S1.3 Electron configurations - IB Chemistry Question Bank

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S1.3 Electron configurations

Practice exam-style IB Chemistry questions for Electron configurations, aligned with the syllabus and grouped by topic.

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Question 1

SL • Paper 1A

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SL • Paper 1A

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A hydrogen discharge tube emits a red line and a violet line in the visible region. How does the photon energy of the violet line compare with that of the red line?

A.

The violet photon has lower energy because it travels at a lower speed.

B.

The violet photon has higher energy because it has a higher frequency.

C.

The violet photon has lower energy because it has a shorter wavelength.

D.

The violet photon has the same energy because both lines are visible.

Question 2

SL • Paper 1A

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SL • Paper 1A

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In the hydrogen emission spectrum, the visible lines are produced by electron transitions ending at which main energy level?

A.

$n = 4$

B.

$n = 1$

C.

$n = 2$

D.

$n = 3$

Question 3

SL • Paper 1A

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SL • Paper 1A

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The maximum number of electrons that can occupy the main energy level with $n = 3$ is:

A.

B.

C.

D. Question 4

SL • Paper 1A

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SL • Paper 1A

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The diagram that best represents an s atomic orbital is:

A.

B.

C.

D.

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Question 5

HL • Paper 1A

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HL • Paper 1A

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The first ionization energy of boron is lower than that of beryllium. The best explanation is:

A.

Boron has a complete outer shell, so removal of one electron requires less energy.

B.

Boron has fewer protons than beryllium, so its outer electron is less strongly attracted.

C.

Beryllium has a paired $2p$ electron, causing greater electron-electron repulsion.

D.

The electron removed from boron is in a $2p$ sublevel, which is higher in energy than the $2s$ sublevel.

Question 6

HL • Paper 1A

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HL • Paper 1A

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The first ionization energy of sulfur is lower than that of phosphorus. The best explanation is:

A.

Sulfur has a paired electron in a $3p$ orbital, so electron-electron repulsion makes removal easier.

B.

Sulfur has a lower nuclear charge than phosphorus, so the outer electron is less attracted.

C.

Sulfur's outer electron is in the $4s$ sublevel, which is farther from the nucleus.

D.

Phosphorus has a paired electron in a $3p$ orbital, so it is harder to ionize.

Question 7

SL • Paper 2

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SL • Paper 2

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A student observes two visible emission lines from a discharge tube. Line A is red and line B is violet.

Compare the wavelength and frequency of line A with line B.

[2]

Explain why the discharge tube produces a line emission spectrum rather than a continuous spectrum.

[2]

Question 8

SL • Paper 2

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SL • Paper 2

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An atom has electrons occupying the first three main energy levels in its ground state.

Calculate the maximum number of electrons that can occupy the third main energy level.

[1]

State the sublevels available in the third main energy level.

[1]

Deduce the period of a main-group element whose highest occupied main energy level is $n = 3$ .

[1]

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Question 9

SL • Paper 1A

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SL • Paper 1A

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The correct orbital diagram for the ground-state electron configuration of oxygen is:

A.

B.

C.

D. Question 10

SL • Paper 1A

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SL • Paper 1A

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The condensed electron configuration of a neutral bromine atom, $Z = 35$ , is:

A.

[Kr] $4s^2 3d^{10} 4p^5$

B.

[Ar] $4s^2 3d^{10} 4p^5$

C.

[Ar] $4s^2 4p^5 3d^{10}$

D.

[Ar] $4s^2 3d^{10} 4p^6$

Question 11

HL • Paper 1A

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HL • Paper 1A

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The convergence limit in the emission spectrum of hydrogen occurs at $\lambda = 91.2\ \text{nm}$
for transitions to $n = 1$ . The first ionization energy calculated from this wavelength is closest to:

A.

$7.88\times 10^5\ \text{kJ mol}^{-1}$

B.

$1.31\ \text{kJ mol}^{-1}$

C.

