Quantitative Chemistry
The good news is, there aren’t that many calculations you have to
do in chemistry. The bad news is that the examining boards like
to give a lot of marks for them, so it’s worth getting used to them.
I hope that this tutorial will give you a feel for doing them and ways to move them
around and get a result.

Target level: easy A level and difficult GCSE.

You WILL need a copy of the periodic table, so go and get that now.

ANSWERS to all of the questions are on the back page. They may even be in
the right order.

If you print this document double-sided, the asymmetrical margins give you
somewhere to staple it together, without losing much visibility.
Reading Formulas
This is so simple that it’s easy to get wrong. Here’s a formula:
H2S2O7
Never mind what it is, what’s in it? There is hydrogen (twice), sulphur, (twice) and
7 oxygens. The numbers go AFTER the element. Easy, but some people get it
wrong.
Try these:
Question 1
how many of each atom is in CH3COOH?
Question 2
how many of each atom is in K2CrO4?
Question 3
how many of each atom is in KMnO4?
Of course, there are other numbers in chemistry; if a chemical reaction uses two
of a particular chemical, then you are going to end up with numbers in front of a
formula, this multiplies everything that follows, up to the end of that formula.
E.g. in 2H2 + O2  2H2O
The first “2” multiplies the H2, so there are 4 hydrogen atoms in total. The oxygen,
as with all reacting gases, comes in pairs (O2), so you end up with two each side
and you need two water molecules on the right (2H2O).
If balancing in general mystifies you, you could perhaps write your own equations
and get them balanced. Even if the products are impossible, it’s still practice.
E.g. In this half-equation, there are charges:
Al3+ + 3e¯  Al
The aluminium starts with a shortage of three electrons, so it has a triple positive
charge. The electrons are replaced in the half-reaction. Notice that the charge is
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at the top right, (where a mathematician would put a power) and the number in
front multiplies everything in the formula after it.
Try these:
Question 4
how many of each atom is in 3H2SO4 ?
Question 5
how many of each atom is in Fe(C2O4)3 ?
Question 6
how many of each atom is in 8S8 ?
Question 7
how many of each atom is in 3Ca(NO3)2 ?
Question 8
how many of each species is in Cu2+(aq) + 2e-  Cu(s) ?
Proton Number (atomic number, also called Z)
This one’s easy too; it’s the smaller number of the two by the element symbol in
the periodic table. Remember, nearly all of the mass of an atom is in the nucleus
and some of that has to be neutrons, so the rest must be protons. Also called
atomic number in old tables and books. The periodic table is in proton number
order because they make more sense than ordering it by atomic masses, which
are occasionally “out of order” and are averages anyway (See “relative atomic
mass” later in this document.
So, try these:
Question 9
What is the proton number of calcium?
Question 10
What is the proton number of iron?
Question 11
What is the proton number of silver?
Mass Number (atomic mass, also called A, but only by his friends)
This is the total number of neutrons and protons in the nucleus. It will be a whole
number. It will describe a particular isotope of an atom, for example, Bromine-79
or bromine-81. However, in the periodic table, you will see relative atomic mass,
which is not the same, because it has been averaged and will usually have
decimals, so you have to round this number up or down in the usual way. Some
simple periodic tables have a whole-number mass quoted.
Try these, what is the mass number of the most common isotopes of:
Question 12
silicon
Question 13
phosphorus
Question 14
carbon
Question 15
magnesium
Examiners like to check that you can work out the neutrons by subtracting the
proton number from the mass number, particularly for isotopes.
Try these, too, how many neutrons in an atom of :
Question 16
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iron-56
To avoid confusion with
isotopes, we put the atomic
mass after the name, or make
it small and put it before:
e.g. 12C or C-12
Page 2 of 16
Question 17
argon-40
Question 18
oxygen-16
Yes, it’s as easy as “big number minus little number”.
Some atoms have isotopes, they have the same proton number, but more or less
neutrons, so they chemically behave exactly like each other, but have different
masses.
Try these:
Question 19
how many neutrons in uranium-235 ?
Question 20
how many neutrons in uranium-238 ?
Question 21
how many neutrons in Bromine-79 ?
Question 22
how many neutrons in 81Br ?
Question 23
how many neutrons in 13C ?
