AQA Chemistry AS Unit 1

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3 Bonding

3.1 The nature of Ionic Bonding

Why Chemical Bonds Form

  • Bonds between atoms involve outer electrons
  • Noble gases have full highest energy levels so are very unreactive
  • When atoms bond they share or transfer electrons to become more stable
  • There are three types of chemical bonds: ionic, covalent and metallic

Ionic Bonding

  • The easiest way for metals to become stable is to lose outer electrons
  • The easiest way for non-metals is to gain electrons
  • So ionic bonding occurs between a metal and non-metal
  • Electrostatic forces: forces of attraction and repulsion between electrically charged particles
  • Ionic compounds exist in a lattice due to the electrostatic forces
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The Nature of Ionic Bonding Cont

Properties of Ionically Bonded Compounds

  • Solid at room temperature
  • They have giant structures so high melting points since energy must be supplied to break the lattice of ions
  • Conduct electricity when molten or aqueous since ions that carry current are free to move
  • Are brittle and shatter easily since lattice of positive and negative ions are moved so like charged ions are in contact so repulsion occurs
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3.2 Covalent Bonding

  • Forms between pair of non-metal atoms
  • A covalent bond is a shared pair of electrons
  • The atoms are held together by electrostatic attraction between nuclei and shared electrons
  • In double covalent bonds four electrons are shared

Properties of Substances with molecular structures

  • They have low melting points since there is only weak attraction between molecules so not much energy is needed to part them
  • Poor conductors of electricity as have neutral overall charge so no charged particles

Co-ordinate Bonding/ Dative Covalent Bonding

  • When one atom provides both electrons
  • The atom that accepts the electron pair is electron deficient
  • The atom donating  has a pair of electrons not used in a bond: Lone Pair
  • Co-ordinate bonds have the same strength and length as ordinary covalent bonds
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3.3 Electronegativity

  • It is the power of an atom to attract the electron density in a covalent bond towards itself
  • Electron density is used to describe the way the negative charge is distributed in a molecule
  • Electronegativity depends on: nuclear charge, distance between nucleus and highest energy electron and the shielding of the nuclear charge
  • The smaller the atom, the closer the nucleus to the highest energy level so greater electronegativity
  • The larger the nuclear charge, the greater the electronegativity

Trends in Electronegativity

  • Going up a group electronegativity increases as atom gets smaller and less shielding
  • Going across a period electronegativity increases as nuclear charge increases and the atom gets smaller
  • The most electronegative atoms are fluorine, oxygen and nitrogen then chlorine 
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Electronegativity Cont

Polarity of Covalent Bonds

  • Polarity is about unequal sharing of electrons between atoms covalently bonded

Covalent Bonds between two atoms of the same

  • The bond with be non-polar as the atoms have the same electronegativity

Covalent Bonds between two different atoms

  • Electrons won't be equally shared as different electronegativities
  • Delta Plus and Delta Minus represent a small charge of less than one electron's worth
  • Delta Plus is given to the less electronegative atom
  • The greater the difference in electronegativity the more polar
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3.4 Metalic Bonding

  • In a metal element the highest energy levels merge
  • Metals consist of a lattice of positive ions in a 'sea' of delocalised electrons
  • The number of delocalised electrons depends on how many electrons lost from each atom
  • Metals have giant structure as metallic bonding spreads through metal

Properties of Metals

  • Good conductors: delocalised electrons move throughout the structure, for heat the sea of electrons partly responsible and vigorous vibrations of closely packed ions
  • Strength depends on charge and size of ion:
    • The greater the charge on ion, the greater the number of delocalised electrons and the stronger the attraction between +ve ions and electrons
    • The smaller the ion the close the electrons to the nucleus and the stronger bond
  • Malleable and Ductile
  • High melting point: due to giant structures, there is strong attraction between metal ions and delocalised electrons so difficult to separate
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3.5 Forces between Molecules

  • There are three types of intermolecular forces:
  • van der Waals, Dipole-Dipoles and Hydrogen bonding

Dipole-Dipole Forces

  • Molecules with polar bonds may have dipole moment
  • With molecule that have more than one polar bond the bonds cancel out so have no dipole moment
  • Dipole-Dipole forces act between molecules that have permanent dipoles

Van der Waals

  • All molecules and atoms have weak electrostatic attractions called Van der Waals
  • Caused by temporary fluctuations in electron clouds that lead to temporary dipoles
  • This induces a temporary dipole in the second molecule
  • Attractions caused between partial charges
  • The size of van der Waals increases with the number of electrons so the larger the mass the stronger the van der Waals 
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3.6 Hydrogen Bonding

  • Hydrogen atom bonded to very electronegative atom e.g. N, O, F
  • The bond is very polar
  • Hydrogen bonds form between the H on the first molecule and a lone pair on the N, O or F atom

Boiling Points of the Hydrides

  • The hydrides are elements bonded to hydrogen
  • As you go down a group the hydride boiling point increases since:
  • Atom gets bigger down the group, the molecule is bigger with more electrons so larger temporary dipoles causing larger induced temporary dipoles and greater electrostatic attraction so stronger van der Waals

Structure and Density of Ice

  • In liquid state H2O hydrogen molecules break and reform easily.
  • When water freezes the molecules are no longer free to move so hydrogen bonds are fixed
  • In ice water molecules are slightly less compact so ice is less dense than water
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3.7 States of Matter