$2.18\times 10^{-18}\ \text{kJ mol}^{-1}$

D.

$1.31\times 10^3\ \text{kJ mol}^{-1}$

Question 12

HL • Paper 1A

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HL • Paper 1A

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A graph of successive ionization energies for an unknown main-group element shows a very large increase between the third and fourth ionization energies. The group of the element is:

A.

Group 2

B.

Group 16

C.

Group 1

D.

Group 13

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Question 13

HL • Paper 1A

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HL • Paper 1A

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The frequency at the convergence limit for an atomic emission spectrum is $3.00\times 10^{15}\ \text{s}^{-1}$ . The first ionization energy is closest to:

A.

$1.99\times 10^{-18}\ \text{kJ mol}^{-1}$

B.

$4.98\times 10^{-42}\ \text{kJ mol}^{-1}$

C.

$1.20\times 10^3\ \text{kJ mol}^{-1}$

D.

$1.20\times 10^6\ \text{kJ mol}^{-1}$

Question 14

SL • Paper 2

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SL • Paper 2

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The diagram shows some electron transitions in the hydrogen atom.

State which set of transitions produces visible lines in the hydrogen emission spectrum.

[1]

Explain why the spectral lines converge at higher frequency.

[2]

State the region of the electromagnetic spectrum for transitions ending at $n = 1$ .

[1]

Question 15

SL • Paper 2

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SL • Paper 2

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The quantum mechanical model describes electrons in atomic orbitals rather than fixed circular paths.

Define the term atomic orbital.

[1]

Describe the shape of an $s$ orbital and the shape and orientations of the $p$ orbitals in a $p$ sublevel.

[2]

State the periodic table block in which an element with its highest-energy electron in a $4p$ sublevel is found.

[1]

Question 16

HL • Paper 2

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HL • Paper 2

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A line emission spectrum has a convergence limit at a frequency of $3.29 \times 10^{15}\ \text{s}^{-1}$ .

State what is represented by the convergence limit at the high-frequency end of the spectrum.

[1]

Calculate the first ionization energy in $\text{kJ mol}^{-1}$ from the frequency of the convergence limit.

[3]

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Question 17

HL • Paper 2

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HL • Paper 2

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The first ionization energy decreases down group 1 from lithium to caesium.

Explain why the first ionization energy decreases down group 1.

[2]

Suggest how this trend is related to the metallic behaviour of group 1 elements.

[1]

Question 18

SL • Paper 1B

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SL • Paper 1B

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A student observed the visible spectra of three gaseous elements using discharge tubes and a diffraction grating. The spectrum of an unknown gas was then compared with reference spectra.

Spectrum	Observed wavelengths / nm
Continuous spectrum	400–700 (all visible wavelengths present)
Reference gas A	410, 434, 486, 656
Reference gas B	405, 436, 546, 577, 579
Reference gas C	447, 502, 588, 668
Unknown gas	405, 436, 547, 578, 579

Distinguish between the continuous spectrum and the line emission spectrum shown.

[1]

Identify the reference gas that is most likely to be present in the unknown sample.

[1]

Explain why the line emission spectrum of an element can be used as a chemical fingerprint.

[2]

Question 19

SL • Paper 1B

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SL • Paper 1B

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A section of a database gives the period number and highest occupied main energy level for several main-group elements.

Element	Period number	Highest occupied main energy level, $n$
H	1	1
He	1	1
Li	2	2
C	2	2
Ne	2	2
Na	3	3
Cl	3	3
K	4	4
Br	4	4
Kr	4	4

Deduce the maximum number of electrons that can occupy the main energy level with $n = 4$ .

[1]

State the relationship shown between the period number and the highest occupied main energy level for these main-group elements.

[1]

Explain why the formula $2n^2$ does not necessarily give the number of electrons present in the highest occupied main energy level of an atom.

