CXC Chemistry: Chemical Bonding

This comprehensive lesson covers the entire Chemical Bonding topic as per the CXC Chemistry syllabus for 2024/2025. It includes detailed explanations, diagrams, examples, and self-assessment questions to help you master this crucial topic.

Introduction to Chemical Bonding

Chemical bonding is the process by which atoms join together to form stable compounds. Understanding chemical bonding is fundamental to understanding chemistry as a whole.

Why Do Atoms Bond?

Atoms bond to achieve a more stable electron configuration, typically by:

Noble Gas Configuration: Most atoms aim to achieve the electron configuration of the nearest noble gas, which has a complete outer shell of electrons.

Types of Chemical Bonds

There are three primary types of chemical bonds:

1. Ionic Bonding

Ionic bonds form between metals and non-metals through the complete transfer of electrons from one atom to another, resulting in positively and negatively charged ions that attract each other.

Characteristics of Ionic Bonds:

Na Na+ Cl Cl-

Fig. 1: Formation of an ionic bond between sodium and chlorine to form sodium chloride (NaCl)

Example: Formation of sodium chloride (NaCl)

Na (2,8,1) → Na+ (2,8) + e- [Loses 1 electron]

Cl (2,8,7) + e- → Cl- (2,8,8) [Gains 1 electron]

The sodium ion (Na+) and chloride ion (Cl-) are attracted to each other by electrostatic forces, forming NaCl.

2. Covalent Bonding

Covalent bonds form between non-metals through the sharing of electrons, resulting in molecules with shared electron pairs.

Characteristics of Covalent Bonds:

H H H H

Fig. 2: Formation of a covalent bond in hydrogen molecule (H2)

Example: Formation of a water molecule (H2O)

Oxygen (atomic number 8) has electronic configuration 2,6

Hydrogen (atomic number 1) has electronic configuration 1

Oxygen needs 2 more electrons to achieve octet (8)

Each hydrogen needs 1 more electron to achieve duet (2)

Therefore, oxygen forms 2 covalent bonds with 2 hydrogen atoms, sharing 2 pairs of electrons.

Types of Covalent Structures:

  1. Simple Molecular Structures: Discrete molecules held together by weak intermolecular forces (e.g., H2O, NH3, CH4)
  2. Giant Covalent/Network Structures: Extended 3D networks of covalently bonded atoms (e.g., diamond, graphite, silicon dioxide)

Fig. 3: Giant covalent structures - Diamond (left) and Graphite (right)

Polar and Non-polar Covalent Bonds

Covalent bonds can be classified as either polar or non-polar:

Cl Cl Non-polar H Cl δ+ δ- Polar

Fig. 4: Non-polar bond (Cl2) vs. Polar bond (HCl)

3. Metallic Bonding

Metallic bonds form between metal atoms, creating a "sea" of delocalized electrons surrounding positive metal ions.

Characteristics of Metallic Bonds:

+ + + + + + + +

Fig. 5: Metallic bonding - Positive metal ions in a "sea" of delocalized electrons

Example: Metallic bonding in copper (Cu)

Copper atoms each contribute one electron to the delocalized "sea" of electrons.

The resulting Cu+ ions are arranged in a regular lattice structure.

The delocalized electrons are free to move throughout the structure, which explains copper's excellent electrical conductivity.

Comparison of Bond Types

Property Ionic Bonding Covalent Bonding Metallic Bonding
Formation Transfer of electrons Sharing of electrons Delocalization of electrons
Between Metal and non-metal Non-metal and non-metal Metal and metal
Physical state at room temperature Usually solid Solid, liquid, or gas Solid (except mercury)
Melting/Boiling points High Typically low (simple molecular)
High (giant covalent)		 Typically high
Electrical conductivity Conducts when molten or in solution Usually poor conductors Good conductors
Solubility in water Usually soluble Usually insoluble (some are soluble) Insoluble
Examples NaCl, CaO, MgCl2 H2O, CO2, CH4,
Diamond, Graphite		 Cu, Fe, Na, Al, Ag

Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules. These are much weaker than the chemical bonds within molecules but are crucial for determining physical properties.

Types of Intermolecular Forces

  1. Van der Waals Forces (Dispersion Forces): Weak attractions resulting from temporary dipoles in molecules
  2. Dipole-Dipole Interactions: Attractions between polar molecules
  3. Hydrogen Bonding: Special case of dipole-dipole interaction involving hydrogen bonded to highly electronegative elements (F, O, N)
H δ+ O δ- H δ+ H δ+ O δ- H δ+ H-bond Cl δ- H δ+ H δ+ Cl δ- Dipole-dipole Strong → Hydrogen bonding > Dipole-dipole > Van der Waals ← Weak

Fig. 6: Intermolecular forces - Hydrogen bonding between water molecules (top) and dipole-dipole interactions between HCl molecules (bottom)

Example: Hydrogen bonding in water (H2O)

In water, the hydrogen atoms have a partial positive charge (δ+) while the oxygen atom has a partial negative charge (δ-).

The hydrogen atom of one water molecule is attracted to the oxygen atom of another water molecule, forming a hydrogen bond.

This explains water's unusually high boiling point compared to other molecules of similar size.

Effects of Intermolecular Forces on Physical Properties

The strength of intermolecular forces affects several physical properties:

Bond Energy and Bond Length

Bond energy is the energy required to break a chemical bond, while bond length is the distance between the nuclei of two bonded atoms.

