reach 1000C (the boiling point of water). Anyways, the
ethanol would turn into a vapour and travel out of the
flask. It would then condense into its liquid form with
the help of the cooling water and fall into the beaker.
Unit 2: Structure and Bonding
Ionic Bonding
Ionic bonding is the bonding in which there has been a transfer of electrons from one atom to another to produce
ions. The substance is held together by strong electrostatic attractions between positive and negative ions. Ions are
formed when it gains or loses electrons. Ones that gain forms negative ions, and ones that lose form positive ions.


A positive ion is called a cation.
A negative ion is called an anion.
You can find the charge of an ion by looking at the group it belongs to. If it belongs to groups 1-4, it has a charge of 14+ (they are positive), whereas, if it belongs to groups 5-0, it has a charge of 3—0. Below is a table containing charges
of common ionic compounds and transition metals.
Ion
Symbol
Charge
Silver
Copper (I)
Ag
Cu
1+
1+
Ammonium
NH4
1+
Ion
Symbol
Charge
Copper (II)
Cu
2+
Cobalt
Co
2+
Nickel
Ni
2+
Hydroxide
Nitrate
Hydrogen Carbonate
OH
NO3
HCO3
111-
Zinc
Zn
2+
Iron (II)
Fe
2+
Carbonate
Sulphate
CO3
SO4
22-
Chromium
Cr
3+
Iron (III)
Fe
3+
Phosphide
Phosphate
P
PO4
33-
Nitride
N
3-
This is an example of a dot and cross diagram. The crosses represent the
electrons on the sodium (anion) and the dots represent the electrons on the
chlorine (cation). In a dot and cross diagram, you must use arrows to show
which electrons are moved from the anion to the cation. On the final
diagram, you mark the new electron(s) on the cation as a cross.
Boiling and Melting Points of Ionic Compounds
Ionic compounds have high boiling and melting points due to strong
intermolecular forces between the atoms. This is because when the ions
are formed during an ionic reaction, one of them would be positive, and one
would be negative. Positive and negative attract and therefore, you get
something like a strong magnet.
 As ionic charge increases, so does the melting/boiling points. Ions
with 2+ and 2- would have stronger attraction because their charges a
stronger, whereas, ions with 1+ and 1- would still have a strong attraction,
but less stronger than 2+- compounds.
Structure of Ionic Compounds
An ionic crystal consists of giant three-dimensional lattices held together by strong electrostatic attractions between
the positive and negative ions.
Structure of Sodium Chloride
This is the basic structure of a sodium chloride crystal. The green is the
chloride and the blue is the sodium. Remember that each sodium is
touched by six chlorides and each chloride is touched by six sodiums.
Look at the middle atoms if unclear. Remember, this structure repeats
itself over and over.
Ionic bonds always produce giant structures.
Ions form closely packed regular lattice arrangement.
They have high melting/boiling points.
The crystals tend to be brittle.
Compounds tend to be soluble in water and insoluble in organic
solvents.
Covalent Bonding
Covalent bonding is formed by sharing a pair of electrons between two
atoms. This is so that both atoms can achieve a full outer shell. It is a strong attraction between the bonding pair of
electrons and the nuclei of the atoms involved. Covalent compounds are only formed when the reactants are nonmetals.
Diagrams YOU Need to Know…
Element
Diagram
H2
Eleme
nt
CH4
Cl2
NH3
HCl(g)
O2
H2O
N2
Diagram
Element
Diagram
CO2
Ethane
Ethene
Simple Molecular Structures
These are gases, liquids or solids with low melting points. Examples include
water, chlorine, oxygen…etc
The covalent bonds between the atoms in a molecule are strong.
However, the forces of attraction between these molecules (inter-molecular
forces) are weak.
They have low melting points, since not a lot of heat is needed to provide the
energy for the molecules to move away from each other, hence, overcome the intermolecular forces
between them.
They tend to be insoluble in water.
They are often soluble in organic solvents.
They do not conduct electricity because the molecules have no overall charge and there are no electrons
mobile enough to move from molecule to molecule.
Giant Covalent Structures
There are no charged ions.
ALL the atoms are joined up to their adjacent atom by extremely strong covalent bonds and packed into
giant regular lattices.
They have very high melting points, since a lot of heat is needed to provide the energy to break apart the
many strong covalent bonds.
They tend to be insoluble in water.
They do not conduct electricity.
Diamond
The diamond is the hardest natural substance. It is a form of pure carbon. Each carbon
atom forms four covalent bonds to the other carbon atoms. They are arranged in a
tetrahedral arrangement. Diamond has a very high melting point, obviously due to very
strong carbon-carbon bonds. It does not conduct electricity because all the electrons in
the outer levels of the carbon atoms are tightly bonded between the atoms. None of them
are free to move around. Diamond is insoluble – like, to both water and other solvents.
Use of Diamond
Saw blades can be tipped with diamonds in high-speed cutting tools used on stone and concrete. The strong
tetrahedral structure makes the diamond hard, making it suitable for this purpose.
Graphite
Graphite is arranged differently – it has a layer structure. Each graphite layer is strong,
but it is easy to separate individual graphite layers. Each carbon atom only forms three
covalent bonds. Graphite conducts electricity because the fourth electron is free to
move around.
Use of Graphite
Because of the layered structure, graphite can be used as a dry lubricant to
lubricate locks.
Metallic Crystals
Metals are giant structures which consist of a regular array of positive ions
in a sea of delocalized electrons. When metal atoms bond together to form
solid, visible metal, their outer electrons are no longer attached to particular
electrons and are free to move around the whole structure.
Metals are able to conduct electricity because the delocalized electrons are
free to move throughout the structure. The energy is picked up by the
electrons and moved around the metals, transferring the electricity
throughout the whole structure. The same goes to heat energy.
Metals are easy to shape because their regular packing makes it simple for
atoms to slide over each other. Metals are said to be malleable.
Introduction to Electrolysis
In metals and carbon, electricity and electric current is simply a flow of electrons or ions. Electrolysis is the chemical
change caused by passing an electric current through a compound which is either molten or in a solution. An
electrolyte is a substance that undergoes electrolysis. It contains ions. It is the movement of the ions, which are
responsible for both the conduction of electricity and the chemical changes that take place. Covalent compounds are
not electrolytes and don’t conduct electricity because they have no free moving electrons. Ionic compounds only
conduct electricity when molten or in a solution because the ions separate and are free to move. These particles can
then carry the electric current.
Experiment to distinguish between electrolytes and non-electrolytes
i.
ii.
iii.
Dissolve substance in water, or if possible, melt it.
Put a conductivity tester into the substance.
If the light bulb lights up, it is an electrolyte.
Explain: When dissolved in water, free moving electrons are able to carry the electric current across from the
cathode to the anode, completing the circuit and lighting the bulb. If the light bulb does not light up, the substance
is obviously not an electrolyte.
But sugar dissolves, why does the bulb not light up? Sugar is a covalent structure.
Diffusion
Diffusion happens when particles spread from higher to lower concentration. It requires a concentration gradient).
Potassium Manganate (VII) Experiment
Diffusion through liquids is very slow if the liquid is totally still. This can be shown but dropping a piece of potassium
manganate (VII) into water. It can take days for the colour to spread
because the gap between each particle is small.
The Bromine Experiment
Showing diffusion in gases can be done by filling a lower gas jar with
bromine gas and topping it with a gas jar filled with air. The bromine
particles and air particles will eventually bounce around to give an even
mixture.
The Ammonium Chloride Experiment
This experiment is used to show that particles in different gases travel at
different speeds. It relies on the reaction between ammonia and
hydrogen chloride gases to give white solid ammonium chloride.
A white ring of ammonium chloride would form near the hydrochloric acid. This shows that ammonia particles have
travelled further to reach the hydrogen chloride gas, showing that it travels faster.
Dilution
Dilution is the reduction of concentration in a solution.
Showing Dilution and Leading to the Idea of Small Sized
Particles
Suppose you dissolve 0.1g of potassium manganate (VII)
in 10cm3 of water to give a deep purple solution. Assume
the smallest drop you can see is 1/1000cm3. The whole
solution will be made up of 10000 drops, each drop
containing 0.00001g of potassium manganate (VII).
Suppose you dilute this down 10 times by taking 1cm 3 of
the solution and making it up to 10cm3 with more water.
Continue doing this until the colour is too faint to see. By
the time of the fifth dilution, each drop will only contain a
billionth of a gram of potassium manganate (VII). If you only needed one ‘particle’ of potassium manganate (VII) per
drop in order to see the colour, the ‘particle’ can’t weigh more than a billionth of a gram.
IS this a good answer? Nowhere near it! A potassium manganate (VII) particle actually weighs about
0.00000000000000000000026g! In reality, you need huge numbers of particles in each drop in order to see the
colour.
Don’t worry I don’t get this either…
Unit 3: Organic Chemistry
Organic chemistry is mainly based around hydrocarbons – compounds made only up of hydrogen and carbon. It is
drawn with lines joining carbons and hydrogen. All carbon bonds have to be bonded to hydrogen – if not something
else. The left picture below shows carbons with all bonds taken up (ethane). The right picture below shows an
incorrect picture of a hydrocarbon because one of the carbons has a free bond.