[1]

Question 20

SL • Paper 1B

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SL • Paper 1B

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The figures show boundary-surface representations of atomic orbitals in two main energy levels.

Identify the orbital with spherical symmetry.

[1]

State the number of orbitals in a $p$ sublevel.

[1]

Deduce the sublevel types available in the main energy level $n = 3$ .

[1]

Explain why an orbital should not be described as a fixed circular path followed by an electron.

[1]

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Question 21

HL • Paper 1A

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HL • Paper 1A

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Successive ionization energy data for a transition element show that the first two electrons are removed at lower energies than the next several electrons. This pattern is best explained by:

A.

Promotion of all $3d$ electrons to $4p$ orbitals before ionization occurs.

B.

Removal of $4s$ electrons before electrons are removed from the $3d$ sublevel.

C.

Pairing of electrons in the $4p$ sublevel before any electrons are removed.

D.

Removal of $3d$ electrons before electrons are removed from the $4s$ sublevel.

Question 22

SL • Paper 2

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SL • Paper 2

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Bromine has atomic number 35.

Write the full electron configuration of a bromine atom.

[2]

Write the condensed electron configuration of a bromide ion, $Br^-$ .

[1]

Explain why the electron configuration of $Br^-$ is the same as that of krypton.

[1]

Question 23

SL • Paper 2

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SL • Paper 2

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The orbital boxes represent the $3d$ and $4s$ sublevels for a ground-state copper atom.

Write the condensed electron configuration of a ground-state copper atom.

[1]

Draw the orbital diagram for the $4s$ and $3d$ electrons in a copper atom.

[2]

State the principle that requires paired electrons in one orbital to have opposite spins.

[1]

Question 24

HL • Paper 2

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HL • Paper 2

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The convergence limit in the emission spectrum of an element occurs at a wavelength of $242\ \text{nm}$ .

Calculate the energy, in J, of one photon at the convergence limit.

[2]

Determine the first ionization energy, in $\text{kJ mol}^{-1}$ , from this convergence limit.

[2]

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Question 25

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HL • Paper 2

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The first ionization energies of period 3 elements generally increase from sodium to argon, but there are small decreases between magnesium and aluminium and between phosphorus and sulfur.

Explain why the first ionization energy generally increases across period 3.

[2]

Explain why aluminium has a lower first ionization energy than magnesium.

[1]

Explain why sulfur has a lower first ionization energy than phosphorus.

[1]

Question 26

HL • Paper 2

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HL • Paper 2

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Successive ionization energy data are shown for an unknown main-group element in period 3.

Identify the position of the largest jump in the successive ionization energies.

[1]

Deduce the group of the element, using the data.

[2]

Explain why successive ionization energies always increase.

[1]

Question 27

SL • Paper 1B

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SL • Paper 1B

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The diagram shows some electron transitions in the hydrogen atom and the corresponding regions of the electromagnetic spectrum.

State the region of the electromagnetic spectrum for transitions that end at the first energy level.

[1]

Describe the relationship between the energy of the emitted photon and the electron transition shown in the diagram.

[2]

Explain how the convergence of lines in the hydrogen spectrum supports the model of discrete energy levels.

[2]

Question 28

SL • Paper 1B

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SL • Paper 1B

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Several arrow-in-box diagrams were proposed for the arrangement of three electrons in a set of degenerate $p$ orbitals.

Identify the diagram that correctly represents a $p^3$ arrangement.

[1]

State the rule that is violated by the diagram showing paired electrons before all three $p$ orbitals are singly occupied.

[1]

State the rule that is violated by the diagram showing two electrons with the same spin in one orbital.

[1]

Explain why the correct diagram is more stable than one with premature pairing.

[1]

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Question 29

HL • Paper 1B

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HL • Paper 1B

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First ionization energies for elements in groups 1 and 17 are plotted against period number.

Compare the first ionization energies of group 1 and group 17 elements in the same period.

[1]

Describe the trend in first ionization energy down each group.

