Foundation Chemistry

Note: From 2008, this modue will be known as 'Atoms, Bonds and Groups'.

This stuff will give you the "foundation" on which you build up the rest of your chemistry knowledge. So, if you don't learn this, hard as it might at first seem, you're snookered later in the course...

Atomic Mass

Look at an atom in a periodic table and you'll see it has a atomic number and a relative atomic mass. The atomic number is just the number of protons. What exactly are protons?

Protons are the positive charges found in the nucleus of an atom. They define what the atom is. For example, any atom with 1 proton is a hydrogen atom. If you change to 2 protons, however, you have a helium atom. In an atom, protons equal electrons.

Electrons are the negative charges that whiz around the nucleus of an atom. These can be knocked off, or more can be added, to produce ions. (charged particles, in other words). They are so small that their weight is negligible. Therefore, atomic mass is just the number of protons and neutrons.

Isotopes: These have the same atomic number (same number of protons) but a different atomic mass. They have a different number of neutrons.

There are quite a few different isotopes occurring naturally. When we look at the atomic mass, we need to take them into consideration so therefore, we take the average of them. This becomes the relative atomic mass. Look at hydrogen's isotopes:

Therefore, it has a relative atomic mass of 1.0027.

To work out relative atomic mass, you need to know the abundance of each isotope. Then it's just a case of multiplying atomic number by abundance, adding together, and then dividing by 100.

Chlorine-35 has an abundance of 75.5% and Chlorine-37 has an abundance of 24.5%. What's the relative atomic mass?

35 x 75.5 = 2642.5
24.5 x 37 = 906.5

2642.5 + 906.5 = 3549

3549 100 = 35.49

Ever wondered how the mass of an atom is measured? It's done in a Mass Spectrometer. The name might sound scary, but in reality, there are only four main stages to it (which funnily enough, you NEED to learn):

  • Ionisation
  • Acceleration
  • Deflection
  • Detection

What it looks like:

Why do the electromagnets deflect atoms of a certain mass? Because it's easier to deflect the smaller ones than the larger ones (naturally). This means, by playing around with the magnetic field, you can make atoms of a particular mass travel towards the detector, while the rest hit the sides of the machine.

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They live underground, they're a brown colour, they burrow holes in fields... Hold on, you're talking about chemical moles? Ah.

So, the question on everyone's lips: what's a mole?

A mole of a substance is the mass required to give 6.023x10²³ atoms or molecules. (Avogadro's Constant).

They didn't just pull 6.023x10²³ out of the air. That's the number of atoms in 1 gram of Hydrogen.

Moles are really useful. Think about a balloon filled with a gram of hydrogen and a gram of lead. Both the same weight, but there are loads more atoms of hydrogen, because each one is smaller than the lead atoms. When you're reacting the same amount of two substances, you CAN'T use their masses.

Therefore, we use moles.

Chemists looked at 1 gram of Hydrogen and calculated that there were 6.023x10²³ atoms in it. Helium is 4 times larger than Hydrogen, therefore it needs 4 times the mass, and so, 1 mole of He = 4g.

However, there were a few problems with measuring against Hydrogen:

  1. It's a gas, and difficult to weigh.
  2. It's explosive.
  3. It doesn't exist as H (H2).

Instead of using Hydrogen as the measuring stick, chemists chose Carbon, a safe, plentiful solid. The relative atomic mass of all elements is relative to Carbon-12, which has a mass of exactly 12 grams.

To work out the number of moles in a solid, we use the following calculation:


How many moles are there in 21 g of Copper?

First step: Look to the Periodic Table and see what the Ar of Copper is. It's 63.55.

Second step: Divide 21 by 63.55

Finally: We get an answer of 0.33 mol.

How many grams are there in 0.32 mol of Carbon?

First step: Look to the Periodic Table to see what the Arof Carbon is: 12.

Second step: Multiply 12 by 0.32.

Finally: We get an answer of 3.84 grams.

Easy! J Remember, whenever you calculate anything a) show your working and 2) include the units (grams, mols, etc.)

(There was a joke just there, did you spot it?)

Now that's all very nice, but what happens when you have a compound, rather than an element?

Well, if you think about it, one molecule of a compound is just the mass of the atoms within it. This is called the relative molecular mass, or the molar mass.

Example: One molecule of CO2 is just 1 carbon and 2 oxygens. We know C has a Ar of 12 and O has an Ar of 16. Therefore the molar mass is 12 + 16 + 16 =
44 g mol-1.

The formula is identical to before but with Mr rather than Ar:

Moles in Gases

Because the particles in gas can move around freely, the same rule can't be applied for working out the moles. However, at a constant pressure and temperature, the number of moles in a particular volume of gas will be the same as in another equal volume of gas. Because of this, chemists found that a mole of gas occupies 24 dm³ (a huge volume: 24,000 cm³).

Using this, we have a formula:

How many moles are there in 203 cm³ of CO2?

