# Ionisation energies

> CIE A-Level Chemistry · Unit 2: Atomic structure
> Source: https://www.owlsprep.com/study/cie-9701-u2-ionisation-energies/

This sub-topic covers definitions of first and successive ionisation energies, factors influencing their magnitude, trends down groups and across periods, and how to deduce electron configurations from ionisation energy data.

**Prerequisites:** [Electronic configuration and sub-shell structure](https://www.owlsprep.com/study/cie-9701-u2-electron-configurations/)

## Learning objectives

- Define first and successive ionisation energies correctly
- Explain how four key factors influence ionisation energy values
- Describe and explain trends in ionisation energies across periods and down groups
- Deduce electron configurations and group number from successive ionisation energy data

## Definitions and Key Factors

**First Ionisation Energy** — The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

*Notation:* I_1

*Example:* For sodium: $\text{Na}(g) \rightarrow \text{Na}^+(g) + e^- \quad \Delta H = I_1$

**Successive Ionisation Energies** — A series of ionisation energies corresponding to the sequential removal of each electron from a gaseous atom/ion after the first electron is removed.

*Example:* Second ionisation energy of sodium: $\text{Na}^+(g) \rightarrow \text{Na}^{2+}(g) + e^- \quad \Delta H = I_2$

Four main factors affect the magnitude of any ionisation energy:
1. **Nuclear charge**: More protons = stronger attraction for outer electrons, increasing ionisation energy.
2. **Atomic radius**: Larger distance between nucleus and outer electron = weaker attraction, decreasing ionisation energy.
3. **Shielding**: Inner electrons repel outer electrons, reducing net nuclear attraction, decreasing ionisation energy.
4. **Sub-shell stability**: Electrons in full or half-full sub-shells have extra stability, increasing ionisation energy.

**Worked example:** Write the balanced equation for the third ionisation energy of aluminium.

1. Recall that the nth ionisation energy removes one mole of electrons from the (n-1)+ gaseous ion. For the third ionisation energy, we start with Al²⁺ (gaseous).
2. Write the full balanced equation:
3. $$\text{Al}^{2+}(g) \rightarrow \text{Al}^{3+}(g) + e^-$$

> **Exam tip:** Always include (g) state symbols for all species in ionisation energy equations — examiners regularly penalise missing gaseous states.

## Trends in First Ionisation Energies

Two key periodic trends are observed for first ionisation energy:
* **Down a group**: Each new element has an extra inner electron shell, so shielding and atomic radius increase. These effects outweigh the increase in nuclear charge, so first ionisation energy *decreases* down a group.
* **Across a period**: Nuclear charge increases, while electrons are added to the same outer shell so shielding stays almost constant. Attraction increases, atomic radius decreases, so first ionisation energy *generally increases* across a period. Small dips in the trend occur between Group 2→13 and Group 15→16.

**Worked example:** Explain why the first ionisation energy of chlorine is higher than that of sulfur.

1. Both elements are in Period 3, so their outer electrons are in the n=3 shell, meaning shielding is identical for both.
2. Chlorine has one more proton than sulfur, so its nuclear charge is higher.
3. Higher nuclear charge creates a stronger attraction between the nucleus and outer electrons.
4. More energy is required to remove an electron, so first ionisation energy of chlorine is higher.

> **tip**
>
> Remember the two common exceptions across a period: the drop in ionisation energy from Group 2 to 13 (3s to 3p) and from Group 15 to 16 (paired electron in p orbital has extra repulsion).

## Successive Ionisation Energies and Electron Configuration

Large *proportional jumps* in successive ionisation energies indicate that the next electron is being removed from a new inner electron shell, which is much closer to the nucleus. This pattern allows us to deduce the number of outer electrons, and therefore the group number of the unknown element.

**Worked example:** The successive ionisation energies (kJ mol⁻¹) for an element are: 738, 1451, 7733, 10540, 13630. Deduce the group of this element.

