Quick Notes - AS Physical Chemistry

Atomic Structure

Structure of an Atom

  • Atoms are the smallest units of an element.
  • Atoms are made of sub-atomic particles:
    • protons (relative charge +1, relative mass 1)
    • neutrons (relative charge 0, relative mass 1)
    • electrons (relative 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; it contains just protons and neutrons, making the nucleus positively charged.
  • Electrons are attracted to the positive charge of a nucleus and exist in orbitals around the nucleus.
  • Atoms can lose or gain electrons to become charged ions.

Electron Orbitals

  • Electrons are negatively charged, so they are attracted to the positive charge of a nucleus (electrostatic attraction).
  • The closer an electron is to a nucleus, the lower its energy.
  • All electrons in an atom want to get as close to the nucleus as possible, but they repel each other.
    • This forces electrons to exist in set regions called orbitals.
    • One orbital can hold two electrons (one pair).
  • Different orbitals have different shapes and energies. Orbital types are labelled with letters – s, p, d and f.
  • Shells describe how far away an electron is from a nucleus, and therefore its energy.
    • Shell 1 is the closest shell to the nucleus, shell 2 the next closest, etc.
  • Due to different shapes, orbitals within a shell can have different energies and form sub-shells.
  • Shells and sub-shells are filled successively – lowest available energy sub-shells and orbitals are filled first.

Electron Configurations

  • Electrons occupy orbitals of lowest energy around a nucleus.
    • Higher energy orbitals are only occupied if all lower energy orbitals are already full.
  • Electron configurations show how electrons are arranged around a nucleus within an atom or ion.
    • The first number shows the shell the electrons are in.
    • The letter shows the shape of the orbital the electrons are in.
    • The second number shows the number of electrons in the orbital(s).
  • Due to electron repulsion and orbital shapes, 4s orbitals are occupied before the 3d orbitals.

Moles (Amounts)

The Mole

  • Atoms and particles are very small, so referring to actual number of atoms in a sample isn’t very practical.
  • The mole is a unit used to refer to an amount of atoms or particles.
  • Just like one dozen = 12 of something, one mole = 6.022 x 1023 of something.
  • The number one mole refers to is known as Avogadro’s constant.

Relative Atomic Mass

  • Relative atomic masses (Ar) are used to compare the masses of all elements on the periodic table.
  • Relative atomic mass (Ar) is the weighted average mass of one atom of each element – compared to 1/12th the mass of one atom of carbon-12.
    • one carbon-12 atom has six protons and six neutrons
    • relative mass of one proton is one
    • relative mass of one neutron is one
    • relative mass of one atom of carbon-12 is twelve
  • The mass of an atom is determined by the number of protons and neutrons it has.
    • Protons and neutrons have (virtually) the same mass.
  • The actual mass of protons and neutrons are too small to use when describing the mass of an atom, instead both particles are assigned a relative mass of one.
  • Relative atomic mass is a weighted average that accounts for the abundance of each naturally occurring isotope of that element.
  • Molar mass is the mass (g) of one mole’s worth of something (g mol-1).

Moles from Mass

  • The number of moles in a sample can be calculated using the formula: moles equation from mass and molar mass showing units
  • The value of relative atomic mass of an element = value of molar mass (in grams per mol).
  • The value of relative formula and molecular mass = value of molar mass (in grams per mol).
  • This means moles can be calculated using: moles equation from mass and relative formula mass

Moles - Solutions

  • The number of moles of a solute in a solution can be calculated using the formula: equation for calculating moles from concentration and volume
  • Volume (of solution) is usually given in cm3 and must be converted into dm3.
    • 1 dm3 is the same as 1000 cm3
  • To convert dm3 to cm3 x 1000
  • To convert cm3 to dm3 / 1000
  • Concentration refers to how many moles of solute are in 1dm3 of solution (units are mol dm-3).