Divide 203 by 24,000 = 0.008 mol

What volume of H2 occupies 0.04 mol?

Times 0.04 by 24 = 0.96 dm³ (960 cm³)

The only thing to remember here is that if you have a cm³ volume, you divide by 24,000 cm³ instead of 24 dm³.

Moles in Solutions

The particles in solutions are freer than solids, but not so much as gases, so another formula is required. We express solutions in the form of concentrations.

The concentration is always in moles per dm³ (mol dm-³ or M). Even though solutions are rarely made up to a volume of 1 dm³, their concentrations are always in the form above.

There are two equations you need to know:


How many moles are there in 250 cm³ of Hydrochloric acid with a concentration of 0.5 mol dm-³?

Step one: Multiply 250 by 0.5 = 125

Step two: Divide 125 by 100 = 0.125 mol

0.4 moles of Hydrochloric acid of concentration 0.8 mol dm-³ reacts with Sodium Hydroxide. What volume of HCl was used?

This one is trickier; it requires you twisting the formula...

Step one: Multiply the 0.4 mols by 1000 = 400.

Step two: Divide 400 by the 0.8 concentration = 500 cm³.

If there is 0.6 mol of NaCl in 230 cm³, what's the concentration?

Step one: Divide 0.6 by 0.230 = 2.6 mol dm-³.

The main thing to watch out for here again is whether your volume is in dm³ or cm³. You may need to convert them from one to the other, so be careful!

That's everything about moles. You must learn all the formulas. There is no question about it. Learn them, and then learn them again. Practice until you can quote them from heart. Then practice some more. They're essential.

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Empirical Formula

You should know all about chemical formulas, that tell us which elements are present in a compound. For example: H2O, H2SO4. Empirical formula is slightly different... The definition is "the simplest integer (whole number) ratio of elements in a compound."

Chemical Formula Empirical Formula
C6H12O6 CH2O
C2H6 CH3

To calculate Empirical Formula (in 5 easy steps!):

  1. Identify the elements present in the compound.
  2. Write down their masses.
  3. Convert the grams to moles. (told ya moles were important!)
  4. Divide by the smallest number of moles.
  5. Ratio up or down to form whole numbers.


What is the empirical formula when 20.7 gram of Lead reacts with 1.6 gram of Oxygen?


20.7 g

mol = 20.7207

= 0.1 mol



1.6 g

mol = 1.616

= 0.1 mol


=> PbO

If you're told the molar mass of a compound, you can then work out it's molecular formula. Take the next example:

A compound has a molar mass of 570 gmol-¹ and has the following composition: Phosphorus 10.88% and Iodine 89.12%. What's the molecular formula?


10.88 g

mol = 10.8831

mol = 0.351


89.12 g

mol = 89.12127

mol = 0.701


Empirical formula = PI2

When you have the empirical formula, you can work out a empirical mass using the Molar Mass = (31 + 2x127) = 285

Then see the Molecular Mass = 570. This is 2x the Empirical mass.

Therefore, the formula is doubled: P2I4.

Molecular Formula is P2I4.

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Balancing Equations

Equations are really useful. They tell us the reactants (what elements react with what), the products (what they produce) and the quantities of each.


Looking at an equation:

  • The reactants are in green. They are separated by a +.
  • An arrow shows a reaction has occurred.
  • The product(s) is in blue. If there's more than one, they are also separated by a +.
  • The state symbols tell us:
    (s) - solid

    (l) - liquid
    - aqueous (dissolved in water)
    - gaseous
  • The formula also tells us the number of moles present. In the example, we have:

See how important moles are? From this, we can calculate that:

(4x23) g of Na + (2x16) g of O2 -----> 2(23x2+16) g of Na2O.

=> 92 g of Na + 32 g of O2 -----> 124 g of Na2O.

We can then ratio up and down, to find out how much of one reactant reacts with another and how much product is produced.

How much Na2O is formed from 8 grams of sodium?

Step 1: 92 g of Na + 32 g of O2 -----> 124 g of Na2O ... divide everything by 92.

Step 2: 1 g of Na + 0.348 g of O2 -----> 15.5 g of Na2O ... times everything by 8.

Finally: 8 g of Na + 2.78 g of O2 -----> 124 g of Na2O

We have an answer: 124 g of Na2O.

In order to do this, though, the equations need to be balanced. Balancing is a nice topic you should remember from GCSE.

Golden Rule: NEVER change the little numbers in the equations.

A classic example:

H2O and H2O2 may look similar


H2O is water and H2O2 is Hydrogen Peroxide. If you had a shower in that, it'd strip away your skin!

Add BIG numbers only, at the beginning of the reactant/product.

Balance the following:

  1. Mg + O2 -----> MgO
  2. CH4 + O2 -----> CO2 + H2O
  3. H2 + I2 -----> HI
  4. Al + O2 -----> Al2O3

See answers here.

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Ionisation Energies

Let's just remind ourselves of the structure of an atom...

