S1.3 Electron configurations
Practice exam-style IB Chemistry questions for Electron configurations, aligned with the syllabus and grouped by topic.
Question 1
What change occurs in photon properties from red light to violet light in the visible spectrum?
A.
Wavelength decreases and energy increases.
B.
Frequency increases and energy decreases.
C.
Wavelength increases and frequency decreases.
D.
Frequency decreases and energy increases.
Question 2
What is the maximum number of electrons in the main energy level with principal quantum number n = 3?
A.
18
B.
8
C.
6
D.
32
Question 3
What distinguishes a line emission spectrum from a continuous spectrum?
A.
It is produced only when white light passes through a prism.
B.
It contains only particular wavelengths emitted by excited atoms.
C.
It always contains wavelengths only in the ultraviolet region.
D.
It contains every visible wavelength without any gaps.
Question 4
What is the shape of an s orbital?
A.
Tetrahedral
B.
Dumbbell-shaped
C.
Trigonal planar
D.
Spherical
Question 5
How many orbitals and how many electrons can a p sublevel contain when full?
A.
3 orbitals and 6 electrons
B.
1 orbital and 2 electrons
C.
7 orbitals and 14 electrons
D.
5 orbitals and 10 electrons
Question 6
What does the convergence limit in an atomic emission spectrum correspond to?
A.
Ionization of the atom from the relevant energy level
B.
Excitation of an electron to the first energy level
C.
Emission of a photon with zero energy
D.
Pairing of electrons in degenerate orbitals
Question 7
Why are successive ionization energies always larger than the preceding ionization energy for the same element?
A.
Each removed electron must come from a lower main energy level.
B.
The same nucleus attracts fewer remaining electrons more strongly after each removal.
C.
The nuclear charge increases after each electron is removed.
D.
The atom becomes larger after each electron is removed.
Question 8
Define a photon.
[1]
Distinguish between a continuous spectrum and a line emission spectrum.
[1]
Question 9
Calculate the maximum number of electrons in the main energy level n = 4.
[1]
State the period number of a main-group element whose highest occupied main energy level is n = 4.
[1]
Question 10
What is the condensed electron configuration of sulfur, Z = 16?
A.
[Ne] 3s² 3p⁶
B.
[Ne] 3s² 3p⁴
C.
[Ar] 3s² 3p⁴
D.
[Ne] 3s¹ 3p⁵
Question 11
An orbital diagram for a p² sublevel shows two electrons paired in one p orbital while the other two p orbitals are empty. Which rule is not followed?
A.
Hund’s rule
B.
Pauli exclusion principle
C.
The 2n² rule
D.
Conservation of charge
Question 12
The convergence frequency for an atom is 1.20 × 10¹⁵ s⁻¹. Which expression gives the first ionization energy in kJ mol⁻¹?
A.
cNₐ/[h(1.20 × 10¹⁵)]
B.
h(1.20 × 10¹⁵)Nₐ/1000
C.
hNₐ/[1000(1.20 × 10¹⁵)]
D.
1000h(1.20 × 10¹⁵)/Nₐ
Question 13
Why does first ionization energy generally increase across a period?
A.
Outer electrons enter higher main energy levels.
B.
The number of occupied main energy levels increases.
C.
Atomic radius increases and shielding decreases.
D.
Nuclear charge increases while shielding changes little.
Question 14
Why is the first ionization energy of aluminium lower than that of magnesium?
A.
Aluminium has a lower principal quantum number for its outer electron.
B.
Aluminium has fewer protons than magnesium.
C.
The electron removed from magnesium is paired in a p orbital.
D.
The electron removed from aluminium is in a higher-energy 3p sublevel.
Question 15
A main-group element has a very large jump between its third and fourth successive ionization energies. What is its group?
A.
Group 18
B.
Group 2
C.
Group 13
D.
Group 15
Question 16
Hydrogen gas in a discharge tube emits several visible lines.
State what happens to an electron in a hydrogen atom when a photon is emitted.
[1]
Explain why the visible lines have different colours.
[1]
Question 17
Describe the shape of an s orbital.
