CAIE A-Level Chemistry AS 1.4 Ionisation Energy Questions

Practise definitions, gas-phase equations, periodic trends and successive-ionisation data using nuclear attraction and shielding.

Exam points

state the full first-ionisation-energy definition and construct gas-phase electron-removal equations
explain trends using nuclear charge, distance, shielding, sub-shell energy and electron pairing
use large jumps in successive values to infer valence-electron number and electronic structure

Covered syllabus points

1.4.0—Scope note 1.4.1—Terms: first ionisation energy, IE 1.4.2—Equations to represent first, second 1.4.3—The trends in ionisation energies across 1.4.4—The variation in successive ionisation 1.4.5—Ionisation energies are due to the attraction 1.4.6—Factors influencing the ionisation energies 1.4.7—The electronic configurations of elements 1.4.8—Periodic Table position from IE data

[Maximum number: 7]

(a)

Define first ionisation energy.

[ 2 ]

(b)

Successive ionisation energies for element A are shown in Table 1.1.

Table 1.1

Use Table 1.1 to deduce the group of the Periodic Table that A belongs to. Explain your answer. Group

[ 1 ]

(c)

Across Period 3 there is a general trend for first ionisation energies to increase due to the increase in attraction between the nucleus and the outer electron.

Explain why the first ionisation energy of sulfur is less than the first ionisation energy of phosphorus.

[ 2 ]

(d)

In an $\mathrm{A} l^{2+}$ ion the nuclear attraction for the outer electron is stronger than in an atom of Na .

Compare the electronic structures of $\mathrm{A} l^{2+}$ and an atom of Na and explain why the third ionisation energy of aluminium is greater than the first ionisation energy of sodium.

[ 2 ]

M1 similarity in electronic structure / shielding of $\mathrm{Al}^{2+}$ and Na both remove electron from (3)s ${ }^{1}$ / single electron in (3)s (sub-level / orbital)

OR $\mathrm{Al}^{2+}$ and Na have same electronic configuration

OR shielding (of outer electron) is the same

M2 greater nuclear charge / number of protons

$\mathrm{Al} l^{(2+)}$ has greater nuclear charge

OR 13p compared to 11p

[Maximum number: 1]

Equations involving four enthalpy changes are shown.

\begin{array}{ll} \mathrm{Na}(\mathrm{~g}) \rightarrow \mathrm{Na}^{+}(\mathrm{g})+\mathrm{e}^{-} & \Delta H=\mathrm{W} \\ \mathrm{Na}(\mathrm{~g}) \rightarrow \mathrm{Na}^{2+}(\mathrm{g})+2 \mathrm{e}^{-} & \Delta H=\mathrm{X} \\ \mathrm{Na}(\mathrm{~s}) \rightarrow \mathrm{Na}(\mathrm{~g}) & \Delta H=\mathrm{Y} \\ \mathrm{Na}(\mathrm{~s}) \rightarrow \mathrm{Na}^{2+}(\mathrm{g})+2 \mathrm{e}^{-} & \Delta H=\mathrm{Z} \end{array}

Which equation represents the second ionisation energy of sodium?

X+Y-W

X-W

Z-W

[Maximum number: 1]

This question is about the first ionisation energies of magnesium and neon.
Which row is correct?

first ionisation energy

from Mg

from Ne

$\mathrm{Mg}>\mathrm{Ne}$

$\mathrm{Ne}>\mathrm{Mg}$

[Maximum number: 2]

Define first ionisation energy.

[Maximum number: 1]

Which process has the largest enthalpy change per mole?

$\mathrm{Al} l^{3+}(\mathrm{g}) \rightarrow \mathrm{Al} l^{4+}(\mathrm{g})+\mathrm{e}^{-}$

$\mathrm{P}^{5+}(\mathrm{g}) \rightarrow \mathrm{P}^{6+}(\mathrm{g})+\mathrm{e}^{-}$

$\mathrm{S}^{6+}(\mathrm{g}) \rightarrow \mathrm{S}^{7+}(\mathrm{g})+\mathrm{e}^{-}$

$\mathrm{Si}^{4+}(\mathrm{g}) \rightarrow \mathrm{Si}^{5+}(\mathrm{g})+\mathrm{e}^{-}$

[Maximum number: 1]

The graph shows the variation of the first ionisation energy with proton number for some elements. The letters used are not the actual symbols for the elements.

Figure for Question 6 — CAIE A-Level Chemistry AS

Which statement about the elements is correct?

P and X are in the same period in the Periodic Table.

The general increase from Q to X is due to increasing atomic radius.

The small decrease from R to S is due to decreased shielding.

The small decrease from U to V is due to repulsion between paired electrons.

[Maximum number: 1]

The second ionisation energy of oxygen is greater than the second ionisation energy of fluorine.
Which factor explains this difference?

The atomic radius of an oxygen atom is smaller than that of fluorine.

The covalent bond in a fluorine molecule is weaker than the bond in an oxygen molecule.

A spin-paired electron is removed from fluorine but not from oxygen.

Fluorine has more electrons in total than oxygen. This causes a greater shielding of the nuclear attraction in fluorine.

[Maximum number: 1]

The fifth to eighth ionisation energies of four elements in Period 3 of the Periodic Table are shown.

Which row refers to chlorine?

fifth

sixth

seventh

eighth

6280

21200

25900

30500

6990

8490

27100

31700

6540

9330

11000

33600

7240

8790

12000

13800

[Maximum number: 1]

Z is a compound of two elements, X and Y .
Element X shows a very large increase between its 5th and 6th ionisation energies. It has the second largest 1st ionisation energy in its group.

Element Y shows a very large increase between its 6th and 7th ionisation energies. It has the largest 1st ionisation energy in its group.

What is compound Z ?

$\mathrm{NO}_{2}$

$\mathrm{PCl}_{5}$

$\mathrm{P}_{4} \mathrm{O}_{10}$

$\mathrm{SF}_{6}$

All question bank results loaded

Covered syllabus points

Question 1

Question (a)

Question (b)

Question (c)

Question (d)

Question 2

Question 3

Question 4

Question 5

Question 6

Question 7

Question 8

Question 9