Combined Science (9204) | Structure, bonding and the properties of matter

Welcome to Structure, Bonding, and Properties of Matter!

Hello future chemist! This chapter is the absolute foundation of understanding how the world around us is built—from the salt you sprinkle on your food to the wires that carry electricity. We are going to explore why atoms stick together and how their "sticking pattern" determines everything about the substance, like whether it melts easily or conducts electricity.

Don't worry if bonding seems tricky at first. We will break it down into simple steps, focusing on why atoms behave like tiny, determined builders trying to achieve maximum stability!

---

1. The Chemistry of Stability: Why Atoms Bond

1.1 The Goal of Every Atom

Remember that atoms are made of protons, neutrons, and electrons. The electrons orbit the nucleus in specific shells.

• The inner shell needs 2 electrons to be full.
• Subsequent shells usually need 8 electrons to be full (this is often called the octet rule).

Atoms are only truly happy (or stable) when their outermost shell (the valence shell) is completely full. The Noble Gases (Group 0/8) are already stable—that’s why they rarely react!

1.2 Forming Bonds

Atoms that are unstable achieve a full outer shell in two main ways:

1. Transferring electrons (happens between metals and non-metals) → Ionic Bonding.
2. Sharing electrons (happens between non-metals) → Covalent Bonding.

Key Takeaway: Bonding happens so atoms can achieve a stable, full outer electron shell.

---

2. Ionic Bonding: The Transfer of Power (and Electrons)

2.1 How Ions are Formed

Ionic bonding involves the complete transfer of electrons from one atom to another, creating charged particles called ions.

1. Metals (Groups 1, 2, 3) like to lose their few outer electrons to get back to the previous, full shell. When they lose negative electrons, they become positively charged cations. (Think: Na → Na⁺)
2. Non-metals (Groups 5, 6, 7) like to gain electrons to fill their shell. When they gain negative electrons, they become negatively charged anions. (Think: Cl → Cl⁻)

Analogy: Ionic bonding is like a "borrowing and lending" scheme. The Metal (lender) gets rid of its excess electron, and the Non-metal (borrower) gains the electron it desperately needs. They are both happier for it!

2.2 The Structure of Ionic Compounds

Once the positive ions (cations) and negative ions (anions) are formed, they are powerfully attracted to each other due to electrostatic forces (opposite charges attract).

These ions pack together in a highly regular, repeating 3D pattern called a giant ionic lattice. (Table salt, NaCl, is a perfect example!)

2.3 Properties of Ionic Compounds

The strong electrostatic forces holding the giant lattice together give ionic compounds very distinct properties:

1. High Melting and Boiling Points: A lot of energy is needed to overcome the incredibly strong electrostatic attraction throughout the entire lattice.
2. Usually Soluble in Water: Water molecules can pull the individual ions away from the lattice.
3. Conductivity:
   • Solid state: They DO NOT conduct electricity because the ions are locked in fixed positions and cannot move.
   • Molten (liquid) or Dissolved state: They DO conduct electricity because the ions are free to move and carry the charge.

---

3. Covalent Bonding: The Act of Sharing

3.1 Forming Molecules

Covalent bonding occurs when non-metal atoms react together. Since both atoms want to gain electrons, they can’t simply transfer them. Instead, they share pairs of electrons so that both atoms feel like they have a full outer shell.

When atoms are held together by covalent bonds, they form a small unit called a molecule. Examples include H₂O (water), O₂ (oxygen gas), and CH₄ (methane).

3.2 Dot-and-Cross Diagrams

We use dot-and-cross diagrams to show which electrons belong to which atom, and which electrons are shared.

• Electrons from one atom are shown as dots (\(\bullet\)).
• Electrons from the second atom are shown as crosses (\(\times\)).
• The shared electrons sit in the middle, forming the covalent bond.

Example: In a chlorine molecule (Cl₂), each Cl atom has 7 outer electrons. They share one pair of electrons, making a single covalent bond. Now, each Cl atom "sees" 8 electrons, achieving stability.

3.3 Properties of Simple Molecular Substances

Molecules (like H₂O or CO₂) are held together by very strong covalent bonds within the molecule. However, the forces between one molecule and the next are very weak. These weak forces are called intermolecular forces.

Analogy: Think of a box of Lego models. The bricks holding one model together are strong (covalent bonds). But the force holding two separate models together (the intermolecular force) is just gravity—very weak!

Because these intermolecular forces are weak, simple molecular substances have specific properties:

1. Low Melting and Boiling Points: Very little energy is needed to break the weak forces between the molecules, so they often exist as gases or liquids at room temperature.
2. Volatility: They evaporate easily (low boiling point).
3. Poor Electrical Conductors: There are no free ions or delocalised electrons to carry charge.

Common Mistake Alert! Always remember the difference: you only need to break the weak intermolecular forces to melt or boil a simple molecular substance, NOT the strong covalent bonds!

---

4. Metallic Bonding: The Sea of Electrons

4.1 Structure of Metals

Metals like to lose their outer electrons. In a solid piece of metal, these outer electrons are not attached to any single atom; they are shared among the entire structure.

Metallic bonding is best described as a regular lattice of positive metal ions surrounded by a mobile 'sea' of delocalised electrons.

4.2 Properties Explained by Metallic Structure

The unique structure of metallic bonding explains why metals have their signature properties:

1. Good Electrical Conductors: The delocalised electrons are free to move throughout the structure and carry electrical charge (current).
2. Good Thermal Conductors: The free electrons can also transfer thermal (heat) energy quickly through the structure.
3. Malleability and Ductility: Metals are malleable (can be hammered into sheets) and ductile (can be drawn into wires). This is because the layers of positive ions can slide over each other without breaking the metallic bond, as the sea of electrons instantly shifts to hold them together again.

Did you know? Gold is so ductile that one gram can be drawn into a wire 2.4 km long!

---

5. Giant Covalent Structures (Allotropes of Carbon)

Some substances based on covalent bonding don't form small, simple molecules. Instead, they form massive 3D structures where every atom is linked to its neighbours by strong covalent bonds. These are called giant covalent structures.

The most important examples in your curriculum are the different forms of Carbon, known as allotropes.

5.1 Diamond (The Ultra-Hard Giant)

• Structure: Each Carbon atom is covalently bonded to four other carbon atoms in a rigid, tetrahedral structure.
• Properties:
  • Extremely hard (used in cutting tools) because of the strong, 3D network.
  • Very high melting point.
  • Does NOT conduct electricity (all outer electrons are used up in the four strong bonds, so none are delocalised).

5.2 Graphite (The Slippery Conductor)

• Structure: Each Carbon atom is covalently bonded to only three other carbon atoms, forming flat, hexagonal layers.
• Properties:
  • Soft and slippery (used as a lubricant or in pencil 'lead'). This is because the forces between the layers are very weak, allowing them to slide easily.
  • Good electrical conductor. Since only three of the four outer electrons are used in bonding, the fourth electron is delocalised (free to move) between the layers.
  • Very high melting point (since breaking the layers still requires breaking strong covalent bonds).

5.3 Graphene

• Structure: Graphene is essentially just a single layer of the graphite structure (a one-atom-thick sheet).
• Properties: It is extremely strong, light, and an excellent conductor of electricity (due to delocalised electrons). It is currently being studied for many high-tech uses.

You have mastered the foundations of chemical structure! Understanding these core concepts—transfer, sharing, and the sea of electrons—will make all future chemistry topics much easier. Great job!