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Contents Page
The creation of this condensed textbook used the Edexcel GCSE
Chemistry specification as a guide to cover all the necessary
specification points, however, the textbook also can be used for other
exam boards. As there is plenty of content overlap that Edexcel has
with AQA (and because AQA is the most common exam board in the
UK) Ken has put a content page for students following the Edexcel
specification and the AQA specification.
For Edexcel GCSE Chemistry
Unit 0 – Formulae, Equations, and Hazards………………….………....4
Unit 1 – Key Concepts in Chemistry…………………………………….7
Unit 2 – States of Matter and Mixtures….…………………………….23
Unit 3 – Chemical Changes, Acids, and Electrolytic Processes………31
Unit 4 – Extract Metals and Equilibria………….……………………...41
Unit 5 – Separate Chemistry I: Transition Metals, Alloys and Corrosion,
Quantitative Analysis, Dynamic Equilibria, Chemical Cells, and Fuel
Cells……………………………………………………………………...48
Unit 6 – Groups in the Periodic Table…………..……………………..67
Unit 7 – Rates of Reaction and Energy Changes………..………….…73
Unit 8 – Fuels, Earth and Atmospheric Sciences……………………...83
Unit 9 – Separate Chemistry II: Qualitative analysis (testing for ions),
Hydrocarbons, Polymers, Alcohols and Carboxylic Acids, Bulk and
Surface Properties of Matter Including Nanoparticles………………..90
For AQA GCSE Chemistry
Topic 1 (Atomic structure and the periodic table) → Unit 1 and Unit 6
Topic 2 (Bonding, structure, and properties of matter) → Unit 1,
Unit 2, Unit 5, Unit 9
Topic 3 (Quantitative chemistry) → Unit 1, Unit 3, Unit 5
Topic 4 (Chemical changes) → Unit 3, Unit 4, Unit 5
Topic 5 (Energy changes) → Unit 1, Unit 5,
Topic 6 (The rate and extent of chemical change) → Unit 1, Unit 4,
Unit 5, Unit 7
Topic 7 (Organic Chemistry) → Unit 1, Unit 8, Unit 9
Topic 8 (Chemical analysis) → Unit 2, Unit 7, Unit 9,
Topic 9 (Chemistry of the atmosphere) → Unit 2, Unit 4, Unit 5, Unit 8
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Foreword
Firstly, I would like to thank Faran Ahmad, James San, Julie Maher, Nirmal Mullagiri (Mary),
Mia Chadwick, Tyisha Rebelo, Asad Chaudhry (for the astounding front cover design),
Daniyaal Anawar, Jammal Othman, Archie Thompson, and Frank G for giving up part of your
summer 2022 to create this resource to students across the UK and internationally.
As you read the words across this page, I hope you realise that people who are 15-17 years
old created this textbook and you are no different. If you are young, I urge you to steal the
time you have now to make the most of your opportunities. Whether you want to be a
doctor, engineer, teacher, lawyer, content creator, rapper, singer, linguist, or politician;
whether you want to be happy, excited, thrilled own two dogs, live in a nice house or have
the money go on a holiday every 8 weeks, I believe in you. You, already actively seeking out
this resource provides evidence that there’s a direction that you want to go in. I wish you
the best!
Good luck,
Ken Tu – Director of this project
U6 2022-2023
I wish you all the best in your GCSEs. By getting this book, it’s clear that you want to succeed.
All you need to do now is put in the effort. Keep going when it feels tough. Make sure that all
the effort you put into your studies, pays off at the end. Go that extra mile to make sure you
don’t regret anything. You’ve got this! – Julie Maher
If you're reading this, you care about your education. That will be key when it comes to
succeeding in your GCSEs. It will take some time to adjust to Year 11, no doubt about that,
but once you get into the rhythm of things, you'll feel more confident than you expect. I hope
this book helps you get into that rhythm! – James San
Seeing as you picked up this textbook, it's clear that you aim to achieve the highest results
you can possibly get for your exams. We encourage you to do your best and aim your highest
in hopes of reaching that goal; you made it this far, after all. Use this guide as much as
possible alongside past papers and exam questions and you're sure to pass with flying
colours. – Nirmal (Mary) Mullagiri
I wish you all the best for the coming year and your GCSEs, the fact that you are currently
reading through this textbook shows that you care about your education and your grades. As
you might already know, year 11 is not a walk in the park. Yes, it will be stressful and giving
up might look a lot easier but really putting in the work will pay off on results day. You’ve
come so far already; you’ll just need to push a little more. Now take a deep breath and go for
it. You’ve got this! -Tyisha Rebelo
Well done for getting this far in your academic career and maintaining a desire to learn and
discover new things along your journey. The fact that you’ve applied yourself and been eager
to pick up this textbook says a lot about you as a student alone. I wish you well with your
studies and good luck for this year! – Mia Chadwick
3
Edexcel GCSE Chemistry: Unit 0 - Formulae, Equations, and Hazards
written by Daniyaal Anawar and Ken Tu
General Word Equations
Below are most of the general word questions from common reactions. This includes general
reactants (e.g. metal and oxygen) and the general product formed (e.g. metal oxide).
Acid and alkali:
acid + alkali → salt + water
E.g. hydrochloric acid + sodium chloride → sodium chloride + water.
Acid and metal oxide:
acid + metal oxide → salt + water
E.g. hydrochloric acid + aluminium oxide → aluminium chloride + water.
Acid and metal carbonate:
acid + metal carbonate → salt + water + carbon dioxide
E.g. hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide.
Acid and metal:
acid + metal → salt + hydrogen
E.g. hydrochloric acid + sodium→ sodium chloride + hydrogen.
Metal and oxygen:
metal + oxygen → metal oxide
E.g. zinc + oxygen → zinc oxide.
Metal and sulphur:
metal + sulphur → metal sulphide
E.g. calcium + sulphur → calcium sulphide.
Metal and water:
metal + water → metal hydroxide + hydrogen
E.g. calcium + water → calcium hydroxide + hydrogen.
Complete combustion of a hydrocarbon:
hydrocarbon + oxygen → carbon dioxide + water
E.g. ethane + oxygen → carbon dioxide + water.
State Symbols
(s) - means the substance is solid.
(l) - means the substance is liquid.
(g) - means the substance is a gas.
Generally for GCSE, it is assumed that most reactions are taking place at room temperature
(25°C) so whilst yes you can have e.g. solid carbon dioxide, carbon dioxide at 25°C is a gas.
(aq) - means aqueous solution so the compound is dissolved in water e.g. NaOH(aq) is sodium
hydroxide but dissolved in water.
Example equation with state symbols: HCl (aq) + CaCO3 (s) → CaCl2 (aq) + H2O (l) + CO2
(g)
Hazards symbols
Anything that has the potential to inflict harm or damage is considered a hazard. The likelihood
that someone or something will suffer injury as a result of being exposed to the hazard is the
associated risk.
Chemical containers have symbols on them to tell you the dangers associated with those
chemicals, and understanding the symbols means that you can work and use the chemicals
safely in the lab.
4
Edexcel GCSE Chemistry: Unit 0 - Formulae, Equations, and Hazards
written by Daniyaal Anawar and Ken Tu
You need to know and be able to recall the hazard symbols, as this recently came up in the July
2022 paper and some were surprised.
Symbol
Meaning
Typical Hazard
Explosive
Possibility to explode in your face I guess?
Flammable
Could start a fire.
Oxidising
Can intensify a fire and increase flammable
range.
Corrosive
Damage clothing and also your skin.
Acute Toxicity
If inhaled/consumed/contacts your skin it
could be fatal.
Hazardous
environment
to
the Damages the ecosystem.
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Edexcel GCSE Chemistry: Unit 0 - Formulae, Equations, and Hazards
written by Daniyaal Anawar and Ken Tu
Health hazard/Hazardous Chemicals could cause health defects e.g.
to the ozone layer
cancer, gene mutation etc.
Serious health hazard
Contact with this substance could have long
term effects on a person’s health.
Gas under pressure
Gas could explode if heated and because it
is pressurised.
Risk Assessments
In chemistry experiments, there are many risks, mainly associated with:
1. The equipment.
2. The chemicals.
When planning an experiment, you need to identify all of the hazards and the risk from each
hazard, including:
1. How likely it could happen?
2. How serious would it be if it did happen?
Creating a plan to reduce such risks is called a risk assessment.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Chemical equations
Chemical equations are used in chemistry to depict what occurs to the reactants in a chemical
reaction. They describe the atoms involved and the changes that occur to the substances
during a reaction. The two main types of chemical equation that you will encounter are:
1. Word equations - simple, less detailed.
2. Symbol equations - complex, more detailed.
Word equations
Here is an example of a word equation, highlighting the reaction between aluminium and
oxygen, forming aluminium oxide - a metal oxide:
aluminium + oxygen → aluminium oxide
The molecules on the left are the reactants as they react, and the molecules on the right are
the products as they are produced.
Symbol equations
Here is the same chemical reaction, however now as a symbol equation:
4Al + 3O2 → 2Al2O3
In chemistry, symbol equations are found to be much more useful,
as they provide us with more information regarding the reactants
and products. Furthermore, the numbers next to the symbols of the
reactants and products are extremely important, so make sure you
don’t mix them up:
1. The large number (coefficient) - how many units there are of
that molecule
2. The small number (subscript) - how many atoms of that
element are in that molecule
Balancing equations
When dealing with symbol equations, they need to be balanced - there must be the same
number of atoms on both sides. This is what the coefficient in a symbol equation is used for.
Take this same equation from before, however now unbalanced:
Al + O2 → Al2O3
The formulae for the compounds are all correct, however, the atoms on the LHS (left hand
side) and the RHS (right hand side) do not match up, and we cannot change the formulae as
this would alter the actual chemical reaction that takes place. Therefore, we must utilise the
coefficients from before. Here are the number of atoms on both sides currently:
Element
No. of atoms in reactants
No. of atoms in products
Al (aluminium)
1
2
O (oxygen)
2
3
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
From looking at the oxygen row, we can see that we can multiply the number of oxygen
atoms in the reactants by 3 and the number in the products by 2 to have 6 (the LCM: Lowest
Common Multiple) each - 2 ✕ 3 = 6, 3 ✕ 2=6. This would result in this following equation
and table:
Al + 3O2 → 2Al2O3
Element
No. of atoms in reactants
No. of atoms in products
Al (aluminium)
1
4
O (oxygen)
6
6
Now, to finish balancing the equation, we can multiply the number of aluminium atoms in the
reactants by 4 to have the final balanced equation and balanced table:
4Al + 3O2 → 2Al2O3
Element
No. of atoms in reactants
No. of atoms in products
Al (aluminium)
4
4
O (oxygen)
6
6
Simple compounds
Compound
Water
Carbon dioxide
Chlorine
Ammonia
Hydrogen
Oxygen
Formula
H2O
CO2
Cl2
NH3
H2
O2
Polyatomic ions
Ion
Ammonium
Nitrate
Sulfate
Hydroxide
Carbonate
Formula
NH4+
NO3-
SO42-
OH-
CO32-
Ionic equations
Ionic equations are used in chemistry to show ionic substances reacting in a solution, with
only the reacting particles included. They can often be quite difficult, so make sure you
practice them a lot. To write an ionic equation:
1. Write out the balanced symbol equation - having it balanced is key.
2. Split up all aqueous molecules so that you write out the equation with all aqueous
ions separated.
3. Cancel out any aqueous ions that are on both sides of the equation.
4. Ensure the charge is the same on both sides.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Take this worked through example:
1. Write equation - AgNO3 (aq) + NaCl(aq) → AgCl(s) + NaNO3 (aq)
2. Split up all aqueous molecules, (notice the solid isn’t split) Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq)
3.
Cancel out any ions on both sides of the equation - Ag+(aq) + Cl-(aq) → AgCl(s)
4.
Ensure the charge is the same on both sides - LHS (1+-1 = 0), RHS (0=0)
5.
Done!
Atomic structure: development of the atomic structure
John Dalton’s (a Chemist in the 18th-19th Century) initial three-part atomic theory
determined that: All substances were made of atoms - small particles which couldn’t be
created, divided or destroyed.
1. Atoms of the same element are identical and, hence, atoms of different elements are
different.
2. Different atoms join together to produce new substances.
JJ Thompson (a Chemist in the 19th-20th Century) then used a cathode-ray tube in 1897 to
conduct an experiment, testing Dalton’s hypothesis. This experiment determined that:
1. Atoms could be divided into smaller particles, therefore
providing evidence to disprove Dalton’s atomic theory.
2. Particles must have a negative charge as the beam moved away
from the negatively charged plate towards the positively charged
plate.
As a result of this, JJ Thompson created the plum-pudding model,
describing atoms as being spheres of positive charge with negative
electrons embedded within them.
Ernest Rutherford then conducted the alpha particle gold foil
experiment in 1909 wherein he shot a beam of positively charged alpha particles at a sheet of
gold foil. This experiment determined that:
1. Most of the atom is empty space as most of the particles went in a straight line
through the atom.
2. The atom has a nucleus which is positively charged as some of the particles were
deflected to the sides.
3. The mass of the atom is concentrated at the nucleus as some particles are deflected
straight back.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
As a result of this, Ernest Rutherford produced the nuclear model of the atom in which:
• The positively charged nucleus is at the centre of the atom.
• The majority of the atom’s mass is concentrated in the nucleus.
• Negative electrons orbit the nucleus.
However, this model caused uncertainty as scientists questioned why the atom didn’t collapse
due to the attraction between the positive nucleus and negative electrons orbiting it.
Niels Bohr contributed to the model’s development after experiments in 1913 in which he
determined that:
1. Electrons orbit the nucleus in fixed shells (or orbitals) a set distance from the
nucleus.
2. Each shell has different energy levels and the shells furthest from the nucleus have
the most energy.
This therefore produced an updated version of the nuclear model.
Following this, in 1918, Ernest Rutherford once again experimented with the atomic model
and he determined that the nucleus could be divided into smaller particles known as protons
which are positively charged.
Finally, in 1932, James Chadwick conducted experiments to prove the existence of neutral
particles in the nucleus called neutrons.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Atomic structure
The modern atomic structure consists of:
•
•
A positively charged nucleus - where the majority of the mass of the atom is
concentrated - in the centre of the atom containing protons and neutrons.
Electrons orbiting the nucleus in shells.
Subatomic Particles
Charge
Mass
Protons
+1
1
Neutrons
0
1
Electrons
-1
1/1840 (You can say very small)
Atoms contain an equal number of protons and electrons. This is because atoms are neutral
and, as protons and electrons have opposing charges, these must be balanced for neutrality to
occur. For example, if an atom has 36 electrons, it must also have 36 protons to counteract
the earlier negative charge. If the atom only had 35 protons, the atom would have a greater
negative charge and, therefore, it would not be neutral.
The nucleus
Atoms are incredibly small (Average size in my
books though! Hehe) and this means their radius is
only around 1 x 10-10 metres! The nucleus, however,
is even smaller! For perspective, the radius of a
nucleus is roughly 10,000 times smaller than that of
an atom and therefore, instead of being distributed
evenly throughout the entire atom, the majority of
the atom’s mass is concentrated within the nucleus.
Term
Meaning
Atomic Number
The number of protons in the nucleus of an atom
Mass Number
The sum of the protons and neutrons within the
nucleus of an atom
What are isotopes?
Isotopes are atoms of the same element with the same number of protons in the nucleus, but
a different number of neutrons. Despite the alternating number of neutrons in the nucleus,
isotopes display the same chemical characteristics as they all have the same number of
electrons in their outer shell and electrons determine how violently chemicals react.
Therefore, the only difference between isotopes is the neutrons in the nucleus (and hence also
the mass of the atom).
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Even if atoms have varying numbers of neutrons in their nucleus, they are still atoms of the
same element. This is because an atom’s element is determined by the number of protons in
its nucleus. Therefore, even if an atom has 40 neutrons instead of the usual 36, it’ll still
remain the same element if the number of protons remains the same.
Relative atomic mass
Relative atomic mass, often written as A , is determined by comparing the average mass of
an element to carbon 12 which is generally used as a standard for comparison.
r
However, due to the existence of isotopes, the relative atomic masses of some elements are
not whole numbers.
1. The relative atomic mass of an element is calculated using the abundance of the
isotopes of an element and finding the average.
2. Isotopes, despite being atoms of the same element, have varying numbers of neutrons
and hence the average can end up not being a whole number.
3. As a result of this, the relative atomic mass is therefore also not a whole number
(Higher Tier Only) The formula for calculating the relative atomic mass of an element is:
Calculating The Number of Subatomic Particles in an Atom
Number of Subatomic
Particle
How to find them?
Protons
No. of Protons is the Atomic Number
Neutrons
Mass number - Atomic number
Electrons
No. of Electrons is the Atomic Number (in an atom NOT
ion)
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
The periodic table: Mendeleev’s periodic table
In 1869, Dmitri Mendeleev developed his periodic table and this soon evolved into the
modern periodic table. Mendeleev’s periodic table was arranged based on:
1. Increasing relative atomic mass
2. Properties of elements (such that elements with similar chemical properties were
placed in the same group)
Mendeleev left gaps where he thought there should be other elements and this allowed him to
predict the properties of these new elements. After the new elements were discovered, it was
realised that Mendeleev’s predictions fitted the properties of the elements very well and his
periodic table was given credibility.
However, although Mendeleev believed he had arranged elements in order of increasing
relative atomic mass, this proved not always to be true and this was due to isotopes not being
taken into account. This is because, although one isotope might make it seem like the mass of
element A is lower than element B, the actual average of all the isotope’s masses (RAM)
might be greater than the average of element B’s isotopes and therefore element A should
actually be placed first in the periodic table.
Modern periodic table
In the modern periodic table, elements
are arranged based on:
1. Increasing atomic number.
2. Element’s chemical
properties.
This arrangement means that elements
with similar properties will be placed
into the same columns (groups) as they
have the same number of electrons on
their outer shell and therefore have
similar chemical properties.
Metals form the majority of the
periodic table’s elements, and they
react to form positive ions. They are positioned on the left side of the periodic table as they
have fewer electrons in their outer shell so they can lose electrons to get a stable outer shell of
8 electrons. This allows the metals to have a stable electronic structure and satisfy its octet
(research this further, it’s not required for GCSE and will never be tested but it’s very
conceptually helpful - Ken).
Non-Metals are found on the right side of the periodic table, and they react to form negative
ions. This is because, as they are on the right side of the table and have a greater number of
electrons, it’s most efficient for them to gain electrons if they wish to get a stable electronic
structure of 8 outer shell electrons.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Electronic configuration
The electronic configuration of an element tells you how many
electrons are in each shell orbiting the nucleus.
So, for example, if oxygen has an atomic number of 8, we know that
it has 8 electrons. The first shell always has a maximum of 2
electrons and the next shells have a maximum of 8 electrons.
Therefore, given that oxygen has 8 electrons, we can infer that there
are 2 electrons in the first shell and 6 in the second shell. This leaves
us with an electronic configuration of [2, 6].
Electronic configuration can also be represented in diagram form as
shown on the right, demonstrating that the first energy shell can have
a maximum of 2 electrons and the shells after can house 8.
The electronic configuration of an element is also related to its position in the periodic table.
1. The group (column) an element is in indicates the number of electrons in the
element’s outer shell.
2. The period (row) an element is in indicates the number of shells an element has, as
well as the shell of the outermost electrons.
For example, if an element is in group 2, period 3, we can tell that the element has 2 electrons
in its outer shell and has three shells in total. From this, we can produce the electronic
configuration of [2, 8, 2] and we can also tell that the element has an atomic number of 12
(2+8+2), and that the element is a metal as it has few electrons in its outer shell. A periodic
table is also provided in the exam and we can therefore also recognise the metal, even just
with the group and period number.
