Electronic Configuration Worksheet With Answers

Electronic Configuration Worksheet: A Comprehensive Guide with Answers

Understanding electronic configuration is fundamental to grasping the behavior of atoms and molecules in chemistry. This worksheet provides a detailed exploration of electronic configuration, covering various aspects from basic principles to complex scenarios. We'll walk through the process step-by-step, offering numerous examples with solutions to solidify your understanding. This comprehensive guide will equip you with the knowledge and skills to confidently tackle any electronic configuration problem. By the end, you'll be able to predict the electronic configuration of various elements and ions, understand the significance of valence electrons, and appreciate the underlying principles of the periodic table.

Introduction to Electronic Configuration

Electronic configuration describes the arrangement of electrons in the different energy levels and sublevels within an atom. It's a fundamental concept in chemistry because it dictates an atom's chemical properties, reactivity, and bonding behavior. Electrons occupy specific orbitals, which are regions of space where there's a high probability of finding an electron. These orbitals are grouped into energy levels (also called shells), denoted by principal quantum numbers (n = 1, 2, 3, etc.), and sublevels (s, p, d, f), each with a specific capacity for electrons.

The Aufbau principle, Hund's rule, and the Pauli exclusion principle are crucial guidelines for determining electronic configuration. Let's briefly review them:

• Aufbau principle: Electrons fill orbitals starting from the lowest energy level and proceeding to higher levels. The order of filling is generally (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p...). Note that there are exceptions to this rule for some elements due to electron-electron repulsion and orbital stability.

• Hund's rule: Within a subshell (e.g., 2p), electrons will individually occupy each orbital with parallel spins before pairing up. This maximizes the total spin and minimizes electron-electron repulsion.

• Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms). This means that each orbital can hold a maximum of two electrons, which must have opposite spins (+1/2 and -1/2).

Step-by-Step Guide to Determining Electronic Configuration

Let's learn how to determine the electronic configuration of an element using the periodic table as a visual aid. We'll illustrate this with several examples.

Example 1: Carbon (C, atomic number 6)

1. Identify the atomic number: Carbon has an atomic number of 6, meaning it has 6 electrons.

2. Follow the Aufbau principle: We start filling orbitals from the lowest energy level:

   • 1s can hold 2 electrons: 1s²
   • 2s can hold 2 electrons: 2s²
   • 2p can hold 6 electrons (three 2p orbitals, each holding 2 electrons): 2p⁴ (We only need 4 electrons to complete the 6 electrons of Carbon)

3. Write the complete electronic configuration: The electronic configuration of carbon is 1s²2s²2p⁴.

Example 2: Oxygen (O, atomic number 8)

1. Atomic number: Oxygen has 8 electrons.

2. Aufbau principle:

   • 1s²
   • 2s²
   • 2p⁶ (We need 6 electrons to reach a total of 8 electrons)

3. Complete electronic configuration: The electronic configuration of oxygen is 1s²2s²2p⁶.

Example 3: Iron (Fe, atomic number 26)

This example demonstrates the filling of d orbitals.

1. Atomic number: Iron has 26 electrons.

2. Aufbau principle: Remember the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d...

   • 1s²
   • 2s²
   • 2p⁶
   • 3s²
   • 3p⁶
   • 4s²
   • 3d⁶

3. Complete electronic configuration: The electronic configuration of iron is 1s²2s²2p⁶3s²3p⁶4s²3d⁶. Note that 4s fills before 3d, according to the Aufbau principle.

Example 4: Copper (Cu, atomic number 29)

This is an example of an exception to the Aufbau principle.

1. Atomic number: Copper has 29 electrons.

2. Aufbau principle (with exception): While the expected configuration would be 1s²2s²2p⁶3s²3p⁶4s²3d⁹, copper exhibits a more stable configuration with a completely filled 3d subshell and one electron in the 4s subshell.

3. Complete electronic configuration: The electronic configuration of copper is 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰.

Electronic Configurations of Ions

Ions are formed by the gain or loss of electrons. To determine the electronic configuration of an ion, we start with the neutral atom's configuration and then add or remove electrons accordingly.

Example 5: O²⁻ (Oxide ion)

Oxygen (atomic number 8) gains two electrons to form the oxide ion, O²⁻.

1. Neutral oxygen: 1s²2s²2p⁴

2. Adding electrons: We add two electrons to the highest energy level (2p).

3. Electronic configuration of O²⁻: 1s²2s²2p⁶

Example 6: Fe³⁺ (Iron(III) ion)

Iron (atomic number 26) loses three electrons to form the iron(III) ion, Fe³⁺.

1. Neutral iron: 1s²2s²2p⁶3s²3p⁶4s²3d⁶

2. Removing electrons: Electrons are removed from the highest energy level first, which is 4s. Then we remove from 3d.

3. Electronic configuration of Fe³⁺: 1s²2s²2p⁶3s²3p⁶3d⁵

Valence Electrons and Chemical Bonding

Valence electrons are the electrons in the outermost energy level (highest principal quantum number). These electrons are primarily responsible for an atom's chemical behavior and bonding capabilities. Elements in the same group (vertical column) of the periodic table have the same number of valence electrons and, therefore, exhibit similar chemical properties.

For example, the alkali metals (Group 1) all have one valence electron, while the halogens (Group 17) have seven valence electrons. This explains their characteristic reactivity and bonding patterns.

Orbital Diagrams

Orbital diagrams visually represent the electronic configuration of an atom, showing the individual orbitals and the electrons they contain. Each orbital is represented by a box, and electrons are represented by arrows. Arrows pointing up and down represent electrons with opposite spins.

Example 7: Orbital Diagram for Carbon (C)

• 1s: ↑↓
• 2s: ↑↓
• 2p: ↑ ↑ ↑

This shows two electrons in the 1s orbital, two electrons in the 2s orbital, and two unpaired electrons in separate 2p orbitals (following Hund's rule).

Frequently Asked Questions (FAQ)

Q1: What are the exceptions to the Aufbau principle?

A1: Some elements, particularly those with partially filled d or f orbitals, show deviations from the Aufbau principle due to increased stability achieved through specific electron configurations. Chromium (Cr) and copper (Cu) are common examples.

Q2: How do I determine the number of valence electrons?

A2: The number of valence electrons is equal to the number of electrons in the outermost energy level (highest principal quantum number). You can also determine the number of valence electrons from the element's position in the periodic table (group number).

Q3: What is the significance of electronic configuration in chemical bonding?

A3: Electronic configuration determines the number of valence electrons, directly influencing an atom's ability to form chemical bonds. Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration, often resembling that of a noble gas (octet rule).

Conclusion

Mastering electronic configuration is a cornerstone of understanding chemical principles. By carefully following the Aufbau principle, Hund's rule, and the Pauli exclusion principle, and utilizing the periodic table as a guide, you can accurately determine the electronic configuration of elements and ions. Understanding valence electrons is crucial for predicting chemical behavior and bonding patterns. This comprehensive guide, along with practice using various examples and worksheets, will solidify your understanding of this vital chemical concept. Remember to practice consistently; the more you practice, the more confident and proficient you will become in determining electronic configurations. Continue to explore and expand your knowledge, and you'll find chemistry becomes increasingly fascinating and rewarding.