Molecular mass
Also called Mr, this is the same as formula mass, but refers to covalent
compounds, ones that have separate molecules, rather than a giant lattice
structure. Anyway, you work them out the same way, by adding-up the atomic
masses, so:
try these; what is the molecular mass of:
Question 24
Water, H2O
Question 25
Methane, CH4
Question 26
Ammonia, NH3
Question 27
Ethene, C2H4
Question 28
Propanoic acid, C2H5COOH
(just count the atoms, even if you don’t know the compound)
Formula mass
Ionic compounds form crystals in which the positive and negative atoms repeat in
3-D, making a “giant structure” or “giant lattice”. There are no separate
molecules, so “molecular mass” isn’t the right words to use. The ratio of the
atoms defines the formula, e.g. MgCl2 magnesium chloride and so on. Just add
up the atomic masses, e.g. for NaCl, the formula mass is:
23 + 35.5 = 58.5
So easy isn’t it? Try these, what is the formula mass of:
Question 29
magnesium oxide, MgO
Question 30
Silicon dioxide, SiO2
Question 31
potassium chloride, KCl
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Relative Atomic Mass
Also called Ar this is the bigger of the two numbers by the symbol in the periodic
table. It’s defined as
The average mass of an atom of an element divided by one-twelfth of the mass of
a carbon-12 atom.
Watch out, this is not just the total of protons and Relative abundance – how
neutrons, you have to work out the average mass much of one atom there is
from the relative abundances (usually percentages) compared to the others.
OR from given masses of isotopes. The important
thing about this is for you to realise that you can't get carbon that is only 12C or
chlorine that is only 35C, you and I have to weigh out samples of real elements with
real isotopic ratios.
The calculation is:
Total of all the (amount of isotope x its mass)
Total of the amounts
(Formula A)
e.g. a question asks:
“A sample of magnesium contains 1.2g of magnesium-23, 94.3g of magnesium-24
and 3.1g of magnesium-25. What is the relative atomic mass of this sample of
magnesium?” Easy or what.
(1.2 x 23) + (94.3 x 24) + (3.1 x 25) =
(1.2 + 94.3 + 3.1)
2363.7
98.6
= 24.47 to 2 d.p.
If they give you percentages, you need to multiply each isotope by its percentage
and then divide by 100. Like this example:
“A sample of carbon contains 99% carbon-12, 0.998% carbon-13 and 0.002%
carbon-14. What is the relative molecular mass of this sample?”
99 x 12 + 0.998 x 13 + 0.002 x 14
100
= 12.01002 or 12.010
Either way, the Ar value you get should be very similar to the mass of the most
plentiful isotope. If it isn’t, go back and do it again.
Try these:
Question 32
a sample of oxygen is shown to be 99.3% oxygen-16 and 0.7%
oxygen-18. What is the Ar of this sample?
Question 33
a sample of chromium is found to be 4 parts chromium-51, 39 parts
chromium-52 and 1 part chromium-53. What is the Ar of this
sample?
Question 34
a 61g sample of gold is shown to be 3g of gold-195, 56g of gold197 and 2g of gold-198. What is the Ar of this sample?
WATCH OUT! If you just bang these into your calculator as, for example,
4x51+39x52+1x53 / 44 then you’ll get the wrong answer because your scientific
calculator will work out 53/44 first and then add it to the others.
(it knows the BODMAS rules and divide is more important than multiply).
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You have to work out the top completely first, then divide it by the bottom, which
you might also have to work out completely first. Just for a change, it is easier to
get the right answer on a 4-function calculator, though don’t use your phone,
because you wouldn’t be daft enough to take that into an exam.
If your calculator has got them, use the brackets buttons to make sure you get
things worked out in the right order.
2 decimal places is enough at GCSE, 3 at A level.
Percentage mass
This is as simple as any other percentage. Not so good at any other percentage,
eh? Try this simple formula:
little number X 100
big number
(Formula B)
Ah, yes, you say, but where do the big and little numbers come from?
Let’s look at a sample question:
Question :
What percentage mass of NH4NO3 is nitrogen?
The little number will come from the mass of the 2 nitrogens: 14 x 2 = 28
The big number will be the total mass of the formula:
14 x 2 + 1 x 4 + 16 x 3 = 80
So you get:
28
x 100 = 35%
80
So it’s
Mass of atoms of the element requested
Mass of all the atoms
(all of them added up)
(formula or molecular mass)
(Formula C)
Try these yourself, to 3 significant figures:
Question 35
What percentage mass of CO2 is oxygen?