Ionic bonding
An ion is an atom which has gained or lost electrons, making it a charged particle with a
positive or negative charge. Positively charged ions are known as cations and negatively
charged ions are known as anions. (I remember the CATions as the PAWsitive ion like a cat
and the anion is just, well, the other one - Ken).
Ions can be formed in reactions between metals and nonmetals, which transfer electrons
between the atoms to form ionic bonds. This can be shown through a dot and cross diagram.
The dots represent the electrons from one atom, and the cross represent the electrons from
another atom. For example, when a sodium atom reacts with a chlorine atom, one electron
from sodium transfers to the chlorine atom,
giving them both a complete outer shell. Since
the sodium atom lost 1 electron, it now has
more protons than electrons and is now a
positively charged ion (cation), with a +1
charge, as it lost one electron. But with the
chlorine atom, it gained 1 electron, so it now
has more electrons than protons, making it a
negatively charged ion (anion), with a -1
charge, as it gained one electron.
14
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
You can also use the atomic number and the mass number of simple ions to calculate:
1. The number of protons (this is simply the atomic number)
2. The number of neutrons (this is mass number - atomic number)
3. The number of electrons (Work out the number of electrons atoms of that element
would usually have, then work out how many electrons have been lost or gained)
Different ions are formed based on the group numbers of the metals and non-metals which
react during ionic bonding. For example, if a group 2 metal reacts, it loses 2 electrons to be left
with a stable outer shell of 8 electrons and this gives it a +2 charge. In terms of configuration,
it goes from [2,8,2] to [2,8]. On the other hand, if a group 7 non metal reacts, it gains an electron
to form a full outer shell of electrons and this gives it a +1 charge. In essence, it goes from [2,7]
to [2,8].
I (Mr Potential) like to use the idea of EBAY to understand charges as they can be rather goofy
at times! If a group 2 metal reacts ionically, it loses 2 electrons and is left with a +2 charge.
This means that the metal is happy (positive) after losing electrons and I see this as EBAY
sellers being happy that they’ve sold a product.
Ionic compounds
Ionic compounds form in a regular structure, known as a lattice. The term “Lattice” in the
field of Chemistry simply means a regular and repeated arrangement of atoms, ions or
molecules. In this case it’s a regular arrangement of atoms and ions. These are held together
by strong electrostatic forces of attraction between the oppositely charged ions and the
force acts in all directions in the lattice. Ionic compounds have the following properties:
• High melting as lots of energy is required to break the many strong bonds which
have strong forces of electrostatic attraction.
• Only conducts electricity in molten/aqueous form as ions are free to move (usually
in fixed position at room temperature).
• Often dissolve in water to form an aqueous solution.
Ionic compounds only conduct electricity in molten/aqueous form as it has delocalised ions in
those states, which can move between the layers and carry a charge, which wouldn’t be possible
in its solid form, as the ions are being used to maintain its regular lattice structure. Also, its
high melting and boiling point is due to its strong ionic bonds, so a lot
of energy is needed to overcome these bonds and change an ionic
compound from solid to liquid/gas.
An example of an ionic compound is the classic Sodium Chloride
(NaCl) shown on the figure to the right.
Another thing that you’ve got to know is how the endings of the
compound names affect the compound. This sounds a bit odd, but it’s
really not that deep. Simply put, if a compound ends in -ide, the
compound contains 2 elements, one of which is a non-metal ion. If a
compound ends in -ate, it contains at least 3 elements, one of which is
Oxygen.
15
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Deducing formulae of ionic compounds
You need to be able to write out the chemical formula of a compound given only its ions. This
is the method that I (Frank) personally found easiest and quickest, shown by an example with
calcium and nitrate to make calcium nitrate:
1. First, write out your ions and their valency in a table - valency is just how many
electrons are gained or lost when the element becomes stable, i.e the charge without the
plus or minus.
2. Then, swap the values as shown in the table.
Ion
Ca2+
NO3-
Valency
2
1
“Swapped” Valency
1
2
After this, just use these numbers as the amount of ions there are, in this case there is 1 calcium
ion and 2 nitrate ions, so the formula is Ca(NO3)2 and that’s the final answer.
However, there is also another, more traditional method where you first find the ratio of charges
between calcium and nitrate. Then, you find the amount of ions needed to balance them, since
of course the compound in the end must be neutral, i.e having 0 charge. For example, since
there are 2 positive charges, you need 2 negative charges to get an overall 0 charge (2+-1+-1 =
0), so therefore you need 1 calcium (Ca2+) and 2 nitrates (NO3-), the same result as before.
Covalent bonding
Covalent bonds occur in most non-metals and non-metal compounds. Covalent bonding
occurs when atoms share pairs of electrons and these therefore form covalent bonds. As
electrons are being shared, electrostatic forces of attraction are very strong and hence the
bonds between atoms are strong.
Covalent bonding can form:
1. Small molecular substances (like Oxygen)
2. Very large molecules (like Polymers)
3. Giant covalent structures (like Diamond)
The “dot and cross” diagram on the right demonstrates how
electrons are shared between atoms and these form covalent
bonds, allowing outer shells to be full and hence forming stable
compounds.
This bit is pretty obvious, but the exam board explicitly requires
students to know that atoms are always smaller than small molecules given that small
molecules are, you guessed it, made of atoms. You can think of atoms as the legos and small
molecular substances are the toys produced from the legos.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Dot and cross diagrams
Dot and cross diagrams are super kawaii and epic so the exam board requires you to use them
a decent amount which is fairly lucky considering they shouldn’t be too bad once you’ve got
your head wrapped around them! The main thing that you’ve got
to be able to do is explain the formation of some covalent
substances using dot and cross diagrams. So, to make things as
easy as possible, let’s use some examples!
Here we have hydrogen (H2) which, pretty simply, just has one
electron being shared by each atom. Due to this, each atom
technically has two electrons and hence has a full outside shell,
leaving it stable.
A tricker example is methane (CH4) wherein each hydrogen atom makes a
covalent bond with the same carbon atom. As carbon has 4 electrons in its
outer shell, it reacts covalently with 4 other hydrogen atoms and, as they
share one electron each, 4 electrons are shared, therefore leaving a stable
outer shell of 8 electrons and the methane compound is hence born.
Simple molecular compounds
Simple molecular compounds are pretty simple (the joke is that they are called simple
molecular substances so me calling them simple is very funny) and they come up fairly often
(oxygen, water etc) so it’s quite a good idea to know some of their properties.
1. Low melting and boiling points (so usually gases or liquids at room temperature)
a) Substances consisting of small molecules have weak intermolecular forces
between the molecules.
b) This is because intermolecular forces increase with the size of the molecules
c) This means that smaller molecules have weaker intermolecular forces so less energy
is required to overcome these forces (covalent forces don’t need to be overcome!)
d) As a result of this, they have low melting and boiling points
2. They don’t conduct electricity
a) This is because small molecules have no overall electric
charge.
b) However, the caveat is that some small molecules breakdown
in water to form ions which can conduct electricity
3. Most simple molecular compounds are insoluble in water
a) Some are soluble as they can form intermolecular forces with
water
b) These forces are stronger than those between water molecules
or their own molecules
4. They are made up of non-metal elements!
17
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Giant covalent substances
Giant covalent structures are a lot bigger than simple molecular compounds (you could
probably infer this from the name) and so the properties are, of course,
different.
1. Very high melting and boiling points.
a) This is because all atoms in these structures are linked to other atoms
via strong covalent bonds.
b) These covalent bonds are very strong, a lot of energy is required
to overcome them and hence their melting and boiling points are
high.
2. Some conduct electricity and some do not.
3. They’re made up of nonmetal elements.
What are metals?
Metals are composed of giant structures of
atoms, arranged in a regular pattern as seen to
the right. The electrons on the outer shells of
metal atoms are delocalised and this means they
are free to move through the entire structure of
the metal. As a result of this, the strong
electrostatic force of attraction between the
sea of delocalised electrons & metal cations
forms strong metallic bonds.
The following chemical and physical properties are the following:
1. High melting and boiling points.
a) This is due to the aforementioned strong metallic bonds and giant structures of
atoms.
2. Conducts heat and electricity.
a) This is because the delocalised electrons carry heat/charge throughout the entire
structure.
3. Malleable
a) The layers of atoms in metals are all the same size therefore can easily slide over
each other.
b) This means the metals can be easily bent and shaped.
4. Insoluble in water
1. Some metals react instead
Allotropes of carbon
Allotropes is just a fancy name - it basically just means different forms of carbon. The two
allotropes of carbon you’ve gotta know of are diamond and graphite and they are both examples
of giant covalent structures.
Diamond
• Structurally, each carbon atom is covalently
bonded to 4 other carbon atoms.
• It’s very hard due to the aforementioned 4 other
carbon atoms.
• It also has a very high melting point due to the 4
covalent bonds.
18
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
It doesn’t conduct electricity as there are no delocalised electrons to conduct
electricity.
• Diamond can be used in cutting tools as it is very hard due to the rigid structure.
Graphite
• Structurally, each carbon is covalently bonded
to 3 other carbon atoms and this forms layers
of hexagonal rings with no covalent bonds
between the layers.
• As there are weak intermolecular forces and
no covalent bonds BETWEEN layers, they
can easily slide over each other and this makes
graphite very soft and slippery. This means it
can be used as an effective lubricant.
• Graphite can also conduct electricity as one
electron from each carbon atom is
delocalised as carbon only makes 3 covalent
bonds. This means it can be used as an electrode.
Graphene
• Graphene is a single layer of graphite.
• It has properties that makes it useful for use in electronics and composites.
• Some of these include being able to conduct electricity.
Fullerenes
• Fullerenes are molecules of carbon atoms with hollow shapes.
• They usually are hexagonal rings of carbon atoms, but can occasionally
be pentagonal or heptagonal.
• Fullerenes can be formed from a variety of carbon atoms, but the first
one discovered (and main one) is the Buckminsterfullerene (C ) which
is spherical.
• A use of fullerenes is delivering drugs to the body.
Carbon Nanotubes
• They are fullerenes in cylindrical form.
• They have very high length to diameter ratios which makes them useful
in nanotechnology, electronics and materials.
• Some examples of uses include reinforcing materials in stuff like tennis
rackets.
•
60
Polymers
•
•
•
•
Polymers are long chains of repeating units called monomers.
They have very large molecules and “n” is used to
show that there are many repeating units.
The atoms in the polymer molecules are linked via
strong covalent bonds.
They have high melting and boiling points and are
solid at room temperature as the intermolecular forces
between polymer molecules are strong
19
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Conservation of mass
In a chemical reaction, no atoms are destroyed, and no atoms are created - mass is always
conserved. Therefore, in a balanced chemical equation, there are the same number and types
of atoms on each side.
This can clearly be demonstrated through a reaction in a closed system, such as a precipitation
reaction. This is where two solutions react and an insoluble solid (a precipitate) is formed in
the solution, leaving no mass change as no reactants or products can escape.
Why can mass increase in an experiment?
If the mass increased in an experiment, the most likely explanation is that one of the reactants
reacted with a gas in the air, most likely to be oxygen.
A good example is iron in an unsealed container reacting with the oxygen from the air to rust,
increasing the mass from the original piece of iron:
iron(s) + oxygen(g) → iron oxide(s)
Why can mass decrease in an experiment?
If the mass decreased in an experiment, the most likely explanation is that gas escaped from
a vessel that was not enclosed, meaning that as the gas is no longer in the reaction vessel, you
can no longer measure its mass, making the total mass decrease.
An example is the thermal decomposition of sodium carbonate to form sodium oxide and
carbon dioxide, a gas, making the mass seem to decrease:
sodium carbonate(s) → sodium oxide(s) + carbon dioxide(g)
Relative formula mass AKA relative molecular mass - M
r
The relative formula mass/relative molecular mass of a compound is the sum of all of the
relative atomic masses (Ar) of all of the atoms in its formula, so essentially just adding up the
mass of everything in that compound. In reactions when calculating the Mr you ignore the
coefficient. E.g. The Mr of 2H2O is just 1+1+16 NOT 2 (1+1+16).
Here’s a worked through example where we find the relative formula mass of methane:
1. First, we find the Ars of carbon and hydrogen, as we know methane is CH4 - Ar of C =
12, Ar of H = 1
2. Next, we find the Mr by multiplying and adding each atom accordingly - 12 + (4x1) =
16.
3. Therefore we are done and the Mr of methane is 16! (not 16 factorial)
Empirical formula
The empirical formula of a compound is the simplest integer (whole number) ratio of atoms
in that compound. Here is a worked through example of finding the empirical formula of
fructose:
1. Find the formula of the compound - Fructose is C H O , the same formula as glucose!
2. Find a number that you can divide the number of atoms by to get whole numbers - Here
that number is 6 as seen here:
6
Element/Atom
Carbon (C)
12
6
Hydrogen (H)
20
Oxygen (O)
Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Chemical Formula
6
12
6
Empirical Formula
1 = 6/6
2 = 12/6
1 = 6/6
3.
Therefore, the empirical formula of fructose (and glucose) is CH O!
2
What is a mole?
The mole is an extremely important concept in chemistry, so make sure you grasp it well as
you’ll use it in a lot of questions in the actual exam. Essentially, one mole is just an amount of
atoms or particles. Specifically, it’s Avogadro’s constant, 6.02✕1023 , a very large number:
602,000,000,000,000,000,000,000 particles.
The reason why this is so useful is that one mole of atoms or molecules has the same mass in
grams as the relative atomic mass of that respective atom or molecule. All this really means is
that:
• One mole of hydrogen weighs exactly 1g - Ar of H is 1.
• One mole of carbon weighs exactly 12g - Ar of C is 12.
• One mole of sodium weighs exactly 23 g - Ar of Na is 23.
However, be careful as you need to take into account how many atoms there are, for example:
• One mole of hydrogen gas weighs exactly 2g (H2 = 2✕1)
• One mole of nitrogen gas weighs exactly 28g (N2 = 2✕14)
• One mole of oxygen gas weighs exactly 32g (O2 = 2✕16)
But, even though all of these have different weights, they all have the same number of particles
- one mole, or precisely, 6.02✕1023 particles.
A key formula that you need to know is:
moles = mass (in grams) /Mr
This is a very common one and you will most likely be using it a
lot.
Concentration
Concentration is a measurement for how much mass there is in a
certain volume - a more concentrated solution will have more
particles per unit volume and a more dilute solution will have less particles in a given volume.
This formula is:
concentration (g dm-3) = mass of solute (g) / volume of solution (dm3)
Here, the solute is just the solid that you dissolve in the solution, and 1dm3 = 1000cm3.
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Edexcel GCSE Chemistry: Unit 1 - Key Concepts in Chemistry
written by Daniyaal Anawar, Frank G, Archie Thompson, Jammal Othman, edited by Ken Tu
Empirical formulas from experiments
You need to be able to describe an experiment to calculate the empirical formula of a simple
compound, with the named example being magnesium oxide (MgO). Here are the steps to do
so:
1. Clean a crucible by heating it until it is red hot and
allow it to cool, reweighing it (in grams) with a lid
using a mass balance.
2. Place a magnesium ribbon inside a clean crucible
and measure its total mass in grams using a mass
balance.
3. Heat the crucible with the magnesium with the lid
on, leaving a small gap for oxygen.
4. Heat it until no further mass change occurs and all of
the magnesium has turned white. Allow the
crucible to cool.
5. Reweigh the crucible in grams using a mass balance.
6. The mass of magnesium oxide is the difference between the final reading and the
initial reading of just the crucible and the lid.
7. Using the mass of oxygen (difference between magnesium and magnesium oxide
readings), we find that the ratio of moles between magnesium and oxygen are 1:1,
highlighting the fact that the empirical formula of magnesium oxide is MgO.
Table to show calculations using example data:
Magnesium
Oxygen
Mass
1.08
1.80-1.08 = 0.72
Mr
24
16
Moles
1.08/24 = 0.045
0.72/16 = 0.045
Mole ratio
1
1
Therefore, as the mole ratio between magnesium and oxygen in the reaction is 1:1, the
empirical formula of magnesium oxide is MgO!
Limiting reactants
The limiting reactant of a reaction is the reactant that gets used up first in a chemical reaction,
limiting how much product can be formed, hence the name. Using a balanced chemical
equation, you can find out the limiting reactant in a reaction. This is done by:
1. Find the moles of each reactant using moles = mass (g) / Mr.
2. Find the mole ratio between each reactant.
3. Compare the ratios from the moles and the actual reaction coefficients, the one with the
smaller numbers is the limiting reactant and is not in excess.
Balancing Equations with Masses
You need to be able to balance an equation given the masses of the reactants and products. You
can do this by:
1. Finding the moles of each substance using moles = mass (g) / Mr
2. Find the mole ratio in the reaction and make them all integers (whole numbers).
3. You now have the balanced symbol equation.
22
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
What are the states of matter?
There are three states of matter:
• Solid
• Liquid
• Gase
These can be modelled using the particle model of
matter- a theory that explains the properties of the
different states of matter and how they behave differently.
In the particle model, each particle is represented by using
a small solid sphere.
Forces between each particle can be weak or strong and
the properties of each state of matter are dependent on
these forces.
Properties of solids:
1. Strong forces of attraction between the particles.
2. Held in fixed positions in a regular lattice arrangement.
3. Keep a definite shape and volume.
4. The particles in the solid have little energy in comparison to the particles in a liquid
or gas.
5. Hardly move- only vibrate about a fixed point. Hotter the solid the more they
vibrate. This causes solids to expand slightly when heated.
Properties of liquids:
1. Some force of attraction between the particles.
2. Free to move past each other but tend to stick together.
3. Don’t keep a definite shape but keep the same volume.
4. Particles have more energy than particles in solid state but less energy than gas state.
5. Particles are constantly moving in random motion. Hotter the liquid the faster they
move. This causes the liquid to expand slightly when heated.
Properties of gases:
1. No/negligible force of attraction between the particles.
2. Particles are free to move.
3. Particles travel in straight lines and only interact when they collide.
4. Don’t keep a definite shape or volume - will always fill any container, exerting a
pressure on the walls.
5. Particles have more energy than the particles in the solid state and liquid state.
6. Particles move constantly with random motion. The hotter the gas the faster they
move. This causes the gas to either expand when heated or their pressure increases.
How can a physical state change occur?
A state change is where a substance changes from one state to another e.g. when ice melts
into water it goes from a solid to a liquid.
1. When a solid is heated - its particles gain more energy.
2. The particles now vibrate more rapidly, this weakens the forces that hold the solid
together and so the solid expands.
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Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
3. When they reach a certain temperature the solid has enough energy to break free
from their fixed position and the solid turns into a liquid. This is called melting.
4. Liquid is heated - its particles gain
more energy.
5. Particles now have more energy
which overcomes the intermolecular
forces that hold the liquid together.
6. When the liquid reaches a certain
temperature the particles have enough
energy to break their bonds. The
liquid now turns to a gas and this is
called evaporation.
7. When a solid changes state to a gas
without becoming a liquid first it is
called subliming.
The amount of energy required to change state depends on the strength of the forces between
particles in a substance. The stronger the forces between the particles, the higher energy
required to overcome these forces of attraction and by extension, the higher the melting and
boiling point of the substance.
2.4) At temperatures below the melting point, the substance is solid. At temperatures above
the melting point but below the boiling point, the substance is liquid. At temperatures
above the boiling point, the substance will be a gas.
What are chemical changes?
Chemical changes are when a new product has been formed when reactants react. Chemical
changes happen during chemical reactions - when bonds between atoms break and the
atoms change positions. Chemical changes are harder to reverse than physical changes.
E.g. in photosynthesis:
carbon dioxide + water → glucose + oxygen
6CO2 + 6H2O → C6H12O6 + 6O2
As you can see the chemicals on the right-hand side (RHS) completely change.
