🌟 Welcome to the World of Covalent Bonding! 🌟

Hello future chemists! This chapter is part of your journey through the Principles of Chemistry, and it’s one of the most fundamental concepts you’ll learn: how atoms stick together to form almost everything around you.

Don't worry if this topic seems tricky at first. We will break down complex molecules into simple diagrams, and by the end of these notes, you’ll be drawing atoms sharing electrons like a professional!

What you will learn in this chapter:

  • What a covalent bond is and why atoms form them.
  • How to draw 'dot-and-cross' diagrams for diatomic, inorganic, and simple organic molecules.
  • The difference between simple molecules and giant covalent structures.
  • How structure dictates the properties of materials like diamond, graphite, and \(C_{60}\) fullerene.

1. The Basics of Covalent Bonding

1.1 What is a Covalent Bond?

Remember the driving force behind all chemical reactions? Stability! Atoms are happiest when their outer electron shells are full, usually with eight electrons (this is called the Octet Rule, or two electrons for hydrogen).

When non-metal atoms react together, they often cannot transfer electrons (like in ionic bonding). Instead, they find a clever solution: they share electrons.

The Definition:

A covalent bond is the strong electrostatic attraction between the nuclei of the bonded atoms and a shared pair of electrons between them.

Analogy: Sharing is Caring!

Imagine you and a friend each have one shoe, and you need a pair to go running. Neither of you will give up your shoe entirely (that would be ionic bonding), so you agree to share the two shoes between you. You both get to wear the "shared pair" of shoes. In chemistry, the "shared shoes" are a pair of electrons.

1.2 Who Forms Covalent Bonds?

  • Covalent bonds form primarily between non-metal atoms.
  • Examples: Carbon (C), Hydrogen (H), Oxygen (O), Chlorine (Cl), Nitrogen (N).

When atoms share electrons, they often form a group called a molecule. A molecule is the smallest unit of a simple covalently bonded substance. Example: Water, \(H_2O\), is a molecule.

Key Takeaway

Covalent bonds occur when non-metal atoms share pairs of outer shell electrons to achieve full outer shells, held together by electrostatic attractions between nuclei and shared electrons.

2. Drawing Covalent Molecules (Dot-and-Cross Diagrams)

Dot-and-cross diagrams are the essential way we visualise covalent bonding. They show outer-shell electrons (the valence electrons) and which ones are being shared.

2.1 Step-by-Step Guide to Drawing Diagrams

To draw a diagram, use the group number of each element to know how many outer shell electrons it has.

Step 1: Count Outer Electrons. Determine the number of valence electrons for each atom (e.g. Hydrogen has 1; Oxygen is Group 6, so 6; Carbon is Group 4, so 4).

Step 2: Assign Symbols. Draw the atoms overlapping where electrons will be shared. Use dots (β€’) for electrons from one atom and crosses (x) for electrons from the other.

Step 3: Share Electrons. Place electrons in the overlapping region so that every atom achieves a noble gas configuration (2 for H; 8 for C, N, O, halogens).

Step 4: Place Unshared Electrons. Put any remaining outer-shell electrons (lone pairs) around the respective atoms.

2.2 Examples of Covalent Molecules

Example A: Hydrogen Chloride (\(HCl\)) - Single Bond

Hydrogen (\(H\)) needs 1 electron. Chlorine (\(Cl\)) has 7 outer electrons and needs 1 more.

They share one pair of electrons (1 from H, 1 from Cl), forming a single covalent bond.

Example B: Oxygen (\(O_2\)) - Double Bond

Oxygen is in Group 6, so it has 6 outer electrons. Each atom needs 2 more to make 8.

They share two pairs of electrons (four electrons total), forming a double covalent bond.

Note: Carbon dioxide (\(CO_2\)) also contains two double bonds (\(O=C=O\)).

Example C: Nitrogen (\(N_2\)) - Triple Bond

Nitrogen is in Group 5, with 5 outer electrons. Each atom needs 3 more to make 8.

They share three pairs of electrons (six electrons total), forming a triple covalent bond.

! Common Mistake Alert !

Always double-check that every atom in your diagram has achieved a full shell! Count the electrons in the shared region for BOTH atoms. For Chlorine, the total must be 8 (6 non-bonding + 2 shared).

Quick Review: Types of Covalent Bonds

Single Bond: 1 shared pair (e.g. \(H_2\), \(HCl\), \(H_2O\), \(NH_3\), \(CH_4\))
Double Bond: 2 shared pairs (e.g. \(O_2\), \(CO_2\), \(C_2H_4\))
Triple Bond: 3 shared pairs (e.g. \(N_2\), \(C_2H_2\))

3. Properties of Simple Molecular Substances

Many covalently bonded substances, such as water (\(H_2O\)), ammonia (\(NH_3\)), methane (\(CH_4\)), and carbon dioxide (\(CO_2\)), exist as small, separate molecules.