 Hydrocarbon – compounds that contain carbon and hydrogen only.
 Homologous series – family of compounds with similar properties
because they have similar bonding. They show a graduation in physical
properties (mpt/bpt) and similar chemical properties such as the general
formula. Alkanes are the simplest.
 Saturated – when carbon cannot take anymore bonds – single
carbon-carbon bonds.
 Unsaturated – presence of a carbon-carbon double bond.
General formula – The formula of different homologous series of carbons.
Isomers – molecules with the same molecular formula but different structural formulae.
Learning the Code
Do you have to remember the formula for propane, butane,
ethane…? No! You can work it out yourself! The first part of the
name tells you how many carbons there are in the longest chain
(not necessarily in total). By the way you have to learn these – at
least the first five. It helps.
For example: propane (left) has three carbons. Butane (right) has
four carbons.
Alkanes and
Code
Number of Carbons
Meth
Eth
Prop
But
Pent
Hex
Hept
Oct
Non
Dec
1
2
3
4
5
6
7
8
9
10
Alkenes
Alkanes and Alkenes are two homologous series.
Pentene has a five carbon chain with a double bond.
Ending
Meaning?
‘ane’
All carbon bonds are
filled with hydrogen –
i.e. they are
saturated
hydrocarbons.
There is a double
carbon-carbon bond
– they are
unsaturated
hydrocarbons.
‘ene’
Isomers
We know what isomers are.
Coding for Double Bonds
For things like pentene and butene, there are many places you can put the double bonds
in.
Pent-1-ene means pentene with the double bond on the first carbon-carbon bond (right).
Pent-2-ene means pentene with the double bond on the second carbon-carbon bond and so on…
But wait! What about pent-4-ene and pent-5-ene? Those don’t exist. Why? Because pent-4-ene is pent-2-ene flipped
over, and pent-5-ene is pent-1-ene flipped over!
Flip: C-C-C=C-C and you get C-C=C-C-C!
Methyl and Ethyl Groups
If the hydrocarbon has a methyl or ethyl group, these two come first,
before the coding for the number of carbons in the chain. But before
even the ‘methyl’ or the ‘ethyl’ there is a number and hyphen to show
which carbon has the methyl or ethyl branch.
For example, this is 2-methylbutane
Code
Meaning
Methyl
Has a branch of CH3
coming off one of the
bonds.
Has a branch of
CH3CH2 coming off.
Ethyl
As you can see, there is a methyl group branching off the second carbon. The rules
are similar to double bonds though, there is no such thing as 3-methylbutane
because that is basically 2-methylbutane flipped over.
But wait! There are five carbons! Why isn’t it 2-methylpentane? Because remember,
these names are based on the longest carbon chain in the hydrocarbon and the
longest carbon chain there is 4, hence, butane. This means that 2-methylbutane is
an isomer of pentane C5H12.
Some Isomers of Butane C4H10
Some Isomers of Pentane C5H12
Alkanes
Alkanes are a homologous series of saturated hydrocarbons. The first five are methane, ethane, propane, butane and
pentane.
The general formula for alkanes is: CnH2n + 2
For example, ethane:
Ethane has two carbons, so n=2.
The formula of ethane must be C2H2(2) +
2
= C2H6.
Complete Combustion of Alkanes
If there is enough oxygen, alkanes will burn in oxygen completely to give carbon dioxide and water. The general
equation for combustion:
Hydrocarbon + Oxygen  Carbon dioxide + Water
The combustion of methane would be: CH4(g) + 202(g)  C02(g) + 2H20(l)
Note: Balancing combustion equations can be annoying. An easy way would be to balance them in the order of
carbon, hydrogen then oxygen.
Incomplete Combustion
If there isn’t enough oxygen, you get incomplete combustion, in which carbon monoxide and water are produced
instead. Carbon monoxide is a colourless, odourless and poisonous gas. It is dangerous because it can combine to
our haemoglobin and stop it from carrying oxygen. As a result, you get ill or even die because oxygen cannot travel to
all parts of your body.
Reaction with Bromine
Alkanes react with bromine under the presence of ultra-violet light. One hydrogen from the hydrocarbon would be
replaced by a bromine atom. This is known as a substitution reaction. Bromine can be used as an indicator for
alkanes and alkenes without UV light. Adding bromine water to alkanes produces no colour change. Reacting
bromine water to alkenes make it turn from brown to colourless.
However, if the mixture of bromine and methane is reacted under UV light, it loses its colour, a mixture of
bromomethane and hydrogen bromide gases is formed.
CH4(g) + Br2(g)  CH3Br (g) + HBr(g)
Alkenes
Alkenes have double bonds, making them unsaturated hydrocarbons.
Alkenes have the general formula of CnH2n – the first four being ethene, propene, butene and pentene.
Combustion
Like alkanes, alkenes burn in oxygen or air to give carbon dioxide and water.
Reaction with Bromine
Alkenes undergo addition reactions, in which part of the double bond breaks
and is used to join other atoms onto the two carbon atoms. When added to
alkenes, and the test tube is shook, the brown of the bromine would be
decolourised, making it suitable as a test for alkenes.
The product of reacting ethene to bromine gives 1, 2-dibromoethane and is a colourless liquid.
CH2=CH2(g) + Br2(aq)  CH2BrCH2Br(l)
Ethanol
All alcohols contain an –OH group attached to a carbon chain. Ethanol is C2H5OH.
What is Needed:
Production of Ethanol