Question 36
What percentage mass of FeCl2 is Iron? (Use 35.5 for Cl)
Question 37
What percentage mass of CH3COOH is hydrogen?
Empirical formulas
These are defined as:
the simplest ratio of one element to another in a compound,
e.g. 1:3 or 2:1:4 are ratios and they lead to formulas, 1:3 could be FeCl3, for
example, 2:1:4 could be Na2CrO4. They don’t have much actual use until after
GCSE, but one form is used to describe the general formula of compounds, e.g.
the alkenes CnH2n. There is one disadvantage to empirical formulas, you can’t
always tell which compound you’ve actually got. For example, C2H4 and C3H6
have the same ratio of C to H, 1:2, so their empirical formula is the same. The
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good thing about empirical formulae is that they are only percentage mass
reversed, you get given the percentages and you have to work out how many of
each element there is. Actually, sometimes you’re given the masses, but it’s the
same method.
The basic calculation is just simply dividing each element’s amount by its
individual atomic mass, which gives you a coarse ratio of how many atoms there
are, then juggle the numbers so they come out at a simple ratio. Let’s face it,
most elements are in a simple ratio in a compound.
e.g. The question says “22g of a certain compound contains 6g of carbon and
16g of oxygen. What is the empirical formula?”
First divide each amount by the appropriate atomic mass:
6
12
This gives us a ratio of
:
16
16
0.5 : 1
Ratio is now
1:2
multiply by 2
one C and 2 O
Therefore the empirical formula is CO2
Of course, the real formula could actually be C10O20, you can’t tell, all you’ve got is
a ratio, but that’s what you wanted anyway.
It’s difficult enough that you’ll probably need another example:
e.g. 11.2g of a compound is found to be 5.84g of carbon, 1.48g of hydrogen and
the rest is oxygen. What is its empirical formula?
Ok, so it’s sneaky, you need the oxygen first by subtraction from the total:
Oxygen = total – (carbon + hydrogen)
= 11.2 – (5.84 + 1.48) = 3.92
The ratio is therefore:
5.84 : 1.48 : 3.92
12
1
16
= 0.4866 : 1.48 : 0.245
This is where your ability with numbers is tested, because it should be fairly
obvious that 1.48 is about 3 times 0.4866 and 0.245 is about half of it.
OR that 0.4866 (call it 0.49) is about twice 0.245 and 1.48 is about 6 times 0.245
so the ratio is 2:6:1 and the formula is C2H6O
Phew! Not easy are they, but if they were easy, there would be only one mark for
them, whereas there’s often 3 at GCSE, or 2 at AS level.
Question 38
A compound is found to be 135g carbon and 45g hydrogen, what is
its empirical formula?
Question 39
A compound is found to be 12%C, 48%O and 40% calcium, what is
its empirical formula?
Question 40
20g of a compound is found to be 12g of C, 5.333g of O and
2.667g hydrogen, what is its empirical formula?
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Notice, the ratios might not be exact and you have to be good at spotting them.
Common ratios are 0.3333:1 or 1:3, 0.25:1 or 1:4 and so on, so don’t be too
precise about your numbers, for a change, on the other hand, 0.33333 is not 0.3.
If you get a ratio like C198O405 then you have taken the numbers too literally and
the compound is probably CO2.
Reacting quantities
These come from the empirical formula (the ratios of the elements in the reaction
formula). For example:
C + O2  CO2
(Formula D)
That’s easy, clearly each carbon involved has to be given 2 oxygens. Remember
the little number below the O says there’s two of them.
Now, if there’s, say, 12g of carbon, you’ll need 32g of oxygen. Can you see why?
Simply because each atom of carbon has a mass of 12, but each oxygen atom
has a mass of 16 and there’s two of them to be included. If you chose equal
masses of the chemicals, you’d be short of oxygen, because it’s heavier than the
carbon. This will sound weird, because you may not have considered oxygen to
be heavy before now.
So, if you are thinking about joining different atoms, you have to count the
different numbers, not masses, and to do that you have to divide everything by the
amount it weighs.
So if you are asked “6g of carbon react with oxygen, what mass of oxygen is
needed?” then you have to think this sort of way:
“6g of C divided by 12 is the same as A grams of oxygen divided by 32. As an
equation this is:”
6
A
(Formula E)
12

32
Now, if you change this around, (multiply both sides by 32) you
can have:
(Formula F)
So A must be 16 grams
6 x 32
A
12
Try these:
Question 41
if 88g of carbon dioxide is created, in the reaction above, how much
carbon was used?
