Chemical bonding structure and properties

Key definitions:

Ion: charged species that are formed by the loss or gain of electrons

Compound: Species made up of more than 1 type of atom

Ionic compound: Compound formed by the electrostatic attraction of oppositely charged ions

Brittle: Something hard that will snap when force is applied

Intermolecular forces: Forced acting between molecules

Delocalised: not in a fixed position

• Ionic Bonding:

  • Formed by electron transfer from a metal to a non-metal.

  • Creates oppositely charged ions.

  • Ions are charged atoms or groups of atoms formed by electron loss or gain.

  • Metals lose electrons (MALE: Metals Are Losers of Electrons) → form positive ions (cations).

  • Non-metals gain electrons → form negative ions (anions).

  • Bonding involves strong electrostatic attraction between the oppositely charged ions.

  Formulae of Ions:

  • Atoms lose/gain electrons to get a full outer shell.

  • The group number shows the number of electrons in the outer shell.

  • The ion’s charge usually matches the group number:

Dot and cross diagrams:

Dot and cross diagrams are used to show the transfer of electrons between metal and non-metal atoms to form positive metal and negative non-metal ions.

Formation of sodium chloride NaCl:

Sodium loses one electron, from the Na+ ion. This electron is transferred to a chlorine atom forming a chloride ion.

Formation of potassium oxide K2O:

Each potassium atom loses an electron, forming 2 K+ ions. These electrons are transferred to one oxygen atom forming an oxide ion.

• Structure and Properties of Ionic Compounds:

  • Ionic compounds form a giant ionic lattice of alternating positive and negative ions.

  • Held together by strong electrostatic forces between oppositely charged ions.

  • Ionic formulae show the simplest whole-number ratio of ions (empirical formula).

  Properties:

  • High Melting and Boiling Points:

    • Strong ionic bonds in the lattice need a lot of energy to break, so melting/boiling points are high.

  • Electrical Conductivity:

    • Solid: Do not conduct – ions are fixed in place.

    • Molten or in solution: Do conduct – lattice breaks down, ions are free to move and carry charge (electrolytes).

  • Magnesium Oxide vs. Sodium Chloride:

    • MgO has higher charged ions (Mg²⁺, O²⁻) than NaCl (Na⁺, Cl⁻), so stronger bonds.

    • More energy needed to break these bonds → higher melting point for MgO.

  • Brittleness:

    • When the lattice is disturbed, like-charged ions can align.

    • This causes repulsion, making the crystal split easily.

Electrolysis of lead (II) Bromide:

Safety: This demonstration will be done in a fume cupboard as poisonous vapours are produced.

Method:

• The circuit was set up as shown in the diagram below. Notice whether the bulb is lit.

• The lead (II) Bromide was heated until it was molten. Notice whether the bold is lit

• The power was turned off, the burner removed and the electrodes lifted out of the molten PbBr2.

Conclusion:

• In solid PbBr₂, Pb²⁺ and Br⁻ ions are fixed in a lattice and cannot move, so no conduction.

• When molten, ions become mobile and can carry charge.

• Passing an electric current through molten PbBr₂ completes the circuit – bulb lights up

• Pb2+ ions move to the cathode, pick up 2 electrons and get deposited as neutral lead atoms (a silvery grey liquid)

  

• Br ions move to the anode, lose an electron and get deposited as neutral bromine atoms. 2 bromine atoms pair up, to form bromine molecules (brown gas)

  

• Molten PbBr2 undergoes decomposition by an electric current into its elements, Pb and Br 2

  

Covalent Bonding:

• Involves sharing a pair of electrons between two non-metal atoms.

• Only outer shell electrons are involved.

• A covalent bond is the electrostatic attraction between the shared electrons and both positive nuclei.

• One shared pair = single bond.

• Two pairs = double bond (e.g. O₂), three pairs = triple bond (e.g. N₂).

Covalent Bonding in Non-Metallic Elements:

• This type of bonding leads to the formation of simple molecules.

Examples:

Structure and Properties of Covalent Bonds:
There are two types of covalently bonded structures:

1) Simple Molecular Structures

• Examples: H₂, O₂, N₂, C₆₀ fullerene, HCl, H₂O, NH₃, CH₄

• Bonding: Strong covalent bonds within molecules

• Structure: Simple molecules with weak intermolecular forces between them

• Properties:

  • Low melting/boiling points – little energy needed to overcome weak forces

  • Do not conduct electricity – no free electrons or ions

Properties of Simple Molecular Substances:

• Low melting and boiling points:

  • Caused by very weak intermolecular forces between molecules.

  • These forces require little energy to overcome.

  • In chemistry, you overcome forces, not break covalent bonds.

• Solubility:

  • Usually insoluble in water, but soluble in non-polar solvents like petrol or trichloroethane.

  • Exceptions (e.g. HCl, CO₂, SO₂, sugar) may dissolve in water due to interactions with water molecules.

• Electrical conductivity:

  • Do not conduct electricity in any state.

  • No charged particles, and electrons are fixed in bonds – they cannot move to carry current.

Example questions:

Explain why substances with a simple molecular structure are gases liquids or low m.p/b.p solids:

• There are very weak intermolecular forces, between the molecules

• Which require little energy to overcome

Explain why the m.p/b.p of simple molecular structures increase, in general, with increasing relative molecular mass:

• The larger the relative molecular mass, the bigger and heavier the molecules

• Molecules have a greater area of contact, and stronger intermolecular forces

• More energy is needed to overcome these forces

2. Giant covalent structures e.g. diamond, graphite, sand

   Bonding: Strong covalent bonds between all atoms

   Structure: Giant covalent with only covalent bond between all atoms in the structure

Properties of giant covalent compounds:

• Extremely high melting and boiling points

  • Many strong covalent bonds need to be broken

  • Within the giant covalent structure

  • This requires a lot of energy

• Insoluble in all solvents

Examples of giant covalent structures:

Metallic bonds:

Metallic Bonding:

• Structure: Giant metallic lattice – positive metal ions in a regular pattern, surrounded by a sea of delocalised electrons.

• Bonding: Strong electrostatic attraction between metal ions and delocalised electrons.

• Each metal atom loses outer electrons, forming positive ions, while electrons become free to move.

Properties of Metals:

• High melting/boiling points:

  • Strong metal–electron attraction needs a lot of energy to break.

• Good conductors of electricity:

  • Delocalised electrons move freely and carry charge.

• Malleable and ductile:

  • Layers of ions can slide, so metals can be bent or drawn into wires.

• Good conductors of heat:

  • Mobile electrons transfer thermal energy quickly.

• Strong and sonorous:

  • Can withstand force and produce a ringing sound when struck.

Bonding summary:

Ionic bonding:

• The strong electrostatic attractions between the oppositely charged ions e.g. NaCl.

• The ionic bond forms due to the transfer of one or more electrons from a metal atom to a non-metal atom which results in oppositely charged ions

Covalent bonding:

• The strong electrostatic attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond

• Covalent substances are formed by the sharing pairs of electrons between atoms of the same and/or different non-metal elements

Metallic bonding:

• The strong electrostatic attractions between the positively charged metal ions and the delocalised electrons e.g. Cu