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Quizzes > Science

Electron Configuration Test: Practice Notation and Periodic Patterns

Quick, free quiz with electron configuration practice questions. Instant results.

Editorial: Review CompletedCreated By: Freek-Jan VernhoutUpdated Aug 27, 2025
Difficulty: Moderate
2-5mins
Learning OutcomesCheat Sheet
Paper art illustration of a periodic table and electron orbits on a coral background for a chemistry skills quiz

This quiz helps you practice electron configuration, write correct orbital notation, and spot periodic trends. For more targeted drills, try the electron configuration quiz or check your understanding with an orbital notation quiz. For quick lookups while you work, keep the quantum number calculator open.

What is the electron configuration of hydrogen?
1s1
1p1
1s2
2s1
Hydrogen, with atomic number 1, has a single electron which occupies the lowest energy orbital, the 1s orbital. According to the Aufbau principle, the 1s orbital is filled first before any higher orbitals. Since hydrogen has only one electron, its configuration is 1s1. .
How many electrons can the 2p subshell hold?
6
4
2
8
Each p subshell has three orbitals (px, py, pz), and each orbital can hold a maximum of two electrons with opposite spins. Therefore, the p subshell can hold 3 × 2 = 6 electrons. This applies to the 2p subshell as well as higher p subshells. .
Which orbital comes next after 4s?
5s
3d
4p
4f
The Madelung rule dictates the order in which orbitals are filled based on the sum of the principal quantum number (n) and azimuthal quantum number (l). After 4s (n + l = 4), the 3d subshell (n + l = 3 + 2 = 5) is next with the lowest n + l value. Even though 4p also has a sum of 5, the 3d orbital has a lower n, so it fills before 4p. .
What is the maximum number of electrons in the third energy level?
32
8
20
18
The third energy level consists of the 3s, 3p, and 3d subshells. The 3s can hold 2 electrons, 3p can hold 6, and 3d can hold 10, giving a total of 2 + 6 + 10 = 18 electrons. Therefore, the third energy level can accommodate up to 18 electrons. .
Determine the electron configuration of calcium with atomic number 20.
[Ar] 4s2
[Kr] 5s2
[Ne] 3s2 3p6
[Ar] 3d2
Calcium has 20 electrons which fill the 1s through 3p subshells equivalent to the argon core ([Ar]). The next available subshell is 4s, which holds the remaining two electrons. Therefore, the ground-state electron configuration is [Ar] 4s2. The 3d subshell remains empty because 4s is lower in energy for calcium. .
Which element has the electron configuration [Ne] 3s2 3p3?
Sulfur
Phosphorus
Silicon
Aluminum
This configuration adds up to 15 electrons: 10 in the neon core, 2 in the 3s orbital, and 3 in the 3p orbital. Element number 15 is phosphorus. Sulfur would have four electrons in 3p (3p4), silicon would have two (3p2), and aluminum would have one (3p1). .
What is the correct order of subshell filling up to the 3s orbital?
1s, 2s, 3p, 2p
1s, 3s, 2s, 2p
1s, 2s, 2p, 3s
1s, 2p, 2s, 3s
The Aufbau principle states that electrons occupy the lowest energy subshell available. This leads to the sequence: 1s, then 2s, then 2p, and then 3s. Each step reflects the increasing principal and azimuthal quantum numbers. Deviations from this sequence do not match observed energy levels. .
How many unpaired electrons does a ground-state oxygen atom have?
4
1
2
0
With six valence electrons (2s2 2p4), oxygen fills the 2p orbitals such that one electron enters each orbital before pairing. This leaves two of the p orbitals with single electrons and one p orbital with a pair. According to Hund's rule, this arrangement minimizes electron repulsion. Hence, oxygen has two unpaired electrons. .
What is the electron configuration of Fe3+ (iron(III) ion)?
[Ar] 3d6
[Ar] 3d5
[Ar] 4s1 3d5
[Ar] 4s2 3d3
Neutral iron (Fe) has the configuration [Ar]4s2 3d6. When forming Fe3+, three electrons are removed: two from the 4s subshell and one from the 3d subshell. This results in [Ar]3d5 for Fe3+. The 4s electrons are lost before the 3d electrons due to their higher energy once occupied. .
Which element has a ground-state configuration ending in 4f5 6s2?
Cerium
Samarium
Neodymium
Promethium
Lanthanides fill the 4f subshell after the xenon core and 6s electrons. Promethium (Z=61) has five electrons in the 4f orbitals and two in 6s, giving [Xe]4f5 6s2. Other nearby lanthanides have different counts: Ce has one 4f electron, Nd has four, and Sm has six. .
What is the term for the stabilization effect observed in half-filled and fully filled subshells?
Aufbau principle
Hund's rule
Pauli exclusion principle
Exchange energy
Exchange energy refers to the stabilization due to parallel spins in degenerate orbitals, where electrons with the same spin can exchange positions without altering the system. This leads to increased stability for half-filled (e.g., d5) and fully filled (d10) subshells. Hund's rule describes that electrons will occupy separate orbitals to maximize total spin, but exchange energy quantifies the added stabilization. .
Identify the anomaly in the electron configurations of transition metals: Which element among these has the configuration [Ar]3d10 4s1 instead of the expected [Ar]3d9 4s2?
Copper
Gallium
Nickel
Zinc
Copper deviates from the expected Aufbau filling order to achieve a completely filled d subshell, which is energetically favorable. Consequently, its electron configuration is [Ar]3d10 4s1 rather than [Ar]3d9 4s2. This anomaly is driven by additional stability from a filled d orbital. .
What is the ground-state electron configuration of lawrencium (element 103)?
[Rn] 5f14 7s2 7p1
[Rn] 5f14 6d1 7s2
[Rn] 5f14 7s1 6d2
[Rn] 5f13 7s2 6d2
Lawrencium (Lr), with atomic number 103, completes the 5f subshell with 14 electrons and has two electrons in 7s and one in 7p according to recent spectroscopic studies. This leads to the configuration [Rn]5f14 7s2 7p1, differing from the expected 6d occupation. The assignment reflects deviations in subshell energy levels at this high atomic number. .
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Study Outcomes