[1]
Describe the shape and orientations of the three p orbitals in a p sublevel.
[1]
Question 18
Phosphorus has atomic number 15.
Write the full electron configuration of a phosphorus atom.
[1]
State the number of unpaired electrons in the 3p sublevel of phosphorus.
[1]
Name the rule used in part (b).
[1]
Question 19
The figure shows the visible spectra from two light sources.
Identify which spectrum is continuous.
[1]
State one feature of the other spectrum that shows it is a line spectrum.
[1]
Suggest why the line spectrum can be used to identify an element.
[1]
Question 20
The figure shows an orbital diagram for a neutral atom up to the 3p sublevel.
Deduce the electron configuration represented.
[1]
Deduce the period of the element.
[1]
Identify the rule shown by the single occupation of degenerate p orbitals before pairing.
[1]
Question 21
A student measured the wavelengths of visible emission lines from an unknown gaseous element and compared them with reference values.
| Line | Unknown / nm | Hydrogen ref. / nm | Helium ref. / nm | Neon ref. / nm |
|---|---|---|---|---|
| 1 | 410.3 | 410.2 | 447.1 | 540.1 |
| 2 | 434.1 | 434.0 | 471.3 | 585.2 |
| 3 | 486.2 | 486.1 | 501.6 | 640.2 |
| 4 | 656.4 | 656.3 | 587.6 | 703.2 |
Identify the reference element that best matches the unknown.
[1]
State one quantitative measurement collected by the instrument.
[1]
Suggest one reason why using a spectroscope is more reliable than observing the colour of the discharge tube by eye.
[1]
Question 22
What is the ground-state electron configuration of a copper atom?
A.
[Ar] 4s² 3d⁹
B.
[Ar] 4s² 3d¹⁰
C.
[Ar] 4s¹ 3d⁸
D.
[Ar] 4s¹ 3d¹⁰
Question 23
Why is the first ionization energy of oxygen lower than that of nitrogen?
A.
Oxygen has one paired 2p orbital, causing extra electron–electron repulsion.
B.
Nitrogen’s outer electron is in a higher main energy level.
C.
Oxygen’s first electron is removed from a 2s sublevel.
D.
Oxygen has a lower nuclear charge than nitrogen.
Question 24
What electron is removed first when Fe forms Fe²⁺ from Fe, [Ar] 4s² 3d⁶?
A.
A 4s electron, because 4s electrons are removed before 3d electrons
B.
A 4p electron, because it is highest in energy
C.
A 3d electron, because 3d fills after 4s
D.
A 3p electron, because it is in the noble-gas core
Question 25
Write the condensed electron configuration of Br⁻.
[1]
Explain why this ion has the same electron configuration as a noble gas.
[1]
Question 26
An orbital diagram for the 2p³ sublevel is shown with one box containing ↑↓, a second box containing ↑ and the third box empty.
Identify the rule not followed.
[1]
Draw or describe the correct arrangement of the three 2p electrons.
[1]
Explain why the correct arrangement is more stable.
[1]
Question 27
Write the condensed electron configuration of a chromium atom.
[1]
State why chromium is an exception to the simple Aufbau filling order.
[1]
Write the condensed electron configuration of Cr³⁺.
[1]
Question 28
The convergence limit for a series in the emission spectrum of gaseous atoms occurs at a frequency of 2.06 × 10¹⁵ s⁻¹.
Calculate the energy, in J, of one photon at this frequency.
[1]
Calculate the corresponding first ionization energy, in kJ mol⁻¹.
[1]
State the meaning of the convergence limit.
[1]
Question 29
State the equation for the first ionization energy of an element X, including state symbols.
[1]
Explain why first ionization energy generally decreases down a group.
[1]
Question 30
Successive ionization energies for an unknown main-group element show a very large jump after the second electron is removed.
Deduce the group of the element.
[1]
Explain your answer.
[1]
Question 31
A spreadsheet database of successive ionization energies is used to compare elements.
State one advantage of using a database for this task.
[1]
Suggest why plotting log(successive ionization energy) rather than successive ionization energy may make trends easier to compare.