1. Look for the largest proportional jump between successive values. The jump between the second (1451) and third (7733) ionisation energy is a 5x increase, much larger than any other change.
2. This large jump confirms that after removing two electrons, the next electron is removed from an inner shell closer to the nucleus.
3. This means there are two outer electrons, so the element belongs to Group 2 of the periodic table.

> **info**
>
> Always look for proportional change, not absolute change. A jump from 1500 to 7700 is far more significant than a jump from 10500 to 13600, which is only a 30% increase.

## Exam Phrasing

**Exam command terms**

Common command terms for this topic have specific marking expectations in CIE 9701:

- **Explain the trend** — You must reference nuclear charge, atomic radius and shielding to earn full marks *(Explaining the trend down Group 2)*

- **Deduce the group** — You must identify the large jump and link it to the number of outer electrons *(Using a table of successive ionisation energies)*

**Check your understanding**

Test your understanding:

1. Which of the following describes where the first large jump occurs for a Group 15 element?

   - After 3 electrons
   - After 5 electrons
   - After 15 electrons
   - After 2 electrons

   *Why:* Correct: A Group 15 element has 5 outer electrons, so the big jump comes after all 5 outer electrons are removed.

## Common pitfalls

- **Wrong:** Forgetting to include (g) state symbols in ionisation energy equations
  - Why it fails: CIE examiners always penalise missing state symbols, as ionisation energy is only defined for gaseous species
  - Correct: Always write (g) after every species in any ionisation energy equation
- **Wrong:** Explaining decreasing IE down a group by only mentioning increasing nuclear charge
  - Why it fails: Nuclear charge does increase, but the effects of increased shielding and atomic radius are larger and dominate the trend
  - Correct: Always state that increased shielding and atomic radius outweigh the increase in nuclear charge when explaining trends down a group
- **Wrong:** Removing inner shell electrons first when analysing successive ionisation energies
  - Why it fails: Electrons are always removed from the highest energy (outermost) shell first
  - Correct: Outer electrons are removed first, so large jumps occur when moving to an inner shell, after all outer electrons are removed
- **Wrong:** Attributing the IE drop between Mg and Al to increased nuclear charge of Al
  - Why it fails: While nuclear charge is higher, the key factor is the energy level of the outer electron
  - Correct: Explain the drop by noting Al's outer electron is in the higher energy 3p sub-shell, further from the nucleus, so less energy is required for removal

## Cheatsheet

| Concept | Key Point | Trend Summary |
| --- | --- | --- |
| First IE | Energy to remove 1 mol e⁻ from gaseous atoms |  |
| Successive IE | Always increase; big jump = new inner shell |  |
| Down a group | I₁ decreases | Shielding ↑, radius ↑ > nuclear charge ↑ |
| Across a period | I₁ generally increases | Nuclear charge ↑, shielding constant, radius ↓ |
| Exception 1: Group 2 → 13 | I₁ drops | Outer e⁻ in higher energy p sub-shell |
| Exception 2: Group 15 → 16 | I₁ drops | Paired p electron has extra repulsion |

## What's next

Ionisation energies are a core foundational concept for CIE A-Level Chemistry, heavily tested in both multiple choice and structured questions. Understanding how ionisation energy trends work will help you explain other periodic properties like electronegativity, atomic radius, and reactivity, and supports the model of electron sub-shell structure that underpins all of chemical bonding and reactivity. This topic connects directly to the study of periodicity as a whole, which is a major topic in the second half of your AS Level course. Mastering ionisation energy explanations now will help you access full marks on many common exam questions later.

- [Chemical bonding](https://www.owlsprep.com/study/cie-9701-u3-overview/)
- [Covalent and coordinate bonding](https://www.owlsprep.com/study/cie-9701-u3-covalent-and-coordinate-bonding/)
- [Metallic bonding](https://www.owlsprep.com/study/cie-9701-u3-metallic-bonding/)

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