Moles - Gases

  • The ideal gas equation is used to calculate the moles of gas present in a system if the volume, pressure and temperature are known.
    • pV = nRT
      • n = moles
      • R = gas constant = 8.31 Jmol-1K-1
      • V = volume (m3)
      • p = pressure (Pa)
      • T = temperature (K)
  • The equation assumes that all gases behave in the same way and are not interacting with each other, they are described as ‘ideal gases’.
  • At room temperature and pressure (293K and 101kPa), the ideal gas equation shows that 1 mole of any gas occupies a total volume of 24000cm3 (24dm3).
  • At standard temperature and pressure (273K and 101kPa), the ideal gas equation shows that 1 mole of any gas occupies a total volume of 22400cm3 (22.4dm3).



  • Bonds between atoms 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 Bonding

  • A covalent bond is the sharing of one pair of electrons between two atoms.
    • Two atomic orbitals from two atoms overlap and create a new area of electron density between the two atoms, called a molecular orbital.
    • The positively charged nuclei of both atoms get attracted to the shared pair of negatively charged electrons and this pulls them closer together.
    covalent bond electrons attraction
  • The distance between the centre of each nuclei to each other is called a ‘bond length’.
  • covalent bond length hydrogen atom distance between nucleus
  • Non-metals form covalent bonds with other non-metals.
  • As covalent bonds occur between atoms, they are atomic bonds.


  • Electronegativity is the ability or 'tendency' of an atom to attract a shared pair of electrons in a covalent bond towards itself.
  • If two atoms with different electronegativities are covalently bonded together, the bond will be polar.
  • electronegativity polar bond partial negative positive charge
    • The atom with the higher electronegativity will attract the bonded electrons towards itself and will end up with a partial negative charge (δ-).
    • The atom with the lower electronegativity will have the bonded electrons pulled away from it and will end up with a partial positive charge (δ+).
  • The greater the difference in electronegativity between two bonded atoms, the more polar the bond.
  • Electronegativity increases across a period in the periodic table and decreases down a group.
  • Electronegativities are relative and the Pauling Scale is often used to compare the electronegativities of different elements - Francium has the lowest electronegativity (0.7) and Fluourine the highest (4.0).

Ionic Bonding

  • Ionic bonding describes the electrostatic attraction that holds oppositely charged ions together.
  • The strength of an ionic bond is based on the charge and size of the ions involved:
    • Greater the charge = stronger the bond.
    • Smaller the ionic radius = stronger the bond
  • To get a stable electron configuration, atoms lose or gain electrons.
    • Metals lose electrons to get a full outer shell and become positively charged ions.
    • Non-metals (usually) gain electrons to get a full outer shell and become negatively charged ions.
  • Oppositely charged ions arrange themselves in a solid to maximise attraction of opposite charges and minimise repulsion to like charges, forming a repeating lattice structure.
  • 2D ionic lattice showing attraction and repulsion

Metallic Bonding

  • In metallic bonding, positive metal ions are attracted to negatively charged delocalised electrons, forming a dense and rigid structure.
  • metallic bonding structure showing positively charged metal ions and delocalised electrons

  • Positively charged metal ions form because outermost electrons are held weakly by the metal atoms nucleus and are able to drift out of their orbitals, becoming delocalised and leaving a positively charged metal ion behind.
    • Delocalised electrons from metal atoms have 'nowhere to go' and stay attracted to the positive charge of metal ions.
    • Electrons are not held by the metal atom but are free to move – they are delocalised.
    • Free movement of delocalised electrons enable metals to conduct electricity.

Intermolecular Forces

  • Intermolecular forces hold molecules together in molecular substances.
  • Temporary induced dipole-dipole forces (London dispersion forces) exist between all molecules.
    • Unequal electron distribution around a molecule creates an instantaneous dipole that can induce a dipole on a neighbouring molecule.
    • Two opposite dipoles from different molecules are attracted to each other and a weak force occurs.
    • They are the weakest type of intermolecular force.
  • The bigger the molecules, the greater the intermolecular forces that can arise between them, giving larger molecules higher melting and boiling points.
  • Partial charges are shown using δ+ (partially positive) and δ- (partially negative).