Question 42
What mass of oxygen is needed to burn 48g of sulphur atoms,
making sulphur dioxide?
Question 43
What mass of iron would you get from 320g of Fe2O3, if you got rid
of the oxygen using a blast furnace?
Question 44
What mass of chlorine is needed to make 50g of magnesium
chloride, MgCl2?
Question 45
Many people should reduce their sodium ion intake to avoid high
blood pressure. What mass of sodium is in 2g of salt, NaCl?
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Moles
These are a way to count the number of atoms that are in a reaction and make
sure they match the formulas involved. You have already met them without
knowing it, in the previous section.
Think about this, 2 grams of hydrogen are burnt in air.
The obvious unbalanced equation is:
H2 + O  H2O.
(Formula G)
This has the usual problem, oxygen should be in
a pair, so the balanced equation is:
2H2 + O2  2H2O.
(Formula H)
Every one oxygen atom joins with 2 hydrogen
atoms. Do atoms of oxygen and hydrogen have
the same mass?
Moles used to be known as a
gram-atom or the gram-atomic
weight of an element or the
gram-molecular weight of a
compound. This terminology
helps you to understand what it
means, but isn’t very easy to
write or abbreviate. “Moles” is
easier to write, is the same
word for the same amount of
atoms or molecules and makes
no sense as a name at all.
You’re right, they don’t. Hydrogen atoms have a
mass of 1, oxygen atoms have a mass of 16,
clearly much heavier. If we added, say, 2 grams of oxygen to the 2 grams of
hydrogen, there would be too little oxygen, because each oxygen atom is 16 times
heavier than a hydrogen atom and 2g, although it looks the same, will be far too
few oxygen atoms.
So, the equation is really telling us that our 2g of hydrogen reacts with 16g of
oxygen, because if 2 hydrogen atoms react with one oxygen atom, then 2g of
hydrogen reacts with 16 g of oxygen. Are you sure you can see why? If not, stop
here and see a real teacher, not a piece of paper.
Now, in real laboratories, we can’t weigh out atoms one by one, we need billions
of the things, because we are using balances that go down to about 0.01g, not the
10-23g or so that an atom weighs. What we need is a fair number so that we
always weigh out the same number of atoms when we weigh out lots of different
masses.
This is called a mole.
A mole is the mass (in grams) of a certain number (6.022 x1023, Avogadro’s
number) of atoms. It is this number because that makes the real mass of a mole
of any element match exactly the atomic mass of that element. So 1 mole of
hydrogen has a mass of 1g, because hydrogen has a mass of 1. 1 mole of
oxygen has a mass of 16g and so on. One mole of something is Avogadro’s
number of that thing. This suggests that you can have a mole of London buses
and indeed you can, as long as you don’t mind how heavy that’s going to be
(about 3 x 1024 tonnes, or about 500 times the mass of the Earth).
Try these:
Question 46
What is the mass of 1 mole of sulphur atoms?
Question 47
What is the mass of 1 mole of sodium atoms?
Question 48
What is the mass of 1 mole of lead atoms?
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Question 49
What is the mass of 1 mole of iodine atoms?
Question 50
What is the mass of 20 moles of vanadium atoms?
Question 51
What is the mass of 0.1 moles of nitrogen atoms?
The most useful thing about this is the way mass calculations are simplified, as
every atom represented in the equation for the reaction means a mole, so looking
back at formula D, each mole of oxygen atoms needs 2 moles of hydrogen atoms
to react with. (Otherwise the formula of water won’t be H2O)
Watch out for this, though, they occasionally ask you about moles of molecules,
which are more than one atom. For example, oxygen, hydrogen, nitrogen, and
the halogens all come in pairs (they’re diatomic), and sulphur comes in rings of 8,
so read the question carefully!
Try these:
Question 52
What is the mass of 1 mole of hydrogen molecules (H2)?
Question 53
What is the mass of 10 moles of oxygen molecules (O2)?
Question 54
What is the mass of 1 mole of sulphur molecules (S8)?
Question 55
What is the mass of 0.1 moles of iodine molecules (I2)?
Moles and molecular mass
Then, of course, they may ask you about moles of molecules of compounds. This
is as easy as atoms, but you have to know how to work out the molecular mass
from the formula, see molecular mass earlier in this document.