You also need to know basic vocabulary in Chemistry. You’ll often hear the words “pure
substance”, “compound”, and “mixture” quite a bit!
A compound means more than one type of atom in a chemical. E.g. Ammonia, NH or
calcium carbonate, CaCO as all of these have multiple elements in them.
3
3
A pure substance means completely made of a single element or a single compound. E.g.
Sodium or lithium oxide. E.g. “inside the evaporating basin are pure copper sulfate crystals”.
A mixture means more than one compound or element present but not chemically
combined. E.g. The air in the atmosphere is a mixture of gases 21% oxygen, 78% nitrogen,
0.9% argon, 0.03% carbon dioxide and other gases.
24
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
How can we test for impurities within a substance?
1. Every pure substance has a very specific and
sharp melting point/boiling point.
2. Mixtures/Impure substances melt gradually
over a range of temperatures and so their
melting point is not exact.
Therefore, we can use the melting and boiling point of
a substance to determine whether it is a pure
substance or a mixture.
Next, we will go over a series of methods of
separating and purifying substances.
Simple distillation
Simple distillation separates out a solute from a
solution. For example, pure water from seawater.
1. Can only be used to separate things with
different boiling points (the solute has a
very different boiling point from the
solvent).
2. Steps of distillation:
• Heat the distillation flask gradually.
• The liquid with the lowest boiling
point evaporates first.
• The vapour then passes into the
condenser where it cools. The vapour
is then condensed and is then
collected in a beaker. This liquid is
called the distillate.
• The remaining solution becomes more
concentrated as the amount of solvent
in it decreases.
Fraction distillation
Fractional distillation is used to separate a mixture
of liquids with similar boiling points. For example,
ethanol from a mixture of ethanol (bp: 78°C) and
water (bp: 100°C), and for separating fractions in
crude oil.
1. The different liquids in fractional distillation
will have different boiling points, so they will
evaporate at different temperatures.
2. The fractional distillation column is hotter at
the bottom and cooler at the top.
3. The vapour rises through the column and
condenses when they reach the part of the
column that is below the temperature of their
boiling point.
25
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
4. You can obtain the different liquids from the column by:
• Collecting different liquids from different parts of the column (lowest
boiling point at the top and highest boiling point at the bottom).
• Continuing to increase the temperature in the column by heating the mixture
(the liquid with the lowest boiling point
is collected first).
Filtration
Filtration is used to separate an insoluble solid from a
liquid. For example, sand and water, and excess
reactant from a reaction mixture.
1. The filter paper has tiny holes/pores in it.
These are large enough to let small molecules
such as dissolved ions or soluble particles
pass through but not the larger undissolved
solid particles. The solid is therefore collected
in the filter paper.
Crystallisation
Crystallisation separates a soluble solid from a solution. For example, copper sulfate
crystals from copper sulfate solution.
1. When the solution is heated in a dish/evaporating basin over a roaring bunsen
burner flame.
2. The solvent (often water)
evaporates.
3. The dish is then left next to a
windowsill (as the sun is warming it
up to continue evaporating the
solvent) to cool.
4. This forms solid crystals. The
crystals are then filtered out of the
solution and left to dry in a warm
place like an oven or they could be
patted with filter paper.
26
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
Paper chromatography
Paper chromatography is used to separate mixtures of soluble substances (often coloured
inks/dyes).
1. Paper chromatography contains two
phases:
• Stationary phase - where the
molecules can’t move, this could
be a solid or a really thick liquid.
In paper chromatography this is
the absorbent paper.
• Mobile phase - where the
molecules can move, this could
be a liquid or a gas. This is the
solvent that moves through the
paper carrying different soluble
substances with it.
Given that you now know the different
methods, you can identify the types of substances you have in a mixture and choose a
technique which is most appropriate in a given question.
The different components in the mixture separate out as the mobile phase moves over the
stationary phase. This happens as the different dissolved substances are attracted to the two
phases in different proportions and so they move at different rates through the paper.
How to interpret a paper chromatography
1. Separation by chromatography produces a chromatogram. This can be used to
distinguish between pure and impure substances. A pure substance produces only
one spot on the chromatogram. An impure substance/mixture produces two or
more spots.
2. A paper chromatogram can also be used to identify substances by comparing them
with known substances. Two substances are probably the same if they contain the
same number of spots in the same colour and if the spots travel the same distance
up the paper. To identify the substances by calculation of Rf values, the Rf values
can be calculated and then compared with the Rf (retention factor) values of known
substances.
• Rf = distance travelled by solute/ distance travelled by solvent
Retention factor (Rf) is basically a measurement of how soluble a substance is in the form
of a ratio. If a solvent moves 5cm up the chromatography paper, how much of that 5cm will
the substance move up with the solvent, 4cm, 3cm etc. A higher retention factor most likely
means that the chemical is more soluble in that solvent.
27
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
Paper chromatography core practical
Aim: To investigate the composition of inks using chromatography.
Equipment:
• Beaker
• Ruler and pencil
• Distilled water or suitable
solvent
• Chromatography paper
• Pipette or capillary tube
Method:
1. Use a pencil to draw a line across
the chromatography paper, 2 cm
from the bottom.
2. Use a capillary tube to add small
dots of different inks to the line on
the chromatography paper. Make sure not to place the dots too close.
3. Place the paper in a beaker containing 1 cm3 of solvent and leave the solvent until it
has moved 2/3 of the way up the paper.
4. Remove the chromatogram from the solvent. Draw a horizontal line across the
chromatography paper (at the solvent front) at this point to measure where the
solvent has reached .
5. Leave the chromatography paper until the solvent has dried.
6. Measure the distance travelled by the solvent front and the pencil line.
•
Key Points:
1. Use a pencil to draw the line on chromatography paper as it is insoluble and
won’t travel up the paper.
2. Ink spots should be placed above the level of solvent (2cm) in the beaker to
prevent them dissolving in the solvent and being washed away.
3. The chromatography paper should be removed from the solvent before the
solvent front reaches the top to allow the Rf values to be calculated.
4. Spots with the same Rf values with the same mobile phase and stationary
phase are likely to be the same substance.
The Rf values can be used to identify the substances in the dye. There is an Rf database
which has a list of chemicals and the Rf values calculated in the database can be used to find
out what substance was in the dye. Another thing to note is that several chemicals may have
the same Rf value in the same solvent therefore to narrow down the chemicals in the dye,
you may wish to repeat the experiment with different solvents.
28
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
Simple distillation in chromatography
This video explains everything about the core practical with distillation in chromatography.
https://www.youtube.com/watch?v=HOS-Z5aqDwA
Aim: To investigate the composition of inks using simple distillation
Equipment:
•
•
•
•
•
•
•
•
•
•
Beaker
Condenser with rubber tubing
Fractionating column
Connectors
Round-bottomed or pear-shaped flask
Bunsen burner
Clamp
Stand
Cold water tap
Thermometer
Method:
1. Add a small volume of ink to a conical flask and connect the flask to a condenser,
you may also add a thermometer. Connect the condenser to a tap and place a beaker at
the end of the opening.
2. Using a Bunsen burner, heat the flask slowly so that the ink bubbles gently. Do not
heat the flask too strongly. Move the bunsen burner away if it starts boiling
vigorously.
3. Collect a sample of the distilled solvent. Turn the Bunsen burner off when finished.
• Key Points:
1. Make sure the condenser is horizontal and the water enters from the bottom
and leaves from the top.
2. Make sure the flask is not heated strongly as this could cause the water to boil
over into the beaker collecting the distillate.
• Analysis:
1. The temperature recorded on the thermometer can be used to identify the
solvent present in the ink by comparing it to the boiling point of possible
known solvents.
2.
How can waste and groundwater be made potable?
1. Potable water is water that is suitable for drinking.
2. Potable water requires:
• Low levels of microbes.
• Low levels of contaminating
substances.
• It is not the same as pure water but it
is safe to drink.
3. Groundwater: found in aquifers (rocks that
trap water underground)
Wastewater: water that has been
contaminated by a human process.
4. Groundwater and wastewater are purified in
water treatment plants using three
processes:
29
Edexcel GCSE Chemistry: Unit 2 - States of Matter and Mixtures
written by Tyisha Rebelo, edited by Ken Tu
•
•
•
Filtration: A wire mesh is used to filter out large objects etc. and then gravel
and sand beds filter out other solid bits present in the water.
Sedimentation: iron/aluminium sulfate is added to the water making finer
particles clump together and these clumps then settle at the bottom.
Chlorination: Chlorine gas is bubbled through the water to kill microbes and
harmful bacteria.
Making sea water potable by using distillation
1. Sea water is found in lakes, rivers and reservoirs.
2. Sea water is made potable by using distillation as sea water contains too much salt
for human consumption:
• Filter the seawater
• Boil the seawater
• Let the water vapour cool and condense
3. However, distillation requires a lot of energy making this method extremely
expensive.
Water without dissolved salts
1. When water is used in analysis during experiments like mixing or dissolving
something in water, deionised water should be used.
2. Deionised water/distilled water is water that has the ions that are present in normal
tap water removed.
3. Although these ions are present in small amounts in tap water, they can interfere
with reactions and give an experiment the false result.
30
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
What are acids & bases?
Acids in a solution dissociate into H+ ions and alkalis in a solution dissociate into OH- ions.
E.g. HCl acid would dissociate into H+ and Cl- where the HCl provided the solution H+ ions.
For a chemical to dissociate means for a compound to break up into simpler constituents, and
in this case simpler constituents ions H+ and Cl-. Something like NaOH (aq) would dissociate
into Na+ and OH- ions.
Neutral solutions are a 7 on
the pH scale. The more acidic a
solution is, the lower its pH
and the more alkaline it is, the
higher the pH. To the right is a
picture of how universal
indicators would behave in the
presence of certain solutions.
Chemical indicators can give away whether the solution is acidic or basic (alkaline). Below is
a table of different indicators you are required to know and how they change colour.
Indicator
Acidic Solutions
Neutral Solutions
Alkaline Solutions
Universal Indicator
Red
Green
Purple
Litmus
Red
Purple
Blue
Phenolphthalein
Colourless
Colourless
Pink
Methyl orange
Red
Orange
Yellow
The more acidic a solution is, the higher the concentration of H+
ions is and the lower its pH. The more alkaline a solution, the
higher the concentration of OH- ions in it, and the higher its pH
is.
Alkalis are bases that are soluble in water. All alkalis are bases but
not all bases are alkalis.
Another thing to know is as pH of a solution decreases by 1, H+
concentration increases by a factor of 10. E.g. If the pH of our
solution is initially 6 then after a reaction the new pH is 5, the H+
ions concentration in the solution has increased by 10 or tenfold.
Furthermore, if the initial pH is 7 and after a different reaction the
new pH is 4, the H+ concentration has increased by 10 ✕ 10 ✕ 10 which is 1000 times.
31
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
Core practical: investigate the change in pH on adding powdered calcium
hydroxide or calcium oxide to a fixed volume of dilute hydrochloric acid
(https://youtu.be/pIk6IpeADq4)
Aim: To investigate the change in pH when calcium oxide and calcium hydroxide are added
to a fixed volume of dilute hydrochloric acid.
Apparatus:
• Measuring cylinder
• Beaker
• Spatula
• Universal indicator paper / pH probe
• Glass rod
• Calcium hydroxide
• Calcium Oxide
• Dilute hydrochloric acid
• pH calibration colour chart
Method:
1. Use the measuring cylinder to measure out a fixed volume of hydrochloric acid and
add it to the beaker.
2. Use a clean glass rod to transfer a few drops from the beaker onto some universal
indicator paper. Wait for 30 seconds and then compare to the pH colour chart. Record
the pH on a suitable table.
3. Add a spatula of calcium hydroxide to the beaker, mix thoroughly and repeat step 2
to record pH.
4. Repeat step 2 until 10 spatulas have been added.
5. Carry out the same experiment again but this time with calcium oxide.
Conclusion:
Use your results to plot a graph. From this graph explain the relationship between the pH and
the number of spatulas added and what this tells you about the presence of H and OH ions in
the solution.
+
32
-
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
What is the difference between the strength and concentration of a solution?
Firstly, what is a solution? A solution forms when a solute is dissolved in a solvent.
Acid/alkali strength is associated with
the number of H+/OH- ions it readily
dissociates in an aqueous solution. The
more H+/OH- ions that are dissociated, the
stronger the acid and vice versa. An acid
which dissociates only a few H+/OH- ions
in a solution would be considered a weak
acid.
Acid concentration is to do with the
number of H+ ions in a solution
compared to the volume of solvent. The
more H+ ions there are in a given volume
of solvent, the more concentrated it is
and vice versa. A solution with very few
H+ ions in a given volume of solvent
would be called dilute.
Some simple chemical tests
You’re required to know how to test for
the presence of some basic chemicals.
Hydrogen: Place a lit splint into a test
with hydrogen gas. If Hydrogen is
present, then there will be a squeaky pop
sound .
Carbon Dioxide: Bubble the gas through lime water. If carbon dioxide is present, then the
lime water will turn cloudy/milky.
Oxygen: place glowing splint into a test tube with oxygen gas. If oxygen is present, the splint
will relight into a flame.
Chlorine gas: damp blue litmus paper turns red, then bleaches and goes white.
33
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
How can we prepare soluble salts from an acid and insoluble reactant?
Examples of insoluble reactants include metals, metal oxide and carbonates. An excess of the
insoluble reactant e.g. copper oxide is added to ensure all of the acid has reacted.
The excess insoluble reactant is filtered out with a funnel and filter paper to ensure that only
the water in the filtrate off to leave behind a pure salt.
Core practical: investigate the preparation of pure, dry hydrated copper
sulfate crystals starting from copper oxide including the use of a water bath
Aim: to obtain pure, dry and hydrated copper sulfate crystals from copper oxide. Hydrated
copper sulfate crystals should be blue and take up a shape that is regular.
Practical Video: https://youtu.be/mNebkJ7_48s
Apparatus:
• Measuring cylinder
• Beaker
• Glass rod
• Spatula
• Evaporating basin
• Bunsen burner
• Water bath
• Conical flask
• Filter Funnel (with filter paper)
• Watch Glass
• Sulfuric Acid
• Copper Oxide
Method:
1. Add some sulfuric acid to the conical flask and use the water bath to heat it.
2. Take a spatula of copper oxide and add it to the conical flask.
3. Use the glass rod to stir the solution.
4. Repeat steps 2 & 3 until the copper oxide is in excess. You can tell as the will be
copper oxide left in the conical flask.
5. Setup the beaker with a funnel (lined with filter paper) and filter the solution to remove
the excess copper oxide.
6. Take the filtrate and add it to the evaporating basin.
7. Use the bunsen burner to heat the evaporating basin until half of the water has
evaporated.
8. Pour the remaining solution into a watch glass and leave until the rest of the water has
evaporated off.
9. Record your observations of the crystals formed by describing their physical properties
such as size and colour.
34
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
Preparing soluble salts from an acid and soluble reactant
For this reaction, a titration must be used. This is because there is NO insoluble substance that
can be removed through filtration and hence, we must only use a set amount of reactants to
ensure that no excess is left in the product mixture if we want the sample to be pure. Titrations
of an acid-alkali reaction to form a soluble salt can be further explained in Unit 5 - Separate
Chemistry I.
What are the solubility rules?
With regards to the different substances in water, some are soluble in water, and some are not.
You must recall which substances are soluble and which ones aren't, which is an unfortunate
memory test IMO :(.
Soluble (in water)
Insoluble (in water)
All sodium, ammonium and
potassium salts
(SAP like the tree sap for maple
syrup)
N/A
All nitrates
N/A
Most chlorides
Except silver and lead chloride (CSL - this sounds
like a Call of Duty clan, for Chlorides no Silver &
Lead)
Most sulfates
Lead, barium and calcium sulfate (S + LBC - Sulfate
+ Lead, Barium, Calcium. LBC is a radio station that
stands for Leading Britain’s Conversation; they talk
about politics, current affairs in the UK.)
Sodium, ammonium, and
Most carbonates
potassium carbonates (SAPC; SAP
C sounds like “sexy” - kind of…)
Sodium, ammonium, and
potassium hydroxide (SAPH; SAP
H reminds me of “slap Hilary” idk
why don’t even ask.
Most hydroxides
35
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
What is electrolysis?
Electrolysis causes the decomposition of electrolytes using electrical energy from a direct
current (DC) power source. This happens due to the free electrolytes being attracted to the
electrodes from the power supply. Electrolytes are ionic compounds that are either dissolved
in water or in a liquid state. When this happens, the ions within the compound are able to
move freely. “Electro” means like electricity or charge and “lyte” comes from Greek meaning
loose
or
dissolved.
1. When the electrolyte makes contact with the two electrodes of the DC circuit, the
circuit is complete.
2. At this point, the electrons drift from the negative terminal of the battery to the
positive terminal.
3. As a result of this, the anode, like in the
diagram, becomes more positively
charged (rather less electrons are
present there) and attracts all the
negative ions in the electrolyte.
4. In contrast, the cathode becomes more
negatively charged (as there’s a buildup of electrons due to the electrons
drifting there) and attracts all the
positive ions in the electrolyte.
5. As a result of this, oxidation and
reduction can take place at the anode
and cathode respectively. Because at the
anode the anions lose electrons but at the
cathode
cations
gain
electrons
ALWAYS!
36
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
Cations are positively charged ions. They are attracted to negatively charged electrodes
called cathodes. Anions are negatively charged ions. They are attracted to positively charged
electrodes called anodes.
When the cations reach the cathodes and the anions reach the anodes, they gain/lose electrons
and discharge to form neutral atoms or molecules. Cations in the electrolyte gain electrons
whilst anions in the electrolyte lose electrons.
What if there are multiple ions in the electrolyte?
In aqueous solutions e.g. NaOH (aq), there
are multiple ions, from the water itself (H2O
can partially dissociate into H+ and OHions) and from the dissolved substances
(ions of the electrolyte in this case is Na+
and OH-). Hence, during electrolysis (at the
cathode and anode) these ions compete in
order to react and the ones that actually
form are based on their reactivity. The
readily discharged an ion is the more likely
that specific atom will form.
How readily discharged the ions are can be
found in the electrochemical series. You do
not need to know them but it’s something
worth pondering.
At GCSE you just need to know them :(. A further explanation can be linked here:
https://www.chemguide.co.uk/inorganic/electrolysis/solutions.html
37
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
Electrolysis of Water: Water contains two ions, H+ & OH-. H+ is attracted to the negative
cathode and gains an electron to form hydrogen. OH- is attracted to the positive anode where it
loses electrons to form oxygen. The volume of hydrogen to oxygen is 2:1 as the molar ratio is
2:1.
2H2O → 2H2 + O2
These are just the products you need to know.
Solution aqueous
Product formed at
Anode
Cathode
Copper chloride, CuCl2(aq)
Chlorine (Cl2)
Copper (Cu)
Sodium chloride, NaCl(aq)
Chlorine (Cl2)
Hydrogen (H2)
Sodium sulfate, Na2SO4(aq)
Oxygen (O2)
Hydrogen (H2)
Sulfuric acid, H2SO4(aq)
Oxygen (O2)
Hydrogen (H2)
Molten lead bromide - the lead ions (Pb2+) gain electrons to become lead atoms (Pb) and the
bromide ions (Br-) lose electrons to become bromine atoms (Br2).
Electrolysis redox half equations
Half equations represent the gain or loss of electrons from a substance.
Half Equation rules:
• Atoms must be balanced.
• Electrons represented by e-.
• Charges on both sides must be balanced.