Permanent Dipole-Dipole

  • Permanent dipole-dipole interactions exist between polar molecules.
  • The partially negatively charged area of one molecule is attracted to a partially positively charged area 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 specific type of intermolecular force.
  • Molecules that can form hydrogen bonds have higher melting and boiling points than similiar sized molecules that are unalbe to.
  • Hydrogen bonds exist between molecules that have a N-H, O-H or F-H bond.
    • Nitrogen, oxygen and fluorine are the most electronegative elements, so the covalent bonds between them and hydrogen are highly polar.
  • Highly polar bonds result in strong attractive forces between molecules.
    • The partially negative nitrogen, oxygen or fluorine forms hydrogen bonds with the partially positive hydrogen atoms from other molecules.
  • Substances that can form hydrogen bonds have higher melting points than similar-sized substances that cannot.

Energy and Enthalpy

Enthalpy Changes

  • During chemical reactions energy is exchanged between reacting substances and their surroundings.
  • Exothermic reactions – products are lower in energy than the reactants and the temperature of the surroundings increases.
  • Endothermic reactions – products are higher in energy than the reactants and the temperature of the surroundings decreases.
  • Enthalpy is a measure of heat energy with the units kJmol-1.
    • The actual amount of energy in a compound or atom is not easily measured in chemistry, so it is the change in heat energy during a reaction (enthalpy change, ΔH) that is measured.
    • Exothermic reactions have negative enthalpy changes (-ΔH).
    • Endothermic reactions have positive enthalpy changes (+ΔH).
  • Standard enthalpy change refers to the thermal energy change that occurs during a reaction, where 1 mole of product is formed in standard conditions (101 kPa and 298K).

Bond Enthalpies

  • The overall enthalpy change that occurs in a reaction can be found by subtracting the sum of all the bond enthalpies of bonds formed from the sum of all bond enthalpies of bonds broken.
  • A bond enthalpy is the amount of energy released or absorbed when a particular bond is made or broken.
    • Energy is released by two atoms when a bond is formed between them.
    • That same amount of energy is absorbed by the two atoms when the bond is broken.
  • Mean bond enthalpies are used to account for the different environments a bond might be in.


  • Calorimetry is a technique used to find the heat energy change that occurs during a reaction.
    • Heat energy change is found by measuring how the temperature of the substance surrounding the reaction changes.
  • Specific heat capacity (c) refers to the energy required to raise the temperature of 1g (or 1kg) of a substance by 1oC.
  • The temperature change of a substance can be used with its specific heat capacity to find the energy it has absorbed or released.
    • Change in energy = mass x specific heat capacity x change in temperature
      • Q = mcΔT
  • Energy is often lost from the substance, whose temperature is being measured, to the surroundings.
    • To increase the accuracy of calorimetry, heat loss must be minimised (extra insulation can help).

Hess' Law

  • Hess’ Law states that the energy change that occurs when a substance is changed into a product is the same, regardless of the route taken.
    • There is often more than one way to go from a reactant to a product, meaning there is more than one ‘route’ to convert a substance to a particular product.
  • Hess cycles are ways of drawing out the energy changes that occur when a substance is turned into a different substance via more than one route.


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, Ea).
    • Collisions between particles that lead to a reaction are called ‘successful collisions’.
  • 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

  • Maxwell-Boltzmann distribution curves show how available energy is distributed among molecules in a gaseous system.
    • At a given temperature, only a certain amount of energy is available for all molecules of gas within a system.
    • Individual molecules have different energies and are constantly colliding, transferring energy with other molecules – either gaining or losing energy.
  • The 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 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 has been favoured compared to the other before equilibrium is reached.
  • Changing the conditions of an equilibrium system can cause the position of equilibrium to move.
  • Le Chatelier’s Principle states that the position of equilibrium will move to oppose a change in conditions.
    • Increasing the temperature of a system forces the position of equilibrium to move to favour the endothermic reaction.
    • Increasing the concentration of a reactant forces the position of equilibrium to move to favour the reaction that uses up that reactant.
    • Increasing the pressure of a system forces the position of equilibrium to move to favour the reaction that produces the fewest moles of gas.