Covalent Bonding
Explore how atoms share electrons to form covalent bonds, understand single, double and triple bonds, and learn about molecular shapes using VSEPR theory.
Sharing Electrons
A covalent bond forms when two non-metal atoms share one or more pairs of electrons. Each atom contributes one electron to the shared pair, allowing both atoms to achieve a stable outer shell configuration.
Types of Covalent Bonds
Single Bond
1 shared pair (2 electrons)
Example: H2, H-H
Double Bond
2 shared pairs (4 electrons)
Example: O2, O=O
Triple Bond
3 shared pairs (6 electrons)
Example: N2, N≡N
Bond strength: Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. N≡N bond energy: 945 kJ mol-1, O=O: 498 kJ mol-1, H-H: 436 kJ mol-1.
Lewis Dot Structures
Lewis dot structures (or electron dot diagrams) show how valence electrons are arranged in a molecule. Bonding pairs are shared between atoms; lone pairs (non-bonding pairs) belong to only one atom. Drawing Lewis structures helps predict molecular geometry.
Steps to Draw Lewis Structures
Step 1: Count valence electrons
Add up all valence electrons from every atom
Step 2: Draw single bonds
Connect atoms with single bonds (2 electrons each)
Step 3: Complete octets
Add lone pairs to outer atoms first, then central atom
Step 4: Form multiple bonds if needed
Convert lone pairs to bonding pairs if central atom lacks an octet
Exception: Hydrogen only needs 2 electrons (duet) to be stable. Some elements in Period 3 and beyond can have an expanded octet (more than 8 electrons) due to available d orbitals.
Molecular Shapes (VSEPR Theory)
The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes. Electron pairs around a central atom repel each other and arrange themselves as far apart as possible.
Linear (2 bonding, 0 lone pairs)
Bond angle: 180 degrees. Example: CO2, BeCl2.
Trigonal Planar (3 bonding, 0 lone pairs)
Bond angle: 120 degrees. Example: BF3, CH2O.
Tetrahedral (4 bonding, 0 lone pairs)
Bond angle: 109.5 degrees. Example: CH4, CCl4.
Bent (2 bonding, 1-2 lone pairs)
Bond angle: approx. 104.5 degrees. Example: H2O (2 bonding, 2 lone pairs).
Polarity of Molecules
Non-polar molecules: Symmetrical shape, bond dipoles cancel out. Examples: CO2, CH4, O2.
Polar molecules: Asymmetric shape or lone pairs create a net dipole moment. Examples: H2O, NH3, HCl.
Electronegativity difference between bonded atoms determines bond polarity. Larger difference = more polar bond.
Key Vocabulary
Covalent Bond
A chemical bond formed by the sharing of one or more pairs of electrons between two non-metal atoms.
Lone Pair
A pair of valence electrons not involved in bonding, belonging to a single atom. Lone pairs affect molecular shape.
Electronegativity
A measure of an atom's ability to attract shared electrons in a covalent bond. Fluorine is the most electronegative element.
VSEPR Theory
Valence Shell Electron Pair Repulsion theory predicts the 3D shape of molecules based on the repulsion between electron pairs.
Worked Examples
Draw the Lewis structure of water (H2O).
Step 1: Valence electrons: O = 6, H = 1 each. Total = 6 + 1 + 1 = 8.
Step 2: Draw O in centre with single bonds to each H. This uses 4 electrons (2 bonds).
Step 3: Remaining 4 electrons go on O as 2 lone pairs.
Answer: O has 2 bonding pairs and 2 lone pairs, giving a bent shape with bond angle ~104.5 degrees.
Explain why CO2 is non-polar despite having polar bonds.
Step 1: Each C=O bond is polar (O is more electronegative).
Step 2: CO2 is linear (180 degrees) -- no lone pairs on the central C atom.
Answer: The two bond dipoles point in exactly opposite directions and cancel each other out, resulting in a non-polar molecule overall.
Predict the shape of ammonia (NH3).
Step 1: N has 5 valence electrons. Total with 3 H: 5 + 3 = 8 electrons = 4 pairs.
Step 2: 3 bonding pairs (N-H) and 1 lone pair on N.
Answer: 4 electron pairs give a tetrahedral electron arrangement, but the molecular shape is trigonal pyramidal (bond angle ~107 degrees) because the lone pair compresses the bond angles.
Knowledge Check
Select the correct answer for each question. Click "Check Answer" to see if you are right.
Question 1
Covalent bonds form between:
Question 2
How many shared electron pairs are in a double bond?
Question 3
What is the shape of methane (CH4)?
Question 4
Which molecule is polar?
Question 5
What effect do lone pairs have on bond angles?
Key Concepts Summary
- ●Covalent bonds form by sharing electron pairs between non-metal atoms.
- ●Single, double, and triple bonds involve 1, 2, and 3 shared pairs respectively. More pairs = stronger, shorter bond.
- ●Lewis structures show bonding pairs and lone pairs, helping predict molecular geometry.
- ●VSEPR theory predicts molecular shapes based on electron pair repulsion.
- ●Polarity depends on both bond polarity and molecular shape -- symmetrical molecules can be non-polar even with polar bonds.