Cathode half equations examples:
Na+ + e- → Na
2H+ + 2e- → H2
Anode half equations examples:
2Cl- → Cl2 + 2e2O2- → O2 + 4eThe half equations (of an electrolysis reaction) can be combined to create an overall ionic
equation. Ions which remain the same on both sides of the equation can be cancelled out as
they are known as spectator ions.
Reduction is the gain of electrons. Reduction happens at the cathode. Oxidation is the loss
of electrons. Oxidation happens at the anode. The most common mnemonic is OILRIG.
Oxidation Is Loss (of electrons). Reduction Is Gain (of electrons).
Video explanation of electrolysis with copper electrodes:
Student Sheet of electrolysis with copper electrodes:
38
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
Core practical: investigate the electrolysis of copper sulfate solution with
inert electrodes and copper electrodes
Practical Video: https://youtu.be/AiEpVTySk70
Student Sheet: https://www.pearson.com/content/dam/one-dot-com/one-dotcom/uk/documents/subjects/science/GCSE-core-practical-sheets/Chemistry/cp4-electrolysiscopper-sulfate.pdf
Aim: To electrolyse copper sulfate solution using inert (graphite) electrodes and copper
electrodes.
Apparatus:
• DC power supply (e.g. battery)
• Connecting leads
• Crocodile clips
• Timer
• Graphite rods x2
• Copper sulfate solution
• Weighing scale
• Beaker
• Test tubes
Method:
1. Add the copper sulfate solution to the beaker.
2. Place the graphite rods in the solution and attach them to the power supply using
connecting leads and crocodile clips.
3. Fill two test tubes with the copper sulfate solution and place over the rods.
4. Turn the power supply on and then take note of observations.
5. Test the gases that form in the test tubes and record observations.
Experiment with copper electrodes:
Apparatus:
• DC power supply
• Connecting leads
• Timer
• Graphite rods x2
• Copper rods/foil x2
• Weighing scale
• Connecting leads
• Crocodile clips
• Beaker
• Test tubes
Method:
1. Add the copper sulfate solution to the beaker.
2. Place the graphite rods in the solution and attach them to the power supply using
connecting leads and crocodile clips.
3. Fill two test tubes with the copper sulfate solution and place over the rods.
39
Edexcel GCSE Chemistry: Unit 3 - Chemical Changes, Acids, and Electrolytic Processes
written by Faran Ahmad, edited by Ken Tu
4. Turn the power supply on and then take note of observations.
5. Test the gases that form in the test tubes and record observations.
6. Repeat the experiment but this time measure the mass of the copper foil and then add
it to the beaker.
7. At the end of the experiment, measure the mass of the copper foil again and calculate
the difference in mass.
40
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
What is reactivity?
The reactivity series is a list of metals in order of their reactivity going from the most to least
reactive.
Metal
Reaction with water
Reaction with
dilute acid
Likelihood to form
cations
Potassium
These elements react with cold
water to form hydrogen and a
metal oxide
Violent reaction
The ability to form
positive ions increases
as you go up the table
Sodium
Calcium
Magnesium These elements react very
slowly with cold water but react
with steam to form hydrogen
Aluminium
and a metal oxide
React to form
hydrogen and a
salt solution
Zinc
Iron
Copper
Do not react with cold water or Do not react
steam
Silver
Gold
An example of the cold water reaction would be as follows:
2Na (s) + 2H2O (l) → 2NaOH (aq) + H2 (g)
An example of the steam reaction would be as follows:
2Al (s) + H2O (g) → 2MgO (s) + H2 (g)
An example of the reaction with dilute acid is as follows:
Fe (s) + 2HCl (aq) → FeCl2 + H2
In these reactions bubbles of gas are formed. The more bubbles formed, the higher up the
table the metal is. The metal atoms lose electrons in all reactions above, forming cations
(positive ions). If a metal has a higher tendency to form cations, it will be more reactive.
41
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Displacement reactions
The reactivity series allows us to predict whether reactions will take place. Metals will react
with compounds of the metals that have a lower reactivity.
When zinc is put into a copper sulfate solution, a copper coating forms onto
the surface of the zinc. This is because the zinc takes the place of the
copper in the compound, forming a zinc sulfate solution.
Zn (s) + CuSO4 (aq) → Cu(s) + ZnSO4(aq)
This is called a displacement reaction. The zinc displaced the copper in the compound.
Displacement reactions only happen if the metal being introduced is more reactive than the
metal already in the compound. For example, Copper cannot displace zinc from its
compounds because copper is less reactive.
Displacement reactions are also redox reactions. The reaction between zinc and copper can
also be written as an ionic equation.
Zn + Cu2+ → Cu + Zn2+
The sulfate ions do not need to be included since they are the same on both sides of the
equation- they are spectator ions. The zinc atoms lose two electrons to form zinc ions. This
can be shown in a half equation where ‘e’ represents an electron:
Zn → Zn2+ + 2e
This is oxidation - the loss of electrons.
The copper ions gain two electrons to form copper atoms:
Cu2+ + 2e → Cu
This is reduction - the gain of electrons.
One substance has been oxidised and another has been reduced. This is called a redox
reaction. Redox can be remembered through the acronym OILRIG - Oxidation is Loss (of
electrons), Reduction is Gain (of electrons). Oxidation is also the gain of oxygen and
reduction is the loss of oxygen.
What are ores?
Most metals are reactive enough to have reacted with other elements to form
compounds in rocks e.g. iron oxide. The process of obtaining a metal from these
compounds is called extraction. There are also some very unreactive metals,
such as gold and platinum, that are found naturally in their native state (as
uncombined elements).
An ore is a rock that contains enough of a certain metal to be extracted to a pure metal for
profit. Haematite is an ore containing iron oxide. Iron can be extracted by heating the iron
extract with carbon. Carbon is more reactive than iron so carbon displaces the iron in iron
oxide and forms carbon dioxide.
iron oxide + carbon → iron + carbon dioxide
This method is used for compounds of metals below carbon in the reactivity series.
42
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Malachite is an ore containing copper carbonate. It is heated to produce copper oxide,
which is then heated with carbon to produce copper.
Metals more reactive than carbon must be extracted using electrolysis. This is when
electricity is passed through a molten ionic compound to break it into its elements. For
example, the bauxite ore contains aluminium oxide which can be electrolysed to form
aluminium. A lot of energy is needed to keep the metal oxides molten for electrolysis which
makes it expensive. Electrolysis is only used to extract very reactive metals that cannot be
obtained by heating their oxides with carbon.
Metal
Method of extraction
Potassium
Electrolysis of a molten compound
Sodium
Calcium
Magnesium
Aluminium
(Carbon)
Zinc
Heat an ore with carbon
Iron
Copper
Silver
Found as the uncombined element
Gold
Biological methods of metal extraction
Copper is traditionally extracted by heating copper sulfide. However, copper ores are running
out so ores with smaller amounts of copper are being used to obtain copper.
Bioleaching uses bacteria grown on a low-grade ore. The bacteria produce a leachate, which
is a solution containing copper ions. Scrap iron is used to obtain the copper by displacement.
Then it is purified by electrolysis. This method is also used for nickel, cobalt and zinc.
43
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Phytoextraction involves growing plants that absorb metal compounds which can then be
burnt to form ash, from which the metal is extracted.
Process
Advantages
Disadvantages
Bioleaching
Does not require high
temperatures
Toxic substances and sulfuric acid
can be produced by the process, and
damage the environment
Phytoextraction
Can extract metals from
contaminated soils
It is more expensive that mining
certain ores
Weather conditions can largely sway
how many plants can be grown
Both bioleaching and
phytoextraction
No harmful gases are
produced (e.g. sulfur
dioxide)
Causes less damage to
the landscape than
mining
Conserves supplies of
higher grade ores
Slow
Oxidation and reduction
Many metals are extracted from metal oxide ores. The oxygen must be removed in order to
obtain the metal from its oxide. The loss of oxygen means that the compound has been
reduced.
Oxidation is the gain of oxygen (and loss of electrons) and reduction is the loss of oxygen
(and gain of electrons). Oxidation and reduction always happen together in redox reactions.
Iron oxide can be heated with carbon to obtain iron.
Iron Oxide + Carbon → Iron + Carbon Dioxide
Aluminium is obtained by removing oxygen from aluminium oxide by electrolysis.
During electrolysis the positive aluminium cations are attracted to the negative cathode to
gain electrons while the negative oxide
anions are attracted to the positive anode
where they lose electrons to form oxygen.
The graphite (carbon) anodes can react with
the oxygen under these high temperatures
to form CO2 (carbon dioxide)
The reactions at the electrodes are as follows:
At the cathode: Al3+ + 3e → Al (reduction)
At the anode: 2O2- → O2 + 4e (oxidation)
44
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Corrosion
Corrosion occurs when a metal reacts with
oxygen (becomes oxidised), making the metal
weaker over time. For iron water is also
required for corrosion alongside oxygen and
that is called rusting.
A metal of higher reactivity will corrode faster.
For example, silver and gold will have little to
no corrosion at all while Sodium and Calcium
will have large amounts of corrosion. Some
metals, however, are an exception to this such
as Aluminium because it forms a protective
oxide layer (or a tarnish), which prevents
further reaction.
Life cycle assessment and recycling
The main advantages of recycling include:
• Natural reserves of metal ores will last longer
• Reduced need to mine ores which is advantageous as mining can damage the
landscape and cause noise and dust pollution
• Less waste metal ends up in landfill sites
• Reduction in pollution due to mining and contamination such as the production of
sulphide ores
• Many metals need more energy to be extracted from their ore than they need to be
recycles
The main disadvantages of recycling include:
• The cost of collecting
• The energy using in collecting, transporting and sorting metals to be recycled
• It can sometimes be more expensive/require more energy to recycle
Life cycle assessment
The life cycle assessment is used to work out the environmental impact of a product. It also
helps people to decide whether it is worthwhile to manufacture and recycle a product. LCAs
can also be used to see the effects of certain materials in products.
45
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Dynamic equilibrium
In some chemical reactions the products react to form the reactants. These are called
reversible reactions. These use a double arrow (⇌) to show that the reactants can form the
products and the products can form the reactants
Example:
Ammonium Chloride ⇌ Ammonia + Hydrogen
The concentrations of the reactants and products change during the reversible reaction.
Eventually, the concentrations of the reactants and products present becomes constant.
Although there is no net change, there is still movement of substances at a constant rate. This
is called a dynamic equilibrium. A dynamic equilibrium can only occur in a closed system,
where there is no loss of reactants or products. In an open system, gases could escape so
equilibrium would not be achieved.
The Haber process
The Haber process is a method to produce ammonia as a reversible reaction between nitrogen
(extracted from the air) and hydrogen (from natural gas) that can reach a dynamic
equilibrium.
N2+3H2 ⇌ 2NH3
The equilibrium position (ratio of the concentrations of reactants and products) can change
by changing the reaction conditions. Since the Haber process is used to produce Ammonia,
the reaction conditions that will form the most product (ammonia) as cheaply as possible are
favoured. In the Haber process, the optimal conditions are a temperature of 450°C, a pressure
of 200 atmospheres and an iron catalyst.
Factors affecting the position of a dynamic equilibrium
The equilibrium position can be affected by changes in temperature, pressure and
concentration. The position of equilibrium will shift to counteract (reduce the effects of) any
changes to the system.
Changes in temperature:
• An increase in temperature: the endothermic reaction is favoured (position of
equilibrium shifts in the endothermic direction) in order to reduce the temperature (by
taking in energy from the surroundings to cool them down)
• A decrease in temperature: the exothermic reaction is favoured (position of
equilibrium shifts in the exothermic direction) in order to increase the temperature (by
releasing energy to the surroundings to heat them down)
Changes in pressure:
• An increase in pressure: the equilibrium position shifts towards the reaction that
forms fewer gas molecules as this reduces the pressure
• A decrease in pressure: the equilibrium position shifts towards the reaction that forms
more gas molecules as this increases the pressure
46
Edexcel GCSE Chemistry: Unit 4 - Extracting Metals and Equilibria
written by Julie Maher, edited by Ken Tu
Changes in concentration:
• An increase in concentration: the equilibrium position shifts towards the direction of
the reaction that uses up the substance added (in order to reduce its concentration)
• A decrease in concentration: the equilibrium position shifts towards the direction of
the reaction that forms more of the substance that has been removed (in order to
increase its concentration)
47
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Structure of metals
All metals have metallic bonding. Metallic bonding is where the
positively charged metal ions are attractive to the negatively
charged sea of delocalised electrons. Delocalised means that the
electrons are free-flowing and are able to move freely in the
structure. Notice how the arrangement of atoms is layer by layer
and in a regular lattice structure. Lattice in the field of
Chemistry simply means a repeated 3D arrangement of atoms or
ions or molecules. In this case it’s an arrangement of atoms.
Transition metals
Most metals are transition metals and they can be found in the
middle block of the periodic table. In
the image to the right, within the red
ring lies all the transition metals. This is
to help you get an idea of where the
transition metals are located in the
periodic table.
Recall that most metals and those which
are transition metals have the following
properties:
• High melting point (because of
the strong electrostatic force of
attraction between the positive
ions and the sea of delocalised electrons, this requires high energy to overcome).
• High density.
• Form coloured compounds with certain other chemicals.
• Transition metals and transition metal compounds can be used as catalysts.
Some common examples of where transition metal compounds are coloured are blue copper
sulphate crystals (right image) which are often created when you do an experiment
practising the method of simple distillation, or red iron (III) oxide (left image).
48
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Corrosion of metals and rusting
Recall that metal + oxygen → metal oxide. Most metals undergo oxidation when oxygen is
present. As a result, a metal’s innate physical properties like strength can become weaker.
Iron is a common example of this, and you
are required to know this for GCSE
Chemistry Edexcel. There are two conditions
for the rusting of iron:
1. Presence of oxygen.
2. Reaction with water.
Corrosion is the first step. When the metal is
exposed to oxygen, it oxidises. The more
reactive the metal, the faster it corrodes. The second step is called rusting, and this is where
water is introduced and literally reacts with the metal oxide. Oftentimes, when corrosion
occurs, an outer layer of tarnish forms. Tarnish is essentially the thin layer of metal oxide.
This prevents oxygen from entering further into the structure and reacting with the atoms
inside of the structure.
In the case of iron, overall, this is the word equation:
iron + oxygen + water → hydrated iron(III) oxide
Hydrated iron (III) oxide is the orange-brown residue/substance you often see on rusted
objects.
Prevention of rusting
There are 3 main ways to prevent rusting:
1. Exclusion of oxygen.
This can be achieved by storing the metal in an inert/relatively
unreactive gas e.g. argon/nitrogen gas.
2.
Exclusion of water.
Often water vapour can exacerbate rusting so desiccant powder can be
used to absorb water vapour and prevent it from reacting with the
iron/metal. Desiccant powder is a hygroscopic substance which
basically means it readily stores/attracts water from its
surroundings. You may have encountered this in the form of
silica gel packets.
49
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
3.
Sacrificial protection.
Metals have different reactivities. So basically
sacrificial protection is where a more reactive
metal coats the actual (less reactive) metal
you would like to protect. This technique
prevents either air or water from reacting with
the metal you wish to protect. E.g. imagine we
wish to sacrificially protect a block of iron
from rusting. We would coat the iron with zinc
and the zinc would oxidise first because it is
more reactive than iron and it forms the
exterior of the block.
Electrolysis in the form of
electroplating
The premise of electroplating is to make an
object look more attractive or increase the
resistance to corrosion. E.g. silver/gold can
be electroplated onto cheaper ‘based
materials’, or chromium (which is resistant
to corrosion) can be electroplated onto iron
to make it resistant to corrosion for structural
purposes.
Assuming you already understand the
principle of electrolysis, this is what happens
in a step by step fashion to electroplate (coat)
the copper ring (diagram in the right) with
silver.
1. The electrolyte used is silver nitrate
solution (AgNO3(aq) composed of
Ag+ ions and NO3- ions).
2. When the direct current flows
through the electrodes, and the
electrolyte, the silver ions in the
electrolyte move to the copper ring.
3. The silver ions gain an electron (the silver ions are reduced): Ag+ + e- → Ag.
4. The silver atoms are deposited onto the copper ring leaving a silver finish.
5. However, at the silver anode, the silver atoms lose an electron, (the silver atoms are
oxidised into the electrolyte): Ag → Ag+ + e-. As the silver is ‘leaving’ the anode, you
will notice the anode getting thinner and losing mass. In contrast, the copper ring will
gain mass.
6.
Galvanising
As aforementioned, sacrificial protection can be used to coat metals with a more reactive
metal such that the metal below is prevented from oxidising. Galvanising is where zinc is
used to coat iron or steel objects. This can be done via electroplating or simply dipping the
object into molten (melted) zinc.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
What is an alloy?
An alloy is a mixture of two or more metals. Alloys are stronger/harder and less malleable
than pure metals.
Pure metals have the following
properties:
• High malleability
• All the atoms are the same size
• Less hard than an alloy
Alloys have the following properties:
• Low malleability
• The atoms are of different
sizes
• Harder than pure metals and
often stronger than pure metals
too.
Steel is an alloy of iron. In pure iron, the layers of iron atoms are arranged in a regular
structure and are all the same size so can easily slide over each other. However, in steel, the
atoms are different sizes so the regular layers are disrupted and this prevents the layers
from easily sliding over each other.
You also need to know about some metals and their alloy uses below. I have listed them out
in a table.
Metal
Alloy
Use of compound/metal
Aluminium Magnalium is an alloy of
aluminium and magnesium.
Magnalium is used in aircraft parts.
Copper
Brass is an alloy of copper
and zinc.
Copper is a strong electrical conductor so is used
in wires. Brass is used for locks, values, sockets
etc…
Gold
24 Carat or 24K is pure
gold
Gold is used in jewellery.
Stoichiometry/quantitative analysis
Stoichiometry is the field of Chemistry involving calculating the quantities of reactants and
products in any chemical reaction. Questions you can be asked are often to do with finding
out the concentration of an acid used in titration or to find the volume of gases produced in
this *insert random chemical reaction* etc…
Titration core practical
More specifically an acid-alkali titration is used to find the exact volume of an acid (often
known concentration) that will neutralise a specific volume of an alkali (of an unknown
concentration) or vice versa. Often, we then use this information to find the concentration of
the alkali. Additionally, as both the acid and alkali are colourless, an indicator is used to
show when alkali has been neutralised.
51
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Aim: By using the neutralisation reaction of hydrochloric acid (HCl of known concentration)
and sodium hydroxide (NaOH of known volume but unknown concentration), find the
concentration of the alkali NaOH. We will do this by first finding the volume of HCl used to
titrate the NaOH and find the moles of HCl. The chemical equation is below:
HCl + NaOH → NaCl + H2O
Apparatus:
• Volumetric pipette (25cm3)
• Conical flask
• White tile
• Burette
• Clamp stand
• Set the equipment as shown in the diagram
•
A good thing to note is that during titration, the
alkali is not placed in the burette as the glass
would etch and get scratched by the alkali. This
would change the volume in the volumetric
burette.
Method:
1. Using a volumetric pipette measure 25cm3 of NaOH solution and pour this into the
conical flask.
2. Fill the burette with HCl acid of a known concentration. Be sure the burette is
closed first and wear safety goggles.
3. Place the conical flask on the white tile.
4. Read the initial volume of the acid on the burette at eye-level/meniscus and record
the volume.
5. Add your indicator to the NaOH solution in the conical flask. You can use
phenolphthalein indicator which in alkali solutions to neutral pH will turn from
pink to colourless. Alternatively, the methyl orange indicator from alkali to neutral
pH will turn from yellow to orange.
6. Add the acid from the burette. Swirl the conical flask until you notice the
corresponding colour change.
In exams, they can often ask you what you should do to get a more accurate titre. The titre is
the total volume of the solution in the burette used after the titration is complete. For this you
can say the following:
•
•
•
Read the initial volume from the meniscus.
