Hybridisation (Cambridge International A-Level Chemistry)

Orbitals & Hybridisation in Covalent Bonding

Bond overlap in covalent bonds

• A single covalent bond is formed when two nonmetals combine

• Each atom that combines has an atomic orbital containing a single unpaired electron

• When a covalent bond is formed, the atomic orbitals overlap to form a combined orbital containing two electrons

  • This new orbital is called the molecular orbital

• The greater the atomic orbital overlap, the stronger the bond

• Sigma (σ) bonds are formed by direct overlap of orbitals between the bonding atoms

• Pi (π) bonds are formed by the sideways overlap of adjacent above and below the σ bond

σ bonds

• Sigma (σ) bonds are formed from the end-on overlap of atomic orbitals
• S orbitals overlap this way as well as p orbitals

Forming sigma bonds

Sigma orbitals can be formed from the end-on overlap of s orbitals

• The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond
• The pair of electrons is found between the nuclei of the two atoms
• There is an electrostatic force of attraction between the electrons and nuclei which bonds the atoms to each other

π bonds

• Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals
• The two lobes that make up the π bond lie above and below the plane of the σ bond
• This maximises the overlap of the p orbitals
• A single π bond is drawn as two electron clouds one arising from each lobe of the p orbitals
• The two clouds of electrons in a π bond represent one bond containing two electrons

Forming pi bonds

π orbitals can be formed from the end-on overlap of p orbitals

Examples of sigma & pi bonds

• Hydrogen

  • The hydrogen atom has only one s orbital
  • The s orbitals of the two hydrogen atoms will overlap to form a σ bond

Sigma bonding in hydrogen

Direct overlap of the 1s orbitals of the hydrogen atoms results in the formation of a σ bond

• Ethene

  • Each carbon atom uses three of its four electrons to form σ bonds
  • Two σ bonds are formed with the hydrogen atoms
  • One σ bond is formed with the other carbon atom
  • The fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a π bond
  • This means that the C-C is a double bond: one σ and one π bond

Pi bonding in ethene

Overlap of the p orbitals results in the forming of a π bond in ethene

Sigma and pi bonding in ethene

Each carbon atom in ethene forms two sigma bonds with hydrogen atoms and one σ bond with another carbon atom. The fourth electron is used to form a π bond between the two carbon atoms

• Ethyne

  • This molecule contains a triple bond formed from two π bonds(at right angles to each other) and one σ bond
  • Each carbon atom uses two of its four electrons to form σ bonds
  • One σ bond is formed with the hydrogen atom
  • One σ bond is formed with the other carbon atom
  • Two electrons are used to form two π bonds with the other carbon atom

Sigma and pi bonding in ethyne

Ethyne has a triple bond formed from two π bonds and one σ bond between the two carbon atoms

• Hydrogen cyanide

  • Hydrogen cyanide contains a triple bond
  • One σ bond is formed between the H and C atom (overlap of an sp C hybridised orbital with the 1s H orbital)
  • A second σ bond is formed between the C and N atom (overlap of an sp C hybridised orbital with an sp orbital of N)
  • The remaining two sets of p orbitals of nitrogen and carbon will overlap to form two π bonds at right angles to each other

Sigma and pi bonding in hydrogen cyanide

Hydrogen cyanide has a triple bond formed from the overlap of two sets of p orbitals of nitrogen and carbon and the overlap of an sp hybridised carbon orbital and a p orbital on the nitrogen

• Nitrogen

  • Nitrogen too contains a triple bond
  • The triple bond is formed from the overlap of the sp orbitals on each N to form a σ bond and the overlap of two sets of p orbitals on the nitrogen atoms to form two π bonds
  • These π bonds are at right angles to each other

Sigma and pi bonding in nitrogen molecules

The triple bond is formed from two π bonds and one σ bond

Hybridisation

• The p atomic orbitals can also overlap end-on to form σ bonds
• In order for them to do this, they first need to become modified in order to gain s orbital character
• The orbitals are therefore slightly changed in shape to make one of the p orbital lobes bigger
• This mixing of atomic orbitals to form covalent bonds is called hybridisation

What is sp³ hybridisation?

• One s orbital and three p orbitals from the same shell mix to form four sp³hybrid 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

4 x sp³ hybrid orbitals are formed from one s orbital and three p orbitals

• 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 sp²hybrid 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

3 x sp2 hybrid orbitals are formed from one s orbital and two p orbitals

• 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

2 sp hybrid orbitals are formed from one s orbital and one p orbital

• 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 C$\equiv$C triple bond containing 1 σ and 2 perpendicular π bonds
    • Linear electron domain geometry
    • Linear molecular geometry
    • A 180° bond angle around each carbon

Exam Hint Carbon forms four bonds. The type of bond depends on how many p orbitals are used in hybridisation:

• sp³

  • 4 - 3 = 1
  • So, the carbon atom forms single bonds

• sp²

  • 4 - 2 = 2
  • So, the carbon atom forms a double bond

• sp

  • 4 - 1 = 3
  • So, the carbon atom forms a triple bond

This page focuses on carbon (a second-period element), but hybridisation also occurs in third-period elements like phosphorus and sulfur.

• These atoms use 3s and 3p orbitals
• They may also use 3d orbitals when forming expanded octets