chapter 3 : chemical bonding

• Ionic bonding

  The strong electrostatic attraction between oppositely charged ions
• Drawing a dot-and-cross diagram for sodium chloride

  1. The sodium atom loses one electron to form an Na+ ion

  2. The chlorine atom gains the electron from sodium, becoming a Cl- ion

  3. Both ions have a full outer shell of electrons
• Drawing a dot-and-cross diagram for magnesium oxide

  1. The magnesium atom loses two electrons to form a Mg2+ ion

  2. The oxygen atom gains these electrons to become an O2- ion

  3. Both ions have a full outer shell of electrons
• Drawing a dot-and-cross diagram for calcium fluoride

  1. The calcium atom loses two electrons to form a Ca2+ ion

  2. Two fluorine atoms each gain an electron, forming two F- ions

  3. All the ions have a full outer shell of electrons
• Covalent bond


  • A bond formed by a shared pair of electrons between nuclei
  • Electrostatic attraction between the positive nuclei of the bonded atoms and the negative electrons between these nuclei holds the atoms together
• Drawing a dot-and-cross diagram of Cl2

  (Outer shell electrons only)
• Drawing a dot-and-cross diagram of CO2

  (Outer shell electrons only)
• Dative covalent bond

  A covalent bond whereby both electrons in the shared pair are donated by one of the bonding atoms only
• Drawing a dot-and-cross diagram of NH4+

  1. Each hydrogen atom is covalently bonded to the nitrogen

  2. There is one dative covalent bond
• Describing the bonding in Al2Cl6

  1. Al2Cl6 is made from two AlCl3 molecules

  2. The Al from one molecule forms a dative covalent bond with a Cl from the other molecule

  3. The same happens with the Al on other molecule

  4. Between the Al atom and the 3 Cl atoms in each AlCl3 molecule, there are covalent bonds
• Sigma (σ) bond

  • The strongest type of covalent bond
  • Formed from the head-on overlap of orbitals
• Pi (π) bond

  • Weaker than a σ bond
  • Formed from the sideways overlap of orbitals
  • A carbon-carbon π bond is formed from the sideways overlap of p-orbitals above and below the plane of the carbon atoms
• How hybridisation to form sp3 orbitals occurs

  1. A 2s orbital electron has been promoted to a 2p orbital

  2. Electrons then rearrange themselves via hybridisation into four identical orbitals called the sp3 orbitals

  3. These four orbitals can then take part in bonding with hydrogen to form methane, CH4
• How hybridisation to form sp orbitals occurs

  1. A 2s orbital electron has been promoted to a 2p orbital

  2. Hybridisation occurs for 2 out of the four orbitals (with a 2s and a 2p orbital)

  3. These two orbitals, sp hybrids, are now identical
• How hybridisation to form sp2 orbitals occurs

  1. A 2s orbital electron has been promoted to a 2p orbital

  2. Hybridisation occurs for three out of the four orbitals (2s and 2 x 2p)

  3. These three orbitals, sp2 hybrids, are now identical
• Electron repulsion theory

  • Electron pairs repel each other meaning they position themselves as far apart as possible
  • All bonding electron pairs repel each other equally
  • Lone pairs offer more repulsion than bonded pairs
• Molecule with 2 bonding pairs

  • Linear, 180°
• Molecule with 3 bonding pairs

  • Trigonal planar, 120°
• Molecule with 4 bonding pairs

  • Tetrahedral, 109.5°
• Molecule with 5 bonding pairs


  • Trigonal bipyramidal, 90° and 120°
• Molecule with 6 bonding pairs

  • Octahedral, 90°
• Molecule with 2 bonding pairs and 2 lone pairs


  • Non-linear/bent, 104.5°
• Molecule with 3 bonding pairs and 1 lone pair

  • Pyramidal, 107°
• Ammonium ion


  • Tetrahedral shape, bond angle of 109.5°
• Electronegativity


  The ability of an atom to attract the bonding pair of electrons in a covalent bond
• Hydrogen bonding

  • The strongest type of intermolecular force
  • A type of permanent dipole-dipole interaction
  • Occurs when a hydrogen atom is bonded to a very electronegative atom (nitrogen, oxygen or fluorine), which is close to another electronegative atom that has a lone pair of electrons
• Drawing a hydrogen bond between two water molecules

  Hydrogen bond
• Bond energy


  The measure of the strength of a chemical bond
• Bond length

  The average distance between the nuclei of two bonded atoms
• Bond polarity

  When two different atoms are joined by a covalent bond, the electronegativities of these atoms will be different meaning the electrons will be drawn towards the atom with the greatest electronegativity. This atom will have a slightly negative charge
• Bond polarity

  When two different atoms are joined by a covalent bond, the electronegativities of these atoms will be different meaning the electrons will be drawn towards the atom with the greatest electronegativity. This atom will have a slightly negative charge while the other will be slightly positive. This charge difference is a dipole. If a bond is polar, it has a dipole.
• Non-metal oxides undergo hydrolysis

  Oxygen is very electronegative. As a result, a permanent dipole forms across the covalent bond and the atom that oxygen is bonded to becomes partially positive. When the oxide is added to water, lone pairs on oxygen in the water are attracted to the partially positive atom in the oxide causing hydrolysis.
• Reactivity of a covalent bond and bond length

  As bond length increases, bond strength decreases. This is because there is less electrostatic attraction between the two nuclei and the shared pair of electrons between them. Reactivity increases.
• Reactivity of a covalent bond and bond strength

  The stronger the bond, the more difficult it is to break (requires more energy) and hence the less reactive the covalent bond.
• Reactivity of a covalent bond and bond polarity

  Generally: the greater the bond polarity, the more reactive the molecule.
• Permanent dipole

  A permanent difference in the partial charges of covalently bonded atoms. This occurs when there is a significant difference in electronegativities of the bonding atoms because the more electronegative atom has greater ability to attract the bonding pair of electrons meaning it has a slight negative charge. The other atom has a slight positive charge.
• Induced dipole-dipole interaction

  The random motion of electrons means that at any one point in time, there may be an uneven charge distribution. This causes an instantaneous dipole to be established between two atoms which can then induce dipoles in neighbouring atoms / molecules.
• Why is bromine liquid at room temperature

  Although bromine only has London forces between molecules, Br2 molecules contain lots of electrons meaning these temporary dipoles are quite strong.
• Why can group 18 elements become liquid despite the fact that they exist as single atoms

  The random movement of electrons within the atoms means that temporary dipoles within the atoms. Temporary dipoles may induce dipoles in neighbouring atoms. If the temperature is low enough, there will not be enough energy to overcome these weak London forces between the atoms meaning the gas will condense.
• Why does the boiling point of group 18 elements increase down the group

  The number of electrons and the atomic radius increases meaning that there are stronger temporary dipole and stronger London forces between the atoms. These stronger forces require more energy to overcome meaning a higher temperature is required to boil the liquid to turn it into a gas.