AS-Level Atomic Structure

Structure of an Atom

  • Atoms are the smallest particles of an element.

  • Atoms are made of sub-atomic particles:

    • ​protons (charge +1, relative mass 1)

    • neutrons (no charge, relative mass 1)

    • electrons (charge -1, relative mass 1/1840)

  • The number of protons an atom has is called its ‘atomic number’.

  • The centre of an atom is called a nucleus and contains just protons and neutrons.

  • Electrons are attracted to the positive charge of a nucleus and exist in orbitals.

  • Atoms can lose or gain electrons to become charged ions.


Electron Orbitals

  • Electrons are attracted to the positive charge of a nucleus (electrostatic attraction).

  • The closer electrons are to a nucleus, the lower their energy.

  • Electrons want to get close to a nucleus but they repel each other.

    • They can only exist in set regions called orbitals.

  • Different orbitals have different shapes and therefore different energies.

  • Sub-shells describe how far away an electron (and its orbital) is from a nucleus.

  • Electron sub-shells are filled successively – lowest energy sub-shells and orbitals are filled first.

The Mole

  • Atoms and particles are very small, so referring to actual numbers of atoms in a sample would require very large numbers that would be hard to work with.

  • The mole is just a unit used to refer to the amount of atoms or particles in a substance.

  • Just like one dozen = 12 of something, one mole = 6.022x10  of something.

  • The number one mole refers to is known as Avogadro’s constant.

Relative Atomic Mass

  • The mass of an atom is determined by the number of protons and neutrons it has.

  • Different elements have different masses. 

  • An element’s ‘relative atomic mass’ is the mass (in grams) of one mole’s worth of its atoms.

  • One mole of carbon-12 (12C) has a mass of 12g. 

    • Each atom of carbon-12 has 6 protons and 6 neutrons.

      • The mass of each must be 1g per mole.

  • Therefore, atomic masses of elements are described as being ‘relative to 1/12 the mass of an atom of carbon-12’.

Moles by Mass

  • The number of moles in a sample can be calculated using the formula:

    • number of moles = mass (g) / molar mass

  • The molar mass (relative molecular mass or relative formula mass) of a compound is calculated by adding up all the relative atomic masses of the atoms within the compound’s formula.

Moles in a Solution

  • The number of moles of a solute in a solution can be calculated using the formula:

    • number of moles = concentration (mol dm  ) x volume (dm ) / 1000

  • 1 dm  is the same as 1000 cm

  • To convert dm to cm x 1000

  • To convert cm to dm / 1000 

  • Concentration refers to how many moles of solute are in 1dm  of solution (units are mol dm  )

Moles of a Gas

  • The ideal gas equation is used to calculate the moles of gas present in a system if the volume  (m ), pressure (Pa) and temperature (K) are known.

    • pV = nRT

      • n = moles

      • R = gas constant, 8.31 Jmol K 

  • The equation assumes that the same number of moles of all gases occupies the same volume at a set temperature and pressure.

  • At room temperature (293K and 101KPa), it is assumed that 1 mole of any gas occupies a total volume of 24000cm (24dm ).
















AS-Level Bonding


  • Atoms can bond together to create larger species.

  • Bonds that hold atoms together are called atomic bonds. 

  • Forces between molecules are called intermolecular forces.

    • Intermolecular forces are broken and formed when molecular substances change state (solid, liquid and gas)


Covalent Bonds

  • Atoms are most stable when they have a ‘full’ outer shell of electron (two electrons for hydrogen, eight electrons for other elements).

  • To have access to full outer shells, atoms ‘share’ electrons between orbitals. 

  • A covalent bond is the sharing of one pair of electrons between two atoms.

Ionic Bonding

  • To get a full outer shell, atoms lose or gain electrons.

  • Metals lose electrons to get a full outer shell and become positive ions. 

  • Non-metals (usually) gain electrons to get a full outer shell and become negative ions.

  • Ionic bonding refers to the attraction between oppositely charged ions.

  • The greater the charge of the ions, the greater the attraction and the stronger the ionic bonding between the ions.

Metallic Bonding

  • Metal atoms are more stable if they lose electrons to have a full outer shell.

  • Electrons lost by metal atoms in solid metal have nowhere to go and stay attracted to the positive charge of the metal ions.

  • Electrons are not held by the metal atom but are free to move - they are delocalised.

  • In metallic bonding, positive metal ions are attracted to the negative charges of the delocalised electrons and this gives metals a dense, rigid structure.

Intermolecular Forces

  • Intermolecular forces exist between molecules

  • When molecular substances change state, it is the arrangement of intermolecular forces that change (not the covalent bonds inside the molecules) 

  • Temporary induced dipole-dipole forces (Van der Waals, London forces, Dispersion forces) exist between all molecules

    • Unequal electron distribution around a molecule creates an instantenous dipole that can induce a dipole on a neighbouring molecule.