  • Atoms are spherical.
  • Most of an atom is empty space.
  • You have the tiny nucleus in the centre made up of protons and neutrons where most of the mass is concentrated.
  • Electrons occupy energy levels around the nucleus (also known as shells).

In GCSE, you were taught that electron orbitals went in the following pattern: 2,8,8... This is WRONG. You were lied to, and even at A-level, we don't get to know the full story (hey, but if anyone does know the truth out there, I would really like to know!). However, our version is closer to the truth than GCSE.

We get proof for electron arrangement from ionisation energies. This as you might have guessed, is something to do with making ions.

The first ionisation energy is the energy required to knock 1 mole of electrons from one mole of gaseous atoms, to give 1 more of gaseous ions. (It must be in the gas state).

That might sound complicated (you have to learn it, lucky you) but if we write it as a formula it appears so much simpler.

Mg(g) ----> Mg(g) + e

In other words, how much energy is needed to remove the electron and make an ion?

For the second ionisation energy, we are simply removing another mole of electrons from the gaseous ion:

Mg(g) ----> Mg²(g) + e

And naturally the third ionisation energy is:

Mg²(g) ----> Mg³(g) + e

And so on...

When chemists looked at this, they saw patterns emerging. Look at the ionisation levels of sodium:

There is a jump in energy after the first electron, and then after 8 electrons, a second jump to the final two. This is clear evidence for the 2,8 theory. So what are the factors that determine how much energy is required to knock off electrons?

  1. Nuclear charge - the larger the element, the stronger the positive pull of the protons in the nucleus.
  2. Distance - The closer an electron is to the positive pull of the nucleus, the harder it is to knock away. (easier the further the electron is from the nucleus).
  3. Electron shielding - complete shells of electrons mask the positive pull of the nucleus, making it easier for electrons to be knocked away.

A key thing to remember here is that 2 and 3 far out weight 1.

GCSE in dispute

So far, it's all sounded pretty good for the old 2,8,8 theory. However, when chemists compared the first ionisation energies of elements in the periods of the periodic table, they found some strange results...

As you can see, the general trend increases as the mass increases. The size of the atom actually decreases though, and this is why it gets harder to knock off the electrons. (Because the nuclear charge increases, the outer electrons - which are all in the same shell - are drawn closer).

But the ionisation level dipped slightly on Boron and Oxygen, and in other periods, they found the same dip on elements in the 3th and 6th group. This couldn't be explained.

They decided it had to be something to do with the second energy level of the atom. The electrons couldn't be in a simple 8 orbital.

Instead, they split it into two subdivisions, one that held 2 electrons, one that held 6. The configuration now looks like this:

Because Boron has a single electron in the 6-sublevel, it is slightly easier to remove. However, once another electron enters, it increases again.

With Carbon, things get even more complex. In the 6-sublevel, elements naturally try to get as far away from each other as possible. Everything is okay with the first three elements that fit into it, but when a fourth comes along, it has to double up with one of the others. Because it's doubled, it is slightly easier to remove (as its electron partner is repelling it too). This occurs on the 10th electron, or in group 6.

Just learn the explanation, no matter how unconvincing it may seem. It'll get you through your exams, right... and if you want to know the truth, you'll need the A-Level stepping stone for university...

To describe the new arrangement of electrons we have completely random letters (as usual).

  • 2 electron shell = s
  • 6 electron shell = p
  • 10 electron shell = d
  • (14 electron shell = f)

They look like this:

The main thing to notice is that 4s comes before 3d, because electrons, in reality, don't circle in nice circles.

When writing out electron configuration like this, you need to know how many electrons the element has. It's a matter of filling up each shell. If it has 8 electrons, you can fill up 1s with 2 electrons, so you write: 1s², you can fill 2s with 2 electrons so you write that as: 2s². Finally you have 6 electrons remaining, which all go into 2p6. All written together you have: 1s² 2s² 2p6.

To write out the new configurations of some well known elements:

H = 1s¹
He = 1s²
Li = 1s² 2s¹
Be = 1s² 2s²
B = 1s² 2s² 2p¹
Ne = 1s² 2s² 2p6
Ca = 1s² 2s² 2p63s² 3p64s²
Fe = 1s² 2s² 2p63s² 3p64s² 3d6
Kr = 1s² 2s² 2p63s² 3p64s² 3d10, 4p6

The exam might ask you to draw electron orbitals with arrow-boxes. They look like this:

Each arrow displays an electron. Here you see the 1s22s23p64s2.

The s orbital is just a standard circle, but the p looks very different, as you can see below. Make sure you learn these shapes, in case they ask you to draw them in the exams.

Notice the fuzzy lines? That's because electrons travel incredibly fast, close to the speed of light.

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Ionic Compounds

There are two types of bonding, covalent and ionic. Ionic is where atoms gain or lose electrons and thus become ions.