[1]
Question 32
State how low first ionization energy is related to metallic behaviour.
[1]
Explain, using ionization energy, why metallic character decreases across a period.
[1]
Question 33
The diagram shows selected electron transitions in a hydrogen atom and the spectral regions in which their emitted photons occur.
Identify the final energy level for the transitions producing visible light.
[1]
State which set of transitions produces the highest-energy photons.
[1]
Explain why lines in a series converge at higher frequency.
[1]
State the relationship between photon energy and frequency.
[1]
Question 34
A table compares proposed electron configurations for several species with Z ≤ 36.
| Species | Proposed configuration | Outer orbital diagram |
|---|---|---|
| Mg | [Ne] 3s² | 3s [↑↓] |
| N | [He] 2s² 2p³ | 2s [↑↑] 2p [↑] [↑] [↑] |
| K | [Ar] 3d¹ | 4s [ ] 3d [↑] [ ] [ ] [ ] [ ] |
| Cl | [Ne] 3s² 3p⁵ | 3s [↑↓] 3p [↑↓] [↑↓] [↑] |
| Cu | [Ar] 4s¹ 3d¹⁰ | 4s [↑] 3d [↑↓] [↑↓] [↑↓] [↑↓] [↑↓] |
Identify one configuration that violates the Aufbau principle.
[1]
Identify one configuration that violates the Pauli exclusion principle.
[1]
Explain why the configuration of Cu shown as [Ar] 4s¹ 3d¹⁰ is acceptable.
[1]
Question 35
The graph shows first ionization energies for consecutive elements in period 3.
Describe the general trend shown.
[1]
Identify one discontinuity in the trend.
[1]
Explain the discontinuity between magnesium and aluminium.
[1]
Question 36
The bar chart shows the first six successive ionization energies for an unknown main-group element.
Identify where the largest jump occurs.
[1]
Deduce the number of valence electrons in the atom.
[1]
Deduce the group of the element.
[1]
Explain why the ionization energies increase successively.
[1]
Question 37
First ionization energies do not increase smoothly across period 2.
Explain why the first ionization energy of boron is lower than that of beryllium.
[1]
Explain why the first ionization energy of oxygen is lower than that of nitrogen.
[1]
Question 38
Vanadium has the electron configuration [Ar] 4s² 3d³.
Write the electron configuration of V²⁺.
[1]
Write the electron configuration of V³⁺.
[1]
Explain why these ions can form by successive ionization.
[1]
Question 39
The convergence wavelength for a spectral series is 91.2 nm.
Convert this wavelength to metres.
[1]
Calculate the energy of one photon at this wavelength.
[1]
State how this photon energy is converted to an ionization energy in kJ mol⁻¹.
[1]
Question 40
The graph shows spectral lines in a series plotted as frequency against line number. The convergence limit is indicated by extrapolation.
Read the convergence frequency from the graph.
[1]
Calculate the energy of one photon at the convergence frequency.
[1]
Calculate the first ionization energy in kJ mol⁻¹.
[1]
Explain why this limit corresponds to ionization.
[1]
Question 41
A database was used to plot successive ionization energies for two unknown elements, once on a linear scale and once on a logarithmic scale.
| Electron removed | X IE / kJ mol^-1 | Y IE / kJ mol^-1 | log10(X IE) | log10(Y IE) |
|---|---|---|---|---|
| 1 | 738 | 1012 | 2.868 | 3.005 |
| 2 | 1451 | 1907 | 3.162 | 3.280 |
| 3 | 7733 | 2914 | 3.888 | 3.464 |
| 4 | 10540 | 4964 | 4.023 | 3.696 |
| 5 | 13630 | 6274 | 4.135 | 3.798 |
| 6 | 17995 | 21270 | 4.255 | 4.328 |
| 7 | 21700 | 25430 | 4.336 | 4.405 |
| 8 | 25660 | 29870 | 4.409 | 4.475 |
State why a database is useful for constructing these plots.
[1]
Compare the visibility of the major jump on the two plots.
[1]
Deduce which unknown has more valence electrons.