Hydration of Ethene
Ethanol can be made by reacting ethene with steam (because it contains more energy)
– a process known as hydration.
Ethene and steam
3000C
60-70 atmospheres
Phosphoric acid
catalyst
CH2=CH2(g) + H2O(g)  CH3CH2OH(g)
Only a small portion of ethene reacts. The ethanol is condensed as a liquid and the unreacted ethene is recycled.
Explaining the Choice of Temperature
Reversible reactions happen in two ways – while ethene is being converted into ethanol, ethanol is also being
converted back into ethene. Reversible reactions can also shift the equilibrium – or ‘alter’ the reaction. Since the
reaction is exothermic – the reaction produces lots of heat. If you increase the temperature, the reaction won’t like it
because it is already producing heat, therefore, it would ‘adapt’ to the conditions by making more ethene so less heat
will be produced. On the contrary, if you decrease the temperature, the reaction would ‘adapt’ to this by increasing
back the temperature; by producing more ethanol – in other words, push the equilibrium to the favourable/forward
reaction. However, making the temperature too low would mean super slow reaction, although more ethanol would be
produced. 300 degrees is therefore, a compromise temperature producing an acceptable yield of ethanol in a short
time.
Explaining High Pressure
In the equation, you have two moles (one mole of ethene and one mole of water) on the left, and one mole (of
ethanol) on the right. Increasing the pressure would mean the equilibrium would be shifted forwards. Why? The
reaction would ‘adapt’ to the conditions by producing more ethanol because you only get one mole of ethanol – which
takes less space than two moles of ethene and water.
Also, there’s the collision theory. Increasing the pressure means that there’d be less space for the atoms to move.
The atoms would also move with more force. This increases the frequency of collisions.
The problem: it’s expensive and ethene might polymerise and turn into polyethene.
Fermentation
Yeast is added to a sugar or starch solution at 300C for several days in the absence of air for anaerobic respiration.
Enzymes in the yeast lower the activation energy, increasing the rate of conversion of the sugar into ethanol and
carbon dioxide. However, they first have to break the sugars into smaller sugars like glucose. In fact, ethanoic acid is
produced and then converted into ethanol.
For example, sucrose:
C12H22O11(aq) + H20  C6H12O6(aq) + C6H12O6(aq)
sucrose + water  glucose + fructose
C6H12O6(aq)  2C2H5OH (aq) + 2CO2(g)
glucose/fructose  ethanol + carbon dioxide
The yeast then gets killed in the mixture, which means that the ethanol produced is impure. To purify it, the alcohol
must undergo fractional distillation.
Comparing the two methods…
Use of Resources
Type of Process
Rate of Reaction
Quality of Product
Reaction Conditions
Fermentation
Uses renewable resources – sugar beet or
sugar cane, corn and other starchy
materials.
A batch process – everything is mixed and
left for several days. It is then removed and
a new reaction is set up – quite inefficient.
Slow, takes several days.
Produces impure ethanol that needs further
processing.
Uses gentle temperatures and ordinary
pressure – relying on anaerobic respiration
of yeast.
Hydration
Uses non-renewable resources – once oil
gets used up, they’re screwed.
A continuous flow process – a stream of
reactants is constantly passed over the
catalyst – more efficient.
Rapid
Produces much purer ethanol.
Uses high temperatures and pressures,
needing a high input of energy – expensive.
Common Question: Which method would poorer places like Brazil use and why? [3 marks]
Answer: Fermentation, because Brazil has the weather conditions to grow large yields of sugar cane and they don’t
have access to crude oil.
Dehydration of Ethanol into Ethene
Dehydration of ethanol produces ethene and water, using hot aluminium oxide as a catalyst.
CH3CH2OH(g)  CH2=CH2(g) + H2O(l)
Crude Oil
Crude oil is a mixture of hydrocarbons. These chains can be super long or super short.
The Trend in Boiling Point and Viscosity


Viscosity means how runny something is
Volatile means how easy it turns into vapour at room temperature
As the number of carbon atoms in molecules increases and gets bigger, intermolecular attractions also increase,
making it more difficult to pull one molecule away from neighbouring ones. As they get bigger, these changes occur:




Boiling point increases – the larger the molecule, the higher the boiling point due to stronger intermolecular
attractions.
Liquids become less volatile – the bigger the hydrocarbon, the more slowly it evaporates in room
temperature. This is again, due to strong intermolecular attractions.
Liquids become more viscous (flow less easily) – Small hydrocarbons are runny, but large ones are much
stickier and gooey (and viscous) because of intermolecular attractions.
Bigger hydrocarbons do not burn as easily, meaning they are less useful.
The Fractionating Column
Crude oil is separated in fractionating
column. This process is fractional
distillation, and splits crude oil into various
fractions depending on their boiling points
and size.
Note: forget about Naphtha
Fraction
Refinery
gases
Uses
Gasoline
Kerosene
Diesel oil


Fuel oil
Bitumen



A mixture of
methane, ethane,
propane and
butane.
Commonly used for
domestic heating
and cooking.
Cars
Used as fuel for jet
aircraft.
As domestic
heating oil.
As ‘paraffin’ for
small heaters and
lamps.
For buses, lorries,
some cars and
railway engines.
Some is cracked to
produce more
petrol.
For ships
Industrial heating
Residue from the
bottom which can
be used for roads.
Combustion and Incomplete Combustion


Combustion of hydrocarbons produces carbon dioxide and water – exothermic.
Incomplete combustion of hydrocarbons produces carbon monoxide and water – in which carbon monoxide
is dangerous because it can bind to haemoglobin and prevent it from carrying oxygen.
In car engines, the temperature reached is high enough to allow nitrogen and oxygen from the air to react, forming
nitrogen oxides. This contributes to smog and causes irritation to human mucus membranes. As well as that, nitrogen
oxides can react with water in the atmosphere and from nitric acid – or acid rain.
Cracking
The Crude Oil Problem
Amounts of each fraction you get depend on the proportions of various hydrocarbons in the original crude oil. Far
more petrol is needed, than something like bitumen. In other words, fractional distillation of crude oil produces more
long-chain hydrocarbons than can be used directly, and fewer short-chain hydrocarbons than required.
The solution? Cracking! Cracking is a useful process in which large hydrocarbon molecules are broken into smaller
ones. Most of the hydrocarbons found in crude oil are long-chain alkanes. Cracking can convert these into alkenes
and shorter alkanes. It is an example of thermal decomposition.
How it Works
The fraction is heated to give a gas and is passed over a catalyst of silica or alumina with
a temperature of 600-700oC.

Long alkane  alkene + alkane


Sometimes you may get more than one type of alkene/alkane.
Make sure the numbers of carbon and hydrogen are balanced.
In an equation, this would read:
What is Needed:
hexane  butane + ethene

Alumina/Silica as
catalyst
600-7000C
C6H14  C4H10 + C2H4
Polymers
Alkenes can be used to make polymers. Polymers are big long molecules of single units called monomers. Molecules
containing carbon-carbon double bonds can be joined together. Part of the double bond is broken and used to join to
other monomers. Joining up lots of monomers to make a polymer is called addition polymerisation.
How to Draw a Polymer
It’s like drawing a hydrocarbon, except the ends are left blank (so it can join more).
As for the repeating unit (which is the unit that keeps repeating itself), you show the
alkene (or the monomer) with its double bond opened up. You then enclose it with brackets and
put an ‘n’ to its right.
Polymers to Know
Polymer
Polyethene
Repeating Unit
How it looks together
Uses
Plastic bags
Plastic bottles
Polypropene
Ropes
Crates
Polychloroethene
PVC for drainpipes
or windows
Electrical insulation
Nylon – Condensation Polymer
In condensation polymerisation, when two monomers combine, a small molecule such as water or hydrogen chloride
is lost. Nylon is made through condensation polymerisation.
The two monomers that make up nylon:
Hexanedioic acid
From a family of compounds called dicarboxylic acids.
1,6-Diaminohexane
From a family known as diamines.
Joining them together…The lost of a Water Molecule
As a block diagram (where the (CH2)6 and (CH2)4 become blocks to make it look easier)
End Note:
Sometimes you may be
given ClOCCH2 CH2 CH2 CH2COCl instead of hexanedioic acid. In this case, just do the same thing, with the lost of
hydrogen chloride HCl.
Unit 4: Analytical Chemistry and Kinestics
Tests for Ions and Gases
Flame Tests: Taking a piece of nichrome, make loop at the end and dip into salts containing ions
Ion
Colour
Li+
Crimson red
Na+
Yellow orange
K+
Lilac
Ca2+
Brick red/orange red
Using Sodium Hydroxide solution
Ion
Colour of Precipitate
Cu2+
Blue
Fe2+
Sludgy Green (or just green)
Fe3+
Orange Brown (rust)
For Ammonium Ions (NH4+)
Heat gently and add sodium hydroxide solution. It will give off a distinctive smell of ammonia (NH3). Ammonia can be
tested by holding a damp red litmus paper. Since it is alkaline, it will turn damp red litmus paper from red to blue.
Using Dilute Nitric Acid and Silver Nitrate Solution
Ion
Colour of Precipitate
Cl-
White
Br-
Pale Cream
I-
Yellow
For Sulphate Ions (SO42-)
Using dilute hydrochloric acid solution and then adding barium chloride solution to form a white precipitate of barium
sulphate.
For Carbonate Ions (CO32-)
Using dilute hydrochloric acid to react with the carbonate, to produce carbon dioxide gas which can be tested by
bubbling through limewater, turning it from colourless to cloudy, milky white.
Tests for Gases
Gas
Test
Result
Hydrogen
Hold a lit splint in presence of hydrogen gas.
Produces a squeaky pop.
Oxygen
Hold a glowing splint in presence of oxygen gas.
Glowing splint relights.
Carbon Dioxide
Bubble through limewater.
Ammonia
Hold damp red litmus paper in ammonia gas.
Turns limewater from colourless to cloudy,
milky white.
Turns damp red litmus paper blue.
Chlorine
Hold damp blue litmus paper in chlorine gas.
Bleaches or turns blue litmus paper white.
Solubility Patterns


All nitrates are soluble.
All sodium, potassium and ammonium compounds are soluble.