  1. Understand Electron Configuration Principles -

    Gain a clear grasp of Aufbau's principle, Hund's rule, and the Pauli exclusion principle to accurately construct electron configurations.

  2. Write Electron Configurations -

    Apply systematic rules to write full and shorthand electron configurations for any element on the periodic table.

  3. Analyze Periodic Table Trends -

    Interpret how electron configurations explain trends in atomic radius, ionization energy, and electron affinity across periods and groups.

  4. Predict Chemical Behavior -

    Use electron configuration insights to predict valence electrons, oxidation states, and likely chemical reactivity of elements.

  5. Classify Elements by Orbital Types -

    Identify and group elements based on their s, p, d, and f orbital electron arrangements.

  6. Evaluate Configuration Exceptions -

    Recognize and explain common exceptions to expected electron configurations, such as those found in transition metals.

Cheat Sheet

  1. Aufbau Principle & Orbital Filling -

    The Aufbau principle dictates that electrons occupy the lowest-energy orbitals first, following the pattern 1s → 2s → 2p → 3s, etc. A popular mnemonic is "1s two, 2s two, 2p six," which helps you remember the order through the first few orbitals. This concept is backed by quantum mechanics and detailed on many university chemistry department sites.

  2. Pauli Exclusion Principle -

    Wolfgang Pauli's rule states that no two electrons in the same atom can have identical sets of quantum numbers, meaning each orbital holds a maximum of two electrons with opposite spins (↑↓). Visualizing this with up/down arrows in diagrams ensures you don't overfill an orbital. This principle is fundamental in textbooks like Atkins' Physical Chemistry.

  3. Hund's Rule of Maximum Multiplicity -

    Hund's rule tells us to distribute electrons singly across degenerate orbitals (same energy) before pairing them, minimizing electron-electron repulsion. For example, in the p sublevel (px, py, pz), place one electron in each orbital first, then pair. This rule is highlighted in NRC and ACS educational resources.

  4. Periodic Table Blocks & Electron Configuration -

    The periodic table is divided into s, p, d, and f blocks, corresponding to the orbital being filled. For instance, Group 2 elements end in ns2, while transition metals fill (n - 1)d orbitals. Recognizing these blocks helps you quickly predict valence electron patterns, as shown in IUPAC guidelines.

  5. Common Configuration Exceptions -

    Some elements like chromium (Cr) and copper (Cu) break the expected filling order to achieve half-filled (d5) or fully filled (d10) stability, e.g., Cr is [Ar] 3d5 4s1 not 3d4 4s2. Memorize these key exceptions by noting that half- and full-filled subshells lower the overall energy. These anomalies are discussed in peer-reviewed chemistry journals and major university curriculum notes.

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