⚛️ CORE Chemistry Study Notes: How Bonding and Structure Relate to Properties ⚛️
Hello future scientist! Welcome to the section where chemistry really makes sense!
Have you ever wondered why metals can be bent, but salt crystals shatter? Or why water boils at a low temperature, but diamond is one of the hardest substances on Earth? The answer lies entirely in how the atoms are bonded and arranged.
In this chapter, we will learn that the type of bond (ionic, covalent, metallic) and the structure (simple or giant) determine all the properties of a substance—like its melting point, hardness, and whether it conducts electricity. Let's dive in!
1. Structures and their Foundations
We classify substances into four main structural groups. The term Giant Structure means the atoms or ions are bonded together in a massive, repeating network, often called a lattice, with no set number of particles.
Quick Review: The Two Main Structure Types
- Simple Molecular Structures: Made up of small, discrete molecules (like H₂O or CO₂). They have strong bonds *inside* the molecules, but weak forces *between* the molecules.
- Giant Structures: These are huge networks (lattices) where all atoms or ions are held together by strong forces. This includes Giant Ionic, Giant Metallic, and Giant Covalent structures.
2. Properties of Giant Ionic Structures
Giant Ionic structures are formed when metals react with non-metals (e.g., Sodium Chloride, NaCl).
Key Feature: Strong Electrostatic Attraction
In these structures, positive ions (cations) and negative ions (anions) are held in a massive, regular lattice structure by very strong forces of attraction (like powerful magnets).
A. Melting and Boiling Points (High!)
Since the forces holding the ions together are so strong (called electrostatic forces), a huge amount of energy is needed to break this lattice and turn the solid into a liquid or gas.
- Property: High melting points and boiling points.
B. Conductivity (The Mobile Charge Rule)
To conduct electricity, a substance needs mobile charge carriers (either free-moving ions or delocalized electrons).
- Solid State: The ions are locked in fixed positions in the lattice. They cannot move.
Result: Do not conduct electricity when solid. - Molten (Liquid) or Dissolved State: When the substance melts or dissolves in water, the lattice breaks down. The ions are now free to move and carry charge.
Result: Conduct electricity when molten or in solution.
💡 Common Mistake to Avoid: Students often think ions can move in the solid. Remember the lattice keeps them tightly fixed!
C. Hardness and Brittleness
Ionic crystals are hard, but they are also brittle (they shatter easily).
Step-by-step: Why Ionic Solids Shatter
- The structure contains alternating positive and negative ions.
- If you hit the crystal, a layer of ions shifts slightly.
- When shifted, ions with the same charge are forced next to each other (positive next to positive, negative next to negative).
- These like charges strongly repel, causing the crystal to break apart (shatter).
Key Takeaway for Giant Ionic Structures: They are strong, melt high, and only conduct electricity when the ions are free to move.
3. Properties of Simple Molecular Structures
These substances are formed by non-metals sharing electrons (covalent bonding), creating small, specific molecules (e.g., O₂, I₂, H₂O, Methane).
The Key Distinction: Strong Bonds vs. Weak Forces
Imagine a Lego house:
The covalent bonds holding atoms together inside one molecule (the individual Lego bricks) are very strong.
The forces holding one molecule to its neighbours (the weak attraction between the finished Lego houses) are very weak. These are called intermolecular forces (IMFs).
A. Melting and Boiling Points (Low!)
When you melt or boil a simple molecular substance, you are not breaking the strong covalent bonds inside the molecule. You are only breaking the weak intermolecular forces between the molecules.
- Property: Low melting and boiling points (often liquids or gases at room temperature).
- Analogy: It takes very little effort to pull the Lego houses apart from each other.
B. Conductivity (Zero!)
Simple molecular substances do not contain any ions, nor do they have any delocalized (free) electrons.
- Property: Do not conduct electricity in any state (solid, liquid, or gas).
Key Takeaway for Simple Molecular Structures: They are held together by weak forces between molecules, leading to low melting points and no conductivity.
4. Properties of Giant Metallic Structures
Metals are made purely of metal atoms (e.g., copper, iron, gold) bonded together.
Key Feature: The Sea of Delocalized Electrons
In metallic bonding, the outer shell electrons (valence electrons) are stripped away from the positive metal ions. These electrons are free to move throughout the entire structure. We call this the sea of delocalized electrons.