Question 56
What is the mass of 3 moles of copper sulphate molecules
(CuSO4)? (assume copper = 64)
Question 57
What is the mass of 0.1 moles of ethanol molecules (C2H5OH)?
Question 58
What is the mass of 8 moles of hexane molecules (C6H14)?
Question 59
What is the mass of 0.4 moles of chlorophyll molecules
(C55H72O5N4Mg4)?
Working with Moles
 Remember Avogadro’s number is 6.022 x 1023
 Remember 1 litre is 1 dm3 (one cubic decimetre)
 Here I’ve used “atomic mass(es)” to mean atomic,
formula or molecular masses.
Just moles:
Avogadro’s number x moles = number of atoms
mass in
grams
number of atoms / moles = Avogadro’s number
number of atoms / Avogadro’s number = moles
moles
atomic
mass(es)
e.g.
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0.05 moles of carbon dioxide contains
0.05 x 6.022 x 1023 molecules
= 3.011x1022 molecules
Moles and real masses:
Moles x atomic mass = a mass in grams
therefore grams / atomic mass = miles
and grams / moles = an atomic mass or whatever
Solutions and moles:
Conc. = amount / volume e.g. 0.5 M/dm3
therefore concentration x solvent volume = moles
moles
and moles / concentration = solvent volume
e.g.
200 ml of 0.22 M CuSO4 solution contains
concentration
solvent
volume
(dm3)
0.2 x 0.22 moles of CuSO4 = 0.044 moles
e.g.
3.4 moles of HCl dissolved in 5 dm3 of water makes a solution of concentration
3.4 / 5 = 0.68 Moles / dm3 , that is, 0.68 moles per litre
e.g.
the volume of 2.1 Moles/dm3 solution that contains 1.3 moles is
1.3 / 2.1 = 0.619 dm3 or 619 ml
The difficult thing about most mole-based calculations is that you start off in
grams, so you have to get moles of one thing and then work with it to get moles of
something else and then convert that back to grams. Phew! The good thing is
that at GCSE, they use simple multiples of the mass, like 4.6g of sodium or 0.7g
of lithium. At A level it’s a different matter, but the principle is still the same and
you can usually get a rough figure in your head if you know the principle of
working with the simple stuff properly.
The other difficult thing is the chemistry. You may get a mark for doing grams to
moles and another for going back again, but it’s working out what’s going on in
between that’s tricky for most people. Look at a typical question – it’s going to ask
you to do something based on an equation. So read the equation to yourself:
C + O2  CO2
Means “An atom of carbon joins to two of oxygen making a molecule of carbon
dioxide.” But it also means “A mole of carbon atoms joins to two moles of oxygen
atoms making a mole of carbon dioxide molecules.”
At its most basic this will give you12g of C + 32g of O2 makes 44g of CO2
OR
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it might be half as much..... 6g of C + 16g of O2 makes 22g of CO2 and so on.
So all you would have to do, given the equation, is try to see how many moles of
everything there is, they’ve probably told you the mass of one anyway.
e.g.
Given: 4Na + O2  2Na2O, what is the mass of oxygen needed in
burning 46g of sodium?
Well, start with the moles of sodium, that’s 4.6 ÷ 23 = 2 moles.
Say “Each mole of sodium reacts with ¼ mole of O2” because that’s what the
equation says. Then you know that 2 moles of Na should need 0.5 moles of O2,
so you need 0.5 * 32 = 16g
Try this:
Question 60
What is the mass of oxygen needed in burning 2.4g of magnesium,
making MgO?
Moles of gases
A mole of any gas at room temperature occupies 24 litres (dm3). Half a mole will
fill half this and so on. A level students prove this in unit 1.
Try this:
Question 61
What is the volume of 2 moles of oxygen?
Question 62
What is the volume of 14 moles of nitrogen?
Question 63
What is the volume of 0.02 moles of helium?
Heat energy change
This is the energy used or needed to heat up a certain amount of a certain
substance by a certain amount. You can work it out using the formula:
Q = mcT
where Q is the heat energy, m is the mass, c is the heat capacity of the material
and T is the change in temperature. T and m depend on your experiment and
c depends on the material, water, for example needs more energy to heat up than
glass does. c is called the specific heat capacity of the material or sometimes just
heat capacity. This assumes that there is no phase change.
