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Atomic structure
 The nucleus contains protons (positively charged) and
neutrons (neutrally charged i.e. no charge).
 The atomic number (proton number) is equal to the
number of protons in the atom’s nucleus.
 The mass number is the total number of protons and
neutrons in the nucleus.
 Ions do not have the same number of electrons as
protons, and so have an overall charge.
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Isotopes and relative masses
 Isotopes are atoms having the same number of
protons but different numbers of neutrons.
 The relative atomic mass is the weighted mean mass
of an atom relative to 12C, so that carbon is exactly 12
on this scale.
 The average relative atomic mass is equal to the sum
of each isotope’s mass for an element x its relative
abundance.
 The relative formula mass of a compound is equal to
the sum of the individual relative atomic masses.
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The mole
 A mole is the S.I. unit for amount of substance and
has units of mol.
 One mole of a substance is simply the relative formula
mass for a compound, or relative atomic mass for an
element in grams.
 The empirical formula is the simplest whole-number
ratio of atoms of each element present in a
compound.
 The molecular formula is the actual number of atoms
of each element in a molecule.
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Calculations using the mole
 Mass calculations: calculate the amount of substance
and then use the chemical equation to deduce the
moles of required substance.
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Acids and bases
 An acid is a hydrogen ion (H+) or proton donor in
solution, whereas a base is a hydrogen ion or proton
acceptor in solution.
 Hydrochloric acid (HCl), sulfuric acid (H2SO4) and
nitric acid (HNO3) are common acids.
 Bases include metal oxides (e.g. MgO), metal
hydroxides (e.g. NaOH) and ammonia (NH3).
 Alkalis are soluble bases and form hydroxide ions,
OH-, in solution.
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Reactions of acids and bases
 Salts are formed when a hydrogen ion from the acid is
replaced by a metal ion, or an ammonium ion.
 Acids react with bases to form a salt and water only;
they react with metal carbonates to form a salt, water
and carbon dioxide gas.
 Metals react with acids to form a salt and hydrogen
gas.
 Salts may chemically combine with water as water of
crystallisation in hydrated salts. (Without water in
anhydrous salts.)
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Oxidation numbers
 An oxidation number indicates the formal charge of a
chemically combined particle in a compound.
 The oxidation number of metals usually equals the
group number (as a positive value) and minus (8 –
group number) for non-metals.
 An element has been oxidised if the oxidation number
increases, and reduced if the oxidation number
decreases.
 When they react, metals are normally oxidised (they
lose electrons), whereas non-metals gain electrons
and are reduced.
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Electronic structure
 Electrons occupy energy levels around the nucleus of
the atom, where each shell has a principal quantum
number.
 For principal quantum number, n = 1, the number of
electrons is 2; for n = 2, the number is 8; then 18; then
32 electrons for n = 4.
 Main energy levels are sub-divided into sub-shells and
these consist of orbitals called s, p and d-orbitals.
 Elements have an electronic configuration that can be
shown in s, p or d notation, for example, sodium is
1s2, 2s2, 2p6, 3s1.
 Gas calculations: 1 mol of any gas occupies 24 000
cm3 or 24 dm3 at room temperature.
 Solution calculations: the amount of substance
dissolved is equal to the concentration x the volume of
solution (in dm3).
 A dilute solution consists of a small amount of
dissolved solute. A concentrated solution consists of a
large amount of solute.
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Chemical bonding
 Ionic bonding takes place when positive ions and
negative ions are attracted in a giant ionic structure.
 Covalent bonding is the sharing of electron pair(s)
between nuclei of atoms.
 The covalent bond and ionic bond are both very
strong chemical bonds.
 A dative covalent bond is one formed in which both
electrons are donated from the same atom.
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Molecular shapes
 The shape of a molecule is determined by the
repulsion between bonded electrons and non-bonded
electrons (lone pairs).
 Lone electron pairs repel more than bonded pairs of
electrons and give rise to distorted shapes.
 By deducing the number of bonded electron pairs and
lone pairs of electrons, the shape of a molecule may
be predicted.
 BF3 is trigonal planar; CH4 and NH4+ are tetrahedral;
SF6 is octahedral; H2O is non-linear (V-shaped/bent);
CO2 is linear and ammonia, NH3, as pyramidal
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Intermolecular forces
 An intermolecular force exists between molecules and
may include hydrogen bonding, dipole-dipole or van
der Waals’ forces.
 Electronegativity is the ability of an atom in a covalent
bond to attract a bonded pair of electrons towards
itself.
 Hydrogen bonding arises in molecules in which a
hydrogen atom is bonded to either an N or O atom.