Towards the end of the titration, add the acid drop by drop.
Use a white tile so you can see the colour change more obviously.
7. Repeat steps 1-6 until you have concordant results. Concordant results means the
titre obtained is within one decimal place of each other e.g. your titres are 25.10cm3,
25.60cm3, 25.65cm3 you stop repeating as the last titre are within 0.1cm3 of each other.
Additionally, you would also count the first one as an anomaly so you would omit that in
your average of the titres obtained.
52
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
By the end of you experiment, you should have data that looks something like this:
Vol of acid
Rough
Attempt 1
Attempt 2
Attempt 3
Final Volume (cm3)
26.40
25.95
26.20
26.00
Initial Volume (cm3)
1.10
1.20
1.30
1.20
Total Titre (cm3)
25.30
24.75
24.90
24.80
The concordant result was 24.90cm3 and 24.80cm3. The average of the two concordant results
is 24.85cm3. So it takes 24.85cm3 of HCl of a known concentration to neutralise 25cm3 of
NaOH solution.
For Chemistry GCSE, you need to know how to carry out simple calculations using the
results of titrations to calculate an unknown concentration of a solution or an unknown
volume of solution required. So let's do that with our results from above and say that the
concentration of the HCl acid is 0.100 mol/dm3.
Remember that 1dm3 =
1000cm3 therefore, 1cm3 =
0.001dm3 or I like to do this
1cm3 = 1 ✕ 10-3 dm3
We need to calculate the
concentration of the NaOH solution!
HCl + NaOH (we’re trying to calculate the concentration of this!) → NaCl + H2O
Step 1:
Calculate the number of moles of the chemical with the known concentration and the titre
achieved. In our case the HCl concentration is 0.100mol/dm3 and its average titre is
24.85cm3.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Step 2:
Using the balanced equation, find the molar ratio and use this to calculate the moles of the
alkali.
HCl + NaOH → NaCl + H O
2
Step 3:
Lastly, calculate the concentration
(mol/dm3) of the alkali as you already know
its volume (25cm3).
You may also be asked to do similar
experiments but to calculate the volume of one
of the solutions instead.
In questions like these, you may also be asked
to find the concentration of mol/dm3 in g/dm3.
Below is the formula relating moles, mass,
relative molecular mass.
My Chemistry teacher always used to make us remember the
mnemonic for “No More Mr (nice guy)” to remember this
equation. M is the symbol for relative molecular mass.
r
54
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
So if our concentration is 0.0994mol/dm3, we just need to
convert the moles to grams. We can do this by dividing the
0.0944 by the relative formula mass of our chemical,
NaOH, which is 23.0 + 16.0 + 1 = 40.0. If g/dm3 is
confusing you think of it like this. Anhydrous NaOH
(without water) has a certain mass and it’s also soluble in
water. Now imagine you dissolve a certain mass of the
NaOH in water. That is where the concept of concentration
in g/dm3 comes from.
55
Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
More stoichiometry with atom economy
Atom economy is a measure of how many reactants form into useful products and this is
given in a percentage e.g. 89% atom economy.
2020 Chemistry Edexcel Q4)b)iv). Below is the formula to calculate atom economy.
Bang 4 marks.
Free marks.
Also if only one product is formed from a reaction, the atom economy is automatically
100% because all the reactants are used to produce one product.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Yield
Actual yield is the mass of the product after a chemical reaction has occurred taking into
account all the imperfect nature of a chemical reaction. Theoretical yield is the mass of the
product should everything react perfectly so the maximum possible amount of product that
could be formed in a chemical reaction. So in all cases the actual yield is less than the
theoretical yield.
There are 3 main reasons why the actual yield is less than the theoretical yield (this comes up
all the time in marks schemes:
• The reaction is incomplete so all the reactants have not finished reacting.
• There are practical losses during the reaction e.g. chemical residue left on the side
of the container during the reaction which hasn’t reacted.
• Unwanted, competing side reactions occurred which do not follow the desired
reaction pathway.
Let’s have a look at a calculating yield example. Below is the equation for percentage yield.
Percentage yield is a ratio in the form of the percentage of actual yield over theoretical
yield.
The question below is a follow up
question from the previous chemical
reaction.
So here is some context to the question: KCl is soluble in water. They are calculating the
percentage yield of KCl crystals which are in the evaporating basin.
You may have noticed that the percentage yield
is over 100% which does not make any sense. How can the actual yield of the KCl be greater
than the theoretical yield? This is true and this is because during the experiment H2O is also
produced and they are placed together in the evaporating basin but the H2O has not all been
evaporated off. Though most questions will not be tricky and do this.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
More stoichiometry with volume of gas & mass
Earlier we learned about chemical reactions with yield so we need to do the same with mass
and volume. Our aim is to find out how much mass of product will be produced from a given
chemical reaction and how much volume of product can be produced (assuming the product
is a gas).
Mass Calculation
Let’s go through an example.
Step 1: Calculate the moles of CaCO3.
Step 2: Calculate the molar ratio
between the moles of this chemical and the
product you must find. In this case CaCO3
and Na2CO3. The molar ratio of CaCO3 :
Na2CO3. Which is 1:1 as shown by the
chemical reaction, 1 mole of CaCO3
produces 1 mole of Na2CO3.
Step 3: Calculate the mass of the
product. In this case the
Na2CO3.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Volume calculation
Below we will go through an example where instead of finding the mass of a product, we are
asked to find the volume of a product. Before that I must explain the formula for the volume
of a gas.
Molar gas volume/molar volume is the
volume occupied by one mole of any gas.
E.g. N2 or O2 or NH3 gas, one mole of any
of these will occupy a certain volume. At
room temperature (278K) and room
pressure (100kPa) (RTP) the molar
volume is 24dm3 or 24,000cm3. So the
new formula triangle for gases at RTP is
this: see formula triangle on the right.
Let’s go through an example question.
Avogadro’s law
If pressure and temperature are the
same for gases in a reaction, equal
volumes of all gases have the same
number of molecules. E.g. if in the
Haber process:
N2(g) + 3H2(g) ⇌ 2NH3(g), 400cm3 of
N2 reacts with excess H2, then 800cm3 of
NH3 is produced as the molar ratio of
N2:NH3 is 1:2 or N2 and O2 if both are at
RTP, they will both occupy the same
volume: 24dm3.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
What is dynamic equilibrium?
Most chemical reactions only react in one direction e.g. Al2O3 + 6HCl → 2AlCl3 + 3H2O.
Dynamic equilibrium is a state where in a reversible reaction, e.g. the Haber Process
(production of ammonia, NH3 from N2 & H2) N2(g) + 3H2(g) ⇌ 2NH3(g), the rate of
forward reaction is equal to the rate of backward reaction. This reaction occurs
simultaneously. Additionally at dynamic equilibrium, the concentration/amount of each
reactant or product is constant. Dynamic equilibrium can only occur in a closed system so
no chemicals can escape from the reaction vessel.
Below shows two graphs for the reaction. The left one, the vertical axis is the rate of reaction
and as you can see, the rates of reaction are the same. The right one, the vertical axis shows
the concentration of reactants and products. The concentration of both the reactants and
products plateaus because dynamic equilibrium is reached: the rate of forward and backward
reaction are the same.
In industry, so like the actual Chemical engineers who make your products like shampoo,
dish soap etc, we may want to maximise the yield of a product. Imagine you’re a Chemical
engineer and run the management for TREsemmé (hair care product brand) and you need to
find a way to source the ingredients. You want your products to be produced with lower
costs and for the yield of your desired ingredients to be as high as possible. Next, we are
going to talk about the factors that affect the equilibrium position* and how industries tend
to compromise on factors to maximise yield of their desired chemical and cost.
Factors affecting dynamic equilibrium
The equilibrium position is which side the reversible reaction is more likely to favour.
Recalling from Unit 4, there are 3 factors that affect the equilibrium position and the catalyst
simply speeds up the rate at which this equilibrium position is achieved: temperature,
pressure, concentration of reactants. The rate at which dynamic equilibrium is achieved is
also known as the rate of attainment of equilibrium.
Just a recap:
Changes in temperature:
•
•
An increase in temperature: the endothermic reaction is favoured (position of
equilibrium shifts in the endothermic direction) in order to reduce the temperature (by
taking in energy from the surroundings to cool the reaction vessel down).
A decrease in temperature: the exothermic reaction is favoured (position of
equilibrium shifts in the exothermic direction) in order to increase the temperature (by
releasing energy to the surroundings to heat the reaction vessel down).
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Changes in pressure:
•
•
An increase in pressure: the equilibrium position shifts towards the reaction that
forms fewer gas molecules as this reduces the pressure in the reaction vessel.
A decrease in pressure: the equilibrium position shifts towards the reaction that forms
more gas molecules as this increases the pressure in the reaction vessel.
Changes in concentration:
•
•
An increase in concentration: the equilibrium position shifts towards the direction of
the reaction that uses up the substance added (in order to reduce its concentration)
A decrease in concentration: the equilibrium position shifts towards the direction of
the reaction that forms more of the substance that has been removed (in order to
increase its concentration)
Let’s look at an example question and how you should phase your
answer.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
The manufacturer wants to maximise the yield of SO . In 6 mark questions, we are often
looking to talk about 3 characteristics and give two points for each of those factors.
3
Step one: Read whether the reaction is exothermic or endothermic, then resolve a conclusion
for temperature and its drawbacks.
1. As the forward reaction is exothermic and reaction pathway B has a lower
temperature than A, the equilibrium position would be shifted further to the right
hence a higher yield of SO3.
2. This does mean that the rate of reaction to produce SO3 would decrease because
the lower temperature in B means less frequent collisions as more molecules have
less energy.
Step two: Talk about our second point - pressure.
3.
The product has fewer moles/volume than the reactants (when the reaction occurs
volume is directly proportional to the moles of gas in specified temperature and pressure) and
because the pressure in B is higher than A, the yield of SO3 is greater in reaction B.
4.
Because the pressure is higher in B, the rate of reaction increases as the particles are
closer together therefore the yield of SO3 is obtained faster because B has higher pressure
than A.
Step three: Talk about our last point - the catalyst.
5.
As B has a catalyst and A does not, the activation energy for the reaction to occur
decreases so equilibrium is reached faster in B than A/catalyst reduces the need for the
temperature to be higher.
6.
However, the catalyst does not cause the yield of SO3 to be affected.
Sometimes you may want to take into account that higher temperatures will be more costly.
So in different questions a drawback of high temperatures can be the increased cost of
production.
Fertilisers in the Haber process
Fertilisers help plants grow as they are mineral ions. They are made from elements, nitrogen,
phosphorus and potassium otherwise known as (NPK fertilisers). Below are some fertilisers
you need to know the formation of. There is a short cut diagram at the end :).
Ammonium nitrate (NH4NO3) is a soluble fertiliser. Below is the reaction equation:
ammonia + nitric acid → ammonium nitrate (a soluble salt)
NH3(aq) + HNO3(aq) → NH4NO3(aq)
The nitric acid is manufactured from ammonia too. I hope this highlights the great use of
ammonia. See reaction below:
ammonia + oxygen → nitric acid + water
NH3(g) + 2O2(g) → HNO3(aq) + H2O(l)
Ammonium sulphate is another soluble fertiliser. Below is the reaction equation:
ammonia + (dilute) sulfuric acid → ammonium sulphate
2NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq)
Ammonium sulphate is produced in a laboratory via titration and then crystallisation. We
also need to be able to compare the similarities and differences between ammonium sulphate
production in a laboratory and industrially.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Production of ammonium sulphate
Laboratory
Industrial
Batch process
Continuous process
Small scale
Large scale
Ammonia as a reactant is solution in the
laboratory process
Ammonia as a reactant is a gas in the
industrial process
Some similarities:
• Both processes are neutralisation reactions.
• Both use sulfuric acid.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
What are chemical cells and fuel cells?
A cell is something that provides a voltage (a potential difference). Using some physics
knowledge, charge flows from an area of high potential to an area of low potential. Electrons
are charge carriers. Ions are charge carriers too. When charged particles move from one area
to another, a current is produced. Both chemical cells and fuel cells are employed to
produce a voltage and by extension electricity.
Chemical cell (Daniell cell) AKA galvanic/voltaic cells
A simple chemical cell produces a voltage using reactants. One of the simplest versions of a
chemical cell is called a Daniell Cell and it requires the following:
• Two different metals, and by extension of different reactivity. The greater the
difference in reactivity the greater the potential difference (voltage) produced.
• A ‘salt bridge’ which permits dissolved, charged ions to pass from one solution to
another. In the Daniell cell, the salt bridge is filter paper soaked with concentrated
potassium nitrate solution so KNO3(aq) or the ions K+ and NO3-.
Both the ZnSO4 and CuSO4 solutions as well as the salt bridge act as an electrolyte enabling
the flow of charged ions.
At GCSE you are not required to know about the specific details of how a chemical cell
operates.
The overall reaction in the Daniell cell is this: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s).
As you can see it is a displacement reaction and generally displacement reactions are
exothermic. This means that energy is released to the surroundings. In this case, energy is
released to the surroundings in the form of electricity rather than heat.
Additionally in non-rechargeable chemical cells, the reactions stop and voltage stops being
produced when all the reactants are used up. However, in rechargeable chemical cells, the
chemical reaction can be reversed when sufficient external electrical current is supplied.
Just like a chemical cell, chemical batteries can exist where more than one chemical cell is
joined in series.
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Edexcel GCSE Chemistry: Unit 5 – Separate Chemistry I: Transition Metals, Alloys
and Corrosion, Quantitative Analysis, Dynamic Equilibria, Chemical Cells and Fuel
Cells
written by Ken Tu
Fuel cells
Whilst chemical cells store their reactants, fuel cells are supplied with reactants externally.
Fuel cells will continually provide a voltage for as long as fuel is being supplied.
Hydrogen-oxygen fuel cells
Hydrogen-oxygen fuel cells produce a
voltage.
1) A hydrogen fuel tank supplies the
reaction vessel with H2 gas.
2) The hydrogen atoms lose electrons
and become H+ ions.
3) The electrons from the reacted
hydrogen atoms travel through the
external circuit.
4) Once the H+ ions pass through the
membrane they react with O2 to
produce water.
5) Electrons are conducted through the
circuit hence charge is carried and a
current is produced.
Below I’ve listed pros and cons of using the hydrogen-oxygen fuel cell.
Pros
Cons
Once set up, fuel cells require very little maintenance.
Hydrogen is a hazard because
it’s highly flammable.
Water is produced as a waste product. If this fuel cell is
used on spacecrafts, the water could be used as drinking
water.
Storage of hydrogen is
difficult because it is a gas.
Fuel cells operate as long as reactants are supplied.
Fuel cells are expensive to
manufacture
Below I’ve listed pros and cons of using chemical cells.
Pros
Cons
Chemical cells operate as long as
reactants are supplied.
Chemical cells will need to be replaced / chemical
cells have a limited lifetime
Once used chemical cells cannot be used again or
need recharging
If on a spacecraft, chemical cells take up valuable
space.
65
Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
written by Mia Chadwick, edited by Ken Tu
Introduction to groups in the periodic table
Elements are positioned as a result of the chemical properties they hold - these chemical
properties are linked to the electronic configuration of the element. Hence, elements with the
same number of electrons in their outermost shell, and by extension therefore similar
chemical properties, are placed in the same group. This lets us use the table to predict the
chemical properties an element would demonstrate based on its position in the periodic table,
and we can classify these elements into specific groups, for example, Group 1 elements are
recognised as ‘The Alkali Metals’.
Group 1 - the alkali metals
The elements of group 1 are known as ‘The Alkali Metals’. They hold this name due to
the fact that once they react with water they form alkali solutions. The alkali metals all
have one electron each in their outer shell, hence they have similar chemical
properties. The group contains the elements shown in the image to the right.
They are soft metals and can be cut easily with a knife or blade - they become softer as you
go down the group. They are shiny and silver when first cut but then tarnish quickly from
exposure to air.
1. They have relatively low melting points - the melting and boiling points of the metals
decrease as you travel down the group.
2. This is due to the atoms getting bigger as their period number increases therefore, they
have more electron shells as we go down the group.
3. In the metallic lattice, the nuclei of the positive ions are further from the delocalised
electrons as the element’s period number increases therefore there is a weaker
electrostatic force of attraction so less energy is required to mechanically break apart
the metal lattice.
66
Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
written by Mia Chadwick, edited by Ken Tu
Reactions of the alkali metals with water
All of the group 1 metals react with water in the same way: they produce a metal hydroxide
and hydrogen, as follows:
group 1 metal + water ⟶ metal hydroxide + hydrogen
2X (s) + 2H2O (l) ⟶ 2XOH (aq) + H2 (g)
The important difference between the reactions of different alkali metals is how quickly they
happen. As you go down the group the more rapidly the reaction occurs, due to increased
reactivity. Sodium’s reaction with water is the most typically asked about and requires
remembering in detail due to the fact that it is in between both lithium and potassium (these are
the three alkali metal + water reactions you must memorise).
SODIUM:
sodium + water → sodium hydroxide + hydrogen
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
LITHIUM:
lithium + water → lithium hydroxide + hydrogen
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
POTASSIUM:
potassium + water → potassium hydroxide + hydrogen
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
What is the pattern in reactivity of group 1 metals?
As we have previously discovered, the metals become more reactive the further down the
group you travel. This is clearly demonstrated by the reactions of each metal with water
described above, as the further down the group, the more vigorous the reaction. Therefore, we
can predict that rubidium and caesium react even more violently than potassium, yet still
produce the corresponding metal hydroxide (rubidium hydroxide and caesium hydroxide).
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Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
written by Mia Chadwick, edited by Ken Tu
Pattern of reactivity explained
But what is the explanation for this pattern? When these metals react the metal atoms are losing
electrons and therefore form metal ions. The different ways in which these reactions occur are
dependent on how easy it is for the outer electron of the alkali metal to escape from the shell i.e. how strongly attracted to the nucleus of the original atom it is.
1. Moving down the group the amount of total electron shells increases. This means the
atoms with a larger number of electron shells are physically larger than those with lesser
shells in total, i.e. lithium is smaller in size than potassium.
2. As the atoms increase in size the outer electron is further from the positive nucleus
therefore is less strongly attracted by its nucleus
3. So can be lost more readily. As a result, the alkali metals get more reactive as you
descend the group.
Group 7 - the halogens
The elements of fluorine, chlorine, bromine, iodine and astatine all have 7 electrons in their
outer shell, therefore we place these elements in group 7 of the periodic table. As we previously
discovered, the number of electrons in an atom's outer shell determines how it will react
therefore these elements (the halogens) will react in a similar way to one another. The Halogens
are non-metals and are diatomic molecules e.g F2, Cl2, Br2 etc.
Physical state at room temperature colour
Fluorine gas
yellow
Chlorine gas
green
Bromine liquid
Red-brown liquid, orange/brown vapour
Iodine
Grey solid, purple vapour
solid
What are the physical properties of group 7 elements?
The melting and boiling points increase down the group, illustrated by the states at room
temperature of the elements shown in the table above.
1. The halogens are all covalent molecular substances and both the melting points and
boiling points increase as the molecule’s relative molecular mass increases.
2. As we descend down Group 7 (G7), the elements have more protons and therefore more
neutrons are found in the nucleus and as a result, the M increases accordingly.
r
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Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
written by Mia Chadwick, edited by Ken Tu
3. Due to more positive protons in the nucleus, the intermolecular forces of attraction
become stronger.
4. As a result, more energy must be put in to overcome the attractive forces and melt or
boil the elements.
5. The halogens are non-metallic elements and therefore poor conductors of heat and
electricity.