Conveniently, 1cm3 (or ml) of water has a mass of 1g
You need to know that the heat capacity of water is 4.2 J ºC-1 g-1 commonly said
as 4.2 joules per degree per gram.
The heat capacity of copper is 0.386 J ºC-1 g-1 which means copper is easier and
cheaper to heat up than water, though at the same temperature as water, it holds
less heat energy and won’t stay hot for so long.
Question 64
What is the heat energy absorbed when 20g of water rises in
temperature by 10ºC?
Question 65
What is the heat energy absorbed when 200g of copper rises in
temperature by 120ºC?
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Question 66
What is the heat energy absorbed when 2kg of water and a copper
kettle of mass 800g both rise in temperature by 70ºC?
Reaction energy per mole
This is just the energy you use when bonds are broken minus the energy you get
when bonds are made.
e.g. in burning methane (CH4 + 3O2  CO2 +
2H2O), you have to break 4 C-H (single) bonds
and 3 O=O (double) bonds, then make two C=O
bonds and four O-H bonds. Notice you have to
know the structures of all the molecules involved in the reaction.
The examiners will give you the numbers for the energy values of each bond, all
you have to do is add the broken ones and the made ones individually and then
subtract the total of made bonds from the total of broken ones.
For example, in burning hydrogen, the bonds broken are two H—
H bonds and an O=O bond. There are four O—H bonds made.
2H2 + O2  2H2O
498
O—H 366
mol-1
The secret is to do the following steps:
1.
2.
3.
4.
5.
6.
7.
8.
H—H 436
O=O
So the total energy change was:
(436 x 2 + 498) – (366 x 4) = 1370 – 1464 = –94 kJ
Some bond
energies:
Make sure you have a balanced equation
Draw all the reactant molecules
Mark their bonds – these will be broken
Get the numbers and add them up, call them A
Draw all the product molecules
Mark their bonds – these will be made
Get the numbers and add them up, call them B
Subtract B from A
Question 67 What is the energy change for this reaction?
C=O
805
Units are kJ
per mole or
kJ mol-1
C + O2  CO2
If the result is negative, you have an exothermic reaction, energy has
been lost from the chemicals. Remember, we are chemists, so the
energy is lost from the chemicals, but we can’t tell, except by
measuring temperature rises (heat energy changes) during the
reaction.Appendix A - elements you should know
(and things you should know about them):
You should know the symbols and names of these elements, or it will be hard to
do GCSE chemistry. I mean it, if you don’t know what a brick is, it’s hard to build
a house, even in Lego.
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co in the ion column means you usually find it covalently bonded.
Symbol
Name
G
Comment
ion
Al3+
Al
Aluminium
3
So reactive it has an inert layer of oxide all the time so
it seems unreactive
Ar
Argon
0
Noble gas, used in light bulbs, 1% of the air
Br2
Bromine
7
Brown poisonous diatomic liquid
Ca
Calcium
2
Metal, usually looks grey/white because it is covered
with oxide
Ca2+
C
Carbon
4
The only conductive non-metal, used for electrodes as
it doesn’t react - has 2 forms(allotropes), diamond and
graphite
co
Cl2
Chlorine
7
Greenish poisonous diatomic gas, though you usually
don’t see the colour because it’s not concentrated
enough.
Cl-
Cu
Copper
T
Orange-coloured metal, good conductor, makes useful
alloys, Cu2+ is most common ion at GCSE
Cu2+
F-
Br-
Cu3+
F2
Fluorine
7
Yellowish poisonous diatomic gas, not found in school
laboratories, terrifically dangerous.
He
Helium
0
Lightest inert
breathing mix
I2
Iodine
7
Black (with a hint of purple) poisonous solid, diatomic
Fe
Iron
T
Hard metal, rusts easily if water and oxygen are
present, steel is mostly iron
Au
Gold
T
Excellent conductor, soft, very dense, expensive
Pb
Lead
4
Soft metal, usually looks dark grey from the oxide coat
Pb2+
Li
Lithium
1
Reacts with (and floats on) water to give an alkaline
solution and hydrogen, yellow flame, kept under oil.
Less reactive than sodium
Li+
Mg
Magnesium
2
Soft metal, quite reactive.
Ne
Neon
0
Noble gas, red glow when electricity passes through it
N2
Nitrogen
5
Diatomic gas, unreactive under normal conditions used to fill food packaging and in the Haber process for
making ammonia, 79% of air.