  • Ionic bonds form between metals and non-metals. (Metals lose electrons, non-metals gain 'em).
  • They are usually non-molecular (They don't have a beginning or an end).
  • They exist in an ionic lattice (regular structure).
  • They conduct when molten or in solutions because of free ions.
  • They have high melting and boiling points because of their lattice structure and the strong electrostatic attraction of bonds. A large amount of energy is requires to break all the bonds in the huge lattices.
  • They are usually hard, but brittle, solids.
  • Often dissolve in water. E.g. salt.

Ever wondered where the little numbers come from in ionic compounds... such as Ag2O? You should have done it in GCSE... Ionic compounds are all about positive ions balancing negative ions.

First off, let's have a look at the following table, showing common ions and their charges. You really need to learn these...

Sodium Na   Hydroxide OH
Lithium Li   Nitrate NO3
Potassium K   Fluoride F
Ammonium NH4   Chloride Cl
Magnesium Mg²   Bromide Br
Calcium Ca²   Iodide I
Barium Ba²   Sulphate SO4²
Silver Ag   Sulphite SO3²
Iron (II) Fe²   Oxide
Iron (III) Fe³   Carbonate CO3²
Lead Pb²   Sulphide S²
Copper Cu²   Phosphate PO4³

From these, we can make all kinds of ionic compounds. You just need the positive charges to equal the negative charges.

What would be the formula of Magnesium Oxide?

Mg ion has a charge of 2+ and O a charge of 2-. These are the same, and so therefore, the formula is: MgO.

What is the formula of Iron(II) Iodide?

Fe(II) has a charge of ² and Iodide a charge of . To balance the charges, we need 2 iodides for every Iron. This makes the formula: FeI2.

What's the formula of Sodium Hydroxide?

Na has a charge of and OH of . They are equal so the formula is: NaOH.

Three general rules for dealing with ionic compounds:

  1. Swap the numbers around.
  2. Ignore any 1s.
  3. Ignore equal numbers.

ALWAYS remember, the little numbers are written at the bottom of the element, not above it.

A tricky one might require brackets:

What's the formula of Copper Nitrate?

Cu has a charge of ² and NO3 has a charge of . We need two NO3 for every Cu. Since NO3 is a compound already, it needs brackets: Cu(NO3)2.

Positive ions are formed when atoms lose electrons, and obviously negative ions are formed when atoms gain electrons. Why does it occur? Well, as usual, the atoms are trying to gain a full electron shell, so they can be stable.

The usual way of drawing this is with a dot-cross diagram:

As you can see for sodium chloride, Na no longer has an electron in its outer shell, but Cl now has a full electron shell.

If you're doubt about charges on ions, look to the periodic table. The groups tell us how many electrons an atom needs to lose or gain to fill it's shell. e.g. Mg is in group II and needs to lose 2 outer most electrons; Cl is is group VII and needs to gain 1 electron.

Ionic Equations

Sometimes chemists shorten a full equation to its ionic equation. This is just the part of the equation where something is happening.

For example:

NaOH(aq) + HCl(aq) -----> NaCl(aq) + H2O(l)

Step one: Split into ions, leaving the covalent molecules:

Na + OH+ H + Cl ------> Na + Cl + H2O

Step two: Cross out any ions that appear on both sides of the equation. This leaves:

OH+ H-----> H2O

Half Equations

Occasionally, they might give you an equation and ask you to write its half equations. This is all about the changes in charge that have occured to each element in the equation. For example, for Mg + F2 -----> MgF2 the half equations would be:

Mg -----> Mg2++ 2e- (Mg has lost two electrons to become 2+ positively charged).

F2 + 2e- ----> 2F - (F2 has gained two electrons to become 2 lots of -1 charged floride).

Distortion of the Ionic Bond

Just to be awkward, sometimes you can get compounds that aren't sure whether they want to be ionic or covalent. They have properties of both. This occurs under certain conditions:

  • The catoin needs a high charge density (in english, this means that the positive ion needs to be small, and the more positive, the better).
  • The anion needs a low charge density (more translation: the negative ion needs to be big and the more negative the better).

Let's have an example for the befuddled masses. Compare NaCl to AlCl3:

Both Na and Cl are roughly the same size and have the same positive and negative pull. This is called a pure ionic bond.
Because Cl is big in comparison to Al, the negative electron in the outer shell is drawn to the very strong pull of Al. (less attraction from its own nucleus). This causes distortion, or even an overlap, as you see here. Partial covalent bonding.

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Covalent Bonding

We've already come across ionic bonding, where elements lose or gain electrons to fill their electron shells. Covalent bonding is a second type of bonding that requires elements to share their electrons.

  • Covalent bonds only occur between non-metals.
  • Covalent compounds are usually molecular (they have a definite beginning an end).
  • Low melting and boiling points because only Van der Waal forces hold them together.
  • They can't conduct electricity because there are no free ions or electrons.

There are three ways of showing covalent bonding:







Lone pairs are the electron pairs that aren't included with bonding. We always draw them on, however.

Shared pairs are then naturally pairs included in the bonding, with a x and a .