Astatine is below iodine in Group 7. As shown above, the colour of the halogens gets darker
as you go down the group. Iodine is a grey solid, and astatine is black which we could’ve
predicted from the above table and remembering to look for patterns in properties of groups of
elements.
Chemical tests for chlorine: To test for the presence of chlorine, we use damp litmus paper.
Chlorine gas will bleach the blue litmus paper white, although the paper might turn red briefly
before it’s bleached because acids are produced when chlorine comes into contact with water
and quickly forms HCl acid which is red in litmus paper but then gets neutralised.
Halogens + metallic elements
The halogens react with metals to form ionic compounds. An additional negatively
charged electron from the metal atom is donated to and attached to the outermost shell
of the halogen atom, therefore providing the halide ion with an overall negative charge
of -1.
However, the ionic compound formed will have varying numbers of halogen atoms, depending
on the valency of the metal. The reaction is less
vigorous as you descend the group, yet all halogen +
metal reactions form metal halide salts eg, NaCl or
MgBr .
2
Halogen + hydrogen
The halogens react with hydrogen to form hydrogen
halides: hydrogen chloride, hydrogen bromide and
hydrogen iodide. Hydrogen halides are all poisonous,
acidic gases and they are covalently bonded. In
addition, they are very soluble in water, reacting with
it to produce acid solutions. An example would be the popular hydrochloric acid. Hydrochloric
acid is formed when hydrogen chloride is dissolved in water.
How can displacement reactions in G7 explain the reactivity pattern?
A more reactive halogen can
displace a less reactive halogen in
an aqueous solution of its salt. If
you were to add chlorine solution to
a solution of colourless potassium bromide, you would observe a colour change from colourless
to orange as bromine is formed. This is due to the fact that chlorine is more reactive than
bromine and has displaced it from the potassium solution. This is because when a substance
is more reactive it will have a greater tendency to react in order to form an overall more
stable compound and no longer stand in its unreacted form. Although, if something is less
reactive it's more likely to either go back to or remain in its unreacted form - being the element.
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Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
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If we carried out the same experiment but used iodine instead of bromine we could observe a
colour change again, yet this time we would see a brown solution of iodine as chlorine would
displace it in the potassium compound. From this given information we can conclude that
chlorine is the most reactive out of the three, however which element comes next? To test this
we would compose the same experiment again but instead, use the colourless potassium iodide
and bromine water. Once the bromine water is added to the solution, a brown colouring will be
observed as the iodine is displaced by the more reactive bromine.
Chlorine solution, Cl2 Bromine solution, Br2
Iodine solution, I2
Potassium
chloride,
KCl
No change - the same
element
No change - the less
reactive bromine cannot
displace chlorine
No change - the less reactive
iodine cannot displace
chlorine
Potassium
bromide,
KBr
Solution changes from
colourless to orange
No change - the same
element
No change - the less reactive
iodine cannot displace
bromine
Potassium
iodide,
KI
Solution changes from
colourless to brown
Solution changes from
colourless to brown
No change - the same
element
From our above findings we can conclude that reactivity of
the halogens decreases as we descend the group.
Therefore, if we were to predict the reactivity of astatine we
could assume that it would be displaced by all of the above
elements, as they are all more reactive.
Reactivity of the halogens explained
1. The halogens react by gaining an electron in their
outer shell. We can therefore explain this decline in
reactivity as you travel down the group through many factors.
2. The outer shell becomes further away from the nucleus, therefore there is a lesser
attraction between the positively charged nucleus and the negatively charged electron
attempting to attach itself to the atom.
3. The electron shielding increases, again weakening the attractive force between the
nucleus and the incoming electron.
4. In conclusion, electrons are gained less easily as you descend the group so therefore
halogens become less reactive.
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Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
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Displacement reactions as redox reactions
The reactions we just studied where the colour of the solution changes
as a less reactive element is displaced by one of greater reactivity are
known as redox reactions as they are often discussed in terms of
oxidation and reduction.
If we were to look at the reaction between potassium bromide and
chlorine water we would be able to realise that the chlorine is acting as
an oxidising agent as chlorine is gaining electrons as it displaces the
bromine. The chlorine atoms gain an electron from the potassium
when they react and form the potassium chloride compound, with the
chloride ions carrying a -1 charge and the potassium ions carrying a +1
charge. Here the chlorine oxidises the Br- ions by taking electrons
away from them, whilst being reduced from Cl2 to 2Cl- ions.
Ionic equation:
Cl2(aq) + 2Br-(aq) → 2Cl-(aq) + Br2(aq)
Balanced half equation:
Cl2(aq) + 2e- → 2Cl-(aq), (reduction)
2Br-(aq) → Br2(aq) + 2e-, (oxidation)
Group 0 – the noble gases
The noble gases are known as group 0 on the periodic table
and they each contain 8 electrons within their valence
shell, meaning it is full - except helium which is in period
1 therefore only having 1 shell and as a result only requires
2 electrons in order to form a complete outer shell. The
noble gases are stable elements due to this property and are
therefore chemically inert (unreactive). They are all nonmetallic, exist as single atoms (monoatomic), colourless,
non-flammable gases at room temperature.
What are the functions of the noble gases?
The noble gases offer a variety of functions despite being
chemically inert, for example helium; because it is less
dense than air and is non-flammable, it is used to fill
balloons and airships since it is much less dense than air balloons filled with it float
upwards. In advertising signs, argon, xenon, and neon are all used. Argon is used to fill
electric light bulbs and provide an inert environment for welding. Low energy light bulbs use
argon in addition to traditional filament lamps; it shares the peculiar characteristic of the other
noble gases in that it glows brightly when a large potential difference is given to it while under
low pressure.
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Edexcel GCSE Chemistry: Unit 6 - Groups in the Periodic Table
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Patterns in the physical properties of the noble gases
There are patterns in the noble gases' physical characteristics, just like there are in the other
groups we've previously studied.
1. Firstly the melting and boiling points of the noble gases are very low, hence why they
are all gases at room temperature.
2. However, they exhibit an increase in boiling point as we move down the group due to
an increase in relative atomic mass.
3. This is because the atoms become larger as you move down the group - therefore a
greater number of protons and neutrons (both of which hold mass).
4. This then causes an increase in intermolecular forces between atoms, which raises the
energy required to overcome intermolecular forces in order to change state.
5. In addition to their low melting and boiling points, group 0 elements have low densities
because their individual atoms are widely spread apart as they have low boiling
points.
6. Therefore at room temperature, noble gases are typically in the gaseous state, but they
become denser as you descend the group.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
written by Ken Tu
What are rates of reaction and energy changes?
In this topic, you will learn about the different factors (concentration, surface area,
pressure, temperature, using a catalyst) that affect how fast a chemical reaction takes place.
You will also learn about exothermic and endothermic reactions that occur when bonds in
the reactants are broken down and when bonds are made in the form of the products.
Collision theory
For a reaction to take place, there must be sufficient activation energy between the two
chemicals/particles as they must collide with enough energy to break bonds. There are 5
conditions that affect the rate of reaction and they are the following: concentration, surface
area, pressure (only in reactions involving gaseous molecules), temperature, and catalyst.
Imagine this displacement reaction taking place, Mg(s) + H2SO4(aq) → MgSO4(aq) + H2 (g)
where in the solution, the hydrogen gas produces bubbles! In an exam if hydrogen gas is
produced and the question says what is observed say ‘bubbles form’
or you see ‘effervescence’ which means a fizz.
Concentration
The concentration of reactant is directly proportional to the rate of
reaction. If we increase the solution concentration this is what
happens to the rate of reaction.
1. As the concentration increases, the number of particles in
reaction per given volume increases. As a result, there are
more frequent collisions as particles are more likely to
collide.
2. The more frequent collisions lead to more successful
collisions (a successful collision is when the reactants react
with sufficient activation energy, and in the correct
orientation to react).
3. Overall, the reaction occurs faster and the reactant is used
up faster too.
Another thing to note is how a reaction (non-reversible) would look
like on a graph. The diagram on the right shows how as a reaction
occurs, the concentration of the reactants
decreases and concentration of the products
increases. The vertical axis is concentration.
Capital ‘M’ is shorthand for mol/dm3.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Temperature
The temperature in which a reaction occurs is directly proportional to the rate of reaction.
1. This increases the average speed of each of the particles in the reaction.
2. This leads to more frequent collisions and more successful collisions in a given
time.
3. Overall the reaction occurs faster in higher temperatures.
Surface area
Surface area is directly proportional to the rate of
reaction. E.g. in the reaction
Mg(s) + H2SO4(aq) → MgSO4(aq) + H2 (g), if the
Mg(s) is refined into powder, it would have a
higher surface area to volume ratio and the
reaction would occur faster. In this reaction
specifically you can notice this as the rate at which
the bubbles of hydrogen gas are produced in the
solution is faster. Or the solution will effervesce
more vigorously.
1. The greater the total surface area of the
chemical, the more available space there is
for reactants to collide with the reactant.
2. There are more sites for the reactions to
take place.
3. More frequent collisions, more
successful collisions in a given time.
4. Overall the rate of reaction increases
5. The opposite is true if you decrease the
surface area to volume ratio.
Pressure
With regards to the rate of a reaction (ignoring yield), a change in pressure is similar to how a
change in concentration will affect the reaction. Pressure is directly proportional to the rate
of reaction.
1. There are more particles in a
given volume
2. This leads to more frequent
collisions, and more successful
collisions
3. Overall the rate of reaction
increases.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Catalyst
A catalyst is a chemical used to speed up a
reaction, by reducing the activation
energy required for the reaction to take
place. The catalyst itself remains unchanged
chemically and also the catalyst does not
lose its mass. Activation energy means the
minimum amount of energy required for a
reaction to take place. When particles
collide, it requires a certain amount of
energy to break bonds. These broken bonds
will produce different chemicals and then
form new bonds which in turn create a new
chemical (the product).
The diagram on the right shows a reaction
profile in which a catalyst is used. As you
can see, the activation energy decreases. This means that it is more likely that the more
particles have sufficient energy to successfully react and as a result the rate of reaction
increases.
Additionally in Biology, you may have come across the term, biological catalyst. An example
of a biological catalyst is an enzyme and these are used in fermentation (the production of
alcohol by using yeast as the enzyme).
Exothermic and endothermic Reactions
After a reaction has taken place, the reaction can end up warmer or cooler than it initially
was. This is because when a reaction takes place, bonds are broken which require energy, and
then bonds are made which releases energy in the form of heat. Depending on whether the
total energy required to break the bonds is greater than or less than the energy released, the
overall reaction is called exothermic or endothermic.
Endothermic reactions occur when heat energy is taken in. Endothermic reactions are
cooler after the reaction occurred.
Exothermic reactions occur when heat energy is released/given out. Exothermic reactions
are hotter after the reaction occurred.
When reactions occur, bonds must be broken and then be made. This is further explained
using the reaction example below.
Mg(s) + H2SO4(aq) → MgSO4(aq) + H2 (g)
For more context, the H2SO4 molecule looks like the image on the right.
The two hydrogen atoms in the H2SO4 must break its bond with the
oxygen atom. Then bonds are MADE between those two broken
hydrogen atoms. This explanation is not completely thorough but I hope
it explains the concept of how reactants react and how the product is then
formed. The act of breaking bonds is endothermic whilst the act of making bonds is
exothermic. A mnemonic to memorise this is Bendo Mexo.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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The diagrams on the right show the reaction
profile for an exothermic and endothermic
reaction. A reaction profile models the change
in chemical energy (energy stored between
bonds) after a reaction takes place. Notice the
following in the types of reactions:
Endothermic reactions - the chemical energy
of the products is greater than the chemical
energy of the reactants, so the temperature
decreases.
Exothermic reactions - the chemical energy of
the products is less than the chemical energy of
the reactants, so the temperature increases.
In the endothermic and exothermic reactions,
the vertical axis represents energy, but it should
be more precise and say chemical energy.
Energy cannot be created or destroyed as a
universal law in the First Law of Thermodynamics (this part is quite pretentious but I want to
be sure you understand) so the vertical axis represents the energy stored between the
bonds of the reactants or the products. Hence why in exothermic reactions, yes the
solution gets hotter but the ‘energy profile’ shows lower in the products because energy
stored between the bonds in the product is lower than in the reactants.
Core practical: investigating how the conditions of a reaction affect its rate
There are two core practicals you should be aware of in GCSE Edexcel Chemistry. One
method alters the surface area of a reactant method using marble chips (calcium carbonate,
CaCO3) and hydrochloric acid (HCl). The second method alters the temperature in fixed
intervals of the reaction using sodium thiosulfate (Na2S2O3) and hydrochloric acid (HCl).
Method 1: marble chips and HCl
Aim: To investigate how a change in surface area of marble chips affects the rate of reaction.
The rate of reaction will be measured using a measuring cylinder. The reaction will produce
carbon dioxide bubbles and the volume of the bubbles produced will be measured.
calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
Apparatus:
• Conical flask
• Delivery tube
• Measuring cylinder
• Water trough
• Clamp
• Stopwatch
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Method:
1. Set the apparatus as shown in the diagram below. The measuring cylinder is placed
upside down but inside the water trough and the dilute HCl can be added to the
conical flask with the marble chips. The additional test tube looking-thing is
unnecessary. To start the reaction you will need to add you dilute HCl and close the
bung immediately after.
Measure a known volume (e.g. 50cm3) of dilute hydrochloric acid.
Add a known mass (e.g. 5g) of marble chips to the conical flask.
Add the dilute HCl to the conical flask and immediately place the bung on top.
Start the stop watch and read the total volume of gas produced on the measuring
cylinder in a fixed time interval (e.g. reading every 30 seconds) until no more new
bubbles are produced.
6. Repeat steps 1-5 but use marble chips that are smaller. (Smaller marble chips should
produce bubbles faster.
Now if we drew a graph of the volume of bubbles produced vs time graph what would that
look like?
2.
3.
4.
5.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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The diagram illustrates how much
volume of CO2 is produced per
minute. Note the following things:
1. The marble chips with a higher
surface area reach the
maximum volume produced
faster than the marble chips
with a lower surface area.
2. The maximum volume of
carbon dioxide produced for
both reactions is the same.
This is because the
mass/volume of reactant used
is the same in each reaction. So
the same amount (volume) of
product (CO2) is produced
each time.
During experiments like this one, you could be asked to find the
rate of reaction at a specific time. To do this you have to draw a
tangent to the curve and calculate the gradient of this tangent. As
represented by the grey line in the diagram on the right.
Sodium thiosulfate and HCl
Aim: To investigate how an increase in temperature affects the
rate of reaction. The idea is that we will warm our sodium
thiosulfate solution in a water bath altering the temperature of it.
We will place this solution on a white tile with a black cross on it.
The solution is clear and beneath the conical flask we can see the black ‘X’ cross. Once we
add the HCl, it reacts and sulphur is the product. Sulphur is a yellow precipitate and is not
clear so we will no longer be able to see the black cross. We will measure the time passed
until we can no longer see the black ‘X’ cross.
sodium thiosulfate + hydrochloric acid → sodium chloride + water + sulphur dioxide + sulphur
Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + SO2(g) + S(s)
Apparatus:
• White tile
• Beaker
• Black whiteboard marker
• Water bath
• Conical flask
• Thermometer
• Stopwatch
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Method:
1. Place a known volume (e.g. 50 cm3) of sodium thiosulfate solution into a conical
flask.
2. Measure a known volume (e.g. 5 cm3) of dilute HCl acid in a test tube.
3. Warm both solutions to the same temperature in a water bath (e.g. 20°C). Using a
thermometer, record this temperature.
4. On a white tile, draw a ‘X’ cross using the black marker.
5. Take the conical flask out and place it on top of the white tile so that the cross is
visible.
6. Pour the HCl into the sodium thiosulphate solution and start the stopwatch.
7. Once you can no longer see the black cross, stop the stopwatch and record the time
taken.
8. Repeat steps 3-7 with different temperatures in fixed intervals (e.g. 5°C intervals)
In the end you should have data that looks similar to this:
Average temperature (°C)
Time for the cross to disappear (s)
20
184
25
169
30
130
35
112
40
90
Heat energy changes in common reactions
You need to know whether a reaction is exothermic or endothermic for the following
reactions:
a. Salts dissolving in water. A salt is an ionic compound consisting of oppositely
charged ions. These are soluble in water.
E.g. NaCl(s) + H2O(l) → Na+(aq) + Cl-(aq) remembering that (aq) means dissolved in
water!
b. Neutralisation reactions - reaction between an acid and a base:
e.g. hydrochloric acid + sodium hydroxide → sodium chloride + water.
c. Displacement reactions - reaction between a metal and a compound of a less reactive
metal:
e.g. magnesium + copper sulphate → magnesium sulphate + copper. Or a reaction of a
halogen and a compound of a less reactive halogen:
e.g. bromine + potassium iodide → iodine + potassium bromide.
d. Precipitation reactions - where an insoluble product forms from two solution:
E.g. barium chloride(aq) + sodium carbonate(aq) → barium carbonate(s) + sodium
chloride(aq).
•
•
Both displacement and neutralisation reactions are always exothermic.
Precipitation reactions & salts dissolving in water can be either exothermic or
endothermic depending on the chemical or salt used.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Calculating enthalpy change (energy change)
Just think of enthalpy change as the fancy word for energy change. You are also required to
know how to calculate the energy change of a reaction. What is the chemical energy change
of the bonds before the reactions vs after the reaction. In stoichiometry, the fancy word for
the area of Chemistry involving the quantitative stuff e.g. mass, volume etc… between the
reactants and products, to calculate energy change, we use kilojoules per mol (kJ/mol or
kJmol-1). A mole is 6.02 ✕ 1023 amount of something. Bond energy, measured in kJ/mol, is
the total amount of energy stored between the bonds of one mole of a substance. E.g. in the
example question below C-H is having a bond energy of 412 kJmol-1 means one mole of C-H
bonds has an energy store of 412 kJ.
Example question
Energy change = sum of the bonds broken - sum of the bonds made.
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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Step 1: Calculate the energy required to break the all bonds (or the left hand side)
Step 2: Calculate the energy required to make the all bonds (or the right hand side)
Step 3: Energy Change = bonds broken - bonds made
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Edexcel GCSE Chemistry: Unit 7 - Rates of Reaction and Energy Changes
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I hope that the concept of energy calculations was clear. Additionally, it is important to note
that in this reaction, the energy change overall was negative. This means that the reaction
was overall, exothermic because more energy was released to the surroundings than taken in.
If the reaction shows a positive energy change, then the reaction was overall endothermic
because less energy was released to the surroundings as more energy was taken in from the
surroundings.
Again ‘energy change’ SPECIFICALLY means chemical energy change so the energy is held
in bonds.
Another cheeky thing I should add is that this concept of calculating energy change is also
done at A level. And this was an A level question but exactly the kind of question the GCSE
could ask of you - though probably with less chemicals.
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Edexcel GCSE Chemistry: Unit 8 - Fuels, Earth and Atmospheric Sciences
written by Mary Mullagiri, edited by Ken Tu
Hydrocarbons, fuels and crude oil
Hydrocarbons are chemical compounds that only contain carbon and hydrogen, hence the
name, and are known as organic compounds because carbon is the foundation for all living
material. These hydrocarbons are found in a complex mixture known as crude oil.
Crude oil is a sticky black oil made of lots of hydrocarbons, mostly alkanes and alkenes,
where the C atoms are all either in chains or rings at different lengths and come from
mining into the earth. This is why it’s such a finite resource (it’s used up faster than it’s
formed) and also a very important one: it provides fuels and feedstock for various
industries. Feedstock is another way of saying raw material which provides reactants in
industrial reactions. Crude oil is also a non-renewable resource.
Because it’s a mixture, chemists need to separate these hydrocarbons into their different
lengths in order to obtain the fractions which are higher in demand e.g. more people in the
world need petrol for fuel so demand for this resource is high hence why crude oil is cracked.
