# Bonding and physical properties

> CIE A-Level Chemistry · 9701
> Source: https://www.owlsprep.com/study/cie-9701-u3-bonding-and-physical-properties/

This core subtopic connects different types of chemical bonding and crystal structure to observable physical properties including melting point, boiling point, solubility and electrical conductivity, and is regularly tested in both multiple choice and written papers.

**Prerequisites:** [Types of chemical bonding](https://www.owlsprep.com/study/cie-9701-u3-types-of-chemical-bonding/); [Intermolecular forces](https://www.owlsprep.com/study/cie-9701-u3-intermolecular-forces/)

## Learning objectives

- Relate structure and bonding type to key physical properties of substances
- Predict melting and boiling points based on bonding and structure
- Explain trends in physical properties for homologous series
- Distinguish between effects of intermolecular and intramolecular forces on properties

## Classification of Structures by Bonding

**Four main solid structure classes** — All solid substances are classified into one of four groups based on their bonding and lattice arrangement, each with distinct physical properties.

| Structure Type | Bonding within lattice | Forces between particles |
| --- | --- | --- |
| Giant ionic | Ionic (electrostatic between ions) | Strong ionic bonds throughout the lattice |
| Simple molecular | Covalent (within discrete molecules) | Weak intermolecular forces between molecules |
| Giant covalent (macromolecular) | Covalent | Strong covalent bonds throughout the entire lattice |
| Giant metallic | Metallic (between cations and delocalized electrons) | Strong metallic bonds throughout the lattice |

**Worked example:** Classify each of the following substances by their structure type: (a) magnesium oxide (b) diamond (c) iodine (d) copper

1. Magnesium oxide is formed from a metal cation ($Mg^{2+}$) and non-metal anion ($O^{2-}$), so it forms a **giant ionic lattice**.
2. Diamond is an allotrope of carbon where every carbon atom is covalently bonded to 4 other carbon atoms extending through the structure, so it is a **giant covalent structure**.
3. Iodine is made of discrete $I_2$ molecules held together by weak London dispersion forces, so it is a **simple molecular structure**.
4. Copper is a metal, so it forms a **giant metallic lattice**.

> **Exam tip:** Always state both the structure type and the type of bonding/force when answering explanation questions to gain full marks.

## Melting and Boiling Point Trends

Melting and boiling point depend on the amount of energy required to overcome the attractive forces holding particles together in the solid/liquid lattice. Stronger forces require more energy, leading to higher melting and boiling points.

**Melting Point** — The temperature at which the solid and liquid phases of a substance are in equilibrium, when sufficient energy is available to overcome forces holding the solid lattice together.

> **warning**
>
> Covalent bonds do NOT break when simple molecular substances melt or boil. Only weak intermolecular forces between molecules are overcome.

**Worked example:** Explain why sodium chloride has a much higher melting point than solid chlorine.

1. First classify the structure of each substance: Sodium chloride is a giant ionic lattice, solid chlorine ($Cl_2$) is simple molecular.
2. In sodium chloride, strong ionic bonds (electrostatic attractions between oppositely charged ions) extend throughout the entire lattice. A large amount of energy is required to break these strong bonds.
3. In solid chlorine, only weak London dispersion forces between discrete $Cl_2$ molecules need to be overcome for melting. Covalent bonds within the $Cl_2$ molecules remain intact.
4. The much stronger forces in sodium chloride result in a far higher melting point.

## Electrical Conductivity and Solubility

A substance can only conduct electricity if it contains charged particles that are free to move through the structure. Solubility follows the rule 'like dissolves like': polar/ionic solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents.

**Worked example:** Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does.

1. In solid sodium chloride, $Na^+$ and $Cl^-$ ions are held in a fixed, regular lattice by strong ionic bonds. The charged ions are not free to move, so the solid cannot conduct electricity.
2. When sodium chloride is melted, the ionic lattice breaks down, and the $Na^+$ and $Cl^-$ ions become mobile.
3. Mobile charged particles are now present, so molten sodium chloride can conduct electricity.