A dative covalent bond is a bond between two elements where both electrons come from only one atom. (the other is an ion).

How do you figure out how to draw them?

  1. Look at what group the element is in in the periodic table. If it's in group five, there are five electrons in its outer shell, group six, six... etc.
  2. How many more electrons does each atom need to be complete? If both need one, obviously, they'll just share one covalent bond.
  3. If they don't seem to bond easily, you may need more of one atom.

Carbon dioxide, for example is:

, O=C=O, CO2

Shapes of Molecules

We all know that electrons repel each other and try to arrange themselves as far apart of possible. In covalent bonds, the electron pairs determine the shape of the molecule. The electron repulsion theory may sound like a horrible mouthful, but in reality, it's fairly simple:

  1. Electrons repel each other and try to arrange themselves as far apart as possible. This means that molecules will move into 3d if necessary.
  2. Lone pairs repel more than bonding pairs. (this is why it's important to draw on the lone pars).
  3. Double or triple bonds repel the same as single bonds.

Okay, so 3d, hmm? Take CH4 (Methane) as an example. In a book, the shape is drawn like this:

This isn't correct though. In reality, it looks like this:

Its shape is called a tetrahedral and has bonding angles of 109.5° -- larger than the 90° angles it would have had if it were 2d. (Bond angles are just the angles between two bonding arms).

So, how do we draw a quick hand of a 3d shape?

The shows a bonding arm that is coming towards you, while the shows a bond arm going behind. Easy enough, if you can visualise it like the picture above.

There are a few examples you should really learn, because the examiners love asking about them.

Molecule Dot-Cross Shape Sketch Bond Angles Shape Name
CH4 109.5° Tetrahedral
NH3 107° Pyramidal
BeCl2 180° Linear
H2O 104.5° Non-linear
BF3 120° Triganol Plainer


  • No lone pairs: Tetrahedral (4 bonding pairs), Linear (2 bonding pairs) or Triganol Plainer (3 bonding pairs).
  • 1 lone pair: Pyramidal
  • 2 lone pairs: Non-Linear
  • 3 lone pairs: Lineal

Giant covalent bonding

While covalently bonded elements are usually molecular (with the definitive beginning and end) there are of course exceptions. In giant molecular bonding, they share electrons throughout the substance, and are thus arranged in a lattice, like in ionic bonding.

Some common examples are:

  • Diamond
  • Graphite
  • Silicon Dioxide (sand)
<- The structure of Silicon Dioxide.

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Metallic Bonding

A third type of bonding is called metallic bonding and, surprise, surprise, it occurs in metal elements (all the same atoms). In metals, the outer electrons are only held very loosely, and thus become delocalised. They move freely through the element, acting as a "cement" to hold the positive nuclei (or they can be called positive ions) together.

  • Positive nuclei are arranged in a regular lattice.
  • A "sea" of delocalised electrons surrounds the positive nuclei.
  • Because the delocalised electrons are free to move (and carry a charge), metals are great conductors, in solid or molten state.
  • The boiling and melting points are generally high because of the electrons "cementing" the nuclei together.
    • The more electrons in the outer shell, the stronger they're held together.
    • The closer the delocalised electrons are to the nuclei, the stronger the attraction.

A nice example:

That's really all there is to it. Make sure you can draw that diagram and label the positive lattice of nuclei and the sea of electrons!

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Intermolecular Bonding

Ever wondered why a solid's a solid or a liquid's a liquid or a gas is... well just gas? It's all about the intermolecular bonding, basically the bonds that hold one molecule to another. We look at three types of intermolecular forces, Van der Waals forces, dipole-dipole forces, and hydrogen bond. They might sound like a mouthful, but they're relatively straightforward.

Van der Waals forces

It's a well known fact that even the simplest molecules attract one another. (Don't look at me so blankly!)

Scientists started experimenting with the very simplest, Helium (not even a molecule, but an atom) and saw that by cooling it enough, it would turn from gas to liquid. However, this liquid formed a puddle at the bottom of the container, and this made them conclude that there must be something holding the He atoms together.

They didn't have a lot to choose from, because all He has are: 2 electrons, 2 protons and 2 neutrons.

One bright spark called Van der Waal came up with a solution. He believed that the random movement of electrons whizzing in the outer shell, would occasionally cause the atom to become charged at one side (if all electrons went to the left of the atom, the right would be slightly positive). He called this a temporary dipole. This then caused a knock on affect on an adjacent atom, as its electrons would be attracted to the positive pull. This was called an induced dipole.

And wow, for a split second, let there be attraction! And so, these forces became known as Van der Waals forces. They only last a split second, but every time one is over, another is happening somewhere else. This holds molecules together.

We show these charges on the molecules with the greek letter delta: δ (delta). Remember, the charges don't make them ions.
  • Van der Waals forces are the weakest type of intermolecular forces.
  • The more electrons in a molecule, the stronger the Van der Waals.
  • The shape of the molecule also affects the Van der Waals. The closer a molecule can get to others, the stronger the forces.