[1]
Suggest one limitation of using only the first few successive ionization energies to identify an element.
[1]
Question 42
The table gives successive ionization energies for a transition element and lists possible electron configurations for some of its ions.
| Section | Entry | Value |
|---|---|---|
| Neutral atom | M | [Ar] 4s² 3d⁶ |
| Ionization energy | IE₁ / kJ mol⁻¹ | 763 |
| Ionization energy | IE₂ / kJ mol⁻¹ | 1562 |
| Ionization energy | IE₃ / kJ mol⁻¹ | 2957 |
| Ionization energy | IE₄ / kJ mol⁻¹ | 5290 |
| Ionization energy | IE₅ / kJ mol⁻¹ | 7240 |
| Ionization energy | IE₆ / kJ mol⁻¹ | 9560 |
| Ionization energy | IE₇ / kJ mol⁻¹ | 12060 |
| Ionization energy | IE₈ / kJ mol⁻¹ | 14580 |
| Ionization energy | IE₉ / kJ mol⁻¹ | 22540 |
| M²⁺ option | A | [Ar] 4s² 3d⁴ |
| M²⁺ option | B | [Ar] 3d⁶ |
| M²⁺ option | C | [Ar] 4s¹ 3d⁵ |
| M³⁺ option | D | [Ar] 3d⁵ |
| M³⁺ option | E | [Ar] 4s¹ 3d⁴ |
Identify the first two electrons removed from the neutral atom.
[1]
Deduce the configuration of the 2+ ion.
[1]
Explain how the successive ionization data support variable oxidation states for this element.
[1]
Question 43
A discharge tube containing a low-pressure gas emits coloured light when a high voltage is applied.
Outline how the emitted light can be separated and recorded as qualitative and quantitative data.
[1]
Explain how line emission spectra arise and why they provide evidence for discrete electron energy levels and for the identity of elements.
[1]
Question 44
Hydrogen has emission lines in the ultraviolet, visible and infrared regions.
State the relationships among wavelength, frequency and photon energy.
[1]
Compare and contrast the electron transitions responsible for the ultraviolet, visible and infrared regions of the hydrogen emission spectrum.
[1]
Question 45
The electron arrangement of atoms can be represented by electron configurations and orbital diagrams.
State the maximum number of electrons in s, p and d sublevels.
[1]
Discuss how the Aufbau principle, Pauli exclusion principle and Hund’s rule are used to construct the ground-state orbital diagram and electron configuration of a sulfur atom.
[1]
Question 46
A student proposes the following condensed configurations:
Cr: [Ar] 4s² 3d⁴
Cu: [Ar] 4s² 3d⁹
Fe²⁺: [Ar] 4s² 3d⁴
State the correct configurations for Cr and Cu atoms.
[1]
Evaluate the student’s configurations, explaining the exceptions and the correct formation of Fe²⁺.
[1]
Question 47
First ionization energy varies across period 2.
Define first ionization energy and write the equation for the first ionization of nitrogen.
[1]
Explain the general trend and the two main discontinuities in first ionization energy across period 2.
[1]
Question 48
The convergence limit for a series in the emission spectrum of an atom occurs at a wavelength of 242 nm.
Outline why the convergence limit can be used to determine a first ionization energy.
[1]
Calculate the first ionization energy, in kJ mol⁻¹, from this wavelength.
[1]
Question 49
Successive ionization energy data are used to infer the electron configuration of an unknown main-group element.
State what is meant by successive ionization energies and why they increase.
[1]
Evaluate how a large jump in successive ionization energies is used to deduce the group of a main-group element, including one limitation of the method.
[1]
Question 50
Transition elements can form ions with different charges.
Deduce the electron configurations of Co²⁺ and Co³⁺ from Co, [Ar] 4s² 3d⁷.
[1]
Discuss how orbital energies and successive ionization energies help explain variable oxidation states in transition elements.
[1]
The Fast Track To Your
Best IB Coursework & College Essays
All content on this website has been developed independently from and is not endorsed by the International Baccalaureate Organization. International Baccalaureate and IB are registered trademarks owned by the International Baccalaureate Organization.