A. Conductivity (Excellent!)
Metals are superb conductors because the delocalized electrons are free to move. If a voltage is applied, these electrons rush through the structure, carrying the charge.
- Property: Excellent conductors of electricity and heat.
B. Melting Points (Usually High!)
The attraction between the positive metal ions and the surrounding sea of negative electrons is generally very strong, requiring a lot of energy to break this giant lattice.
- Property: High melting points (though some, like sodium, are lower).
C. Malleability and Ductility
This is where metals differ dramatically from ionic compounds.
- Malleable: Can be hammered into sheets (like aluminium foil).
- Ductile: Can be drawn out into wires (like copper wire).
When you hit a metal, the layers of positive ions can slide over each other. Because the sea of delocalized electrons holds the structure together non-directionally, the layers can shift without causing massive repulsion, meaning the metal changes shape instead of shattering.
Key Takeaway for Giant Metallic Structures: High melting points, great conductivity due to delocalized electrons, and they are malleable/ductile because layers can slide.
5. Properties of Giant Covalent Structures (Macromolecules)
These structures are non-metal atoms bonded together (covalently) into one huge, continuous lattice. They are sometimes called macromolecular structures. The key examples are Diamond and Graphite (both forms of Carbon), and Silicon Dioxide (Silica).
A. General Properties (Very Strong!)
Since every atom is held to its neighbour by extremely strong covalent bonds throughout the entire structure, they are incredibly hard to break.
- Property: Extremely high melting and boiling points (the highest of all structure types).
- Property: Usually very hard (e.g., diamond).
B. Conductivity (Usually Poor, but one big exception!)
Most giant covalent structures (like Diamond and Silicon Dioxide) have no free ions or delocalized electrons, meaning they do not conduct electricity.
Case Study: The Two Faces of Carbon
Diamond and Graphite are both made only of carbon atoms, but they have wildly different properties due to their structure.
Structure 1: Diamond
Each carbon atom is covalently bonded to four other carbon atoms in a rigid, 3D tetrahedral network.
- Hardness: Extremely hard. Used in cutting tools.
- Conductivity: Poor conductor of electricity (all outer electrons are used up in the strong bonds, so there are no free/delocalized electrons).
Structure 2: Graphite
Each carbon atom is covalently bonded to three other carbon atoms, forming flat, hexagonal layers. The layers are held together by weak intermolecular forces.
- Hardness: Soft and slippery. Used in pencils and as a lubricant. Why? Because the weak forces between the layers allow them to slide easily over one another.
- Conductivity: Good conductor of electricity. Why? Since carbon only bonds to three neighbours, the fourth outer electron is delocalized (free to move) between the layers, allowing charge to flow.
🔥 Memory Trick: Graphite has Layers and Random electrons (delocalized) - making it soft and conductive.
Key Takeaway for Giant Covalent Structures: They are extremely robust (high melting point). Diamond is hard and non-conductive. Graphite is soft and conductive due to its layered structure and one delocalized electron per atom.
📝 Quick Review Summary Table 📝
Use this table to quickly compare and contrast the properties—this is essential for exam success!
Type 1: Simple Molecular (e.g., H₂O, CO₂)
- Structure: Small molecules, weak forces between them.
- Melting/Boiling Point: Low.
- Conductivity: No.
- Hardness: Soft.
Type 2: Giant Ionic (e.g., NaCl)
- Structure: Lattice of alternating ions, strong forces.
- Melting/Boiling Point: High.
- Conductivity: Yes, when molten or dissolved (free ions).
- Hardness: Brittle.
Type 3: Giant Metallic (e.g., Fe, Cu)
- Structure: Positive ions in a sea of delocalized electrons.
- Melting/Boiling Point: High.
- Conductivity: Yes, in all states (delocalized electrons).
- Hardness: Malleable and Ductile.
Type 4: Giant Covalent (Diamond, Graphite)
- Structure: Massive network of atoms held by strong covalent bonds.
- Melting/Boiling Point: Extremely High.
- Conductivity: Usually No (Diamond), but Yes for Graphite (delocalized electrons in layers).
- Hardness: Very Hard (Diamond) or Soft (Graphite).
You did a great job reviewing these structures! Understanding the bonding is the key to predicting a substance's behaviour. Keep practicing!