Dipole-dipole bonds

So what exactly is a dipole?

A dipole is a charge separation. Look above at the two δs in Van der Waals. That change in charge is called a dipole.

When you have covalent bonding like those just mentioned, you rarely get pure covalent bonds, such as H2:

Both nucleuses have identical change, and so pull the shared pair of electrons evenly.

Instead, we might have the bond pulled slightly in one direction. Most of the time, a slight pull isn't enough to make any difference. There are only three elements in which it does: Nitrogen, Oxygen and Fluorine. These are called electronegative elements and they create a dipole by pulling the covalent bond away from the element they're sharing with.

A couple of examples, to show how it works:

Here, Nitrogen pulls the electron bonds from each hydrogen, making it slightly negative while the hydrogens are positive. Oxygen pulls the electron pairs towards it, making them slightly positive. Notice how only elements connected to O are affected though.

Molecules containing dipoles are said to be polar.

When you get a group together, the positive deltas of one molecule are attracted to the negative deltas of another. Thus, we have attraction. The bonds are stronger than Van der Waals and are called dipole-dipole, or permanent dipole-dipole bonding.

This shows two molecules, attracted to each other with dipole-dipole attraction. It is usually drawn on as three dots. ...

You need to be careful though, because not all compounds containing electronegative atoms are polarised. A good example of when it doesn't happen is BF3:

As you can see, the electronegative element F is surrounding B. But because the three Fs are pulling equally, they cancel out each others' dipoles.

For it to be polarised, there needs to be a positive end and a negative end.

Hydrogen Bonds

Finally, we come to the strongest type of intermolecular bonding, Hydrogen bonding. As you can probably guess, it involves hydrogen. Hydrogen bonding is basically the same as dipole-dipole bonding, but where the electronegative element (NO or F) is joined to a H.

The most classic example is water. It has oxygen bonded with two hydrogens. Why does this result in something stronger than dipole-dipole?

It's all about hydrogen. Hydrogen has a single electron. When NO or F pull its electron away, it results in an exposed nucleus which is very positive. You have another molecule adjacent with a lone pair (for example in oxygen) and there is a strong attraction.

Here you can see the structure of H2O when bonded with hydrogen bonds. The H nuclei are exposed, causing strong attraction between them and the lone pairs.

What effect do dipole-dipole, or hydrogen bonds cause? Well, for a start they cause the molecules to have a much higher boiling/melting point than expected because the energy required to break the bonds is much greater.

Ice floats

Ever wondered why? It's all about the hydrogen bonds. When water is in liquid state it looks like this:

All the molecules are clustered together, and compact. When they freeze however, they go into the regular structure above with lots of gaps. This makes ice less dense, and so it can float on water. Pretty cool, huh?

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Oxidation and Reduction

Back in the day, they would have told you something like:

Oxidation is gain of oxygen and reduction is loss of oxygen.


Oxidation is loss of hydrogen and reduction is gain of hydrogen.

(or neither if you did single science...)

These definitions are okay, but it goes further than that. Remember the old OIL RIG? Oxidation Is Loss of electrons and Reduction Is Gain of electrons.

But what happens if an atom loses oxygen but gains electrons? See... it can get complex. Fortunately, Chemistry has a way of making things simpler (doesn't happen very often, I know.) It gives everything an Oxidation Number and by seeing how these change, we can determine whether an atom has been oxidised, reduced or stayed the same:

  • Oxidation occurs when the Oxidation Number increases.
  • Reduction occurs when the Oxidation Number decreases.

Sadly, to work out what the ON is, you need to know a load of rules:

  1. The ON of a free element is always zero. e.g. the ON of Mg = 0.
  2. The combined ON of atoms in a compound is also 0. e.g. NaCl2 = 0.
  3. In compounds, ON in Group I = +1, Group II = +2,etc..
  4. In compounds, H = +1 (except in metal hydrates, such as NaH where it would be -1).
  5. In compounds, O = -2 (except in peroxides).
  6. In compounds, F/Cl/Br/I = -1 (except when combined with O or F).
  7. The ON of a singular ion is just its charge. e.g. Na= +1 and Cl = -1.
  8. The sum of all atoms in a polyatomic ion (ion made of more than one atom) is equal to the charge on the ion. e.g. CO3. O has an ON of -2 so 3x-2=-6. To leave a -1 ion, we need C to be +5.

Um, did I say this made it EASIER? Well, it is, once you get your head around the rules.

Let's do a few examples, to show you how it works:

What is the oxidisation number of Ba in BaCl2?

We know from the Periodic Table that Ba is in group 2. Therefore it must have a ON of +2. Easy.

...of P in P2O3?

Compounds have a total ON of 0.

O has a ON of -2. For O3: -2x3=-6.

Therefore, P2: +6

P = +3.

...of P in PO4³?

ON of ion = -3.

ON of O = -2. For O4: 4 x -2 = -8.

Therefore, P must equal +5.