Most carbonates and hydroxides are insoluble except for sodium, potassium and ammonium.
All sulphates are soluble except barium and lead(II) sulphate.
All chlorides are soluble except lead(II) and silver chloride.
Reactions of Metals to Acids
Metals react very similarly to dilute hydrochloric acid and dilute sulphuric acid.
Metals
Magnesium
Aluminium
Zinc
Iron
Reaction to Acid
Rapid fizzing, mixture gets very hot, colourless magnesium sulphate/chloride solution forms.
Is slow due to its coat of aluminium oxide which prevents aluminium from contacting the acid.
On heating, this layer is removed, aluminium will start fizzing rapidly – abit like Mg.
Zinc reacts slowly with cold dilute acid and may produce some effervescence. On heating
however, it fizzes more.
Iron also reacts slowly with cold dilute acid and will produce abit of effervescence when heated.
Combustion of Hydrogen
Bonds are broken in the hydrogen and oxygen molecules. These form new bonds of water molecules. This reaction is
exothermic, and gives out water in the form of steam, before it condenses into a liquid. The reaction is:
2H2(g) + O2(g)  2H2O(l)
Testing for Water
Water turns white anhydrous copper(II) sulphate blue. It’s reaction is
CuSO4(s) + 5H2O(l)  CuSO4•5H2O
Or you can use cobalt chloride paper – which turns pink in the presence of water.
You can check the purity of water by showing that it freezes at exactly 0ºC and boils at exactly 100ºC.
Rates of Reactions
Experiment Setup
To measure the effects of changes in surface area, concentration of solutions, temperature and use of catalyst, you
can react calcium carbonate marble chips with dilute hydrochloric acid and measure the mass of CO 2 produced by
weighing the difference in mass of the reactants and the mass of the products (there won’t be any change in mass
produced, because the initial mass of reaction will equal the final mass, however, since carbon dioxide gas is formed,
this will escape from the flask, and therefore, the amount of mass lost will be
the mass of carbon dioxide produced. Plot the results on a graph with mass
against time and you’ll get an upward curve.
For the reaction to occur, acid particles must collide with the surface of the
marble chips. As the acid particles get used up, the collision rate decreases, so
the reaction slows down.
Changes in Surface Area of Solid
You can repeat the above experiment by keeping the
same mass of marble chips, just using smaller ones to
increase the surface area. The reaction happens
faster. You have to remember the graphs. Notice
however, that in the end, the amount of carbon dioxide
produced is still the same – just that the small chips
experiment happens faster.
Why does it happen faster? Because the surface area in contact with the gas or liquid is much greater. Less marble
chip particles are hidden away from the acid particles.
Changes in the Concentration of
Solutions
Repeat the original experiment but using
only half as concentrated as before. The
something like this (ignore the 80% line)
reaction happens slower and produces
carbon dioxide gas:
hydrochloric acid
graph should look
– in which the
half as much
In terms of collision theory, if you
concentration of reactants, the reaction
because it increases the frequency of
second.
increase the
becomes faster
collisions per
Changes in the Temperature of the Reaction
Do the original experiment again, but this time, at a higher temperature.
Your graph will look like this (ignore the concentration label ‘cause it’s
WRONG unless the lower concentration solution is still in excess).
Increasing the temperature means more kinetic energy for the particles,
which make them move faster, therefore, making them collide more
frequently.
Also, not all collisions make new bonds. Some particles just bounce off each
other. In order for a reaction to happen, particles have to collide with a
minimum amount of energy called activation energy. Increasing the
temperature produces a very large increase in the number of collisions that
have enough energy for a reaction to occur.
In the following diagram, a) shows a fail collision and b) shows a successful
one.
Changing
the Pressure
Changing the pressure of a reaction where the reactants are only
solids or liquids makes virtually no difference, so the graphs remain
unchanged. Increasing the pressure in a reaction where the reactants
are gases does speed the reaction up. This is because it forces the
particles closer together, so they hit each other more frequently.
Catalysts and How
They Work
Catalysts speed up the rate of reactions but aren’t used up in the
process. You can show that manganese (VI) oxide is a catalyst by
simply having two conical fasks containing hydrogen peroxide.
Hydrogen peroxide decomposes to give oxygen and water. Put the manganese (VI) oxide in one of the flasks.
Oxygen would be given off quickly. To check that the manganese (VI) oxide hasn’t been used up, simply filter it out
from the solution and weigh it (remember to weigh it before the experiment too!). The graph should look like the
pressure graph – in which the rate of reaction increases, but the amount you get at the end is still the same.
So how does it work?
Adding a catalyst gives the reaction an alternative route for reactions with a lower activation energy.
Unit 5: Quantitative Chemistry and Energetics
A mole is a measure of the amount of substance. One mole contains 6 x 1023 (also known as the Avogadro Number)
particles (atoms, molecules or formulae) of the substance. For example, 1 mol of sodium contains 6 x 10 23 atoms of
sodium.
Calculating Relative Atomic Mass
Chlorine has two isotopes: chlorine-35 and chlorine-37. A typical sample will be 75% chlorine-35 and 25% chlorine37.
The RAM = (0.75 x 35) + (0.25 x 37) = 35.5g
Calculating Relative Formula Mass
Straightforward stuff.
H2
Ca(OH)2
Hx2
(1 x Ca) + (2 x O) + (2 x H)
2x1=2
40 + (2 x 16) + (2 x 1) = 74
Amount in moles = Mass of Substance (g) /RFM of Element or Compound (g)
Calculations using moles:
The equation for sodium chloride is: 2Na + Cl2  2NaCl
If 2.3g of Na was used:
a)
b)
c)
Find out how many moles of Chlorine was used
Find out the volume of Chlorine used in the reaction
Find out the mass of sodium chloride produced
a)
Firstly, convert the grams of Na into moles:
2.3 / 23g = 0.1 mol
The equation says that 2 moles of Na and 1 mole of Cl (1 mole of a diatomic molecule is always X 2) is needed to
produce 2 moles of NaCl, so if 0.1 mol of Na is used, then half of that is the amount of chlorine used in the reaction in
moles.
So moles of Cl used = 0.1 / 2 = 0.05 mol
b)
One mole of any gas has a volume of 24 dm3 (24000cm3) at room temperature and pressure. This is also
called the molar volume.
Cl2 is a gas and the moles used in the reaction = 0.05 mol
So the volume of Cl2 gas used = 0.05 x 24000 = 1200cm3
c)
The moles of NaCl produced is 0.1 mol (if 2 moles of Na gives 2 moles of NaCl, then 0.1 mole of Na will give
0.1 mole of NaCl). So all you do is:
i)
Find the RFM of NaCl (58.5)
ii)
Multiply that by 0.1 (5.85g)
Molar Concentrations – The Hard Part
Remember that:


Mol/dm3 means moles per litre (e.g. a salt solution of 0.5 mol/dm 3 means 0.5 moles (or 58.5/2 = 29.25g) of
salt was dissolved in a litre of water
It’s all about proportion
20 cm3 of 0.5 mol dm3 sodium hydroxide solution was dissolved with 25cm3 of hydrochloric acid to form a sodium
chloride solution. Calculate the concentration of HCl needed to react with the NaOH
NaOH + HCl  H2O + NaCl
RFM of NaOH = 40g = 1 mole of NaOH
0.5 mol dm3 of NaOH means (40 x 0.5) 20g of NaOH was dissolved in 1000cm 3 of water
The amount of moles in 20cm3 of NaOH solution:
20cm3/1000cm3 x 0.5 moles = 1/50 x 0.5 = 0.01 moles of NaOH
The equation says that 1 mole of NaOH + 1 mole of HCl gives 1 mole of NaCl
So 0.01 moles of NaOH + 0.01 moles of HCl gives 0.01 moles of NaCl
So 0.01 moles of HCl was present in 25cm3 of HCl solution! However, concentration is measured in mol dm 3 so:
1000cm3/25cm3 x 0.01 mol = 0.4 mol dm3 of HCl used.
Calculating the Empirical Formula and Molecular Formula
The empirical formula is the simplest formula and only tells you the
ratio of the various atoms. Suppose 2.4g of magnesium combined
with 1.6g of oxygen, you can use a table to work out the empirical
formula. (Mg = 24 O = 16)
Percentage
Combining Masses
Number of moles
=
Ratio of Moles
Empirical Formula
C
87.5
87.5
85.7/12
7.14
H
14.3
14.3
14.3/1
14.3
1:2
CH2
What about with
percentage
figures?
Combining Masses
Number of moles
=
Ratio of Moles
Empirical Formula
Mg
2.4
2.4/24
0.10
O
1.6
1.6/16
0.10
1:1
MgO
Suppose you had a compound containing 85.7% C, 14.3% H and
you were asked to calculate the empirical formula. Firstly, you
assume that 100% = 100g! (C = 12 H = 1)
However, you know that CH2 does not exist. Remember this is only the ratio. To find the molecular formulae, you
need to know the relative formula mass of the compound. Suppose it was 56g for the above question.
Firstly, find out the RFM of CH2 = 12 + 2 = 14g
Find out how many times 14 goes into 56, so 56/14 = 4 times
Which means the molecular formula is C4H8!
Obtaining Formulae Experimentally
Metal Oxides
Hydrogen can be passed over metal oxides to reduce it to the
metal. To find the formula of copper oxide, the experimental steps
are as follows:
1.
2.
3.
Measure the mass of the empty combustion tube.
Use a spatula to put copper oxide into the tube. Weigh
the tube.
Set up the apparatus as shown. Turn the gas at the jet
to light the excess gas.
4.
5.
6.
7.
Heat the copper oxide until it has all turned into red copper.
Stop heating but leave gas passing through until everything has cooled.
Weigh the combustion tube.
Put masses in a table and calculate empirical formula from there.
Mass of Empty Tube
Mass of tube + Copper Oxide (Before)
Mass of tube + Copper
Mass of Oxygen
Mass of Copper
52.2g
66.6g
65.0
66.6 – 65.0 = 1.6g
65.0 – 52.2 = 12.8g
In the Case of Water of Crystallisation
Combining Masses
Number of moles
=
Ratio of Moles
Empirical Formula
Cu
12.8
12.8/64
0.20
O
1.6
1.6/16
0.10
2:1
Cu2O
When substances crystallise from a solution, water becomes chemically bounded with the salt. This is called water of
crystallisation and the salt is said to be hydrated.
Suppose you had to find the formula of a BaCl2•nH2O (a barium chloride crystal), to find n:
1.
2.
3.
4.
Weight the mass of an empty crucible.
Add barium chloride crystals and reweigh.
Heat the crucible gently (so the barium chloride won’t decompose), and reweigh.
Put masses into a table and calculate the formula from there.
Mass of Empty Crucible
Mass of tube + Crystals (Before)
Mass of tube + Anhydrous Crystals
(After)
Mass of BaCl2
Mass of Water
30.00g
32.44g
32.08g
Combining Masses
Number of moles
=
Ratio of Moles
Empirical Formula
2.08g
0.36g
BaCl2
2.08
2.08/208
0.01
H2O
0.36
0.36/18
0.02
1:2
BaCl2•2H2O
Calculating Percentage Yield
Most of the time, when you do carry out a chemical reaction, you get less than you expect. The rest of it has been lost
in some way perhaps due to spillages or losses when chemicals are transferred.
Suppose you work out that 10g of A will give 500g of the product, but you only get 400g…?
The percentage yield is (400/500) x 100 = 80%
A general formula would be: (mass produced/expected mass to be produced) x 100
Endothermic and Exothermic Reactions

∆H represents the molar enthalpy change for
exothermic and endothermic reactions
Endothermic
Exothermic
Heat energy is
taken in
Breaking of bonds
Heat energy is
given out
Making of bonds
∆H = + N kJ mol-1
∆H = - N kJ mol-1
Energy Calculations
The general formula:
Bonds of all the reactants – Bonds of all the products = Energy change
Example: Methane reacts with chlorine to produce chloromethane and hydrogen chloride. The equation:
CH4 + Cl2  CH3Cl + HCl
You would be given a table with the bonds and the energy
required to break/bond them:
Bond
Energy (kJ mol -1)
C-H
413
C - Cl
346
H - Cl
432
Cl - Cl
243
Reactants:
4 C – H bonds (CH4) = 4 x 413 = 1652kJ
1 Cl – Cl bond (Cl2) = 1 x 243 = 243 kJ
Products:
Total: 1652 + 243 = 1895 kJ
3 C – H bonds = 3 x 413 = 1236 kJ
1 C – Cl bond = 1 x 346 = 346 kJ
1 H – Cl (HCl) = 1 x 432 = 432 kJ Total: 2017 kJ
(Carbon can form 4 bonds. In this case, 3 of them bonds with 3 hydrogen and the last one bonds with chlorine)
Energy Change = 1895 – 2017 = -122 kJ
the reaction is exothermic
Describing Simple Calorimetry Experiments


All these involve measuring a temperature change during the reaction.
Specific heat is the amount of heat needed to raise the temperature of 1g of a substance by 1 0C. For water,
the value is 4.18 J g-1 0C-1 (joules per gram per degree Celsius).
Heat Given Out = Mass x Specific Heat x Temperature Rise
For Neutralisation, Displacement Reactions and Dissolving
Mass of Weighing Bottle + Mg (g)
Mass of Weighing Bottle Afterwards (empty) (g)
Mass of Mg used (g)
10.810
10.687
0.123
They all follow the same method. This example
involves measuring the heat evolved (or energy)
when magnesium reacts with dilute sulphuric acid.
1.
2.
3.
4.
Initial Temp. (0C)
Final Temp. (0C)
Temperature Rise (0C)
17.4
27.5
10.1
Pour an excess of sulphuric acid into a polystyrene cup and measure the temperature of the acid.
Pour some magnesium powder into a weighing bottle and weight it.
Pour the powder into the acid and record the highest temperature.
Weigh the empty weighing bottle.
Let’s say the total mass of the solution and Mg is 50g
Heat evolved when 0.123g of Mg reacts = 50 x 4.18 x 10.1J = 2111J = 2.111kJ
To find out the heat evolved when 1 mole of Mg reacts (Mg = 24g):
(2.111/0.123g) x 24 = 412 kJ
The temperature rose, meaning the reaction is exothermic so:
Mg(s) + H2SO4(aq)  MgSO4(aq) + H2(g)
∆H = -412 kJ mol-1
This is actually smaller than the accepted value, which is around -417 kJ mol-1. A reason for this could be that heat
was lost too quickly. Using a mercury thermometer may give better results.
Combustion
1.
2.
3.
4.
Put 100cm3 of water into a conical flask and record the
temperature.
Fill the spirit burner with alcohol (let’s say ethanol) and weight.
Light the spirit burner and record the temperature of water until
there is say, a 400C increase.
Reweigh the spirit burner.
Volume of water (cm3)
Mass of water being heated (g)
100
100
Mass of burner before (g)
Mass of burner after (g)
37.355
36.575
Mass of ethanol burnt (g)
Original temp. of water (0C)
0.780
21.5
Final temp. of water (0C)
Water temperature increase (0C)
62.8
41.3
Heat gained = 100 x 4.18 x 41.3 = 17260 J = 17.26 kJ
Ethanol is C2H5OH
One mole of ethanol = 46g
Amount of heat produced from 1 mole of ethanol = (17.26/0.780) x 46 = 1020 kJ
KEY POINTS FOR THIS UNIT