3.1 The Key to Properties: Types of Forces

To understand physical properties, distinguish between two forces:

  1. Covalent Bonds: The strong electrostatic attractions within the molecule (holding H to O in \(H_2O\)). These are not broken when substances melt or boil.
  2. Intermolecular Forces: The weak attractions between separate molecules (holding one \(H_2O\) molecule next to another).
Analogy: The Lego Bricks

Think of a molecule as a single, solid Lego brick. The plastic forming the brick (the covalent bond) is very tough. But pulling one brick away from another (breaking the weak intermolecular force) requires far less energy.

When simple molecular substances melt or boil, only the weak intermolecular forces of attraction are overcome, while the strong covalent bonds remain intact.

3.2 Melting and Boiling Points

  • Low Melting/Boiling Points: Simple molecular substances are usually gases, liquids, or low-melting solids at room temperature because little energy is needed to overcome the weak intermolecular forces.
  • Trend with Relative Molecular Mass (\(M_r\)): As relative molecular mass increases, the intermolecular forces of attraction between molecules become stronger. More energy is required to overcome these forces, so melting and boiling points increase as \(M_r\) increases.

3.3 Electrical Conductivity

  • Non-conductors: Simple molecular substances do not conduct electricity in solid, liquid, or gaseous states.
  • Explanation: They have no free ions and no delocalised electrons to move and carry an electric charge; all valence electrons are localised in bonds or lone pairs.

Key Takeaway

Simple molecular substances have low melting and boiling points due to weak intermolecular forces of attraction (which increase with molecular mass) and do not conduct electricity because they lack mobile charged particles.

4. Giant Covalent Structures (Macromolecules) and Allotropes of Carbon

In giant covalent structures, millions of atoms are linked by strong covalent bonds in continuous three-dimensional networks. Because huge amounts of energy are needed to break these strong covalent bonds, giant covalent structures are solids with very high melting and boiling points.

4.1 Diamond (Giant Covalent)

Structure: Each carbon atom forms four strong covalent bonds to four other carbon atoms in a rigid, 3D tetrahedral network.

Properties & Explanations:

  • Very High Melting Point: Breaking the vast network of strong covalent bonds requires massive thermal energy.
  • Extremely Hard: The rigid, tetrahedral giant lattice firmly holds every atom in place. (Used in cutting tools and drill bits).
  • Electrical Insulator: Does not conduct electricity because all four outer valence electrons of every carbon atom are held in covalent bonds (no free or delocalised electrons).

4.2 Graphite (Giant Covalent)

Structure: Each carbon atom forms three strong covalent bonds to three other carbon atoms, creating flat, hexagonal layers. Between the layers, there are only weak intermolecular forces.

Properties & Explanations:

  • Soft and Slippery: The weak forces between layers allow the layers to slide over each other easily. (Used as a solid lubricant and in pencils).
  • High Melting Point: A large amount of energy is still required to break the strong covalent bonds within the hexagonal layers.
  • Conducts Electricity: Because each carbon atom only forms 3 bonds, its 4th outer electron is delocalised and free to move throughout the layers to carry charge. (Used as electrodes).

4.3 \(C_{60}\) Fullerene (Simple Molecular)

Structure: \(C_{60}\) fullerene consists of hollow spheres of 60 carbon atoms bonded together in a soccer-ball shape (pentagons and hexagons). Unlike diamond and graphite, \(C_{60}\) is a simple molecular substance, not a giant lattice.

Properties & Explanations:

  • Lower Melting/Boiling Point than Diamond or Graphite: Although the covalent bonds within each \(C_{60}\) sphere are strong, only weak intermolecular forces exist between individual \(C_{60}\) molecules, requiring much less energy to overcome.
  • Soft / Slippery: The individual spherical molecules can roll and slide over one another due to weak intermolecular attractions.
  • Poor Electrical Conductor: Although each carbon atom forms 3 covalent bonds and has a delocalised electron, the electrons cannot move freely between different \(C_{60}\) molecules, so fullerene is a poor conductor of electricity compared to graphite.

πŸ“š Chapter Review: Summary Comparison

Substance / Structure Bonding / Forces Broken on Melting Melting & Boiling Points Electrical Conductivity Hardness / Texture
Simple Molecular (e.g. \(H_2O\), \(CO_2\), \(CH_4\)) Weak intermolecular forces between molecules Low (increases with higher \(M_r\)) Non-conductor (no mobile charged particles) Usually gases, liquids, or soft solids
Diamond (Giant Covalent) Many strong covalent bonds throughout 3D lattice Very High Non-conductor (all electrons localised in 4 bonds) Extremely hard
Graphite (Giant Covalent) Strong covalent bonds within layers Very High Good conductor (1 delocalised electron per C atom) Soft and slippery (layers slide)
\(C_{60}\) Fullerene (Simple Molecular) Weak intermolecular forces between \(C_{60}\) cages Moderate (much lower than diamond/graphite) Poor conductor (electrons do not easily jump molecules) Soft / slippery solid (spheres slide over each other)