Ionic and covalent bonding: SAQ 2

Study Notes

Chemical bonding occurs when 2 or more elements combine in a chemical reaction, forming a compound.
A compound is a substance made up of two or more different elements chemically combined.
Examples of chemical formulas: H₂O, SO₂, HCl

Hydrated substances contain molecules of water in definite proportions.
Water bound in this way is called water of crystallisation.
Heating removes this water, and compounds without this water attached are called anhydrous.

Noble gases are extremely stable due to their electronic configuration, which includes 8 electrons in their outer energy level.
The Octet Rule states that when bonding occurs, atoms tend to reach an electronic arrangement with 8 electrons in their outermost energy level.
Exceptions to the Octet Rule:
- Elements close to He tend to lose electrons to have 2 electrons in their outermost energy level.
- Transition elements (d-block metals) do not obey the Octet Rule.
- The rule sometimes works for sulphur and phosphorous, but not always.

The valency of an element is the number of bonds an atom of the element forms when it reacts.
Valency can be found by identifying the number of electrons an atom must lose, gain, or share to get a stable electronic configuration (i.e. Nobel Gas configuration).
Examples of valency:
- Sodium (Na): 1 bond
- Oxygen (O): 2 bonds
- Carbon (C): 4 bonds

Ionic bonds are created by the transfer of electrons.
Ionic bonds are formed between the electrostatic attraction (unlike charges attract) between the cation and anion.
Examples of ionic bonds:
- Lithium fluoride (LiF)
- Calcium and flourine
Covalent bonds involve the sharing of electrons.
Covalent bonds are formed when two non-metal atoms share a pair of electrons.
Examples of covalent bonds:
- Hydrogen (H₂)
- Hydrogen chloride (HCl)

High melting and boiling points.
Solids at room temperature.
Dissolve in water.
Conduct electricity.
Examples of ionic compounds:
- Sodium chloride (NaCl)
- Calcium carbonate (CaCO₃)

Covalent bonds can be formed between:
- Atoms of the same element (e.g. H₂)
- Atoms of different elements (e.g. HCl)
Electrons in a covalent bond:
- Bond pairs (shared electrons)
- Lone pairs (unshared electrons)
Examples of covalent bonds:
- Water (H₂O)
- Ammonia (NH₃)

Sigma (σ) bonds:
- Formed by the end-on overlap of orbitals.
- Single pair of electrons shared.
Pi (π) bonds:
- Formed by the sideways overlap of p-orbitals.
- Multiple pairs of electrons shared.
Examples of sigma and pi bonds:
- Oxygen (O₂): 1 σ bond, 1 π bond
- Nitrogen (N₂): 1 σ bond, 2 π bonds

Non-polar bonds:
- Electrons shared equally between atoms.
- Example: Hydrogen (H₂)
Polar bonds:
- Electrons shared unequally between atoms.
- Example: Hydrogen chloride (HCl)
Electronegativity:
- The ability of an atom in a covalent bond to attract the shared electrons.
- Examples of electronegativity values:
  - Hydrogen (H): 2.20
  - Chlorine (Cl): 3.16

Electronegativity values can be used to predict what types of bonds will be formed between atoms.
The greater the difference in electronegativity values, the greater the degree of polarity.
Examples of predicting bonds:
- CS₂: non-polar covalent bond
- HBr: polar covalent bond
- NaCl: ionic bond