Remember to always include the + and - to numbers!

In the redox equation below, calculate the oxidisation numbers of Pb and Sn and state which has been oxidised and which has been reduced:

PbO2 + 4H + Sn² -----> Pb 2 + Sn4+ 2H2O

ON of PbO2 = 0.

ON of O = -2 and so O2: -2 x 2 = -4.

Therefore, ON of Pb must be +4.

ON of Pb2= +2

=> ON has decreased so Pb has been reduced.

ON of Sn2 is same as charge on ion = +2.

ON of Sn4 = +4

=> ON has increased so Sn has been oxidised.

To summarise:

Oxidation Number (or state) is the number of electrons that an atom has lost gained or shared when forming bonds. It can be positive, negative of zero.

Oxidation is when the ON increases, whereas reduction is where the ON decreases.

An oxidising agent oxidises other substances and a reducing agent reduces other substances (obviously).

Disproportionation (what a word!) is where the element is both oxidised and reduced at the same time. It does happen...

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The Periodic Table

I've mentioned the Periodic Table a few times, but what exactly is it? Well, here it is:

The Periodic Table is a very useful tool for chemists. It shows:

  • The divide between metals and non-metals. All metals are on the left hand side, while non-metals are on the right.
  • The groups that elements are in. Elements in the same group have similar properties.
    • Group I is called the Alkali Metals.
    • Group II is called the Alkaline Earth Metals.
    • Group VII is called the Halogens.
    • Group 0 is called the Noble Gases.
  • It shows the periods, which also have repeating patterns in them (which we see from the study of Period Three very soon).
  • It shows all the elements arranged in order of increasing atomic number.
  • It shows us the sub-levels of elements:
    • The s-block contains Group I and II, eg: Li = 1s2 2s1.
    • The p-block contains Group IV to 0, eg: Ne = 1s2 2s2 2p6
    • The d-block contains mostly transition elements, eg:
      Fe = 1s2 2s2 2p6 32 3p6 4s2 3d10.

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Period Three

Period Three is of course referring to the third period of the Periodic Table.

Don't ask me why the exam board decided to look into this period in particular, but we're stuck with it...

So, there are a few facts you need to know, along with why.

As you go along the period, the atomic radius of elements decreases. This is because while you're adding more electrons, they all occupy the same shell. The nucleus gains an extra proton, and so there's an increased nuclear charge. This increases attraction of the nucleus on the outer electrons, and so reduces the size of the atom.

Electrical conductivity changes along the period.

Na, Mg and Al are metals and so have metallic bonding. They are all good conductors, and as you add an extra electron to the outer shell the conductivity increases.

Si is a macromolecular structure, with no free ions or electrons. However, impurities can cause it to be a semiconductor.

P, S, Cl and Ar are all simple covalent molecules. They have no free ions or electrons and so are poor conductors.

Melting point changes too...

It increases from Na up to Al because in metallic bonding, the more electrons in the outer shell, the stronger the attraction holding them together.

Si has the biggest boiling/melting point because it's a macromolecular structure with very strong covalent bonds throughout.

P4, S8, Cl2 and Ar only have Van der Waals holding them together, meaning they all have low boiling points. The boiling point decreases along them, with the exception of S8 (because there are 8 of it, there are lots of electrons to make Van der Waals forces with).

Remember the ionisation pattern of the First Period from earlier? It's identical here:

The ionisation energy increases across the period, with the exception of the third (Aluminum) and sixth (Sulphur) elements, where the ionisation energy dips a little.

For Aluminum, this is because we've moved from the 's' sub-level into the 'p'. The electron is a little further away from the nucleus and so a little easier to be removed. For Sulphur, we have two electrons in the 'p' sublevel pairing up for the first time. This additional repulsion makes the 6th electron that bit easier to be removed.

Check the Ionisation Energies section if you're not sure!

To summarise Period Three: (you don't need to learn all the numbers, just to illustrate...)

Electron Config.
Atomic Radius (nm)
Physical State
(at room temp)

Electrical Conductivity

Melting / Boiling Point
Sodium Na 1s²2s²
0.191 Solid Metal Good 386K
Magnesium Mg 1s²2s²
0.160 Solid Metal Good 922K
Aluminum Al 1s²2s²
0.130 Solid Metal Good 933K
Silicon Si 1s²2s²
0.118 Solid Non-Metal Semi-conductor 1683K
Phosphorus P 1s²2s²
0.110 Solid Non-Metal Poor 317K
Sulphur S 1s²2s²
0.102 Solid Non-Metal Poor 371K
Chlorine Cl 1s²2s²
0.099 Gas Poor 172K
Argon Ar 1s²2s²
2p3s² 3p
0.095 Gas Poor 84K

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Group Two

And to represent the exam syllabus, Group 2 has been selected! Well, at least these dudes have a proper name - as mentioned above, they're called the Alkaline Earth Metals. There are a few facts you need to know:

  • The outer most electron configuration is the same for each element: n s². For example, Mg: 1s2 2s2 2p6 3s2 3p2 4s2.
  • They're all metals, meaning they have metallic bonding and they also bond with non-metals in ionic bonding. (losing their two outer electrons to become +2 ve).
  • Because they're all trying to lose two electrons, they have similar chemical properties.
  • Atomic radius increases as you go down the group. Yes, you have a stronger nuclear charge, BUT with each element down, you're adding another complete shell of electrons. This creates electron shielding, which 'masks' the pull of the nucleus.
  • Ionisation energy decreases because the distance from the nucleus and the electron shielding make it easier for outer electrons to be knocked off.
  • This also means that reactivity increases as you go down the group.