1 mole is the Avogadro constant number of particles
The molar volume is 24 dm3 or 24 000 cm3
Number of moles = mass / RFM
1 dm3 = 1000 cm3
Mol dm-3 = mol per 1000 cm3


G dm-3 = grams per 1000 cm 3
Energy/Heat = mass x specific heat x temperature rise
Unit 6: Chemistry in Society
Electrolysis
RECAP! Electrolysis is a chemical change caused by passing an electric current through a compound which is either
molten or in solution. An electric current (in chemistry terms), is a flow of electrons or ions. An electrolyte is a
substance that undergoes electrolysis. Electrolytes all contain ions. Ionic compounds, for example, are electrolytes.
Electrolytes can only undergo electrolysis when molten or in a solution, where the ions are free to move. Covalent
compounds are not electrolytes because they don’t contain ions.
Electrolysis can form new substances when ionic compounds conduct
electricity. It is set up as so:



The electrodes are usually made of carbon because it is
fairly un-reactive.
The positive electrode is called the anode.
The negative electrode is called the cathode.
A simple example of electrolysis involves molten (melted) lead (II)
bromide:
So what happens?
Molten lead is found at the bottom of the
cathode.
Bromine gas comes out of the anode.
When the power supply is switched off, no more
bubbles are produced and everything else
stops.
What the hell happened?

Since the lead (II) bromide (PbBr2) is molten, its
ions are free to move around.

The bromide ions are attracted to the positive
electrode. The extra electron which makes the
bromide ion negatively charged is deposited into
the anode, thus, turning them back into neutral
bromine atoms. These then covalently bond to join bromine atoms (i.e. bromine gas).

On the other hand, the lead ions gain back to electrons (it has a 2 + charge) and become normal lead atoms.
These fall to the bottom of the container as molten lead.
The half-equation at the anode would be:
2Br-  Br2 + 2eWhat it basically means is that, two bromine ions are formed when
one bromine molecule receives two electrons (to fill its shell). This is
the format for all anions (ions that are negatively charged).
The half equation at
the cathode would
be:
Pb2+ + 2e-  Pb
All cation half-equations are of this form. Half-equations basically
show the gaining and losing of electrons. This basically means that
the lead (II) ions get two electrons to become a neutral lead atom.
With molten substances, the metal will be produced at the cathode
and whatever it’s bonded to will be produced at the anode.
When you electrolyse aqueous solutions (not molten salts), things are much different because you have to consider
the water molecules too. Water is a weak electrolyte but it can ionise to form hydrogen and hydroxide ions.

If the metal is more reactive than hydrogen, then
hydrogen ions from water is discharged instead. These
pair up to form hydrogen gas that escapes as bubbles.

If the metal is below hydrogen, you get the metal
produced.

If you have solutions of halides (chlorides, bromides or
iodides), you get the halogen (chlorine, bromine or
iodine) produced.

With other negative ions such as sulphates, oxygen
would be produced.
The electrolysis of sodium chloride solution (brine) does not give sodium and chlorine! Here’s the electrolysis:

Sodium is higher than hydrogen in the reactivity series, so
hydrogen ions from the water in the sodium chloride solution
is discharged instead at the cathode.
2H+ + 2e-  H2

Chloride ions give up one electron each (chloride ion = 1charge) and become chlorine atoms. These covalently bond
to form chlorine gas and bubbles out of the solution at the
anode.
2Cl-  Cl2 + 2e-

When all the chlorine has been removed from the solution,
only hydroxide (OH-) ions and sodium (Na+) ions are left, as
well as some water. These combine to form sodium
hydroxide solution (NaOH).
The electrolysis of sodium chloride solution is used to manufacture sodium hydroxide solution. The process is slightly
different – it is electrolysed in a diaphragm cell:
The products are kept separated by the diaphragm. If the
chlorine produced were to react to hydrogen, it would cause an
explosion on exposure to sunlight or heat to give hydrogen
chloride. Furthermore, if the chlorine were to react with the
sodium hydroxide solution formed, it would form bleach. Uses of
sodium hydroxide include:



Making bleach
Making soap
Making paper – NaOH breaks the wood down
Uses of chlorine include:

Sterilising water

Making hydrochloric acid

Making bleach
And… the electrolysis of copper sulphate solution:

Copper is lower
than hydrogen and
therefore, a coat of it forms at the cathode.
Cu2+ + 2e-  Cu

Oxygen gas is discharged from the hydroxide ions in the water
because the sulphate ions are more stable.
4OH-  2H2O + O2 + 4eIf you electrolyse the solution for longer, something else happens. The
hydrogen ions are being discharged and remains in the solution. Similarly,
sulphate ions are being discharged either. As a result, the solution turns into sulphuric acid (H 2SO4) and it begins
electrolysing:

Sulphate ions are being discharged from the acid
so oxygen is discharged from the hydroxide ions
instead.
4OH-  2H2O + O2 + 4e-

There are only hydrogen ions arriving at the cathode so they discharge as hydrogen gas.
2H++ 2e-  H2
Common Exam Question: Why is twice as much hydrogen produced than oxygen? For every four electrons that flow
around the circuit, one molecule of oxygen and two molecules of hydrogen are produced.
Electrolysis Calculations
Back to moles! Here are some things to know:



One faraday means one mole of electrons passing around the circuit.
One faraday = 96000 coulombs.
Charge (coulombs) = Current (amps) x Time (seconds)
Example: What mass of copper is deposited on the cathode during the electrolysis of copper (II) sulphate solution if
0.15A flows for 10mins?
The electrode equation is:
Cu2+ + 2e-  Cu
Calculate the coulombs involved:
10mins x 60 = 600 seconds
Charge = 0.15 x 600 = 90 coulombs
The equation says that 1 mole of copper ions + 2 moles of electrons give 1 mole of copper atoms
1 mole of electrons = 96000 coulombs
2 moles of electrons = 192 000 coulombs
2 moles of electrons (192 000 coulombs) give 1 mole of copper (RFM = 64g), so 90 coulombs give:
(90/192 000)
x 64g = 0.03g
(Coulombs worked out/coulombs of electrons) x RFM of element = Mass of element deposited
When involving gases…
Example: During the electrolysis of dilute sulphuric acid, hydrogen is released at the cathode and oxygen at the
anode. Calculate the volumes of hydrogen and oxygen produced if 1.0A flows for 20mins
The electrode equations are:
2H++ 2e-  H2
4OH-  2H2O + O2 + 4eAssume the molar volume of gas to be 24000 cm3
For hydrogen:
2H++ 2e-  H2
2 hydrogen ions + 2 moles of electrons give 1 mole of hydrogen molecule
20mins x 60 = 1200 seconds
1200 x 1A = 1200 coulombs
2 x 96 000 coulombs (2 moles of electrons) = 192 000 coulombs
1 mole of hydrogen gas = 24000 cm3
192 000 coulombs give 1 mole of hydrogen gas (or 24000 cm3 of hydrogen gas)
So 1200 coulombs give: (1200/192 000)
x 24000 = 150 cm3 of hydrogen produced
For oxygen:
4OH-  2H2O + O2 + 4e20mins x 60 = 1200 seconds
1200 x 1A = 1200 coulombs
4 moles of hydroxide ions give 2 moles of water + 1 mole of oxygen gas + 4 moles of electrons (4 x 96000 = 384 000
coulombs)
1 mole of oxygen gas = 24000 cm3
384 000 coulombs give 1 mole of oxygen gas (or 24000 cm3 of oxygen gas)
So 1200 coulombs give: (1200/384 000)
x 24000 = 75 cm3 of oxygen produced
The equation:
(Calculated coulombs/moles of electrons) x (mole of gas x 24000) = amount of gas produced in cm 3
Reversible Reactions and Dynamic Equilibria
Some reactions are reversible. Reversible reactions are indicated by the symbol
Some examples of reversible reactions include:

Copper (II) Sulphate Crystals
Heating the blue hydrated copper (II) sulphate crystals
causes them to lose their water of crystallisation, making
them turn from blue to white – the white copper (II)
sulphate crystals are described as ‘anhydrous’ meaning
‘without water’:
CuSO4•5H2O  CuSO4 + 5H2O
However, this reaction can be reversed by simply adding water to the
crystals. The crystals will become hydrated again:
CuSO4 + 5H2O  CuSO4•5H2O

Heating Ammonium Chloride
When ammonium chloride is heated in a test tube, the white crystals decompose into
hydrogen chloride gas and ammonia gas. These flow upwards and recombine again
further up the test tube:
NH4Cl  HCl + NH3
This later recombines:
HCl + NH3  NH4Cl
Introducing Dynamic Equilibria
Things change when reversible reactions are carried out under ‘closed’ conditions – meaning no substances are
added to the reaction mixture and no substances can escape from it. Heat however, can be given off or absorbed.
In a reversible reaction, you have the forward reaction (the reaction going from left to right) and the back reaction (the
opposite of the forward reaction) happening at the same time. Both rates of reactions will become equal and this point
is the dynamic equilibrium. It is dynamic in a sense that the reactions are still continuing, and equilibrium in a sense
that the total amounts of the various things present are now constant. In other words:
A + 2B
C+D
When you have a reaction like the above, A + 2B (forward reaction) is reacting to produce C + D (back reaction). At
the same time, C + D is reacting to produce A +2B. In the end, you have equal amounts of products and reactants.
Another way to think of is, is to imagining walking down an elevator that goes up, making sure you’re walking at the
same speed as the elevator. You would be going down, but everytime you take one step down, the elevator goes one
step up. In the end, you remain where you are.
So how would you produce more of substance C in a reversible reaction such as the above? You can do this by
altering the position of the equilibrium by either:




Changing the pressure
Changing the temperature
Increasing/Decreasing the concentrations of substances present
Adding a catalyst
If a dynamic equilibrium is disturbed by changing the conditions, the reaction moves to counteract the change.
In other words, the reaction will either go more towards the ‘forward’ direction or the ‘back’ direction in an attempt to
‘adapt’ to the conditions.
A + 2B
C+D
Changing the Concentration
What happens when more of A is added? If you add more A, the reaction will want to remove it. This can only be
done by reacting more A to 2B, and in the end, this gives us more C and D. The conditions, in this case, favour the
forward reaction.
Changing the Pressure
When you increase the pressure, you bring molecules closer together. Increasing the pressure will always help the
reaction go in the direction which produces the smaller number of molecules. In this case, we have 3 molecules on
the left (one A and 2 Bs), whereas, we only have 2 molecules on the right (one C and one D). The reaction can only
reduce the pressure by producing few molecules. This can only be done by producing more C and D.
Changing the Temperature
Suppose the forward reaction was exothermic
A + 2B
C+D
∆H = -100 kJ mol-1
This would mean that the back reaction would be endothermic by the same amount (-100 kJ mol-1).
Suppose the temperature was decreased, the reaction would respond by increasing the temperature back up again.
This can only be done by producing more C and D because the forward reaction is exothermic. Increasing the
temperature will of course, have the opposite effect.
Adding a Catalyst
Adding a catalyst speeds up the forward and back reactions by the same proportion. This means that there is no
change in the position of the equilibrium.
Methods of Extraction
The extraction of metals depend alot on its position in the reactivity series. Costs are also factors to take into account.
For metals up to zinc, the cheapest method is usually heating the ore with carbon or carbon monoxide to reduce it.
For metals more reactive than zinc, electrolysis is usually used.
The Extraction of Aluminium
Aluminium is extracted from an ore called bauxite, which is impure aluminium oxide. Aluminium ions are attracted to
the cathode and are reduced to aluminium:
Al3+ + 3e-  Al
Oxide ions are attracted to the anode and lose electrons to form oxygen gas:
2O2-  O2 + 4e-
There are a few
keep in mind:
things to

Melting aluminium oxide requires extremely high temperatures. Instead, it is dissolved in an aluminium
compound called cryolite.

Because of high temperatures, the carbon anodes will burn with oxygen to form carbon dioxide. This means
the anodes have to be replaced regularly.

The cost of electricity is also a major factor – the cell has currents up to 100 000A so it is expensive.
Uses of aluminium include:
Use
To make aircraft carriers
Saucepans
Property of Aluminium That Makes This Useful

Resists corrosion due to its aluminium oxide coat

Has low density

Strong

Aluminium has a shiny appearance

It is a good conductor of heat

It resists corrosion

It has a low density
The Iron Blast Furnace
Iron is extracted from an ore that contains iron (III) oxide called haematite.
Coke is impure carbon. It burns to form carbon dioxide. This
is a strongly exothermic reaction.
C + O2  CO2
At high temperatures, the carbon dioxide is reduced by
more carbon to give carbon monoxide.
C + CO2  2CO
Carbon monoxide is the main reducing agent:
Fe2O3 + 3CO  2Fe + 3CO2
Carbon may also reduce the iron (III) oxide:
Fe2O3 + 3C  2Fe + 3CO
The heat of the furnace causes the limestone to thermally
decompose to form calcium oxide and carbon dioxide:
CaCO3  CaO + CO2
The calcium oxide reacts with silicon dioxide (one of the
impurities found in haematite) to form calcium silicate, which
melts and trickles to the bottom of the furnace as molten
slag:
SiO2 + CaO  CaSiO3
Uses of Iron (you don’t really need to learn this)
Types of Iron
Iron Mixed With
Some Uses
Wrought Iron
Pure Iron
Decorative work such as gates and railings
Mild Steel
0.25% Carbon
Nails, car bodies, ship building
High – Carbon Steel
0.25 – 1.5% Carbon
Cutting tools
Cast Iron
About 4% Carbon
Manhole covers, guttering, engine blocks
Stainless Steel
Chromium and Nickel
Cutlery, cooking utensils
Preventing the Rusting of Iron (in which iron oxidises into iron oxide Fe2O3)
Using Barriers
Keep water/oxygen away from the iron by painting, coating with oil…etc.
Alloying the Iron
Such as allowing it with chromium and nickel to produce stainless steel
Using Sacrificial Anodes
Galvanising iron by coating it with a layer of zinc. Zinc is more reactive than iron and
will corrode instead. During the process it loses electrons to form ions. These
electrons flow into the iron so any iron atom which has lost electrons immediately
regains them. These means even if the zinc is scratched, the iron won’t rust.
The Haber Process
The Haber process is used to make
ammonia NH3
Uses of ammonia include:



Making fertilisers
Making nitric acid
Making nylon
The equation for the reaction is:
N2 + 3H2



