Medium Electron Configuration Practice Questions

Concept Explanation

Electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals, describing how electrons are arranged around the nucleus based on energy levels and subshells.

To master medium-level electron configuration, you must understand three fundamental rules. First, the Aufbau Principle states that electrons fill lower-energy orbitals before moving to higher ones. Second, Hund's Rule dictates that every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied. Third, the Pauli Exclusion Principle asserts that no two electrons in an atom can have the identical set of four quantum numbers. At this level, we also encounter exceptions in the transition metals, such as Chromium and Copper, where half-filled or fully-filled d-subshells offer increased stability. Understanding these patterns is essential for predicting chemical behavior, similar to how periodic trends help us understand atomic size and electronegativity.

Solved Examples

1. Determine the full electron configuration for Phosphorus (Z = 15).

   1. Identify the total number of electrons: 15.

   2. Follow the orbital filling sequence: 1s, 2s, 2p, 3s, 3p.

   3. Fill the orbitals: 1s² (2 electrons), 2s² (2 electrons), 2p⁶ (6 electrons), 3s² (2 electrons), 3p³ (3 electrons).

   4. Final configuration: 1s² 2s² 2p⁶ 3s² 3p³.

2. Write the noble gas (shorthand) configuration for Iron (Z = 26).

   1. Locate Iron on the periodic table and find the preceding noble gas, which is Argon (Z = 18).

   2. Represent the first 18 electrons as [Ar].

   3. Identify the remaining 8 electrons. After Argon, electrons fill the 4s and then the 3d orbitals.

   4. Fill 4s² (2 electrons) and then 3d⁶ (6 electrons).

   5. Final configuration: [Ar] 4s² 3d⁶.

3. Identify the ground-state configuration for the Copper atom (Z = 29), noting any exceptions.

   1. The expected configuration based on the Aufbau principle would be [Ar] 4s² 3d⁹.

   2. Recognize that a completely filled d-subshell (d¹⁰) provides extra stability.

   3. One electron from the 4s orbital promotes to the 3d orbital.

   4. Final configuration: [Ar] 4s¹ 3d¹⁰.

Practice Questions

1. Write the full electron configuration for the Sulfur atom (Z = 16).

2. Provide the noble gas notation for the Silver (Ag) atom, keeping in mind its position in the transition metals.

3. Determine the electron configuration for the Magnesium ion (Mg²⁺).

4. Which element has the ground-state electron configuration [Kr] 5s² 4d¹⁰ 5p²?

5. Write the electron configuration for the Bromide ion (Br⁻).

6. Identify the total number of p-electrons in a neutral Chlorine atom.

7. Write the noble gas configuration for Chromium (Z = 24), accounting for stability exceptions.

8. How many unpaired electrons are present in the ground state of a Manganese (Mn) atom?

9. Write the electron configuration for the Scandium ion (Sc³⁺).

10. An element has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵. Name the element and its group.

Answers & Explanations

1. 1s² 2s² 2p⁶ 3s² 3p⁴: Sulfur has 16 electrons. Following the Aufbau principle, we fill 1s, 2s, 2p, 3s, and finally place the remaining 4 electrons in the 3p subshell.

2. [Kr] 5s¹ 4d¹⁰: Like Copper, Silver (Z=47) is an exception. It promotes an electron from the 5s to the 4d orbital to achieve a stable, fully-filled d-subshell.

3. 1s² 2s² 2p⁶: Magnesium (Z=12) loses two electrons to form Mg²⁺. The electrons are removed from the outermost 3s orbital, leaving the configuration of Neon.

4. Tin (Sn): Krypton represents 36 electrons. Adding 2 (5s), 10 (4d), and 2 (5p) gives a total of 50 electrons, which corresponds to Tin.

5. [Ar] 4s² 3d¹⁰ 4p⁶: Bromine (Z=35) gains one electron to form Br⁻, completing the 4p subshell and achieving the stable configuration of Krypton.

6. 11: Chlorine (Z=17) has the configuration 1s² 2s² 2p⁶ 3s² 3p⁵. The p-electrons are 6 from the 2p subshell and 5 from the 3p subshell (6 + 5 = 11).

7. [Ar] 4s¹ 3d⁵: Chromium is an exception where a half-filled d-subshell provides more stability than a full 4s subshell and a 3d⁴ configuration.

8. 5: Manganese [Ar] 4s² 3d⁵ has five electrons in the 3d subshell. According to Hund's Rule, each of the five d-orbitals will hold one electron, all unpaired.

9. [Ar] or 1s² 2s² 2p⁶ 3s² 3p⁶: Scandium (Z=21) loses three electrons (two from 4s and one from 3d) to reach a stable noble gas configuration.

10. Bromine (Group 17): The total electron count is 35. This element is in the p-block, specifically the Halogen group. This relates closely to how we determine Lewis structures based on valence electrons.

Quick Quiz

Frequently Asked Questions

What is the difference between ground state and excited state?

The ground state is the lowest energy arrangement of electrons in an atom, following standard filling rules. An excited state occurs when an electron absorbs energy and jumps to a higher energy orbital, leaving a vacancy in a lower shell.

Why does the 4s orbital fill before the 3d orbital?

According to the Aufbau principle, the 4s orbital has a slightly lower energy level than the 3d orbital in neutral atoms. This occurs because the 4s orbital is more penetrating, allowing electrons to spend more time closer to the nucleus despite the higher principal quantum number.

How do you write electron configurations for ions?

For cations, remove the required number of electrons starting from the highest principal energy level (n), specifically the s-orbital before the d-orbital in transition metals. For anions, add electrons to the next available orbital in the filling sequence until the total charge is accounted for.

What are the exceptions to the Aufbau Principle?

Exceptions primarily occur in the transition metals, such as Chromium ([Ar] 4s¹ 3d⁵) and Copper ([Ar] 4s¹ 3d¹⁰). These atoms shift an electron to achieve half-filled or fully-filled d-subshells, which are energetically more stable states than the predicted configurations.

How does electron configuration relate to the periodic table?

The periodic table is organized into blocks (s, p, d, f) that correspond to the subshell being filled by the valence electrons. An element's position in a specific period and group directly reflects its highest occupied energy level and the number of electrons in its outermost shell.

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