Energy Changes on Heating

  • Solid: they vibrate around a fixed point, this increases the distance between particles so the solid expands
  • Solid to liquid: energy has to be supplied to weaken the forces between particles, the energy needed is the enthalpy change of fusion
  • Liquid: the particles move quicker so have more kinetic energy causing the liquid to expand
  • Liquid to gas: energy needed to break intermolecular forces, the energy needed is the enthalpy change of vapourisation
  • Gas: particles gain kinetic energy so move faster making the gas expand
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States of Matter Cont


  • Have regular arrangement and are held together by forces of attraction
  • The strength of the forces of attraction affects physical properties
  • The stronger the force the higher the melting point

Ionic Crystals

  • Have strong electrostatic attractions
  • High melting points so need a lot of energy to break

Metallic Crystals

  • Metals exist in a lattice of positive ions embedded in delocalised electron sea
  • High melting points since strong metallic forces
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States of Matter Cont

Molecular Crystals

  • Held by van der Waals, dipole-dipoles or hydrogen bonds
  • Have weak  intermolecular forces so have low melting temperatures
  • Crystals are soft and break easily
  • Doesn't conduct electricity as no charged particles to charge charge

Macromolecular Crystals

Diamond and Graphite

  • Both made only of Carbon
  • Have different arrangement of carbon
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States of Matter Cont


  • Each carbon atom forms four single covalent bonds
  • The four electron pairs repel
  • Diamond forms a tetrahedron shape with 109.5 degree angle between bonds
  • It forms a 3D lattice shape so has these properties:
  • Hard material, high melting point and doesn't conduct electricity since no free charged particles


  • Has van der Waals and strong covalent bonds: Covalent bonds mean has a high melting point
  • Each carbon atom forms 3 single covalent bonds so form a trigonal planar with 120 degree bond angle
  • Each carbon has a spare electron in p orbital that are delocalised and move through the planar adding strength to the bonding, the electrons cause it to conduct
  • Weak van der Waals mean the layers can slide past each other causing graphite to be soft
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3.9 Shape of Molecules and Ions

Electron Pair Repulsion Theory

  • Each electron pair will repel other electron pairs
  • Electron pairs can be lone pairs or shared pairs

Two pairs of Electrons

  • The shape is linear and the bond angle is 180 degrees

Three pairs of Electrons

  • The shape is trigonal planar and the bond angle is 120 degrees

Four pairs of Electrons

  • The shape is tetrahedral and the bond angle is 109.5
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Shape of Molecules and Ions Cont

Five Pairs of Electrons

  • The shape is trigonal bipyramid and the bond angles are 90 and 120 degrees

Six Pairs of Electrons

  • The shape is octahedral and the bond angles are 90 degrees

Molecules with lone pairs of electrons

  • Lone pairs affect the shape of a molecule: you lose 2 degrees for every lone pair
  • Lone pairs repel more than bonding pairs as lone pairs are only attracted to one positive nucleus whereas bonding pairs stretch over two positive nuclei
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4 Periodicity

The Blocks of the Periodic Table

  • All elements with their highest energy electron in the s sublevel are in the s block
  • P sublevels are in the P block
  • D sublevels are in the D block


  • Elements in the same group form a 'chemical family' so have similar properties
  • They have same number of electrons in their highest energy level


  • In the S Block elements get more reactive down the group
  • To the right elements (non-metals) get more reactive up the group


  • There are trends in physical and chemical properties across a period
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4.2 Trends in properties of Period 3

  • Periodicity: regular recurrence of properties of elements arranged in periodic table order
  • Elements Na, Mg and Al are metals so have giant structures, they lose highest energy electrons to form ionic compounds
  • Si has 4 electrons in its highest energy level that it forms covalent bonds with, it is classed as a semi metal
  • P, S and Cl are non-metals that can form ionic or covalent bonds
  • Ar is a noble gas with a full highest energy level so is unreactive

Trends in Melting and Boiling Point

  • Giant structures have high melting and boiling points
  • Molecular/ atomic structures have low melting and boiling points
  • The metals have increasing mp and bp since strength of metallic bonding increases as the charge on the ion increases meaning more delocalised electrons hold the lattice together
  • The non-metals have molecular structures that depend on the sizes of the van der Waals between the molecules. The MP and BP goes S8>P4>Cl2 as have less bonds to overcome each time. Silicon has a giant structure of covalent bonds so has higher MP
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4.3 More trends in Period 3

Atomic Radii

  • Atomic Radius depends on type of bonds formed
  • It decreases across the period, atoms get larger down a group
  • Decreases as along the period are adding protons to the nucleus and electrons to highest energy level, the increased charge pulls the electrons closer to the nucleus and as there is no extra shielding the size of the atom decreases
  • Increases down a group as gain extra energy level so highest energy level further from nucleus

First Ionisation Energy: X(g)->X+(g) + e-(g)

  • Ionisation energy increases across the period and decreases down the group
  • Increase across period as extra proton in same highest energy level, increased charge means more difficult to remove electron
  • Decrease down group as gain filled energy level which increases shielding also electron is further from the nucleus so easier to remove
  • There is a drop from one period to another as start a new highest energy level, so increase in atomic radius and outer electron further away from nucleus 
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4.4 Ionisation Energies

Drop between Group 2 and 3

  • Al is lower than Mg as its' highest energy electron is in the P sublevel which takes less energy to remove than Mg's electron in the S sublevel

Drop between Group 5 and 6

  • S is lower than P as S has two electrons paired in the P sublevel orbital as they are already repelling each other it is easier to remove it from the S

Successive Ionisation Energies

  • Removing an electron from an atom one at a time gets harder each time
  • This is as the next electrons are being removed from an ion
  • In a graph the ionisation energies show when each energy level begins as a increase in energy is shown
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good, detailed notes, very useful :D

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