∆H = -92 kJ mol-1
2NH3
The forward reaction would be favoured by a low temperature because the forward reaction is exothermic
(so lowering the temperature would cause the reaction to make more NH 3 to heat things up abit more).
450ºC isn’t a low temperature. It is however, a compromise temperature, because if the temperature was
made to be low, the reaction would be so slow that it would take a very long time to produce much ammonia.
Pressure is also another compromised. Because the forward reaction has less molecules than the back
reaction (2 molecules of NH3 as opposed to 1 N2 and 3 H2 molecules), the forward reaction would be
favoured by a high pressure. 200 atm is high, but anything higher would be extremely expensive.
The iron catalyst speeds the reaction up but has no effect on the equilibrium. However, if the catalyst wasn’t
used, the reaction would be too slow.
The Contact Process
1.
Burn sulphur in air to form sulphur dioxide SO2
S + O2  SO2
2. Use an excess of air to react sulphur dioxide to more oxygen to form
sulphur trioxide
2SO2 + O2
2SO3
∆H = -196 kJ mol-1
3. Reacting sulphur trioxide with water will give an uncontrollable fog of
concentrated sulphuric acid. Instead, sulphur trioxide is absorbed in concentrated
sulphuric acid to give fuming sulphuric acid (oleum):
SO3 + H2SO4  H2S2O7
4. This is converted into twice as much concentrated sulphuric acid by
careful addition of water:
H2S2O7 + H2O  2H2SO4
The reversible reaction here is:
2SO2 + O2
2SO3
∆H = -196 kJ mol-1

Because the forward reaction is exothermic, a low temperature has to be used. Again, if a low temperature
is used, the rate of reaction would be too slow, so 450ºC is a compromise.

As for the pressure, a low pressure is needed because the forward reaction contains fewer molecules than
the back reaction.

The catalyst, vanadium (V) oxide (V2O5) speeds up the rate of reaction but has no effect on the equilibrium.
Again, without a catalyst, the rate of reaction would be extremely slow.
Uses of sulphuric acid include:

Making fertilisers – including ammonium sulphate and other substances

Detergents – including hand soaps and shampoos

Paint Manufacture – used to extract titanium dioxide from titanium ores
A Few Extra Bits and Pieces
Indicator Solutions
Note: Methyl orange is orange in neutral
Phenolphthalein
colourless in acid
pink in alkali
solutions, however, these two indicator solutions
are both yes no indicators – meaning, if there is a
Methyl orange
red in acid
yellow in alkali
reaction, the methyl orange (for example) would
turn from red to yellow (or vice versa depending on whether the solution it has been dropped into is acidic or
alkaline). It will not turn orange.
The acidity or alkalinity
is measured in pH.
indicator cam be used to measure the approximate pH of a solution.


Acids are sources of H+ ions
Alkalis are sources of OH- ions
for more info, go to page 10
Reactions Between Metals, Metal Compounds and Acids
Basic things to know:


Hydrochloric acid reacts with metals and metal compounds to form a metal chloride
Sulphuric acid reacts with metals and metal compounds to form a metal sulphate
And some equations:
Metal + Acid  Salt + Hydrogen
Mg + 2HCl  MgCl2 + H2
Mg + H2SO4  MgSO4 + H2
Metal Oxide + Acid  Salt + Water
MgO + 2HCl  MgCl2 + H2O
MgO + H2SO4  MgSO4 + H2O
Metal Carbonate + Acid  Salt + Water + Carbon Dioxide
MgCO3 + 2HCl  MgCl2 + H2O + CO2
Mg CO3 + H2SO4  MgSO4 + H2O + CO2
Preparing Soluble Salts Using Titration
The exact amount of acid needed to neutralise an alkali can be found
by titration. This technique can be used to make pure crystals of
a soluble salt (one that dissolves in water).
In the example below, an acid and an alkali react to
make sodium chloride.
of something
Universal
1.
The burette is filled with hydrochloric acid.
2.
A known quantity of alkali (say 50 cm3 sodium hydroxide) is released from a pipette into the conical flask.
The tap on the burette is turned open to allow the acid to be added drop by drop into the alkali.
3.
The alkali contains an indicator (phenolphthalein).
4.
When enough acid has been added to neutralise the alkali, the indicator changes from pink to colourless.
This is the end point of the titration.
5.
The titration can be repeated using the same amounts of acid and alkali but without the indicator.
6.
Pure salt crystals which are free from indicator can then be crystallised from the neutral solution.
Precipitation Reactions
The process of making a solid come from a solution is called precipitation. The solid itself is called a precipitate.
An insoluble salt (one that doesn't dissolve) can be made by reacting the appropriate soluble salt with
an acid or alkali or another salt.
You are normally asked to prepare a solid from two soluble solutions – so know your solubility rules:





All nitrates are soluble.
All sodium, potassium and ammonium compounds are soluble.
Most carbonates and hydroxides are insoluble except for sodium,
potassium and ammonium.
All sulphates are soluble except barium and lead(II) sulphate.
All chlorides are soluble except lead(II) and silver chloride.
Example: Prepare Silver Chloride
You’ll need:


A soluble silver salt – what about silver nitrate?
A soluble chloride – like magnesium chloride?
You can make it up really. So the equation:
AgNO3 (aq) + MgCl2(aq)  AgCl(s) + Mg(NO3)2(aq)
The Reactivity Series
Metals are arranged based on their reactions in the reactivity series.
They can be deduced by using displacement reactions, in which a less
reactive metal is pushed out of its compound by a more reactive metal. For
example, the reaction between magnesium and copper (II) oxide:
Mg + CuO  MgO + Cu
Displacement reactions are examples of redox reactions, in which oxidation
and reduction occurs in the same reaction.
The reducing agent is a substance that reduces something else. In this
case, it’s magnesium. The oxidising agent is a substance that oxidises
something else. In this case, it’s copper.
Remember OILRIG: Oxidation is gain (of electrons); Reduction is loss (of
electrons)
Let’s look at this in terms of an ionic equation:
STATE SYMBOLS
(s)
solid
(l)
liquid
(g)
gas
(aq)
aqueous solution
(dissolved in water)
Mg + Cu2+ + O2-  Mg2+ + O2- + Cu
This basically looks at what turns into ions and what does not. The magnesium turns into a positive ion because it
loses two electrons to bond with oxygen. The copper ion, because it’s displaced, regains the two electrons it lost to
the oxygen.
Notice in this equation, the oxygen ion does not change? It remains an ion. In this case, the oxygen ion is the
spectator ion. Spectator ions aren’t included in the ionic equation, so the proper ionic equation of this displacement
reaction should be:
Mg + Cu2+  Mg2+ + Cu
As mentioned earlier, the magnesium loses two electrons to bond with oxygen and copper gains two electrons. This
can be written as half-equations:
Mg  Mg2+ + 2eCu2+ + 2e-  Cu
Reactions with substances are different depending on its reactivity:
Metals
Potassium, Sodium,
Lithium
Calcium
Reaction with Dilute Acids
Too reactive to add safely to acids
Zinc or Iron
Reaction with Water
Very vigorous, produces hydroxides and
hydrogen gas
Reacts gently and produces the same products
as above
Reacts with steam to produce magnesium oxide
and hydrogen
Reacts slowly and forms an oxide and hydrogen
Anything Below
No reaction
No reaction
Magnesium
Can be added to very dilute acids – but it’s
going to be violent!
Reacts vigorously – strongly exothermic –
forms a salt and hydrogen
Reacts slowly
Rusting of Iron
This requires two things:


Oxygen
Water
Preventing the Rusting of Iron (in which iron oxidises into iron oxide Fe 2O3)
Using Barriers
Keep water/oxygen away from the iron by painting, coating with oil…etc.
Alloying the Iron
Such as allowing it with chromium and nickel to produce stainless steel
Using Sacrificial Anodes
Galvanising iron by coating it with a layer of zinc. Zinc is more reactive than iron and
will corrode instead. During the process it loses electrons to form ions. These
electrons flow into the iron so any iron atom which has lost electrons immediately
regains them. These means even if the zinc is scratched, the iron won’t rust.
More on sacrificial anodes – if the iron has already rusted, it can still be displaced by the more reactive zinc:
Fe2O3 + 3Zn 2Fe + 3ZnO
And this is the end.
Good luck!