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Chemistry GCSE Definitions

Atomic structure

Proton – +1

Electron – weighs 1/1836 -1

Neutron – changes for isotopes (neutral)

Isotopes do not change the atomic number, but increase the mass number because of the increase in neutrons,

This will change the relative atomic mass depending on spread of isotopes.

Mass spectrometry, sample in gaseous state vaporised, and bombarded with electrons, forming positive ions which are accelerated and passed by a magnet which then splits apart different strengths, forming a graph showing abundance of different species.

Can find isotopes

Relative atomic mass can be calculated

First ionisation energy

The amount of energy required to remove one mole of electrons from each atom in the gas phase to form a singly positive ion.

Second ionisation energy

The amount of energy required to remove the next electron from an atom.

Requires more energy after the first, since this will usually lead to open electrons, or electrons left in unfilled shells.

Jump to between shells requires a lot more energy, and shows large peak on graph

Across periodic table generally increases but then drops when starting a new row because, outer shell contains one electron an so is not held well due to nuclear shielding

*Slight dips occur at between elements which have half filled shells, since the next element will have an extra e- on top of the half filled shell which is out of placed, and the element is more stable without it.

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S shell contains 2

P shell contains 6

D shell contains 10

*the d shell will not fill until the next s shell fills (3p 4s 3d 4p)



Electron affinity

The energy change per mole of gain in electrons to form an anion in the gas phase.

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First affinity generally negative (exothermic)

The second is always (positive) (endothermic)


Ionic equations

Ions separate

Insoluble and covalent as usual

Cross spectator ions out (on both sides)

Moles = mass/Mr(Ar)

Moles = Molarity x volume x 10-3


Calculate moles

Do ratio

Multiply/divide if solution taken from larger amount or diluted

Back-titration (unreacted method)

Calculate moles of substance used on reactants (usually alkali)

Multiply to get value of entire flask etc, e.g. if 25cm used from 250 then x 10

Work amount of moles of (acid) used altogether using Molarity x concentration

Take away amount made used altogether – amount used in reaction

This will give moles of (acid) in the equation to make up solutions, so take ratio if necessary

Calculate mass using moles x mr

Work out percentage of actual/original x 100

Bonding and structure

+ve ions are cations attracted to -ve (metals)

-ve ions are anions attracted to +ve (non-metal)

When dissolving or hydration, six water ligands by dative covalent bonding

Ionic bonding – low Ie energy

– Non-metal has high electron affinity (electron gain energy)

– Metal forms large ions with low charge

– Non-metal forms small ions of low charge

If cation is small/high charge then it is very polarising and has high charge density

If anion is big then it is polarisable

Lattice enthalpy is the measure of strength of an ionic substance, can be used for solubility.

Lattice enthalpy is the energy change per mole for the process

M+ + X- = MX

Ignores polarisation so actual covalence can be calculated

Covalence increases lattice energy

P/S orbs head on lap for sigma bonds

P orbs side lap for pi bond

A covalent bond is polar if there is a large difference in electronegativity

Electronegativity is the measure of how strongly an atom attracts electrons when in a covalent bond.

Two bonded pairs – linear 180

Two bonded two lone – bent linear 104

Three bonded pairs – trigonal planer 120

Three bonded, one lone pair – pyramidal 107

Three bonded, two lone pairs – t shaped two 90 and one 180

Four bonded pairs tetrahedral – 109.5

Five bonded pairs – trigonal bipyramid 3 are 120 and two 90

Six bonded pairs – octahedron 90

Two double bonds linear 180

One double, two single – trigonal planer 120

Two double bonds, two single – tetrahedral 109.5

Two double bonds, one lone pair – bent linear 104

Noble gases increase in temperature for mpt

*Hydrides also

Group 7

Fluorine pale yellow gas

Chlorine greenish gas

Bromine brown volatile liquid

Iodine dark grey lustrous solid

The hydrogen halides are very soluble

Produce strong acids

HF has strongest bond, and decreases down the group

Have high ionisation energies

Produce ppts with Ag+ ions, of which chloride is soluble with NH3, bromide with conc. NH3 and iodide insoluble.

React with conc. sulphuric acid since halogens are reducing agents

– Chlorides will produce –hydrogen sulphates and HCL the chlorine cannot be reduced further

– bromides are stronger reducing agents since it is bigger and so loses its e easier so will produce HBR and –hydrogen sulphates, and the HBR further produces bromine and water and sulphur dioxide

– Iodide is even bigger and so further reduces conc. H2SO4, this produce the above products and then the HI further produces, I2 and H2S and water, I2 and S and water, I2 and SO2 water. I3- can also be produced in the excess of I2

All halogens are strong oxidising agents, and decreases down the group, so Cl can oxidise Br and so on.

oxo-anions – stuff with oxygen like- carbonates or hydroxides

Unit 2 – Organic and energetic


?H is the heat change within a chemical reaction under standard conditions (atmospheric pressure around 298K)

Standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burned in oxygen, under standard conditions

Standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions

Standard enthalpy of neutralisation is

Exothermic graph starts high and goes down visa versa for endothermic

*Graphite more thermodynamically stable than diamond so diamond not in standard state

Specific heat capacity – is the heat required to increase the temperature of 1g of substance by 1K

Heat transfer = mass x specific heat capacity x temperature change

?H = kjmol-1 (after dividing by moles, kj before dividing)

Experiment – solutions make one excess but accurate, and then divide by accurate moles for ?H

First law of thermodynamics – energy cannot be created or destroyed but only changed from on form to another

Hess’s Law – The enthalpy change for a reaction is independent of the route by which the reaction is achieved, provided that the pressure and physical states of the reactants and products are the same in each case

Enthalpy of dissociation energy – enthalpy change when mole of a gaseous substance is broken up into free gaseous atoms

Can also be called and enthalpy for covalent bonds


Carbon can catenate, form bonds with itself

homologous series – similar chemical properties, gradual variation, formula

Nucleophiles – species which seek out positive centres and must have a lone pair if electrons to donate to form a covalent bond

Electrophiles – species which seek out negative centres and must be capable of accepting a lone pair of electrons to form a covalent bond

Aliphatic – normal

aromatic – contain rings do not obey rules

Free radical (can form from breaking bonds like Cl2) homolytic fission

Calculating energy per unit – moles gives g; divide kj by moles to get kgperg then times by density for calorific value (kjcm-3)

A tertiary iodide will be the most reactive because the C-I bond is the weakest and is on the third carbon.

Alcohols CnH2n+1OH

Aldehydes and keytones – CnH2nO (keytone cant have H as second R group)




Kinetic stability means that the reactants are thermodynamically unstable but do not have enough energy to react to react and so are kinetically stable.

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Chemistry GCSE Definitions. (2020, Jun 01). Retrieved from

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