Electron Configuration Test: Practice Notation and Periodic Patterns
Quick, free quiz with electron configuration practice questions. Instant results.
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.
Study Outcomes
- 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.
- Write Electron Configurations -
Apply systematic rules to write full and shorthand electron configurations for any element on the periodic table.
- Analyze Periodic Table Trends -
Interpret how electron configurations explain trends in atomic radius, ionization energy, and electron affinity across periods and groups.
- Predict Chemical Behavior -
Use electron configuration insights to predict valence electrons, oxidation states, and likely chemical reactivity of elements.
- Classify Elements by Orbital Types -
Identify and group elements based on their s, p, d, and f orbital electron arrangements.
- Evaluate Configuration Exceptions -
Recognize and explain common exceptions to expected electron configurations, such as those found in transition metals.
Cheat Sheet
- 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.
- 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.
- 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.
- 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.
- 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.