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Atomic structure and isotopes
Isotopes are atoms of the same element with different numbers of neutrons and different masses
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The
outer electrons are involved in chemical reactions
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Physical properties depend on the mass of an atom
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Relative mass
Carbon-12 is used as the standard measurement of relative masses
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Relative atomic mass is the weighted mean mass of an atom compared with 1/12th of the mass of carbon-12
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1 mole of a substance is the amount of substance which contains as many particles as there are carbon atoms in 12g of
carbon-12 atoms
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02 x 10²³ mol¯¹
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Molar mass has the same numerical value as the relative
molecular mass/relative formula mass
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Determination of formulae
Anhydrous refers to a substance that contains no water molecules/no water of crystallisation
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Water of Crystallisation refers to water molecules that form an essential part of the crystalline structure of the compound
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Acids, bases, alkalis and neutralisation:
Bases are species that are proton acceptors that neutralise acids
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Acids are species that are proton donors
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Alkalis are soluble bases that release OH¯ ions when added to water
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A cation is a positively charged ion, usually a metal ion or ammonium ion in a salt
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Dissociation is the breaking up of a compound into simpler constituents e
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ions
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The same applies for a base with water:
B + H₂O
BH⁺ + OH¯
The strength of an acid, HA, or a base, B, is the extent of its dissociation into H⁺ and A¯ or BH⁺ + OH¯ ions or to what
extent it is ionised in solution
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g
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A strong acid is almost 100% ionised in solution
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The forward reaction is favoured, so nearly all the base dissociates in water and
lots of OH¯ ions are released
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It is written as a reversible reaction
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Energy levels, shells, sub-shells, atomic orbitals, electron configuration
Region
Maximum number
Each shells holds up to 2n²
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An orbital is a region around the nucleus that can hold
nd
2 shell
8
up to two electrons, with opposite spins
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Each atomic orbital can hold up to two
s-orbital
1
electrons
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s-subshell
2
p-subshell
6
A p-orbital has a 3D dumb-bell shape with its centre at the nucleus
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Giant ionic lattice structures contain oppositely charged ions which are strongly attracted to each other in all directions
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The greater the charge, the stronger the electrostatic forces between the ions and the greater the amount of energy
required to break up the ionic lattice during melting
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This breaks down the giant
ionic lattice as the water molecules surround and pull the ions away from the lattice and cause it to dissolve
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When solid, they
can’t conduct electricity because the ions are fixed in position by the strong ionic bonds and aren’t mobile
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The attraction overcomes the repulsion between the two positively charged nuclei resulting in a covalent bond
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Average bond enthalpy is the measurement of covalent bond strength
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The shapes of simple molecules and ions
Name of shape Diagram of shape
Bond angles
No
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5°

2 bond pairs and 2 lone pairs

H₂O

Trigonal planar

120°

3 bond pairs and no lone pairs

BF₃

Trigonal
pyramidal

107°

3 bond pairs and 1 lone pair

NH₃, SbCl₃

Tetrahedral

109
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A lone pair is slightly more electron-dense than a bonded
pair so a lone pair repels more than a bonded pair
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5°
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Electron repulsion theory predicts the molecular shapes
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Electronegativity and bond polarity
Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond
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Electronegativity increases towards the top right of the
periodic table, with fluorine having the most electronegative atoms
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Reactive metals form compounds with the least electronegative atoms
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If a molecule has an equal share of electrons, it is a perfect 100% covalent bond
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A non-polar molecule can have polar bonds with permanent dipoles but still be symmetrical
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A more electronegative element can form a non-polar molecule as it can be symmetrical so the dipoles cancel out
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g
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A polar covalent bond has a difference in electronegativity
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The bonding
electrons are closer to the electronegative atom
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Polar molecules are non-symmetrical because there is a charge difference across the whole molecule
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There are three common types of intermolecular forces: hydrogen bonds, permanent dipole-dipole forces and van der
Waals’ forces
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Permanent dipole of one molecule
attracts permanent dipole in a different polar molecule forming a weak permanent dipole-dipole force
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They are weak
intermolecular attractions between very small, temporary dipoles in neighbouring molecules
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The small induced dipoles attract one another
causing weak intermolecular forces
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An increase in electrons increases the strength of the van der Waals’ forces, the size of the induced dipoles and the
boiling points
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If there are no van der Waals’ forces, it
would be impossible to liquefy the noble gases
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Molecules containing -NH and -OH bonds are polar with strong permanent
dipoles
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In a hydrogen bond, the electron deficient Hδ+ on one molecule attracts a lone
pair of electrons on a Oδ- or Nδ- on a different molecule
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Ice is less dense than water because the molecules of ice are held apart by hydrogen bonds forming an open lattice
structure
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The stronger the intermolecular forces, the higher the melting and boiling points
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They have low melting and boiling points because the intermolecular forces are weak van der Waal’s forces which need a
relatively low amount of energy to break
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They are soluble in non-polar solvent because the van der Waals’ forces form between the simple molecular structure
and the non-polar solvent
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When simple molecules are broken, only the weak van der Waals’ forces break not the strong covalent bonds

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