[1]

Explain the trend down a group.

[2]

Question 30

HL • Paper 2

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HL • Paper 2

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Vanadium has the condensed electron configuration $[Ar] 4s^2 3d^3$ and forms ions with different positive charges.

Write the condensed electron configuration of $V^{2+}$ .

[1]

Explain why the first two electrons removed from vanadium are the $4s$ electrons rather than $3d$ electrons.

[1]

Explain how successive ionization energy data can support the existence of variable oxidation states in vanadium.

[2]

Question 31

SL • Paper 1B

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SL • Paper 1B

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A table gives information for some atoms and ions with atomic numbers up to $Z = 36$ .

Species	Z	Condensed e- configuration
Fe	26	[Ar] 4s $^2$ 3d $^6$
Fe $^{2+}$	26	[Ar] 3d $^6$
Cu	29	[Ar] 4s $^1$ 3d $^{10}$
Br	35	[Ar] 4s $^2$ 3d $^{10}$ 4p $^5$
Kr	36	[Ar] 4s $^2$ 3d $^{10}$ 4p $^6$

Deduce the number of electrons in $\text{Br}^-$ .

[1]

Write the condensed electron configuration of $\text{Br}^-$ .

[1]

Write the condensed electron configuration of $\text{Fe}^{2+}$ and explain which electrons are removed first.

[2]

Write the observed condensed electron configuration of a copper atom.

[1]

Question 32

HL • Paper 1B

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HL • Paper 1B

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The graph shows wavelengths of lines in the hydrogen emission spectrum that converge at the high-frequency limit. Use $h = 6.63 \times 10^{-34}\ \text{J s}$ , $c = 3.00 \times 10^8\ \text{m s}^{-1}$ and $N_A = 6.02 \times 10^{23}\ \text{mol}^{-1}$ .

State what is represented by the convergence limit in this emission spectrum.

[1]

Calculate the first ionization energy of hydrogen in $\text{kJ mol}^{-1}$ using the convergence wavelength from the graph.

[3]

Suggest why the lines become closer together near the convergence limit.

[1]

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Question 33

HL • Paper 1B

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HL • Paper 1B

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The graph shows first ionization energies for consecutive elements in period 3.

Describe the general trend in first ionization energy across period 3.

[1]

Explain the general trend across period 3.

[2]

Explain the decrease in first ionization energy from magnesium to aluminium.

[1]

Explain the decrease in first ionization energy from phosphorus to sulfur.

[1]

Question 34

HL • Paper 1B

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HL • Paper 1B

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Successive ionization energies for two unknown period 3 elements, X and Y, were obtained from a database.

State why successive ionization energies always increase for a given element.

[1]

Deduce the group of element X using the largest jump in its successive ionization energies.

[2]

Deduce the condensed electron configuration of X.

[1]

Deduce the group of element Y.

[1]

Question 35

SL • Paper 2

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SL • Paper 2

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A spectroscope is used to examine light from two lamps. Lamp A contains a hot filament. Lamp B contains a low-pressure gas excited by an electric discharge.

The two spectra are compared.

I.

Distinguish between the spectrum from lamp A and the spectrum from lamp B.

[2]

II.

Explain why the low-pressure gas in lamp B emits only certain wavelengths of visible light.

[2]

Discuss how the spectrum from lamp B could be used to identify the gas.

[3]

A student states that two lamps that appear the same colour to the eye must contain the same element. Evaluate this statement.

[1]

Question 36

SL • Paper 2

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SL • Paper 2

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The principal quantum number, $n$ , labels a main energy level in an atom. Potassium is in period 4 of the periodic table.

The maximum capacity of a main energy level is given by $2n^2$ .

I.

Deduce the maximum number of electrons in the main energy levels $n=3$ and $n=4$ .

[2]

II.

Explain why your answer for $n=4$ does not mean that a ground-state potassium atom has 32 electrons in its fourth main energy level.