The method they use, and you should know, is called fractional distillation.
As in the diagram below:
1. Crude oil is pumped into a giant furnace called a fractionating column which is
being heated from the bottom.
2. Smaller hydrocarbons boil easily in this hot environment, so they become gaseous
and rise up higher.
3. Eventually, they reach a cool enough temperature and condense onto a metal sheet.
They drip through the pipes at different heights and are then collected separately.
4. Each section in this diagram is known as a fraction. It is called such because it is a
fraction of the original oil.
5. Each fraction contains a group of hydrocarbons of similar lengths.
You must also be able to recall the uses of each fraction:
• Gases are used for domestic heating and cooking
• Petrol, or Nafta, is used for car fuel
• Kerosene, or paraffin, is used for fuel for aircraft
• Diesel oil is used in cars and trains
• Fuel oil is for ships and power stations
• Bitumen is used for surface roads and roofs
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Edexcel GCSE Chemistry: Unit 8 - Fuels, Earth and Atmospheric Sciences
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Remember the mnemonic: Good Penguins Keep Diving For Bass
The entire reason why this method works
to separate the different fractions is
because of how many C and H atoms are
in a hydrocarbon. The longer the
hydrocarbon, the higher the boiling
point and amount of energy required to
vaporise it. Essentially, as the length of
the chain increases, you need more
energy to break the many
intermolecular forces. Moreover, it is
easier to ignite a shorter hydrocarbon
than it is a longer one, which is why
gases are so flammable. It is also why
gases are at the top of the fractionating
column and liquids at the bottom. Most
of these hydrocarbons are part of a
specific homologous series called
alkanes.
Homologous series - alkanes and
alkenes
A homologous series is no different to
the binomial naming structure you see in biology - it tells us which ‘species’ a compound is
classified into and each series contains its own specific characteristics. There are four that
you must know for the exam:
1. All compounds in a homologous
series have the same general
formula. It’s like the same
algebraic formula for the
compounds.
2. All alkanes, like the ones shown
to the side, have the general
formula of CnH2n+2 where n
stands for the number of C atoms
which are in the hydrocarbon. All
alkenes have the general formula
of CnH2n.
3. An alkane only has C-C single
bonds, hence the single lines in
the diagram.
4. An alkene will have at least one
C=C double bond.
5. Each compound in the
homologous series will differ by
CH2 from their neighbouring compounds. Basically, this means that CH4 (methane)
and C2H6 (ethane) differ by one CH2 group; if you add a CH2 to methane, you get
ethane.
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6. Alkanes are called saturated hydrocarbons because their molecules are completely
joined by single bonds.
7. Whereas alkenes are unsaturated because they have at least one double bond. This
means alkenes are more reactive because they’re able to have more reactions than
alkanes - this is why you can make polymers, or repeating chains, out of alkenes, but
not out of alkanes.
Trend in physical properties of alkanes
Within the homologous series, you have a trend in physical properties as the chain length
increases. This is most commonly seen in an increase in boiling points when increasing the
chain length - more molecules means stronger intermolecular forces (IMF), which means
you need more energy to overcome IMFs and break the bonds, meaning the boiling point
increases.
Finally, all compounds in a
homologous series share
similar chemical properties
because they are classified in
the same series. For example,
alkenes will turn brown
coloured bromine water
colourless while alkanes won’t
do anything. Both alkanes and
alkenes are flammable and both
are very volatile, becoming
gaseous at room temperature
very quickly.
Complete and incomplete combustion
Complete combustion of a fuel simply means that all the alkanes and alkenes are converted
into CO2 and H2O when they’re fully burnt in a sufficient supply of O2. Energy is also
given out, mostly as heat, some as light. Let’s take the example of propane.
C3H8 + 5O2 → 3CO2 + 4H2O
All the carbon in the propane was converted into carbon dioxide by reacting with the oxygen.
The hydrogens were bumped off to form water when they also reacted with excess oxygen.
On the other hand, the incomplete combustion of hydrocarbons can form carbon and
carbon monoxide along with water. Carbon dioxide is not formed because there is an
insufficient supply of O2. Let’s take another look at that example:
C3H8 + 3O2 → 2CO + C + 4H2O
During incomplete combustion, there isn’t enough O2 to react with the hydrocarbon, forming
the toxic gas CO, or carbon monoxide. This gas is colourless and odourless, and carbon
monoxide binds to haemoglobin in your blood. This new chemical (carboxyhemoglobin) in
your blood doesn’t not carry sufficient oxygen in your body for respiration. This means your
cells don’t get enough oxygen for respiration and could die. This kills the tissue and then
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causes death if you don’t get enough oxygen in time. The carbon gets dispersed as fine black
powder called soot.
Soot can build up and block things we use all the time, like
ventilators, boilers, fireplaces and other appliances. The
problem is that this soot can cause respiratory problems by
disturbing the lining of our airways. It can also blacken
buildings when released into the atmosphere during industrial
processes.
Impurities and the atmosphere
Realistically, crude oil is not just a mixture of hydrocarbons, but also contains impurities as
a result of mining. When burning fuels from fractional distillation in the industry, other
pollutants are formed. The most important ones are sulphur oxides and among these,
sulphur dioxide. When burned, sulphur dioxide rises up and is released as chemical waste
along with soot - you can see this most clearly in industrial chimneys and plumes of black
smoke. The sulphur dioxide reacts with droplets in clouds to create sulphurous acid and due
to the presence of oxygen, this acid fully forms sulphuric acid. You can see the formation in
this series of symbol equations:
S + O2 → SO2
SO2 + H2O → H2SO3
H2SO3 + ½O2 → H2SO4
Once formed, these two acids mix in the clouds and fall as acid rain.
Engines and nitrous oxides
Fuels are also used within cars and engines. When they’re burnt, the engines become hotter
and hotter, increasing the internal temperature and causing nitrogen and oxygen from the
atmosphere to react together. This forms nitrogen oxides which are all pollutants and form
smog in the atmosphere as well as causing acid rain and health problems such as bronchitis
due to their toxic nature.
N2 + O2 → 2NO
As a result, some manufacturers began creating cars which run
on hydrogen instead of petrol, which produces water instead
of nitrous oxides. Of course, the advantages are that hydrogen
can be made in many different ways and is a renewable source
which means there won’t be a danger of it ever running out.
Hydrogen cars are also sustainable and won’t cause damage to
buildings or people, but the downsides are that they’re
incredibly expensive and the technology has been known to
cause serious shock risks as well as the flammability of the
fuel. A hydrogen car works by using the energy produced from
reacting oxygen with hydrogen, meaning using two extremely reactive gases to power a car.
Also, it’s harder to store hydrogen because you don’t get as many atoms per volume with a
gas as you would with a liquid.
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Cracking
The fractions which produce fuels are all
different types. Petrol, kerosene and diesel are
all non-renewable fossil fuels, meaning they
come from mining and are limited in supply.
Methane is also a non-renewable fossil fuel,
albeit a smaller one, but it’s found in natural
gases instead of the ground. If you’ve ever
been near someone who farted, you’ll have a
decent idea of what it smells like.
As you can see, there are various lengths of
alkanes, but there are longer chained ones
than there are short ones. You can make
smaller hydrocarbons from large alkanes,
which is known as cracking. Cracking is the
process where longer alkanes are broken
down, or cracked, into smaller ones, some
of which are other alkanes while others are
unsaturated alkenes, meaning they have a
C=C double bond instead of only single
bonds.
The reason why we do this is because few
small hydrocarbons are formed naturally,
however, these ones are more useful than
longer hydrocarbons. This is necessary in
order to meet the commercial demand for
fuels, since larger ones are only seldom used.
It also produces more alkenes than what
you’d normally get, which are essential for feedstock.
Earth and atmospheric science
Scientists believe that when the earth was formed billions of years ago, the atmosphere was
formed by volcanic activity on the surface. This is because the early atmosphere contained
little to no oxygen while containing lots of carbon dioxide. Water vapour was also
incredibly high while other gases appeared in small amounts.
Scientists analysed rock formations that date back to the earth’s formation and found that
these rocks contain iron compounds. What’s special about these compounds is that they
decompose if oxygen is present and would have only been able to form if there was a lack
of oxygen.
Because of the volcanic activity, they also believe that the atmosphere contained large
amounts of CO2 and water vapour at that time since that’s primarily what volcanoes
happen to release. You could probably compare that early atmosphere to the one we have
today. Currently, the atmosphere is formed of roughly 78% to 80% nitrogen, 21% oxygen
and finally the rest is made of other naturally occurring gases like argon and CO2.
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1. Over time, all the condensation in the world formed the oceans. It did this by
raining a lot since all that steam from volcanoes cooled in the air and came back
down as rain.
2. Carbon dioxide also decreased later on because it
got trapped in the rain and dissolved into the
oceans over time and scientists can refer to this
special blend of ocean, carbon dioxide and
anything else that might have been living as an
organic “soup”.
3. Within this soup, primitive plants, which popped
up sometime after the volcanoes, began to slowly
use CO2 and photosynthesis oxygen, increasing
its abundance gradually. Photosynthesis is the
process where plants make glucose and oxygen
from water, sunlight and CO2 gas. You’ll need to remember the formula as well, so
here it is:
6CO2 + 6H2O → C6H12O6 + 6O2
Side note: Plants only photosynthesis when light is available since that’s the main form of
energy they use to complete the process. They’ll revert to respiration whenever they don’t
have light or enough CO2 in order to survive.
Over time, the accumulation of plant and algae activity caused more oxygen to be released
into the atmosphere and decreased the amount of carbon dioxide therein over another few
billions of years.
Speaking of oxygen, the way you test for it is quite simple. You take a glowing splint and
hold it in a test tube filled with an unknown gas. If the splint reignites, you know it’s
oxygen because O gas is essential for combustion, i.e. burning stuff.
2
The greenhouse effect
Carbon dioxide, methane and water vapour
are all greenhouse gases. A greenhouse gas
is responsible for absorbing heat energy
from the earth and dispersing it evenly in all
directions. This causes the earth to keep
warm and gradually increases its
temperature, known as the infamous
greenhouse effect.
Humans, AKA you, also contribute to this
effect because as we burn fossil fuels and
release large plumes of smoke from our
factories, cars, homes and other such
wonderful inventions, we end up increasing
the amount of CO2 in the atmosphere. We
cause more and more greenhouse gases to be
released into the atmosphere with no real
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place to go and so cause the earth to become more and more unstable in terms of climate.
We don’t even know how bad things are because the measurements taken at these times and
locations are uncertain - we can’t fully trust how valid these sources are since half of them
were aimed at making it seem like a tiny detail to protect our way of life.
Global warming’s impact on climate change
As the diagram shows, ever since we began to increase
our activity, the average temperature has increased in
only the span of ~2000 years, give or take. This
drastically changes the climate by causing glaciers and
ice to melt into the oceans, which raises sea levels. As
sea levels rise, water ends up being in places where it’s
unhelpful which changes the rainfall patterns and
affects flooding and droughts in different areas. All
these factors culminate in habitat changes and end up
collapsing the ecosystems we depend on for various
things.
If we continue to live as we have been and pump out as
much pollution as we do, eventually the earth will not
be able to sustain us. These dominoes will fall down and we won’t be able to come out of it,
but there is hope: International treaties are in place to reduce greenhouse gas emissions
which help reduce the damage long-term. Right now, there are ways to mitigate climate
change, though they are a bit expensive.
In agriculture, we can use crop rotation to plant different crops at different times of the year
to help adapt to newer climates. In cities, we can build irrigation systems to help with
drought season and we can build flood defences to stop flooding from becoming a problem
later down the line. These plans are all very large-scale and can also cause harm to the
ecosystems they’re replacing, but it’s necessary to cope with what we’ve already done.
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Hydrocarbons and homologous series
Organic chemistry is the study of chemical compounds that contain carbon. Carbon is
‘organic’ because it is contained in all living organisms. Nowadays, chemists are finding
ways to make carbon compounds synthetically, such as pharmaceutical drugs or plastic
polymers.
Hydrocarbons are the foundation of organic chemistry. They are molecules that contain only
hydrogen and carbon atoms. Most hydrocarbons consist of a carbon chain, and certain groups
of atoms can attach themselves to the chain. These groups of atoms are called functional
groups and determine the chemical properties of the hydrocarbon. A group of molecules with
the same functional group can be classified as a homologous series. The molecular formulae
of adjacent members in a homologous series differ by a CH2 unit.
Different atoms can form a different number of bonds with other atoms:
• Hydrogen atoms can form up to one bond.
• Oxygen atoms can form up to two bonds.
• Nitrogen atoms can form up to three bonds.
• Carbon atoms can form up to four bonds.
Alkanes
Alkanes are the simplest homologous series and consist of a carbon chain where all of the
carbon atoms are surrounded by hydrogen atoms. Since hydrogen atoms surround the chain,
alkanes do not have a distinctive functional group. In addition, alkanes can be described as
saturated because all of the atoms only form single bonds.
Alkanes have the general formula CnH2n+2, where n denotes the number of carbon atoms.
The reason for this general formula is that each carbon atom is covalently bonded to two
hydrogen atoms on both sides. There are also two hydrogen atoms bonded at the ends of the
carbon chain.
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As with all other organic compounds, alkanes must be given names. The nomenclature of
organic chemistry was devised by IUPAC (International Union of Pure and Applied
Chemistry), an international organisation of chemists that draws up standards so that
chemists throughout the world use the same universal conventions of chemistry.
For alkanes, IUPAC has given the following rules:
• The name must have the suffix -ane, which shows that alkanes only have single
bonds.
• The name must have a prefix, which shows how many carbon atoms are in the chain.
• meth- indicates 1, eth- indicates 2, prop- indicates 3, but- indicates 4, pent- indicates
5, and hex- indicates 6 (for GCSE Chemistry, you are not expected to know prefixes
beyond hex-).
Looking at the diagram on the previous page, you can see the displayed formula for each
alkane with its corresponding name and molecular formula below.
Alkenes
Alkenes are another homologous series and look very similar to alkanes. However, the
fundamental difference between them is that alkenes have at least one double bond between
two carbon atoms in the chain, whereas alkanes only have single bonds. Due to the double
bond, alkenes can be described as unsaturated.
Alkenes are more reactive than alkanes because the double bond can open to form two single
bonds. The two carbon atoms involved in the double bond now form a single bond, and the
other single bond can be formed with a completely different molecule.
The general formula of an alkene is CnH2n, where n denotes the number of carbon atoms.
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For alkenes, IUPAC has given the following rules:
• The name must have the suffix -ene, which shows that alkenes have one double bond.
• The name must have a prefix - the same prefixes used for alkanes can also be used for
alkenes (and all other organic compounds).
Looking at the diagram on the previous page, you might have noticed that there is no
methene, and the alkene series begins with ethene instead. This is because two carbon atoms
are needed for a double bond to form, yet methene would hypothetically only have one
carbon atom and so cannot form the double bond. Therefore, methene does not exist.
You might have also noticed that butene can be displayed in two different ways, as the
double bond can form in two different places in the carbon chain. This phenomenon is called
isomerism - isomers are molecules that have the same molecular formula but have
different displayed formulae.
Underneath the isomers, you can see that they have two different names - but-1-ene and but2-ene. The number in between the names is called a locant and denotes which carbon atom
the double bond starts to form. The order of locants starts from the ends of the carbon chain.
• Looking at the first isomer of butene, the locants would range from 1 to 4 as butene
has four carbon atoms in the chain. The double bond is located at either carbon 1 or
carbon 3, depending on what end of the chain you’re reading from. This means that
the name of the isomer could be either but-1-ene or but-3-ene. However, we use the
smallest locant possible, hence the name is but-1-ene, not but-3-ene.
• In the second isomer, the double bond is located at carbon 2 from both ends of the
chain, therefore the name of this isomer is but-2-ene.
Reactions of hydrocarbons with bromine water
One of the quickest and most common ways of distinguishing
between alkanes and alkenes in the lab is by adding bromine
water. Bromine water is orange in colour.
Alkanes do not react with bromine water as all of the carbon and
hydrogen atoms have formed the maximum number of covalent
bonds. Therefore, bromine water will remain orange when added
to an alkane.
However, alkenes can open their double bond to form two single
bonds, so they can form single bonds with bromine molecules.
Therefore, bromine water will decolourise (become colourless)
when added to an alkene.
The reaction between an alkene and bromine water can be
described as an addition reaction - two smaller molecules react to form one larger molecule
with no other products. The product of this reaction is a dibromoalkane. Bromo- indicates that
bromine is present, di- indicates that two bromine atoms are present, and -ane indicates that
only single bonds are present.
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Let’s look at the reactions of different alkenes with bromine water:
Ethene opens its double bond to form single bonds with two bromine atoms. Since a bromine
atom is bonded to each carbon atom, two locants are needed in the name (hence 1,2dibromoethane).
Propene opens its double bond to form single bonds with two bromine atoms. Since a
bromine atom is bonded to two carbon atoms in the chain, two locants are needed in the name
(hence 1,2-dibromopropane).
Since butene has two isomers, dibromobutane will also have two isomers. But-1-ene will
form 1,2-dibromobutane, whereas but-2-ene will form 2,3-dibromobutane. This is because
the bromine atoms are bonded to the two carbon atoms that formed the original double bond.
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Alcohols
Alcohols are a homologous series of molecules that contain a hydroxyl (-OH) functional
group. They have the general formula CnH2n+1OH, where n denotes the number of carbon
atoms.
Regarding IUPAC nomenclature rules, all alcohols have the suffix -ol and use the same
prefixes. They can also be described as saturated, like alkanes, because they only have single
bonds.
Propanol and butanol have two isomers, as shown below:
Ethanol also has two isomers, but one of its isomers, methoxymethane (or dimethyl ether), is
not an alcohol - it is a different type of organic compound. Because of this reason, you will
not be required to know or draw the displayed formula of methoxymethane for GCSE
Chemistry.
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Dehydration reactions with alcohols
If you heat a mixture of an alcohol and an acid catalyst (usually sulfuric acid), an alkene and
water are formed in a dehydration reaction. It’s ‘dehydrated’ because water is lost from the
alcohol for each alkene molecule formed.
Here is an example with
ethanol:
•
•
•
Propanol would be dehydrated to form propene and water.
Butanol would be dehydrated to form but-1-ene, but-2-ene and water.
Even though methene does not exist, methanol can be dehydrated. However, it does
not form an alkene and water - instead, it forms methoxymethane.
The production of ethanol
Ethanol can be manufactured using a hydration reaction with ethene. In this reaction, ethene
produced by cracking is heated with steam in the presence of a phosphoric acid catalyst. The
equation for the reaction is:
C2H4 (g) + H2O(g) → C2H5OH(g)
This reaction typically uses a temperature of around 300 °C and a pressure of around 60-70
atmospheres.
Notice that ethanol is the only product, so this hydration reaction is also an addition reaction.
The process is continuous - as long as ethene and steam are fed into one end of the reaction
vessel, ethanol will be produced. These features make it an efficient process with a 100%
atom economy. However, ethene is made from crude oil, which is a non-renewable resource.
The reaction also uses a lot of energy, which can make it quite expensive.
Another way of producing ethanol is to break down carbohydrates with yeast. This is called a
fermentation reaction. For example, with glucose:
C6H12O6 (aq) → 2C2H5OH(aq) + 2CO2 (aq)
The carbohydrates can come from any source, but sugar cane and sugar beet plants are often
used. The yeast cells contain an enzyme called zymase, which acts as a biological catalyst
and speeds up the fermentation reaction without being used up.
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The method of fermentation is as follows:
1. Mix yeast and a solution of glucose in a clean container. Seal the container and leave
it in a warm place.