 Water molecules, and other substances consisting of
hydrogen bonding, have anomalous properties as a
result.
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Bonding and physical properties
 Metals consist of a close-packed arrangement of
positive ions, through which delocalised electrons
move.
 Metals are very good electrical conductors as a result
of having mobile electrons.
 Giant structures have high melting and boiling points
due to strong chemical bonds acting throughout the
structure.
 Giant ionic structures conduct electricity when molten,
and when dissolved in water due to mobile ions, not
electrons.
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Periodicity
 When the elements are arranged in order of their
atomic number, there is a regular repetition of
physical and chemical properties.
 Elements in the same group have similar chemical
and physical properties.
 In the Periodic Table, ionisation energies increase
moving across a period from left to right, and
decrease moving down a group.
 Electron structures, atomic radii, melting points and
boiling points all show periodicity.
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Group 2 elements – the alkaline earth metals
 These elements all react with water to form a solution
of the hydroxide and hydrogen gas. These elements
react with oxygen to form the oxide.
 Reactivity increases on descending the group
because the outer two electrons are further from the
nucleus and are less shielded.
 Metal hydroxides are weak alkalis and typically have a
pH between 8 and 11.
 Group 2 carbonates decompose with greater difficulty
as the group is descended, to form the metal oxide
and carbon dioxide gas.
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Group 7 elements – the halogens
 All halogens exist as diatomic molecules in which van
der Waals’ intermolecular forces act between the
molecules.
 Halogens dissolve in organic solvents, like hexane, to
form characteristic colours, for example, iodine forms
a purple solution.
 Halogen atoms gain one electron to form halide ions,
X-, and this ability becomes easier on moving up the
group.
 Halogen atoms become larger on descending the
group, so a gained electron is only weakly attracted
due to greater shielding.
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Reactions of chlorine and halide ions
 Chlorine disproportionates in water to form
hydrochloric acid and chloric(I) acid, the latter being
an oxidising agent.
 Chlorine also disproportionates in cold, aqueous
sodium hydroxide to form sodium chloride, sodium
chlorate(I) and water.
 Chlorine is used to kill germs in water supplies, but is
also toxic to humans at higher doses.
 Halide ions are detected with silver(I) nitrate solution
and the subsequent reaction with ammonia solution.
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Representing organic compounds
 The empirical formula is the simplest whole-number
ratio of atoms of each element present in a
compound.
 The molecular formula is the actual number of atoms
of each element in a molecule.
 The general formula is an algebraic formula for a
member of a homologous series, for example, CnH2n+2
for an alkane.
 Structural, displayed and skeletal formulae are all
used to show organic molecules in different ways.
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Isomerism
 Structural isomers are compounds with the same
molecular formula, but different structural formulae.
 Stereoisomers are compounds with the same
structural formula, but with a different arrangement in
space.
 E/Z isomerism is an example of stereoisomerism in
which lack of free rotation about the C=C bond allows
two different forms to exist.
 Cis-trans isomerism occurs when two hydrogen
atoms, on different carbon atoms, adopt either the
same or opposite sides of the C=C.
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Hydrocarbons from crude oil
 A hydrocarbon is a compound containing hydrogen
and carbon atoms only.
 Crude oil is separated by fractional distillation, in
which the fractions have different boiling points.
 The carbon atom in hydrocarbons, as in alkanes, is
forming four single covalent bonds in a tetrahedral
arrangement.
 As the chain length of the hydrocarbons increases, so
too does the size of the van der Waals’ force,
resulting in a high boiling point.
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Reaction of alkanes
 Alkanes are hydrocarbons, and their molecules are
weakly attracted by van der Waals’ forces.
 Alkanes are saturated hydrocarbons; this means that
they consist of only carbon single bonds.
 Alkanes are mainly used as fuels, when combusted
they provide heat energy.
 Alkanes react with halogens in ultraviolet light to form
halogenoalkanes. This is a free radical substitution
mechanism.
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Alkenes
 Alkenes are unsaturated hydrocarbons; they consist
of one or more carbon double bonds, making them
more reactive than alkanes.
 A double bond consists of a and a bond. The
bond is formed by the adjacent overlap of porbitals.
 The bond is the reactive part of the double bond;
the bond is of lower reactivity.
 The arrangement of the bond about the carbon double
bond is trigonal planar, with an internal angle of 120°.
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Chemical reactions of the alkenes
 Alkenes react with hydrogen (to form alkanes); with
halogens (to form dihalogenoalkanes); and steam (to
form alcohols).
 Alkenes also react with hydrogen halides (to form
halogenoalkanes) and can polymerise to form addition
polymers, e.g. polypropene.