[2]

An atom has the condensed electron configuration [Ne] $3s^2 3p^5$ . Deduce its period and explain your answer.

[2]

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Question 37

SL • Paper 2

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SL • Paper 2

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A model of atomic structure divides each main energy level into sublevels. Some orbital boundary surfaces are shown.

Atomic orbitals and sublevels are used to describe electron arrangements.

I.

Distinguish between an atomic orbital and a sublevel.

[2]

II.

Describe the shape and orientation of the s and p orbitals shown.

[2]

Explain how sublevels are related to the block nature of the periodic table, using period 3 as an example.

[3]

Question 38

HL • Paper 1B

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HL • Paper 1B

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A transition element M has the condensed electron configuration $[\text{Ar}]\ 4s^2\ 3d^6$ . The graph shows its first six successive ionization energies.

State which sublevel loses electrons first when M forms cations.

[1]

Write the condensed electron configuration of $\text{M}^{2+}$ .

[1]

Suggest why the increase from the second to the third ionization energy is larger than the increase from the first to the second.

[1]

Use the data to suggest why this element can form ions with more than one oxidation state.

[2]

Question 39

HL • Paper 1B

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HL • Paper 1B

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A convergence frequency was measured from the emission spectrum of a gaseous element. A database value for its first ionization energy is also shown. Use $h = 6.63 \times 10^{-34}\ \text{J s}$ and $N_A = 6.02 \times 10^{23}\ \text{mol}^{-1}$ .

Observation / unit	Value
Emission line 1 frequency / Hz	$1.01 \times 10^{15}$
Emission line 2 frequency / Hz	$1.09 \times 10^{15}$
Emission line 3 frequency / Hz	$1.15 \times 10^{15}$
Emission line 4 frequency / Hz	$1.20 \times 10^{15}$
Emission line 5 frequency / Hz	$1.22 \times 10^{15}$
Convergence frequency / Hz	$1.24 \times 10^{15}$
Database first ionization energy / kJ mol $^{-1}$	496

Calculate the first ionization energy from the convergence frequency.

[2]

Calculate the percentage difference between the experimental value and the database value.

[1]

Evaluate whether the experimental spectrum supports the database value.

[2]

Question 40

SL • Paper 2

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SL • Paper 2

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The visible line emission spectrum of hydrogen contains a red line and several blue-violet lines. Other hydrogen lines occur outside the visible region.

The line emission spectrum provides evidence about electron energy levels in hydrogen.

I.

Explain why the existence of separate lines supports the idea of discrete energy levels.

[2]

II.

Explain why the lines in a hydrogen emission series converge at higher frequency.

[2]

Compare the photons emitted by transitions ending at the first, second and third main energy levels of hydrogen.

[3]

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Question 41

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SL • Paper 2

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Vanadium forms several ions in compounds. Selenium is a non-metal in period 4.

Electron configurations can be written in full or condensed form.

I.

Write the full electron configuration of a ground-state vanadium atom, $Z=23$ .

[2]

II.

Write condensed electron configurations for $\text{V}^{2+}$ and $\text{Se}^{2-}$ .

[2]

III.

Explain why the configuration of $\text{V}^{2+}$ is not [Ar] $4s^2 3d^1$ .

[1]

Complete an orbital diagram for the outer electrons of a ground-state vanadium atom and explain how the diagram follows the Aufbau principle, Hund's rule and the Pauli exclusion principle.

[3]

Question 42

SL • Paper 2

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SL • Paper 2

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Chromium and copper are exceptions to the simple Aufbau filling pattern for ground-state atoms.

The expected and observed configurations of chromium and copper are compared.

I.

Write the observed condensed electron configurations of chromium and copper atoms.

[2]

II.

Explain why these observed configurations are more stable than the expected configurations [Ar] $4s^2 3d^4$ and [Ar] $4s^2 3d^9$ .

