Structure and Bonding - Introduction

Chemical bonding is one of the most important concepts in chemistry because it explains how atoms join together to form all the materials around us. Understanding bonding helps us predict the properties of substances and explain why materials behave the way they do.

Types of Strong Chemical Bond

All materials are held together by three main types of strong chemical bond. Each bond involves electrostatic attraction (the force between charged particles), but the particles involved and the way electrons behave are different.

Ionic Bonding

Ionic bonding happens when metals react with non-metals - like when sodium reacts with chlorine to make table salt.

• Metal atoms lose one or more electrons to form positively charged ions, for example Na+Na+

• Non-metal atoms gain those electrons to form negatively charged ions, for example Cl−Cl−

• The oppositely charged ions attract each other through a strong electrostatic force that acts in all directions, creating a giant ionic lattice

Think of this like a 3D framework where positive and negative ions are arranged in a regular pattern. Because of this three-dimensional network, ionic compounds such as sodium chloride form crystalline solids with high melting points.

Covalent Bonding

Covalent bonding occurs when non-metals bond together. This happens in elements like oxygen gas (O2O2​) and in compounds like water (H2OH2​O).

• Atoms share pairs of electrons to fill their outer shells

• The shared pair creates an electrostatic attraction between the nuclei and the bonding electrons

Depending on how many atoms join together, covalent substances can exist as small molecules (like water), giant covalent structures (like diamond) or polymers (like plastics).

Metallic Bonding

Metallic bonding is found in pure metals like copper and in alloys such as brass (a mixture of copper and zinc).

• Metal atoms release some of their outer electrons, which become delocalised (not attached to any particular atom) and move freely through the structure

• Positive metal ions are arranged in layers, and the delocalised electrons attract the ions, binding the structure together

This 'sea of delocalised electrons' explains why metals conduct electricity and heat, and why they can be hammered into shape (malleable) without breaking.

How Chemical Bonds Form

Only the electrons in the highest occupied energy level (the outer shell) are involved in bonding. The type of bond that forms depends on what happens to these electrons:

1. Transfer of electrons produces ionic bonds
   Na→Na++e−Na→Na++e−

2. Sharing of electrons produces covalent bonds
   Cl+Cl→Cl2Cl+Cl→Cl2​

3. Pooling of electrons leads to metallic bonds
   M→Mn++ne−M→Mn++ne−

The strength of each bond type comes from the electrostatic force of attraction between opposite charges. Because this force depends on the size of the charges and the distance between them, ionic compounds and metals typically have high melting points, while small covalent molecules may have low melting points and can be gases or liquids at room temperature.

Key terms

Ionic bond - An electrostatic attraction between oppositely charged ions formed when electrons transfer from a metal to a non-metal.

Covalent bond - A strong electrostatic attraction between two nuclei and a shared pair of electrons.

Metallic bond - The attraction between positive metal ions and a sea of delocalised electrons.

Electrostatic force - The attractive or repulsive force between charged particles.

Delocalised electron - An electron that is not attached to one specific atom and can move freely throughout the structure.

Giant ionic lattice - A regular 3D arrangement of positive and negative ions held together by ionic bonds.

Worked example

Question: A magnesium atom reacts with an oxygen atom. Explain the bonding that occurs and predict the formula of the compound formed.

Solution:

1. Identify the groups: Magnesium is in Group 2, oxygen is in Group 6

2. Work out electron transfer:

   • Magnesium loses two electrons: Mg→Mg2++2e−Mg→Mg2++2e−

   • Oxygen gains two electrons: O+2e−→O2−O+2e−→O2−

3. Explain the bonding: The oppositely charged ions Mg2+Mg2+ and O2−O2− attract each other through electrostatic forces to form ionic bonds

4. Determine the formula: Since one Mg²⁺ ion balances one O²⁻ ion, the formula is MgOMgO (magnesium oxide)

Real-world applications of bonding theory:

• Table salt (sodium chloride) conducts electricity when molten because the ions are free to move and carry charge. However, solid salt doesn't conduct because the ions are fixed in position in the crystal lattice.

• Diamond has an extremely high melting point (over 3500°C) because each carbon atom is covalently bonded to four others in a rigid 3D network that's very difficult to break.

• Copper wires conduct electricity because delocalised electrons can move freely throughout the metallic structure, carrying electrical current.

• Aluminium foil can be bent and shaped because the layers of metal atoms can slide over each other while the sea of electrons maintains the bonding.

Investigating electrical conductivity of different bond types

Aim: To compare the electrical conductivity of ionic, covalent and metallic substances and relate this to their bonding.

Apparatus:

• Simple electrical circuit with battery, bulb and connecting wires

• Solid sodium chloride

• Graphite (pencil lead)

• Aluminium foil

• Bunsen burner and crucible

• Tongs

Method:

1. Set up a simple circuit with a battery and bulb, leaving a gap for testing materials

2. Place solid sodium chloride in the gap and record whether the bulb lights

3. Repeat with graphite and aluminium foil

4. Carefully heat sodium chloride in a crucible until it melts

5. Test the molten sodium chloride (using tongs for safety)

Safety:

• Wear safety goggles when heating

• Use tongs when handling hot materials

• Ensure electrical connections are secure

Expected Results:

• Solid NaCl - bulb doesn't light (no conductivity)

• Molten NaCl - bulb lights (conducts electricity)

• Graphite - bulb lights dimly (limited conductivity)

• Aluminium - bulb lights brightly (excellent conductivity)

Conclusion: The results support bonding theory: ionic compounds only conduct when ions are mobile (molten), metals conduct due to delocalised electrons, and most covalent substances don't conduct (graphite is unusual because it has some delocalised electrons).

Comparison table