2. Keep the mixture between 30 °C and 40 °C. Fermentation happens fastest within
these temperatures. At lower temperatures, the reaction slows down. At higher
temperatures, the zymase in the yeast denatures and the reaction would stop.
3. It’s important to keep the mixture in anaerobic conditions, with no oxygen. Oxygen
would convert the ethanol into ethanoic acid, which is what you get in vinegar.
4. When the concentration of alcohol reaches about 10%-20%, the fermentation reaction
stops because the yeast gets killed off by the alcohol. The yeast falls to the bottom of
the container and the ethanol solution is collected from the top.
Fermentation produces a dilute concentration
of ethanol. The fermented mixture can then
be fractionally distilled to produce more
concentrated alcohol. Different types of
alcoholic drinks contain different percentages
of alcohol. The typical ethanol concentration
of beer is around 4% while some spirits, such
as whisky and vodka, have a concentration of
40%.
To make a concentration of alcohol above
20%, ethanol must be concentrated by
fractional distillation of the fermentation
mixture. Remember that fractional
distillation separates two miscible liquids
with similar boiling points. Ethanol has a
lower boiling point than water. This means
that when the fermentation mixture is heated,
ethanol evaporates into the fractionating
column, while the water stays as a liquid. A Liebig condenser is used to condense the ethanol
vapour by cooling it. The concentrated ethanol can then be collected in a separate flask.
Fermentation has several advantages over hydrating ethene. Since it uses plant sugars as
feedstock as opposed to cracking products from crude oil, it is renewable. Moreover, it is less
expensive since lower temperatures and pressures are used (hence less energy is used).
However, fermentation is a batch process and so only certain volumes of ethanol can be
produced at a time. This means that fermentation is not very adaptable to changes in the
demand for ethanol. Since by-products are also produced, fermentation does not have a 100%
atom economy and other ways to use the by-products must be determined to reduce the
amount of waste.
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Investigating the combustion of alcohols
The combustion of alcohols is an exothermic reaction - it releases energy. This energy can be
used to generate electricity, heat homes, power cars etc. and so alcohols can act as fuels, like
the fractions of crude oil.
To test how well a particular alcohol would act as a fuel, the amount of energy it releases
when it is burned must be measured. We can do this by using the following method:
1. Put some alcohol into a spirit burner and measure the mass of the burner and fuel with
a mass balance.
2. Measure 100cm of distilled water into a copper calorimeter.
3. Insulate the calorimeter by using a draught shield, then cover it with an insulating lid
after placing a thermometer inside. This helps to make sure that minimal energy is
lost to the surroundings.
4. Take the initial temperature of the water, then put the burner under the calorimeter
and light the wick.
5. Stir the water throughout using the thermometer. When the heat from the burner has
made the temperature of the water rise by 20 °C, extinguish the flame.
6. Immediately reweigh the spirit burner and subtract the burner and fuel’s original mass
from it – this gives you the mass of fuel that has been used.
7. Repeat the experiment using other alcohols.
3
You can use your results to compare the masses of alcohol needed to increase the temperature
of the water and hence their usefulness as fuels. The lower the mass of fuel needed to raise
the temperature of the water by a given amount, the better the alcohol is as a fuel, as it
releases more energy per gram.
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Carboxylic acids
Carboxylic acids are a homologous series of molecules that contain a carboxyl (-COOH)
functional group. The carboxyl group is formed by the combination of a carbonyl (C=O)
group and a hydroxyl (-OH) group.
Carboxylic acids have the general formula CnH2n+1COOH, where n denotes one less than the
number of carbon atoms. The reason why n is one less for carboxylic acids is that the carbon
atom in the -COOH group counts as part of the carbon chain for the prefix. For example,
ethanoic acid’s molecular formula is CH3COOH.
Regarding IUPAC nomenclature rules, all carboxylic acids have the suffix -oic acid.
Carboxylic acids have the typical properties of acids. For example, they react with metals to
form a metal salt and hydrogen, and they react with metal carbonates to form a metal salt,
water and carbon dioxide. Since they only partially dissociate in water to produce H+ ions,
carboxylic acids are weak acids.
Carboxylic acids can be made by oxidising alcohols with an oxidising agent, such as sulfuric
acid or potassium dichromate. For example, ethanol can be oxidised to form ethanoic
acid:
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Addition polymerisation
Polymers are complex molecules that consist of long chains of smaller molecules called
monomers. Addition polymers are polymers formed in a reaction with no other polymers.
Poly(alkenes) are the most common addition polymers.
To make poly(alkenes), smaller alkene monomers open their double bonds so that the carbon
atoms can join together to form a long chain. The name of the poly(alkene) comes from the
alkene used as the monomer. For example, ethene monomers form poly(ethene) and propene
monomers form poly(propene). The molecular formula of the poly(alkene) is simply the
molecular formula of the alkene in brackets, with a subscript n outside. For example,
poly(ethene)’s molecular formula is (C2H4)n and poly(propene)’s molecular formula is
(C3H6)n.
To draw the displayed formula of a poly(alkene), firstly replace the double bond with a single
bond. Draw the other atoms in their same positions, then draw brackets around the formula,
ensuring that the bonds extend out of the brackets. Finally, place a subscript n outside the
brackets.
You might also be asked to draw the repeating unit of a poly(alkene). To do this, remove the
brackets and the subscript n, then replace the single bond with a double
bond.
Different poly(alkenes) have different properties that make them suitable for different uses:
• Poly(ethene) is flexible, an electrical insulator and cheap. It is used for plastic bags,
bottles and wire insulation.
• Poly(propene) is flexible, strong, tough and mouldable. It is used for crates, furniture
and ropes.
• Poly(chloroethene), also known as PVC, is tough, an electrical insulator and
waterproof. It is used for window frames, water pipes and electrical wires.
• Poly(tetrafluoroethene), also known as PTFE or Teflon®, is unreactive, slippery and
non-stick. It is used for non-stick pans and waterproof clothing.
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Condensation polymerisation
Condensation polymers are polymers formed in a reaction where water molecules are
eliminated for each bond made. The monomers used to make condensation polymers have
two functional groups (they don’t necessarily have to be two of the same group).
One of the most common condensation polymers is the polyester. Polyesters are made from
dicarboxylic acids and diols (di-alcohols). During the condensation reaction, an ester link (COO) is formed between the dicarboxylic acid and the diol, shown in the red circle. Two
water molecules are eliminated, shown in the blue circles. Two different shades are used to
show which atoms form each water molecule.
This means that for any n dicarboxylic acids and diols, 2n water molecules are eliminated to
form the polyester.
There are also several naturally occurring polymers that you need to know for GCSE
Chemistry:
• DNA stands for deoxyribonucleic acid and is found in all living organisms. It consists
of two strands that twist together to form a double helix. Each strand is a polymer,
made of monomers called nucleotides. The order of bases within the nucleotides
creates a unique genetic code for the organism.
• Carbohydrates are organic compounds that contain carbon, oxygen and hydrogen.
They are used by living organisms to release energy. One common carbohydrate is
starch, a polymer made of long chains of glucose molecules.
• Proteins are polymers made up of amino acids. Amino acids contain two functional
groups at their ends - one end has an amino (-NH2) group and the other end has a
carboxyl (-COOH) group. As n amino acids join together to form proteins, n water
molecules are eliminated.
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The economic and environmental sustainability of polymers
When a scientist carries out a life cycle assessment of a plastic, they will likely find issues at
each stage, from raw materials to disposal.
The alkene monomers used to make poly(alkenes) are often produced by cracking, which
uses crude oil. Crude oil is a non-renewable resource, so its supply will deplete and its price
will increase over time. Crude oil is versatile in the fact that its fractions can be used as fuels,
heating, surfacing for roofs and roads etc. Therefore, as the supply of crude oil depletes,
opportunity costs will have to be made as to how we can use the remaining crude oil
optimally.
There are two main methods of disposing of polymers:
• Landfill sites: When polymers are too difficult or expensive to recycle, they are often
placed in landfill. There are two main problems with landfill. Firstly, it takes up
valuable land that could be used for rearing livestock and growing crops for food or
biofuels. Secondly, the polymers thrown into landfill sites are often nonbiodegradable, meaning that they will stay there for many years.
• Combustion: The energy released in burning polymers can be used for heating or
generating electricity. However, there is the possibility that toxic gases can also be
released when burning polymers. For example, when polymers containing chlorine
are burnt, such as poly(chloroethene), hydrogen chloride gas is produced. Carbon
dioxide may also be produced and this is a known greenhouse gas that contributes to
global warming.
To delay the economic dilemma of crude oil, we can recycle polymers. Recycling reduces the
amount of non-biodegradable waste in landfill sites and the volume of toxic or harmful gases
released when polymers are burnt. Recycling polymers will also mean that less crude oil will
need to be used to produce the alkene monomers needed to make them. From an economic
standpoint, recycling centres create jobs, which will be particularly beneficial for areas with
low economic activity.
However, there are also disadvantages to recycling polymers. Each time a
polymer is recycled, its strength decreases. This means that recycled
polymers will eventually have shorter lifespans or lower quality. In
addition, polymers must be separated by type before they can be
remoulded into new polymers, which can be difficult and
expensive.
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Nanoparticles
Nanoparticles are minuscule particles that consist of a few hundred atoms. They typically
have diameters ranging from 1 nm (1 × 10-9 m) and 100 nm (1 × 10-7 m).
Nanoparticles have a very high surface area to volume ratio, which means that they have a
very large surface area in comparison to their volume. The smaller the particle, the larger the
surface area to volume ratio. This property allows a greater proportion of atoms in a
nanoparticle to interact with substances than atoms in a bulk material, meaning that
nanoparticles will typically have different properties than their bulk counterparts.
Nanoparticulate materials
Scientists can use the properties of different nanoparticles to make nanoparticulate
materials. Since nanoparticles have a high surface area to volume ratio, they can be very
effective catalysts for many reactions. A good example is the use of platinum nanoparticles in
fuel cells.
Fullerenes are a simple molecular structure and nanoparticle whose structure allows them to
contain other substances. Fullerenes can be absorbed into the body more easily than other
larger substances, so scientists can place medicines, such as chemotherapy drugs, inside these
fullerenes so that they can be delivered to cells quickly.
Since nanoparticles are so small, they can be added to bulk materials to make them stronger
and more durable without increasing their weight significantly. This makes them suitable for
different sports equipment, like tennis rackets and golf clubs.
Graphene is a nanoparticle that can conduct electricity since one electron is delocalised for
every three covalent bonds that each carbon atom makes. Therefore, graphene can be used for
various electronic devices, such as display screens or memory chips.
Silver nanoparticles have the distinctive property of being antibacterial, as they can penetrate
the cell walls of bacteria and permanently damage cell membranes, causing cytoplasm to leak
out of bacterial cells. This property makes silver nanoparticles suitable for deodorants.
Finally, testing of nanoparticles has shown that they are better at protecting the skin from
ultraviolet rays than substances typically used in sunscreens. Moreover, since they don’t
leave white marks on the skin, nanoparticulate
sunscreens often give better skin coverage.
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The risks of nanoparticles
Nanoscience is a new branch of science and there are still many questions about nanoparticles
that scientists have not yet found answers to. Remember that nanoparticulate properties are
often different to bulk material properties, so even if the bulk material is considered safe, we
cannot assume the same safety for nanoparticles.
One particular question is the use of nanoparticles on or inside the body, such as sunscreens
and medicinal drugs in fullerenes. Some people speculate as to whether
they can damage or accumulate in cells to dangerous levels if they are
not broken down easily. Others suggest that if certain nanoparticles are
breathed in, they might encourage or exacerbate respiratory problems.
Another concern is that when antibacterial-treated fabric or clothing is
washed, some of the silver nanoparticles could travel into water
reservoirs and rivers as they are non-biodegradable. This could change
the conditions of aquatic environments to the point where they become
harmful to wildlife, such as fish.
Polymers as materials
Most polymers are insulators of heat and electricity, they can be flexible and are easily
moulded. They are often cheaper than most other materials, and they also tend to be less
dense.
Many different monomers can be used to make polymers. This means that the properties of
different polymers are varied, resulting in many applications:
• Polyesters are used in clothing due to their ability to stretch and dry quickly, making
them an ideal substitute for cotton in clothing such as T-shirts.
• High-density poly(ethene) [HDPE] is used to make water pipes as it’s strong and
rigid.
• Light, stretchy polymers such as low-density poly(ethene) [LDPE] are used for plastic
bags and bottles.
• Poly(styrene) foam is used in packaging to protect breakable things and as a thermal
insulator.
• Heat-resistant polymers such as melamine resin and poly(propene) are used to make
plastic kettles.
• Condensation polymers can degrade and break down over time, so polymer products
don’t always last as long as those made from other materials.
Ceramics as materials
Ceramics are non-metallic solids with high melting points
that aren’t made from carbon-based compounds. They are
excellent insulators of heat and electricity. They tend to be
very brittle and stiff, but also strong and hard-wearing. They
don’t degrade or corrode like other materials can, so they can
last a lot longer. For example, porcelain plates or bowls are
excellent at insulating heat. This is why you can carry hot food
in them without hurting your hand. However, porcelain
shatters easily if dropped as it’s brittle.
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There are two types of ceramics that you need to know:
• Clay ceramics: Clay is a mineral formed from weathered and decomposed rock. It’s a
soft material when it’s dug out of the ground, so it can be moulded into different
shapes. When fired at high temperatures, it hardens to form a clay ceramic. As clay
can be moulded when wet and then hardened, it’s ideal for making pottery and bricks.
As clay bricks are hard, they can withstand the weight of more bricks on top of them.
• Glass: Glass is generally transparent and strong, can be moulded when hot and can be
brittle when thin. Most glass is soda-lime glass, which is made by heating a mixture
of limestone, sand and sodium carbonate until it melts. When the mixture cools, it
comes out as glass.
Composites as materials
Composites are made of one material embedded in another. Fibres or fragments of a
material, known as the reinforcement, are surrounded by a matrix acting as a binder. The
main disadvantage of most man-made composites is that they tend to be much more
expensive to produce than other materials.
The properties of composites depend on the matrix and the reinforcement used to make them,
so they have many different uses:
• Fibreglass consists of fibres of glass embedded in a matrix made of polymer (plastic).
It has a low density like plastic but is very strong like glass. It is used for objects like
skis, boats and surfboards.
• Carbon fibre composites also have a polymer matrix. The reinforcement is either
made from long chains of carbon atoms bonded together (carbon fibres) or from
carbon nanotubes. Carbon fibre composites are generally very strong and light, so
they can be used in aerospace and sports car manufacturing.
• Concrete is made from aggregate (a mixture of sand and gravel) embedded in cement.
It is strong and rigid and so it is ideal for use as a building material, like in skate
parks.
Metals as materials
Metals are good conductors of heat and electricity. This can be an advantage or disadvantage,
depending on what the material is needed for. They are malleable, so like polymers, they can
be formed into a variety of shapes.
Some metals corrode easily, but products made from corrosion-resistant metals can last for a
very long time. Metals are usually less brittle than either ceramics or polymers, so they’re
likely to deform but stay in one piece where other materials may shatter.
Metals can also be mixed with other elements to form alloys. For example, as aluminium is
malleable and low in density, aluminium alloys are used in the construction of aircraft
fuselages.
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Testing for cations
Ionic substances consist of a cation (a positive ion) and an anion (a negative ion). Scientists
can carry out tests on unknown ionic substances to identify the cations and anions that form
them.
Some cations cause an ionic substance to exhibit a particular colour when burnt. In the lab,
we can conduct a flame test. Dip a nichrome wire loop into dilute hydrochloric acid and then
rinse it with distilled water. Dip the loop into your unknown substance and place it in the
clear blue area of a Bunsen burner flame.
You need to know the colours of the flames produced by the
following cations:
• Lithium ions give a crimson flame.
• Sodium ions give a yellow flame.
• Potassium ions give a lilac flame.
• Calcium ions give an orange-red flame.
• Copper ions give a blue-green flame.
The other common test for cations in the lab is by adding sodium
hydroxide solution. Some cations form precipitates with particular
colours when sodium hydroxide solution is added.
• Calcium ions form a white precipitate.
• Copper ions form a blue precipitate.
• Iron(II) ions (Fe2+) form a green precipitate.
• Iron(III) ions (Fe3+) form a brown precipitate.
• Aluminium ions initially form a white precipitate, then
redissolve in excess sodium hydroxide to form a
colourless solution.
To work out whether a substance contains ammonium ions (NH4+), all you need to do is add
some sodium hydroxide solution to the substance and gently heat it. If ammonia gas is given
off, it means that there are ammonium ions in the substance.
You can test for ammonia gas by holding a piece of damp red litmus paper over it. If the
gas is ammonia, the litmus paper will turn blue. Even though ammonia has a very
distinctive strong smell, do not sniff a mystery gas to identify that it is ammonia. At high
concentrations, ammonia is an irritant and toxic.
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Testing for anions
Remember that when an acid reacts with a metal carbonate, a metal salt, water and carbon
dioxide are produced. The carbon dioxide (collected in a gas syringe) can be bubbled into
limewater. If the limewater becomes cloudy, then carbon dioxide is present and so carbonate
ions were present in the initial reaction with the acid.
Halide ions (ions derived from the Group 7
elements, the halogens) can be identified by adding
dilute nitric acid and silver nitrate solution.
• Chloride ions give white precipitates of
silver chloride.
• Bromide ions give cream precipitates of
silver bromide.
• Iodide ions give yellow precipitates of
silver iodide.
The nitric acid should be added before the silver
nitrate solution so that any carbonate ions (which
would also produce a precipitate with silver nitrate
solution, confusing the results) are removed. It’s
also important to use nitric acid when testing for halide ions rather than hydrochloric acid.
Hydrochloric acid would introduce chloride ions to the solution, so a white precipitate would
be formed regardless of whether the solution originally contained chloride ions or not.
Barium sulfate is a white precipitate and is often formed in the lab to identify sulfate ions. To
make barium sulfate, add dilute hydrochloric acid to your unknown substance to remove any
carbonate ions (which can also produce white precipitates when added with the next
chemical). Then, add barium chloride solution to your unknown substance. If a white
precipitate forms, you’ll know that it is barium sulfate and so sulfate ions were present in the
unknown substance.
Flame photometry
You might have realised that flame tests can only identify one cation. They cannot be used to
identify multiple cations at once because the different colours would mix in the flame. If you
know that an unknown substance contains multiple cations, you can use flame photometry
instead.
Like flame tests, flame photometry involves placing a sample into a flame. As the ions heat
up, they become excited and move to higher
energy levels (or shells). When the electrons
fall back down to their original energy levels,
they transfer energy as light. This light is
passed through a flame photometer, which
detects different wavelengths of light. The
flame photometer produces a line spectrum
of different emissions.
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Each cation produces a different line spectrum since no two cations have both the same
charge and the same number of electrons. Therefore, the wavelengths of light that cations
emit will be different.
The readings on the flame photometer
can correspond to concentrations of the
cation with a calibration curve. Using
the calibration curve on the right, you
can see that a flame photometer reading
of 4 indicates that the concentration of
sodium ions is 0.02 g dm .
-3
Flame photometry is an instrumental
method of analysis - a technique that
involves machines and computers
instead of solely relying on humans (like
flame tests). There are three main
advantages of instrumental methods of analysis rather than chemical tests:
• They are more sensitive since they can detect very minuscule amounts of a substance.
• They are faster since they can be automated.
• They are more accurate since they do not involve human error.
Perhaps the most obvious advantage of flame photometry is that it can be used no matter how
many cations are present in the unknown substance, unlike flame tests. Due to its high
sensitivity, it can produce line spectra of many different cations, which can be accompanied
by a calibration curve to determine accurate concentrations of each cation.
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