O2
Oxygen
6
Reactive diatomic gas, helps things to burn, 21% of the
air
P
Phosphorus
533569709 17/02/2016
5
gas
-
children’s
balloons,
divers’
White or red non-metal, burns spontaneously in air,
kept under water, plants need it in small amounts, so
do we
IFe2+
Fe3+
Mg2+
co
O2and
co
co
Page 13 of 16
K
Potassium
1
Reacts with water to give an alkaline solution and
hydrogen, pink flame, more reactive than sodium, kept
under oil, plants need it in small amounts, so do we
K+
Ag
Silver
T
Silvery metal, mostly unreactive, so used for jewellery.
AgCl is light-sensitive, giving black crystals of silver.
Ag+
Na
Sodium
1
Reacts with (and floats on) water to give an alkaline
solution and hydrogen, yellow flame, kept under oil
Na+
S
Sulphur
6
Yellow non-metal solid. Oxides make acid solutions
and contribute to acid rain. Molecule forms ring of 8, as
S8, though you don’t need to remember that.
S2-
Zn
Zinc
T
Not as reactive as aluminium or magnesium, a
protective oxide layer resists some corrosion and when
plated onto iron/steel, the zinc is a sacrificial anode.
and
co
Zn2+
It’s also worth remembering which are in groups 1, 7 and 0 and that all gases and
halogens (g7) come in pairs.
You could also benefit from running my formulas revision program, available as a
download from www.nicechocolate.plus.com
On the same website is a program for download called atoms, which is useful
revision.
index
A
H
Appendix, 13
atomic mass, 2
atomic number, 2
H = mcT, 11
half-equation, 2
Heat energy change, 11
B
I
balanced equation, 8
BODMAS rules, 4
isotopes, 3
isotopic ratios, 4
C
J
chemistry, 10
Just moles, 9
E
M
Empirical formulas, 5
exothermic, 12
F
Formul
a mass,
3
Formul
as, 1
Mass Number, 2
mole, explained, 8
molecular mass, 9
Molecular
mass, 3
Moles, 8
N
neutrons, working out, 2
Nitrogen, 13
G
P
gases, moles of, 11
Percentage mass, 5
periodic table, 1
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Page 14 of 16
phase change, 11
Proton Number, 2
S
Solutions, 10
R
W
Reacting quantities, 7
Reaction energy, 12
real masses, 10
Relative Atomic Mass, 4
website, 14
Working with Moles, 9
Appendix B Answers to
questions
1)
2xC, 2xO,
4xH
20 )
146
42 )
48g
21 )
44
43 )
224g
2)
2xK, 1xCr,
4xO
22 )
46
44 )
37.36g
3)
1xK, 1xMn,
4xO
23 )
7
45 )
0.786g
24 )
18
46 )
32g (31.998g)
6 x H, 3 x S,
12 x O
25 )
16
47 )
23g (22.99g)
26 )
17
48 )
207.2g
5)
1 x Fe, 6 x C,
12 x O
27 )
28
49 )
126.9g
6)
64 sulphur
atoms
28 )
74
50 )
1.019kg
29 )
40
51 )
1.4g
7)
Ca x 3, N x 6
and 18 x O
30 )
60
52 )
2g
31 )
74.5
(this depends
which periodic
table you use)
53 )
320g
54 )
256g
55 )
25.38g
32 )
16.014
56 )
480g
33 )
51.932
57 )
4.6g
34 )
196.93
58 )
688g
35 )
72.72%
59 )
385.6g
36 )
44.09%
60 )
1.6g
37 )
6.667%
61 )
48 litres
38 )
CH4
62 )
336 litres
39 )
CaCO3
63 )
0.48 litres
40 )
C3H8O
64 )
840J
41 )
24g don’t
forget the unit,
it’s a real
value.
65 )
9264J
66 )
(2000 x 4.2 +
800 x 0.386) x
4)
8)
One copper2+
ion and 2
electrons
making one
copper atom
9)
20
10 )
26
11 )
47
12 )
28
13 )
31
14 )
12
15 )
24
16 )
30 neutrons
17 )
22 neutrons
18 )
8 neutrons
19 )
143 neutrons
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Page 15 of 16
70 = 609,616J
67 )
498 – 1610 =
1112 kJ per
mole
(make C=O twice)
(break O=O oxygen)
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