 The reaction between a halogen and an alkene
involves an electrophilic addition mechanism.
 An electrophile is a lone pair electron acceptor.
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Properties and preparation of ethanol
 Alcohols consist of one or more –OH groups in an
organic molecule.
 Alcohols are all water soluble and have relatively high
boiling points (are not very volatile) as a result of
hydrogen bonding.
 Ethanol can be made industrially by the reaction of
ethane with steam in the presence of a phosphoric(V)
acid catalyst.
 Ethanol for the drinks industry is produced by the
fermentation of sugars, such as glucose.
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Reactions of alcohols
 Alcohols may be classed as either primary, secondary
or tertiary.
 Primary alcohols may be oxidised to form aldehydes
then carboxylic acids, and secondary alcohols to form
ketones.
 Alcohols may react with carboxylic acids, in the
presence of a concentrated sulfuric(VI) acid catalyst,
to form esters.
 Alcohols may eliminate water, in the presence of an
acid catalyst, to form alkenes.
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Halogenoalkanes
 Halogenoalkanes are formed when a hydrocarbon
has one or more hydrogen atoms substituted for
halogen atoms.
 Halogenoalkanes may be hydrolysed to form alcohols,
using hot sodium hydroxide solution.
 The hydroxide ion acts as a nucleophile when
reacting with a halogenoalkane. A nucleophile is a
lone pair electron donor.
 The C–I bond is weaker than the C–Br and C–Cl
bond, and so it breaks faster when attacked by
nucleophiles.
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Uses of halogenoalkanes
 Chloroethene and tetrafluoroethene are used to make
the polymers PVC and PTFE (Teflon) respectively.
 CFCs (chlorofluorocarbons) were manufactured for
use in aerosols and refrigerants because of their low
reactivity and high volatility.
 CFCs have caused the ozone layer to break down,
leading to an increased intensity in damaging
ultraviolet radiation on Earth.
 Biodegradable CFCs and HCFCs are now used as
alternatives to CFCs.
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Modern analytical techniques
 Infrared spectroscopy is used to detect the presence
of certain covalent bonds in an organic molecule by
bond vibration.
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Enthalpy changes
 Exothermic reactions are those that produce heat
energy, while endothermic reactions absorb heat
energy.
 Standard enthalpy of formation is the heat change
involved when one mole of a compound is formed
from its elements under standard conditions.
 Standard enthalpy of combustion is the heat released
when one mole of a compound is combusted in
excess oxygen under standard conditions.
 Average bond enthalpy is the heat energy required to
break one mole of the specified bonds in the gas
phase.
 An absorption of approximately 3000 cm-1 appears in
most organic molecules, as this relates to the C-H
bond.
 Mass spectrometry can be used to detect the
presence of isotopes and their abundances.
 When an organic molecule is placed into a mass
spectrometer, fragments may appear, such as C2H5+
at m/z = 29.
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Rates of reaction
 The collision theory enables us to explain the effect of
changing concentration, surface area, temperature etc
on reaction rates.
 For a reaction, collisions must take place between
particles with sufficient energy before a successful
collision may result.
 A catalyst works by increasing the rate of a chemical
reaction, at the end of which the catalyst is chemically
unchanged.
 A catalytic converter allows pollutant gases like CO
and NO to react on the surface of Pd, Pt and Rh to
form less harmful products.
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Temperature and its effect on reaction rate
 Particles at a fixed temperature: a range of energies –
some have low and some high, but most have a
middling energy.
 Reactions have a certain minimum energy
requirement for a reaction to take place; this is called
the activation energy.
 As temperature increases, the distribution moves to
the right, giving a greater proportion of molecules
exceeding the activation energy.
 As more particles have an energy exceeding the
activation energy, there will be a greater chance of a
successful collision.
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Dynamic chemical equilibrium
 In a reaction at equilibrium, the forward and reverse
reaction rates are equal.
 Le Chatelier’s principle: an equilibrium will oppose a
change in external conditions by shifting to reduce the
effect of the change.
 In the Haber process, a compromise temperature is
used between the rate and yield for ammonia.
 A greater yield of ammonia is formed at high pressure
due to fewer gas molecules existing on the product
side of the equilibrium.
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The chemistry of the air
 The greenhouse effect results from certain gases in
the atmosphere, such as carbon dioxide, absorbing
infrared radiation.
 Bonds like C=O, O–H and C–H bonds in carbon
dioxide, water and methane absorb infrared radiation.
 Ozone in the stratosphere absorbs ultraviolet radiation
from the Sun and oxygen is formed. This is a
reversible process.
 Air pollution can be controlled by reducing hazardous
chemicals from industry and using high atom
economy processes.
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