Hard Electron Configuration Practice Questions

Concept Explanation

Electron configuration is the symbolic representation of how electrons are distributed among the various atomic orbitals of an atom or ion according to their energy levels. This distribution follows fundamental physical principles, primarily the Aufbau Principle, which states that electrons fill lower-energy subshells before moving to higher ones. However, at a hard level of chemistry, we must account for nuanced phenomena such as the stability of half-filled and fully-filled d-subshells, the anomalous configurations of transition metals, and the specific rules for forming cations and anions. Understanding these advanced patterns is essential for predicting chemical reactivity and is closely related to properties like ionization energy and periodic trends.

To master hard electron configuration, you must navigate three core rules:

• The Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.

• Hund’s Rule: Electrons will occupy empty orbitals of the same energy (degenerate orbitals) singly before pairing up to minimize electron-electron repulsion.

• Stability Anomalies: Elements like Chromium ([Ar] 4s¹ 3d⁵) and Copper ([Ar] 4s¹ 3d¹⁰) deviate from the standard filling order because a half-filled or fully-filled d-subshell offers a lower energy state than the expected 4s² 3dⁿ arrangement.

When dealing with transition metal ions, a critical rule applies: electrons are always removed from the highest principal quantum number (n) orbital first. For example, in the fourth period, electrons are removed from the 4s orbital before the 3d orbital, even though the 3d orbital was filled last according to Aufbau.

Solved Examples

Example 1: The ground-state configuration of the Gadolinium (Gd, Z=64) atom.

1. Identify the nearest noble gas preceding Gadolinium: Xenon ([Xe], Z=54).

2. Determine the remaining electrons: 64 - 54 = 10 electrons.

3. Following the filling order: 6s² (2 electrons), then 4f and 5d.

4. Gadolinium is an exception in the lanthanide series; it prefers a half-filled f-orbital for stability. Instead of [Xe] 6s² 4f⁸, it becomes [Xe] 6s² 4f⁷ 5d¹.

5. Final Configuration: [Xe] 4f⁷ 5d¹ 6s².

Example 2: The electron configuration of the Iron(III) ion (Fe³⁺).

1. Start with the neutral Iron atom (Z=26): [Ar] 4s² 3d⁶.

2. Identify the electrons to remove for a 3+ charge: 3 electrons total.

3. Apply the rule: Remove from the highest 'n' shell first. Remove both 4s electrons.

4. Remove the remaining 1 electron from the 3d subshell.

5. Final Configuration: [Ar] 3d⁵.

Example 3: Identifying an element from an excited state configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵ 4p¹.

1. Count the total number of electrons: 2+2+6+2+6+1+5+1 = 25 electrons.

2. In a neutral atom, the number of electrons equals the atomic number (Z).

3. Atomic number 25 corresponds to Manganese (Mn).

4. Note that this is an excited state because an electron has moved from the 4s or 3d orbital into the 4p orbital.

Practice Questions

1. Write the full ground-state electron configuration for the Silver (Ag) atom, accounting for any stability anomalies.

2. Determine the electron configuration for the Ruthenium(IV) ion (Ru⁴⁺).

3. Identify the element that has the following ground-state configuration for its 2+ ion: [Kr] 4d⁸.

Start Learning Smarter Today

Join thousands of students using AI-powered study tools to achieve better results.

Get Started Free

4. Which of the following ions is isoelectronic with Thallium(I) (Tl⁺)? Provide its abbreviated electron configuration.

5. An element in the 5th period has 3 unpaired electrons in its ground state and is a metal. Provide the abbreviated configuration.

6. Explain why the configuration of Tungsten (W) is [Xe] 4f¹⁴ 5d⁴ 6s² and not [Xe] 4f¹⁴ 5d⁵ 6s¹, despite the chromium/molybdenum pattern.

7. Write the configuration for the Lead(IV) ion (Pb⁴⁺) and explain which electrons were lost.

8. Determine the number of unpaired electrons in a ground-state gaseous Nickel(II) ion (Ni²⁺).

9. Write the electron configuration for the Copper(I) ion and explain its diamagnetic or paramagnetic nature.

10. Identify the period, group, and block for an element with the configuration [Rn] 7s² 5f¹⁴ 6d¹⁰.

Answers & Explanations

1. Silver (Ag): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹ 4d¹⁰. Like Copper, Silver promotes one 5s electron to the 4d subshell to achieve a completely filled d-subshell, which is more energetically favorable.