Factors Affecting Bond Energy and Bond Length

Important relationships:

Bond Type Average Bond Energy (kJ/mol) Average Bond Length (pm)
C-C (single bond) 348 154
C=C (double bond) 614 134
C≡C (triple bond) 839 120
C-H 413 109
O-H 463 96
N-H 391 101

Lewis Structures and VSEPR Theory

Lewis structures show the arrangement of valence electrons in molecules, while VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on electron pair repulsions.

Drawing Lewis Structures

  1. Calculate the total number of valence electrons
  2. Draw the skeletal structure (central atom usually has lowest electronegativity)
  3. Complete octets for terminal atoms
  4. Place remaining electrons on central atom
  5. Form multiple bonds if necessary to achieve octets

Example: Lewis structure for CO2

  1. Total valence electrons: C (4) + O (6) + O (6) = 16 electrons
  2. Skeletal structure: O-C-O
  3. Complete terminal atom octets: O:C:O (using 12 electrons)
  4. Remaining 4 electrons go on central C
  5. Form double bonds to achieve octets: O=C=O

VSEPR Theory and Molecular Geometry

VSEPR theory predicts molecular geometry based on the arrangement of electron pairs around the central atom:

Electron Pair Geometry Molecular Geometry (examples) Bond Angle
2 electron pairs Linear (CO2, BeF2) 180°
3 electron pairs Trigonal planar (BF3)
Bent (SO2)		 120°
~119°
4 electron pairs Tetrahedral (CH4)
Pyramidal (NH3)
Bent (H2O)		 109.5°
~107°
~104.5°
5 electron pairs Trigonal bipyramidal (PCl5)
See-saw (SF4)
T-shaped (ClF3)
Linear (I3-)		 90°, 120°
~90°, ~120°
~90°
180°
6 electron pairs Octahedral (SF6)
Square pyramidal (BrF5)
Square planar (XeF4)		 90°
90°
90°
O C O Linear 180° F B F F Trigonal Planar 120° C H H H H Tetrahedral 109.5° O H H Bent 104.5°

Fig. 7: Common molecular geometries predicted by VSEPR theory

Electronegativity and Polarity

Electronegativity is a measure of an atom's tendency to attract electrons in a chemical bond. Differences in electronegativity lead to bond polarity.

Electronegativity Trends

Determining bond type based on electronegativity difference (ΔEN):

Molecular Polarity

A molecule is polar if it has:

  1. Polar bonds (due to electronegativity differences)
  2. An asymmetrical arrangement of these polar bonds

Examples:

Applications and Properties of Different Bonding Types

Applications of Ionic Compounds

Applications of Covalent Compounds

Applications of Metals and Alloys

Glossary of Key Terms

Bond: The force of attraction that holds atoms together in a compound.
Chemical Bond: A lasting attraction between atoms, ions, or molecules that enables the formation of chemical compounds.
Valence Electrons: Electrons in the outermost shell of an atom that participate in bonding.
Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight electrons in their valence shell.
Ionic Bond: A chemical bond formed through the complete transfer of electrons from one atom to another.
Covalent Bond: A chemical bond formed through the sharing of electron pairs between atoms.
Metallic Bond: A chemical bond formed by the attraction between metal cations and delocalized electrons.
Electronegativity: A measure of an atom's tendency to attract electrons in a chemical bond.
Polar Bond: A covalent bond in which electrons are shared unequally, resulting in partial positive and negative charges.
Non-polar Bond: A covalent bond in which electrons are shared equally between atoms.
Polar Molecule: A molecule that has an uneven distribution of charge due to polar bonds and asymmetric structure.
Intermolecular Forces: Forces of attraction between molecules.
Hydrogen Bond: A special type of dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative atom and another highly electronegative atom.
Bond Energy: The energy required to break a chemical bond.
Bond Length: The average distance between the nuclei of two bonded atoms.
Lewis Structure: A representation of covalent bonding using dots and lines to show valence electrons and bonds.
VSEPR Theory: Valence Shell Electron Pair Repulsion theory, used to predict the geometry of molecules.
Alloy: A mixture of a metal with one or more other elements.
Coordinate Covalent Bond: A covalent bond in which both shared electrons come from the same atom.
Dipole Moment: A measure of the polarity of a bond or molecule.

Self-Assessment Questions

Multiple Choice Questions

1. Which type of bond forms between a metal and a non-metal?

  1. Ionic bond
  2. Covalent bond
  3. Metallic bond
  4. Hydrogen bond

Answer: a. Ionic bond

Explanation: Ionic bonds typically form between metals (which tend to lose electrons) and non-metals (which tend to gain electrons). The metal atom donates electrons to the non-metal atom, resulting in the formation of positive and negative ions that attract each other.

2. Which of the following compounds has the highest melting point?

  1. CH4 (methane)
  2. NaCl (sodium chloride)
  3. H2O (water)
  4. C2H6 (ethane)

Answer: b. NaCl (sodium chloride)

Explanation: Ionic compounds like NaCl typically have high melting points due to the strong electrostatic forces between the oppositely charged ions. Simple molecular compounds like methane, water, and ethane have much lower melting points because they are held together by weaker intermolecular forces.

3. Which of the following represents a polar covalent bond?

  1. H-H
  2. Cl-Cl
  3. O-H
  4. Na-Cl

Answer: c. O-H

Explanation: The O-H bond is polar covalent due to the significant difference in electronegativity between oxygen (3.5) and hydrogen (2.1). H-H and Cl-Cl are non-polar covalent bonds because the electronegativity difference is zero. Na-Cl is an ionic bond due to the very large electrone