Ionisation Energy | ICSE Class 10 Chemistry

01e⁻remove

Removing an electron

The positive nucleus attracts negative electrons. To remove one, that attraction must be overcome — so energy must be supplied.

Electron held by nucleus ↓ Energy supplied ↓ Electron removed ↓ Positive ion formed

A neutral atom has equal numbers of protons and electrons. If it loses an electron, it has more protons than electrons and becomes a positive ion, called a cation.

Ionisation energy — textbook definition

The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion.

Student-friendly meaning: ionisation energy tells us how difficult it is to remove an electron from an atom.

M(g) + ionisation energy → M⁺(g) + e⁻

The names ionisation potential (I.P.), ionisation energy (I.E.), first ionisation energy and ionisation enthalpy are used for this idea in the chapter.

Check · 1 of 18

During ionisation, energy is:

Correct: energy is supplied to overcome the attraction of the nucleus.

Write it · 2 of 18

What is the positive ion formed after electron removal called?

Answer: cation.

02(g)meaning

Read the definition precisely

Every word in the definition earns its place. A neutral atom has no overall charge; an isolated atom is considered alone, not joined to another atom; and a gaseous atom is shown by (g).

Units given in the chapter
Unit	Use
`eV/atom`	electron volt per atom
`kJ mol⁻¹`	kilojoule per mole — the SI unit stated in the chapter

Memory hook

Ionisation energy = energy to eject an electron. Both energy and eject begin with E.

Check · 3 of 18

Which equation correctly represents ionisation?

Correct: the atom is gaseous, energy is supplied, and a positive gaseous ion plus an electron are formed.

Write it · 4 of 18

State the SI unit of ionisation energy in words.

Answer: kilojoule per mole. It is written as kJ mol⁻¹.

03↔factors

What controls ionisation energy?

According to the chapter, two main factors matter: atomic size and nuclear charge.

Larger atomic size Greater nuclear charge ↓ ↓ outer electron farther away stronger attraction for electrons ↓ ↓ weaker attraction more energy needed ↓ ↓ lower ionisation energy higher ionisation energy

Think of the nucleus as a magnet. A stronger pull makes the electron harder to remove. But an electron that is farther away feels a weaker pull.

Lithium and sodium

Lithium is 2, 1; sodium is 2, 8, 1. Sodium has an extra shell, so its outer electron is farther from the nucleus. Thus sodium has lower I.E. than lithium.

Chapter values
Element	I.E. (kJ mol⁻¹)
Lithium	520
Sodium	498

Check · 5 of 18

Why does a larger atom generally have lower ionisation energy?

Correct: greater distance weakens the attraction on the outermost electron.

Tap-sort · 6 of 18

Which chain correctly explains greater nuclear charge?

Correct: greater nuclear charge holds the electron more strongly, so more energy is required.

04→period

Across a period

Ionisation energy generally increases from left to right across a period, with exceptions.

Across a period ↓ Nuclear charge increases; atomic size decreases ↓ Outermost electron held more strongly ↓ Ionisation energy generally increases

The word generally is essential. The source tables show exceptions, so do not write that the value always increases smoothly.

Period 2 ionisation energy (kJ mol⁻¹)
Li	Be*	B	C	N*	O	F	Ne
520	899	801	1088	1402	1314	1681	2080

The listed exceptions are Be 899 then B 801, and N 1402 then O 1314. Their causes are not explained in this section, so remember the exceptions without adding an unsupported explanation.

Period 3 ionisation energy (kJ mol⁻¹)
Na	Mg*	Al	Si	P*	S	Cl	Ar
496	737	577	786	1011	999	1256	1520

Here the marked exceptions are Mg 737 then Al 577, and P 1011 then S 999. Noble gases have high values: neon is 2080, argon 1520, and helium has the chapter’s highest value, 2372 kJ mol⁻¹.

Check · 7 of 18

Which is the correct exam statement about I.E. across a period?

Correct: use “generally increases” and mention exceptions.

Write it · 8 of 18

Which element has the highest ionisation energy according to the chapter?

