3.4 Covalent bonding and coordinate (dative covalent) bonding

Define covalent bonding.

The electrostatic attraction between the nuclei of 2 atoms and a shared pair of electrons.

How many electrons are shared in a single, double, and triple covalent bond?

• Single/C–C bond → 2 electrons shared.
• Double/C=C bond → 4 electrons shared.
• Triple/C≡C bond → 6 electrons shared.

How can elements in period 3 expand their octet?

• SO₂: sulfur forms a double bond with each O → sulfur can have more than 8 electrons.
• PCl₅: P forms 5 single bonds → 10 electrons in outer shell.
• SF₆: S forms 6 single bonds → 12 electrons in outer shell.

Define coordinate bonding.

A covalent bond where both electrons come from the same atom.

Give an example of coordinate bonding in ammonium ion.

In NH₄⁺, nitrogen in NH₃ donates its lone pair to bond with H⁺ (which has no electrons, only a proton). This gives the ammonium ion (NH₄⁺) with 4 N–H bonds.

Explain coordinate bonding in Al₂Cl₆.

AlCl₃ is electron deficient (Al has only 6 electrons after bonding). A chlorine atom donates a lone pair to Al, forming a dative bond. Two AlCl₃ units join to form Al₂Cl₆ with 2 dative covalent bonds (Cl → Al).

Define a lone pair.

Pairs of electrons in the outer shell of an atom that are not bonded.

What is the difference between σ and π bonds?

• σ bond: first bond formed between 2 atoms, direct overlap of atomic orbitals, stronger than π bonds, found in all single bonds.
• π bond: second/third bond in double or triple bonds, formed from sideways overlap of p orbitals above and below σ bond plane, weaker than σ bonds, found in C=C and C≡C.

How many σ and π bonds are in single, double, and triple bonds?

• Single bond → 1 σ bond.
• Double bond → 1 σ + 1 π bond.
• Triple bond → 1 σ + 2 π bonds.

Define hybridisation.

Mixing of atomic orbitals to form new orbitals for bonding.

Describe sp³ hybridisation.

1s + 3p orbitals mix → 4 sp³ orbitals, tetrahedral shape (109.5°). Found in CH₄, C–C in ethane.

Describe sp² hybridisation.

1s + 2p orbitals mix → 3 sp² orbitals + 1 unhybridised p orbital. Shape = trigonal planar (120°). Found in C=C bonds (σ from sp², π from unhybridised p).

Describe sp hybridisation.

1s + 1p orbitals mix → 2 sp orbitals + 2 unhybridised p orbitals. Shape = linear (180°). Found in C≡C (alkynes), C≡N in HCN.

Define bond energy.

The energy required to break one mole of a particular covalent bond in the gaseous state.

Define bond length.

The internuclear distance of 2 covalently bonded atoms.

How do bond length and energy relate?

• Shorter bonds → nuclei closer → stronger attraction → higher bond energy → bonds harder to break → less reactive.
• Longer bonds → nuclei further → weaker attraction → lower bond energy → bonds easier to break → more reactiv