The first two ionisation energies of calcium are 590 and 1145 kJ mol^-1. Calculate the energy to convert 0.200 mol of Ca(g) into Ca^2+(g). [2 marks]

ScotlandChemistrySyllabus dot point

Why do the properties of the elements change in regular patterns across the periodic table?

The trends in covalent radius, first ionisation energy and electronegativity across periods and down groups, explained in terms of nuclear charge, number of occupied shells and the screening effect of inner electrons.

An SQA Higher Chemistry answer on periodic trends, explaining how covalent radius, first ionisation energy and electronegativity change across periods and down groups in terms of nuclear charge, the number of occupied electron shells and the screening effect of inner electrons.

Generated by Claude Opus 4.810 min answerUpdated 2026-06-02

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this key area is asking
The three explaining factors
Across a period
Down a group
Ionisation energy as an equation
Worked example: total ionisation energy
Putting it together
Examples in context
Try this

What this key area is asking

The SQA wants you to explain, not just state, the periodic trends. You must account for the changes in covalent radius, first ionisation energy and electronegativity using three ideas: nuclear charge, the number of occupied electron shells, and the screening effect of inner electron shells. "Explain fully" questions here demand that all three factors are weighed against each other, so this key area rewards careful, complete wording.

The three explaining factors

Across a period

Going left to right, protons are added to the nucleus while electrons fill the same outer shell, so screening barely changes.

Covalent radius decreases: the increasing nuclear charge pulls the same shell of electrons closer.
First ionisation energy increases: the outer electrons are held more tightly, so more energy is needed to remove one.
Electronegativity increases: the smaller, more positively charged atom attracts bonding electrons more strongly.

Down a group

Going down a group, each element has an extra occupied shell and more screening inner electrons.

Covalent radius increases: the outer electrons occupy a higher shell, further from the nucleus.
First ionisation energy decreases: the outer electron is further away and more screened, so it is easier to remove.
Electronegativity decreases: the bonding electrons sit further from the nucleus and feel less pull.

Ionisation energy as an equation

Worked example: total ionisation energy

Worked example

The first three ionisation energies of aluminium are $577$ , $1820$ and $2740 \text{ kJ mol}^{-1}$ . Calculate the energy needed to convert $0.250 \text{ mol}$ of gaseous aluminium atoms entirely into $Al^{3+}(g)$ ions.

Step 1: Identify which ionisations are needed

Forming $Al^{3+}$ from $Al$ removes three electrons, so all three ionisation energies apply.

Step 2: Add the per-mole energies

577 + 1820 + 2740 = 5137 \text{ kJ mol}^{-1}

Step 3: Scale by the number of moles

E = 0.250 \times 5137 = 1284 \text{ kJ}

Notice the sharp jump from the first to the second ionisation energy: removing an electron from a positive ion always costs more.

Putting it together

The key insight is that the extra shells and screening down a group outweigh the rising nuclear charge, whereas across a period the rising nuclear charge dominates because no new shell is added.

Examples in context

The periodic trends are not just exam abstractions; they drive real chemistry. The high electronegativity of fluorine (top right of the table, smallest atom) is why hydrogen fluoride forms strong hydrogen bonds and why $C-F$ bonds in materials such as PTFE are so unreactive. At the other extreme, the low ionisation energy of caesium, the easiest stable element to ionise, is exploited in photoelectric cells and atomic clocks, where a tiny amount of energy frees its loosely held outer electron. The steady fall in ionisation energy down Group 1 is also why reactivity with water increases from lithium to potassium: the more easily the single outer electron is lost, the more vigorous the redox reaction.

Try this

Q1. Explain why the first ionisation energy increases across period 3 from sodium to argon. [2 marks]

Cue. The nuclear charge increases while electrons fill the same shell, so the outer electrons are held more tightly and need more energy to remove.

Q2. Account for the increase in covalent radius down Group 7. [2 marks]

Cue. Each element has an extra occupied shell and more screening, placing the outer electrons further from the nucleus.

Q3. The first two ionisation energies of calcium are $590$ and $1145 \text{ kJ mol}^{-1}$ . Calculate the energy to convert $0.200 \text{ mol}$ of $Ca(g)$ into $Ca^{2+}(g)$ . [2 marks]

Cue. $(590 + 1145) \times 0.200 = 1735 \times 0.200 = 347 \text{ kJ}$ .

Exam-style practice questions

Practice questions written in the style of SQA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

SQA Higher 20183 marksExplain fully why the first ionisation energy decreases down Group 1 from lithium to caesium. Refer to nuclear charge, the number of occupied shells and screening in your answer.

Show worked answer →

Markers want all three factors weighed against each other, not just distance.

Going down Group 1, each successive element has an additional occupied electron shell, so the outer electron is held in a shell that is further from the nucleus. There are also more inner electron shells, which screen (shield) the outer electron from the full attraction of the nucleus.

Although the nuclear charge (number of protons) increases down the group, the effects of the extra shell and the increased screening outweigh it, so the outer electron is held less tightly. Less energy is therefore needed to remove it, and the first ionisation energy decreases.

A common mark lost is quoting only "the atom is bigger" without naming the extra occupied shell and the increased screening, or forgetting to acknowledge that nuclear charge rises but is outweighed.

SQA Higher 20224 marksThe first ionisation energy of magnesium is

738 \text{ kJ mol}^{-1}

and its second ionisation energy is

1451 \text{ kJ mol}^{-1}

. (a) Write the equation, with state symbols, for the first ionisation energy of magnesium. (b) Calculate the energy required to convert

0.500 \text{ mol}

of gaseous magnesium atoms entirely into

Mg^{2+}(g)

ions.

Show worked answer →

Part (a) tests the definition as an equation; part (b) is a two-ionisation calculation.

(a) The first ionisation energy refers to one mole of gaseous atoms losing one mole of electrons:

Mg(g) \rightarrow Mg^+(g) + e^-

(b) Forming $Mg^{2+}$ requires both the first and second ionisation energies. The total energy per mole is:

738 + 1451 = 2189 \text{ kJ mol}^{-1}

For $0.500 \text{ mol}$ :

E = 0.500 \times 2189 = 1095 \text{ kJ}

Markers reward the state symbols in part (a) (a frequent loss) and adding both ionisation energies before scaling by the number of moles in part (b).

Related dot points

Sources & how we know this

SQA Higher Chemistry Course Specification — SQA (2018)