Answer: helium. Its listed value is 2372 kJ mol⁻¹.

05↓group

Down a group

Ionisation energy decreases down a group.

Down a group ↓ New shells are added; atomic size increases ↓ Outermost electron is farther from the nucleus ↓ Electron is removed more easily ↓ Ionisation energy decreases

Although nuclear charge increases down a group, the increase in atomic size has the stronger effect in this chapter’s explanation.

Group 1 values (kJ mol⁻¹)
H	Li	Na	K	Rb	Cs
1312	520	498	419	403	375

Among the alkali metals: Li > Na > K > Rb > Cs. Caesium has many shells and a very large atomic size, so its outer electron is comparatively easy to remove.

Group 17 values (kJ mol⁻¹)
F	Cl	Br	I
1681	1256	1143	1008

Thus F > Cl > Br > I. Caesium has the lowest determined ionisation energy in the chapter, 375 kJ mol⁻¹; francium’s has not been correctly determined because it is radioactive.

Tap-sort · 9 of 18

Choose the increasing order of ionisation energy for Li, Na, K and Rb.

Correct: I.E. decreases down the group, so reverse the downward order for increasing I.E.

Write it · 10 of 18

Which element has the lowest determined ionisation energy in the chapter?

Answer: caesium. Its listed value is 375 kJ mol⁻¹.

06M/Ncharacter

Metals, non-metals and I.E.

Low ionisation energy High ionisation energy ↓ ↓ electron lost easily electron difficult to remove ↓ ↓ metallic behaviour non-metallic behaviour

Metals tend to lose electrons and form positive ions, so they usually have low ionisation energy. Non-metals hold their electrons more strongly and do not generally lose them easily, so they usually have high ionisation energy.

For a Group 1 metal: M → M⁺ + e⁻. Lower I.E. means its electron is lost more readily; metallic character and metal reactivity increase.

Check · 11 of 18

Which is expected to lose an electron more easily: Na or Cl?

Correct: low I.E. means easy electron loss, which is characteristic of metals such as sodium.

Tap-sort · 12 of 18

Which comparison is correct?

Correct: ionisation energy decreases down Group 17: F > Cl > Br > I.

07✓mastery

Worked examples and practice

Worked comparisons

1. Which has higher I.E.: lithium or sodium?

Lithium. Both are in Group 1; I.E. decreases down a group, and lithium is above sodium.

2. Which has lower I.E.: sodium or potassium?

Potassium. It lies below sodium, has a larger atomic size, and its valence electron is removed more easily.

3. Which has higher I.E.: sodium or chlorine?

Chlorine. Both are in Period 3 and I.E. generally increases from left to right.

4. Which has the highest I.E. among Li, Be, F and Ne?

Neon. They are in Period 2 and the broad trend rises towards the right.

5. Arrange Li, Na, K and Rb in increasing I.E.

Rb < K < Na < Li.

6. Arrange F, Cl, Br and I in decreasing I.E.

F > Cl > Br > I.

7. Which loses an electron more easily: Na or Cl?

Sodium. It has much lower ionisation energy than chlorine.

Check · 13 of 18

Which has higher ionisation energy: Li or Na?

Correct: I.E. decreases down Group 1, so lithium is higher than sodium.

Check · 14 of 18

Which order is decreasing ionisation energy?

Correct: I.E. decreases down Group 17.

Common mistakes to avoid

Energy must be supplied to remove an electron.

Ionisation energy is required to remove an electron. Electron addition is taken up under electron affinity in the next lesson.

A larger atom has its outermost electron farther from the nucleus, so it is removed more easily and I.E. is lower.

Extra shells increase atomic size. In the chapter’s explanation this stronger effect makes I.E. decrease down a group.

It generally increases, but the textbook tables show exceptions.

Metals usually have low I.E. because they lose electrons relatively easily.

Tap-sort · 15 of 18

Choose the best reason why I.E. decreases down a group.

Correct: atomic size increases enough to overcome the effect of increased nuclear charge.