**Check your understanding**

Test your understanding of solubility:

1. Why is solid iodine insoluble in water?

   - Iodine is ionic and water is polar
   - Iodine is non-polar and cannot form strong interactions with polar water
   - The covalent bonds in $I_2$ are too strong to break
   - Water has a higher boiling point than iodine

   *Answer:* Iodine is non-polar and cannot form strong interactions with polar water

   *Why:* Correct! Like dissolves like: non-polar iodine only dissolves well in non-polar solvents like hexane, not polar water.

## Trends in Homologous Series

For homologous series of simple molecular compounds (e.g. alkanes), boiling point increases with increasing carbon chain length. This is because larger molecules have greater surface area and stronger London dispersion forces between molecules.

**Worked example:** Explain why the boiling point of pentane ($C_5H_{12}$) is higher than that of ethane ($C_2H_6$).

1. Both pentane and ethane are simple molecular alkanes, so only intermolecular London dispersion forces need to be overcome for boiling.
2. Pentane has a larger molecular mass and greater molecular surface area than ethane, so the London dispersion forces between pentane molecules are stronger.
3. More thermal energy is required to overcome the stronger intermolecular forces in pentane, so it has a higher boiling point than ethane.

**Exam command terms**

Common command terms for this topic in CIE exams have specific expectations:

- **Explain** — You must explicitly link structure and bonding to the physical property, not just state the trend *(Explain why diamond has a high melting point requires linking strong covalent bonds to energy needed)*

- **Predict** — Use your knowledge of bonding to deduce the property of an unfamiliar compound *(Predict whether silicon dioxide conducts electricity requires recalling its giant covalent structure with no free electrons)*

## Common pitfalls

- **Wrong:** Saying covalent bonds break when simple molecular substances melt/boil
  - Why it fails: Only intermolecular forces between molecules are overcome; covalent bonds within molecules remain intact
  - Correct: State that weak intermolecular forces between molecules are overcome, and covalent bonds do not break
- **Wrong:** Claiming all giant covalent substances do not conduct electricity
  - Why it fails: Some giant covalent structures have free delocalized electrons
  - Correct: State that graphite and graphene conduct electricity, but diamond and silicon dioxide do not
- **Wrong:** Saying solid ionic compounds conduct electricity because they contain ions
  - Why it fails: Ions in solid ionic lattices are fixed in place and not mobile
  - Correct: State that only molten or dissolved ionic compounds conduct electricity, because their ions are free to move
- **Wrong:** Attributing higher boiling points of alcohols to stronger covalent bonds compared to alkanes
  - Why it fails: This confuses intramolecular covalent bonds with intermolecular forces
  - Correct: Explain that alcohols form hydrogen bonds between molecules, which are stronger than London dispersion forces in alkanes of similar mass
- **Wrong:** Claiming all metals have high melting points
  - Why it fails: Some metals like mercury and group 1 metals have weak metallic bonding and low melting points
  - Correct: State that most metals have high melting points, but note exceptions when answering questions

## Cheatsheet

| Structure Type | Typical mp/bp | Electrical Conductivity | Solubility in Water |
| --- | --- | --- | --- |
| Giant Ionic Lattice | High | Solid: No; Molten/Aqueous: Yes | Usually soluble |
| Simple Molecular | Low | No (unless ionises in water) | Soluble only if polar/H-bonding possible |
| Giant Covalent | Very High | No (except graphite/graphene) | Insoluble |
| Giant Metallic | Mostly high (Hg: liquid RT) | Yes (solid and molten) | Insoluble (usually react instead of dissolving) |

## What's next

Understanding the link between bonding and physical properties is foundational for almost all other topics in A-Level Chemistry. In organic chemistry, you will apply this knowledge to explain trends in boiling points of different functional groups and predict solubility of organic compounds. In inorganic chemistry, it helps explain periodic trends in melting and boiling points of Period 3 elements and their compounds. This topic also underpins further study of crystal structures, and the energetics of bond breaking and forming in thermodynamics. Mastery of this subtopic is essential to score full marks on common explanation questions in both Paper 1 and Paper 2 of CIE 9701.

- [States of matter](https://www.owlsprep.com/study/cie-9701-u4-overview/)
- [Gases and the ideal gas equation](https://www.owlsprep.com/study/cie-9701-u4-gases-and-ideal-gas-equation/)
- [Liquids and Solids](https://www.owlsprep.com/study/cie-9701-u4-liquids-and-solids/)

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