    • The two opposite dipoles from two molecules are attracted to each other and a weak force occurs

    • They are the weakest intermolecular forces

Permanent Dipole-Dipole

  • If a molecular substances is polar (has polar bonds and an overall uneven distribution of electron density) permanent dipole-dipole interactions exist between molecules

  • The slightly negative part of one molecule is attracted to the slightly positive part of another molecule.

  • The attractions are permanent and are stronger than temporary, induced dipole-dipole forces

  • The greater the polarity of a bond or molecule, the stronger the permanent dipole-dipole interactions

Hydrogen Bonding

  • Hydrogen bonds are a type of permanent dipole-dipole force that are very strong

  • They exist between molecules that have a N-H, O-H or F-H bond

  • Nitrogen, Oxygen and Fluorine are the most electronegative elements and the covalent bonds between them and hydrogen are highly polar

  • Highly polar bonds result in strong attractive forces between molecules (the slightly negative N, O or F forms hydrogen bonds with slightly positive hydrogen atoms from  other molecules)

  • Substances that can form hydrogen bonds have higher melting points than those unable to

AS-Level Energy and Enthalpy

Enthalpy Changes

  • During chemical reactions, energy is exchanged with the surroundings.

  • In exothermic reactions, products are lower energy than the reactants and the temperature of the surroundings increases as energy is released overall.

  • In endothermic reactions, products are higher energy than the reactants and the temperature of the surroundings decreases as energy is absorbed overall.

  • The actual amount of energy in a compound or atom is not measured in chemistry but the change in energy during a reaction can be measured and is called enthalpy.

  • Standard enthalpy changes refer to the thermal energy change that occurs during a reaction, where 1 mole of product is formed in standard conditions (100 KPa and 298K).


Bond Enthalpies

  • The energy levels of two individual atoms lower when they bond together.

  • A certain amount of energy is released by two atoms when a bond is formed, and that same amount of energy is absorbed by the two atoms when the same bond is broken.

  • The amounts of energy released or absorbed for a particular bond is called its bond enthalpy.

  • Mean bond enthalpies are used to account for the different environments a bond may be in.

  • The overall enthalpy change that occurs in a reaction can be found by subtracting the sum of all bond enthalpies formed from the sum of all bond enthalpies broken.


  • The energy change that occurs during a reaction can be found by measuring how the temperature of the surroundings changes during a reaction, this is called calorimetry.

  • Specific heat capacity refers to the energy required to raise the temperature of 1kg of a substance by 1°C.

  • The temperature change that a reaction forces on a substance can be used with the specific heat capacity to find the energy that has been exchanged.

  • Change in energy = mass x specific heat capacity x change in temperature

  • Q = mc∆T

  • Energy is often lost to the surroundings from the substance being used to find the temperature change (for example water).

  • To increase the accuracy of calorimetry, heat loss must be minimised (extra insulation can help).

Hess' Law

  • There is often more than one way to turn a substance into a product.

  • The energy change that occurs when a substance is changed into a product is the same, regardless of the route taken. This is called Hess’ Law. 

  • Hess’ Cycles are ways of drawing out the energy changes that occur when a substance is turned into a product via more than one route.

AS-Level Kinetics

Collision Theory

  • Kinetics is the study of how fast reactions occur.

  • How quickly a reaction happens is called its reaction rate.

  • Reactions occur when reactant particles collide together with enough energy (activation energy), creating a successful collision.

  • Increasing the frequency of successful collisions increases the rate of reaction.

  • Increasing the temperature, concentration and pressure of reacting particles increases the rate of a reaction.


Boltzmann Distribution Curves

  • Boltzmann distribution curves show the number of particles with a given amount of energy in a system.

  • Effects of changing concentration, temperature and the use of a catalyst on the energies of particles can be shown.


  • In a reversible reaction, reactants react together to produce products (forward reaction) and those products also react to produce the original reactants (reverse reaction).

  • At equilibrium, the rates of both the forward and reverse reactions are the same – the concentrations of all reactants and products do not change.

  • The position of equilibrium describes how much one reaction is favoured compared to the other.

  • Changing the conditions of an equilibrium system causes the position of equilibrium to move to oppose this change (Le Chatelier’s Principle). 

  • Increasing the temperature of a system forces the position of equilibrium to move to favor the endothermic reaction.

  • Increasing the concentration of a reactant forces the position of equilibrium to move to favor the reaction that uses up the reactant.

  • Increasing the pressure of a system forces the position of equilibrium to move to favor the reaction that produces the fewest moles of gas.