Write it · 16 of 18

Complete: across a period, ionisation energy __________, with exceptions.

Answer: generally increases.

Concept check — answer, then open

1–4: Define I.E.; supplied or released? ion formed? units?

I.E. is the energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positive gaseous ion. Energy is supplied. A positive ion/cation is formed. Units: eV per atom and kJ per mole.

5–8: two factors; across a period; down a group; why a larger atom has lower I.E.?

Atomic size and nuclear charge. I.E. generally increases across a period, with exceptions; it decreases down a group. In a larger atom the outer electron is farther away and attracted less strongly.

9–12: Which has lower I.E., a metal or non-metal? Highest? Lowest determined? F or I?

A metal; helium; caesium; fluorine.

Give reason: why does I.E. generally increase across a period?

Across a period, nuclear charge increases and atomic size decreases. The outermost electron is therefore attracted more strongly, so more energy is required to remove it.

Give reason: why does I.E. decrease down a group?

New shells make the atom larger. The outermost electron is farther from the nucleus and is removed more easily; this atomic-size effect overcomes increased nuclear charge.

Give reason: why does Na have lower I.E. than Li; and why does F have higher I.E. than Li?

Na has one more occupied shell, so its valence electron is farther away and less strongly held. F and Li are in the same period, but F has greater nuclear charge and smaller atomic size, so it holds electrons more strongly.

Give reason: why do metals usually have low I.E. and non-metals high I.E.?

Metals lose valence electrons comparatively easily and form positive ions. Non-metals hold their electrons strongly and do not lose them easily.

Level 1: expand I.E.; definition; equation; positive ion; SI unit; factors; trends.

I.E. = ionisation energy. Definition as above. M(g) + I.E. → M⁺(g) + e⁻. Cation. kJ mol⁻¹. Atomic size and nuclear charge. Across: generally increases with exceptions; down: decreases.

Level 2: small/large atom; nuclear charge; Group 1; metals/non-metals; “generally”; Cs vs Li; F vs I.

Small atom: electron close and strongly held, so high I.E. Large atom: electron far, so low I.E. Greater nuclear charge raises I.E. Group 1 falls because extra shells increase size. Metals have lower I.E. because they lose electrons more easily. “Generally” is needed because exceptions occur. Cs is lower than Li because it has many more shells. F is higher than I because F is smaller and its electron is held more strongly.

Increasing I.E.: Li, Na, K; F, Cl, Br, I; Na, Si, Cl, Ar; Cs, Rb, K, Na.

K < Na < Li; I < Br < Cl < F; Na < Si < Cl < Ar; Cs < Rb < K < Na.

Decreasing I.E.: Li, Na, K, Rb; F, Cl, Br, I; Na, Cl, Ar; H, Li, Na, K.

Li > Na > K > Rb; F > Cl > Br > I; Ar > Cl > Na; H > Li > Na > K.

ICSE: State variation across a period and down a group, with reasons.

Across: generally increases because nuclear charge increases and atomic size decreases; exceptions occur. Down: decreases because added shells increase atomic size and the outer electron is easier to remove.

ICSE pairs: higher I.E. in Li/Na, Na/Cl, F/Br, K/Cs, Cl/Ar.

Li; Cl; F; K; Ar.

ICSE ordering: Na, Mg, Si, Cl in increasing I.E.

Na < Mg < Si < Cl, using the broad Period 3 trend.

ICSE data: 520, 498, 419, 375 kJ mol⁻¹ in one group.

The trend is a decrease down the group. 375 kJ mol⁻¹ belongs to the element lowest in the group, and that element loses its valence electron most easily.

One-minute revision

IONISATION ENERGY = energy needed to remove an electron M(g) + I.E. → M⁺(g) + e⁻ Units: eV/atom; kJ mol⁻¹ Factors: atomic size (bigger → lower I.E.); nuclear charge (greater → higher I.E.) Across a period → generally increases, with exceptions Down a group → decreases Metals → usually low I.E. Non-metals → usually high I.E.

Ionisation potential
or ionisation energy

By the end, you can