Hybridisation (HL) (DP IB Chemistry): Revision Note

Hybridisation

What is hybridisation?

• Hybridisation is the process by which atomic orbitals mix to form new hybrid orbitals used in covalent bonding

• It explains why atoms like carbon form bonds that are identical in strength and direction

• Carbon’s ground state electronic configuration is 1s² 2s² 2p²

  • An orbital spin diagram shows this as:

• This suggests that carbon has only two unpaired electrons available for bonding

• However, carbon typically forms four covalent bonds (e.g. in methane)

  • These bonds are all equivalent

• To account for this, one 2s electron is promoted to a 2p orbital, giving four unpaired electrons

• These orbitals form four identical sp³ hybrid orbitals

  • The new orbitals are sp³ because one s orbital and 3 p orbitals merge

• In Period 3 elements, the 3s and 3p orbitals can also hybridise to form sp³ orbitals

  • This allows atoms like phosphorus and sulfur to form tetrahedral or expanded shapes

The shape of s and p orbitals

• Quantum mechanics shows that:

  • A 1s orbital is spherical

  • A p orbital is dumbbell or figure-of-eight shaped

• There are three p orbitals all at right angles to each other, known as p_x, p_y and p_z

What is sp³ hybridisation?

• One s orbital and three p orbitals from the same shell mix to form four sp³ hybrid orbitals

• These hybrid orbitals have ¼ s character and ¾ p character

  • These orbitals are asymmetric, with a larger lobe similar in shape to a p orbital

• The four sp³ orbitals arrange themselves with tetrahedral geometry

• This hybridisation explains the bonding and shape in molecules like methane and ammonia

• Methane, CH₄:

  • The carbon atom forms four single covalent bonds

  • Each carbon sp³ hybrid orbital overlaps head-on with a hydrogen 1s orbital

  • This results in:

    • Four identical sigma bonds

    • Tetrahedral electron domain geometry

    • Tetrahedral molecular geometry

    • A 109.5° bond angle

• Hybrid orbitals can accommodate both bonding pairs and lone pairs of electrons

• Ammonia, NH₃:

  • The nitrogen atom forms three single covalent bonds

  • Each nitrogen has three bonding pairs and one lone pair in sp³ hybrid orbitals

  • This results in:

    • Three identical sigma bonds and one lone pair

    • Tetrahedral electron domain geometry

    • Trigonal pyramidal molecular geometry

    • A 107° bond angle

What is sp² hybridisation?

• One s orbital and two p orbitals from the same shell mix to form three sp² hybrid orbitals

• These hybrid orbitals have ⅓ s character and ⅔ p character

  • These orbitals are asymmetric, with a larger lobe similar in shape to a p orbital

• The three sp² orbitals arrange themselves with trigonal planar geometry

• This explains the bonding and geometry seen when carbon forms a double bond, such as in alkenes

• Ethene:

  • Each carbon atom forms three sigma bonds and one pi bond

  • The carbon atoms are sp² hybridised

  • Each carbon uses three sp² orbitals to form σ bonds:

    • Two with hydrogen atoms

    • One with the other carbon

  • One unhybridised p orbital to form a π bond with the other carbon

  • This results in:

    • One C=C double bond containing 1 σ and 1 π bond

    • Trigonal planar electron domain geometry

    • Trigonal planar molecular geometry

    • A 120° bond angle around each carbon

• This bonding arrangement also occurs in carbonyl groups, where both carbon and oxygen use sp² hybrid orbitals to form the double bond

What is sp hybridisation?

• One s orbital and one p orbital from the same shell mix to form two sp hybrid orbitals

• These hybrid orbitals have ½ s character and ½ p character

  • These orbitals are asymmetric, with a larger lobe similar in shape to a p orbital

• The two sp orbitals arrange themselves with linear geometry

• This explains the bonding and geometry seen when carbon forms a triple bond, such as in alkynes

• Ethyne:

  • Each carbon atom forms two sigma bonds and two pi bonds

  • The carbon atoms are sp hybridised

  • Each carbon uses two sp orbitals to form σ bonds:

    • One with hydrogen

    • One with the other carbon

  • Two unhybridised p orbitals form two π bonds with the other carbon

  • This results in:

    • One CC triple bond containing 1 σ and 2 perpendicular π bonds

    • Linear electron domain geometry

    • Linear molecular geometry

    • A 180° bond angle around each carbon