[2]

Copper can form both $\text{Cu}^+$ and $\text{Cu}^{2+}$ ions.

I.

Write condensed electron configurations for $\text{Cu}^+$ and $\text{Cu}^{2+}$ .

[2]

II.

Evaluate the statement: $\text{Cu}^{2+}$ is formed by removing two $3d$ electrons because $3d$ is written after $4s$ in the configuration.

[1]

Question 43

HL • Paper 2

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HL • Paper 2

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The convergence limit for a series in the hydrogen emission spectrum has a wavelength of $91.2\ \text{nm}$ in the ultraviolet region.

The convergence limit is used to calculate the first ionization energy of hydrogen.

I.

Explain what the convergence limit represents in terms of the electron in a hydrogen atom.

[1]

II.

Calculate the energy, in J, of one photon at this convergence wavelength.

[2]

III.

Calculate the first ionization energy of hydrogen in $\text{kJ mol}^{-1}$ .

[2]

Discuss why the lines before the convergence limit correspond to excitation within the atom rather than ionization.

[3]

Question 44

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HL • Paper 2

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A graph of first ionization energy against atomic number is obtained for the elements in period 2.

The graph contains both an overall trend and two discontinuities.

I.

Explain the general trend in first ionization energy across period 2.

[2]

II.

Explain the decrease in first ionization energy from beryllium to boron.

[2]

III.

Explain the decrease in first ionization energy from nitrogen to oxygen.

[2]

Evaluate the statement: first ionization energy increases smoothly with atomic number across every period.

[2]

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Question 45

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HL • Paper 2

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The successive ionization energies of an unknown period 3 element, E, are shown.

Use the successive ionization energy data to deduce information about element E.

I.

Identify where the largest jump occurs and deduce the group of E.

[2]

II.

Deduce the valence electron configuration of E.

[1]

III.

Suggest the identity of E.

[1]

Explain why successive ionization energies increase and why a very large jump occurs in this dataset.

[2]

Evaluate the usefulness of plotting these data on a logarithmic scale.

[1]

Question 46

HL • Paper 2

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HL • Paper 2

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The convergence frequencies for emission series that correspond to first ionization are $1.24\times10^{15}\ \text{s}^{-1}$ for sodium and $1.04\times10^{15}\ \text{s}^{-1}$ for potassium.

Use the convergence frequency data to calculate first ionization energies.

I.

Calculate the first ionization energy of sodium in $\text{kJ mol}^{-1}$ .

[2]

II.

Calculate the first ionization energy of potassium in $\text{kJ mol}^{-1}$ .

[2]

III.

Deduce which metal has the lower first ionization energy and relate this to its convergence frequency.

[1]

Discuss why first ionization energy decreases from sodium to potassium.

[3]

Question 47

HL • Paper 2

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HL • Paper 2

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Iron is a transition element. A plot of successive ionization energies for iron shows that several electrons can be removed before a very large jump associated with the argon core.

Electron configurations help interpret successive ionization energy data for iron.

I.

Write condensed electron configurations for $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$ .

[2]

II.

Explain the order in which electrons are removed when iron forms these ions.

[2]

Discuss how the successive ionization energy pattern helps explain variable oxidation states in iron.

[3]

Compare the successive ionization energy pattern of iron with that expected for a main-group element with two valence electrons.

[1]

Question 48

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HL • Paper 2

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A database is used to plot $\log_{10}(IE)$ for successive ionization energies of two period 3 elements, A and B.

Use the logarithmic successive ionization energy plot to identify the electron arrangements of A and B.

I.

Deduce the groups of elements A and B.

[2]

II.

Write condensed electron configurations for A and B.

[2]

III.

Explain why the major jump occurs after the number of electrons identified in (a)(i).

[1]

Evaluate the use of only the first ionization energy, rather than successive ionization energies, to assign an element to a group.

[2]

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S1.2 The atom

S1.4 Counting particles, mass. The mole

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