## 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 ).

23​

-1

-3

3

3

3

3

3

3

3

3

3

-3

3

3

## AS-Level Bonding

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.

Calorimetry

• 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.

Equilibrium

• 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.