2. Ru⁴⁺: [Kr] 4d⁴. Neutral Ruthenium is [Kr] 5s¹ 4d⁷ (an anomaly) or traditionally [Kr] 5s² 4d⁶. When forming a 4+ ion, we remove the two 5s electrons first, then two 4d electrons, leaving 4d⁴.

3. Palladium (Pd). The 2+ ion has [Kr] 4d⁸. Adding two electrons back gives the neutral atom with 46 electrons. Note: Neutral Pd is actually an extreme anomaly [Kr] 4d¹⁰, but the atomic number identifies it as Palladium.

4. Mercury (Hg) or Lead(II) (Pb²⁺). Tl⁺ has 80 electrons. Its configuration is [Xe] 4f¹⁴ 5d¹⁰ 6s². Mercury (Hg) is isoelectronic as it also has 80 electrons.

5. Niobium (Nb) or Rhodium (Rh). Niobium ([Kr] 5s¹ 4d⁴) has 5 unpaired electrons. However, Vanadium's group member Niobium is tricky. Let's look at Rhodium: [Kr] 5s¹ 4d⁸. Detailed orbital diagrams show Rhodium fits common "hard" criteria for specific unpaired counts in transition blocks.

6. Relativistic effects and pairing energy. In heavier elements like Tungsten, the energy gap between s and d orbitals changes, and the stability gained from a half-filled d-shell in W is less than the energy required to promote the s-electron, unlike in Cr or Mo. (Reference: LibreTexts Chemistry).

7. Pb⁴⁺: [Xe] 4f¹⁴ 5d¹⁰. Neutral Pb is [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p². To form Pb⁴⁺, the atom loses the two 6p electrons and the two 6s electrons (the outer shell electrons).

8. 2 unpaired electrons. Ni is [Ar] 4s² 3d⁸. Ni²⁺ is [Ar] 3d⁸. In the five d-orbitals, 8 electrons fill as: 5 up, then 3 down. This leaves 2 orbitals with single (unpaired) electrons.

9. Cu⁺: [Ar] 3d¹⁰. It is diamagnetic. Because all electrons in the 3d subshell are paired, there is no net magnetic moment.

10. Period 7, Group 12, d-block (Copernicium). The highest principal quantum number is 7 (Period 7). The 6d subshell is filled, placing it in the 12th column of the d-block.

Quick Quiz

Frequently Asked Questions

What is the (n+l) rule in electron configuration?

The (n+l) rule, also known as Madelung's rule, specifies that orbitals with a lower sum of the principal quantum number (n) and azimuthal quantum number (l) are filled first. If two orbitals have the same sum, the one with the lower 'n' value is filled first.

Why are there exceptions to the Aufbau principle in transition metals?

Exceptions occur because the energy difference between the ns and (n-1)d subshells is very small. In certain cases, the stability gained from electron-electron repulsion minimization in half-filled or fully-filled d-subshells outweighs the energy cost of shifting an electron.

How do you write electron configurations for anions?

To write the configuration for an anion, you add the number of electrons corresponding to the negative charge to the next available lowest-energy orbital. For example, Chlorine (Z=17) adds one electron to its 3p subshell to become Cl⁻ with a neon-like 3p⁶ ending.

What does it mean for an atom to be isoelectronic?

Isoelectronic refers to two or more molecular entities (atoms, ions, or molecules) that have the same number of electrons and the same electron configuration. A common example is the Na⁺ ion and the Neon atom, both having 10 electrons.

How do you determine if an atom is paramagnetic or diamagnetic?

An atom is paramagnetic if it contains at least one unpaired electron, causing it to be attracted to a magnetic field. It is diamagnetic if all its electrons are paired, resulting in a slight repulsion from magnetic fields.

What is the noble gas shorthand notation?

Noble gas notation is a method of simplifying electron configurations by using the symbol of the previous noble gas in brackets to represent the core electrons. This allows chemists to focus on the valence electrons, which are responsible for chemical behavior.

Enjoyed this article?

Share it with others who might find it helpful.

Share Article