All Group 2 metals burn vigorously with Oxygen to form their oxides (white solids).
e.g. 2Mg(s) + O2(g) -----> 2MgO(s).

They will also react with water to produce hydroxides and hydrogen gas.
e.g. Ca(s) +2H2O(l) -----> Ca(OH)2(s) + H2(g).
Magnesium is an exception though. It reacts with steam to form:
Mg(s) +H2O(g) -----> MgO(s) + H2(g).

They also react with hydrochloric acid. They fizz and produce hydrogen gas, and then vanish into a colourless chloride solution.
e.g. Mg(s) + 2HCl(aq) -----> MgCl2(aq) + H2(g).

Reacting the oxidised group two element with hydrochloric also causes it to disappear into the colourless chloride solution, and water.
e.g. MgO(s) + 2HCl(aq) -----> MgCl2(aq) + H2O(l).

And finally, reacting the carbonate form of the element causes it to fizz (producing carbon dioxide) and then vanish into the colourless chloride solution and water.
e.g. MgCO3(s) + 2HCl(aq) -----> MgCl2(aq) + CO2(g)+ H2O(l).

Calcium Carbonate

Calcium Carbonate... CaCO3... something you might vaguely remember from GCSE... limestone... loads of reactions that you had to learn... And guess what? They return to haunt you at A-Level too. Muahahaha! Lots of "just got to learn 'em" facts coming up. Fortunately, the entire topic can be illustrated in one pretty picture. Learn to draw it by heart:

Group 2 uses:

  • **Ca(OH)2 is spread on soil to neutralise acidic soils.**
  • **Ma(OH)2 is used in indigestion tablets as an antacid (anti-acid).**
  • Mg is used on fireworks or in sacrificial protection.
  • MgO is used as lining for furnaces.
  • CaCO3 is used in cement.
  • CaO is used in mortar.

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Group Seven

Finally, (yes, I really did say that word - we've almost finished Foundation Chemistry!) we have the interesting Group 7 to study. You'll probably remember the good ol' Halogens from GCSE...

  • They're all non-metals.
  • They float around in pairs to become more stable: F2, Cl2, Br2, I2.
  • They all form ionic bonds, trying to gain one outer electron. As you move up the group, reactivity increases because it's easier for the nucleus to pull in electrons from other atoms if there's less distance and less electron shielding. The opposite of group 2.
  • The atomic radius increases as you move down the group (for the same reasons as explained in group 2).
  • Physical appearance is interesting:
    • Fluorine is a yellow gas.
    • Chlorine is green gas.
    • Bromine is a red liquid.
    • Iodine is a purple solid.
  • When either Bromine or Iodine are displaced from a reaction, they appear brownish.
  • Melting point and boiling points increase as you go down the group. Van der Waals forces hold them together. Because Iodine is big, with lots of electrons, it is held strongly together enough to stay as a solid at room temp, while Bromine's intermolecular forces keep it liquid. F and Cl are both gases.
  • Electron configuration always ends with n p5. e.g. Cl: 1s2 22 2p6 3s2 3p5.

Reactions of Halogens

Displacement: Halogens higher up can "kick out" those beneath them and steal their solutions.
For example: Cl2+ 2KBr -----> 2KCl + Br2
It's simpler just to learn the Ionic Equations:

  1. Cl2+ 2Br -----> 2Cl + Br2
  2. Cl2+ 2I -----> 2Cl + I2
  3. Br2+ 2I -----> 2Br + I2

To test for these halide ions, you just add silver nitrate (AgNO3):

  2. Cl-- Ag(aq) + Cl(aq) -----> AgCl(s) White precipiate formed, which is soluble in concentrated or dilute ammonia solution.
  3. Br-- Ag(aq) + Br(aq) -----> AgBrl(s) Cream precipitate formed, which is soluble in dilute ammonia solution only.
  4. I-- Ag(aq) + I(aq) -----> AgI(s) Yellow precipitate formed, which is insoluble in ammonia solution.

Disproportionation reactions:

Cl2 + 2NaOH -----> NaCl + NaClO + H2O (the Na in NaCl is reduced, while the Na in NaClO is oxidised.)

Cl2 + H2O HCl + HClO (Cl in HCl is reduced, while Cl in HClO is oxidised.)

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And that's pretty much everything you need to know!