Hybridisation (Cambridge (CIE) AS Chemistry): Revision Note

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

• 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

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

• 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

Sigma and pi bonding in ethene

• 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

• 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

• 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

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 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