Complete The Full Electron Configuration For Each Element Below

Embarking on the journey to understand the electronic structure of atoms, mastering the art of writing full electron configurations is a fundamental step. An electron configuration describes which orbitals of an atom are occupied by electrons. Understanding and accurately writing electron configurations unlocks insights into an element's chemical properties, reactivity, and its position on the periodic table. Let's delve into how to systematically complete the full electron configuration for various elements.

Understanding the Basics

Before diving into specific examples, let's solidify the underlying principles that govern electron configurations:

• Orbitals and Quantum Numbers: Electrons reside in specific regions around the nucleus called orbitals. These orbitals are described by a set of quantum numbers:

  • Principal Quantum Number (n): Defines the energy level of the electron (n = 1, 2, 3,...). Higher values indicate higher energy levels.
  • Azimuthal or Angular Momentum Quantum Number (l): Defines the shape of the orbital (l = 0, 1, 2,..., n-1).
    • l = 0 corresponds to an s orbital (spherical).
    • l = 1 corresponds to a p orbital (dumbbell-shaped).
    • l = 2 corresponds to a d orbital (more complex shapes).
    • l = 3 corresponds to an f orbital (even more complex shapes).
  • Magnetic Quantum Number (ml): Defines the orientation of the orbital in space (ml = -l, -l+1,..., 0,..., l-1, l).
  • Spin Quantum Number (ms): Describes the intrinsic angular momentum of the electron, which is quantized and referred to as spin. It can be either +1/2 (spin up) or -1/2 (spin down).

• The Aufbau Principle: Electrons first fill the lowest energy orbitals available before occupying higher energy orbitals. This principle provides a roadmap for building up electron configurations.

• Hund's Rule: Within a subshell (e.g., p, d, or f), electrons will individually occupy each orbital before any orbital is doubly occupied. This maximizes the total spin angular momentum and minimizes electron-electron repulsion, leading to a more stable configuration.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of all four quantum numbers. This means that each orbital can hold a maximum of two electrons, and these electrons must have opposite spins.

The Order of Filling Orbitals

The order in which orbitals are filled isn't simply 1s, 2s, 2p, 3s, 3p, etc. Due to the complex interactions between electrons, the energies of orbitals can overlap. A helpful mnemonic diagram (often called the Madelung rule or diagonal rule) assists in remembering the correct filling order:

1. Draw a series of diagonal lines, starting from the top left.
2. Follow each diagonal line, adding the orbitals in the order you encounter them.

This gives the following filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Writing Full Electron Configurations: A Step-by-Step Guide

To illustrate the process, let's work through several examples, each increasing in complexity.

1. Hydrogen (H):

• Atomic number: 1 (meaning 1 proton and 1 electron in a neutral atom)
• The single electron occupies the lowest energy orbital, which is the 1s orbital.
• Full electron configuration: 1s¹

2. Helium (He):

• Atomic number: 2 (2 protons, 2 electrons)
• Both electrons fill the 1s orbital. Since an s orbital can hold a maximum of two electrons, it's now full.
• Full electron configuration: 1s²

3. Lithium (Li):

• Atomic number: 3 (3 protons, 3 electrons)
• The first two electrons fill the 1s orbital (1s²).
• The third electron must occupy the next available orbital, which is the 2s orbital.
• Full electron configuration: 1s² 2s¹

4. Oxygen (O):

• Atomic number: 8 (8 protons, 8 electrons)
• Start filling orbitals according to the Aufbau principle: 1s² 2s².
• We now have 4 electrons left to place. The next orbital is the 2p orbital, which has three sub-orbitals that can each hold two electrons.
• According to Hund's rule, we first fill each 2p orbital individually with one electron before doubling up in any of them. This gives us 2p⁴.
• Full electron configuration: 1s² 2s² 2p⁴

5. Sodium (Na):

• Atomic number: 11 (11 protons, 11 electrons)
• Filling the orbitals: 1s² 2s² 2p⁶. At this point, we've used 10 electrons.
• The next available orbital is the 3s orbital.
• The remaining electron occupies the 3s orbital: 3s¹.
• Full electron configuration: 1s² 2s² 2p⁶ 3s¹

6. Chlorine (Cl):

• Atomic number: 17 (17 protons, 17 electrons)
• Filling the orbitals: 1s² 2s² 2p⁶ 3s². At this point, we've used 12 electrons.
• The next orbital is the 3p orbital.
• We have 5 electrons left to place in the 3p orbitals. Following Hund's rule, we first fill each 3p orbital individually, then double up: 3p⁵.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁵

7. Potassium (K):

• Atomic number: 19 (19 protons, 19 electrons)
• Filling the orbitals: 1s² 2s² 2p⁶ 3s² 3p⁶. At this point, we've used 18 electrons.
• The next orbital according to the Aufbau principle is the 4s orbital.
• The remaining electron occupies the 4s orbital: 4s¹.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

8. Iron (Fe):

• Atomic number: 26 (26 protons, 26 electrons)
• Filling the orbitals: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². At this point, we've used 20 electrons.
• The next orbital is the 3d orbital.
• We have 6 electrons left to place in the 3d orbitals. Following Hund's rule, we first fill each 3d orbital individually, then double up: 3d⁶.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

9. Copper (Cu):

• Atomic number: 29 (29 protons, 29 electrons)
• Filling the orbitals: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹. This is where things get interesting!
• Although the above configuration follows the Aufbau principle, it's not the most stable configuration. Copper is more stable with a completely filled 3d subshell and a half-filled 4s subshell.
• One electron from the 4s orbital moves to the 3d orbital, resulting in a full 3d¹⁰ configuration.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

10. Silver (Ag):

• Atomic number: 47 (47 protons, 47 electrons)
• Following the Aufbau principle, we would expect: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d⁹.
• Similar to copper, silver achieves greater stability by having a completely filled d subshell.
• One electron from the 5s orbital shifts to the 4d orbital to fully occupy it.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹ 4d¹⁰

11. Gold (Au):

• Atomic number: 79 (79 protons, 79 electrons)
• Following the Aufbau principle would lead to a configuration ending in ... 6s² 4f¹⁴ 5d⁹.
• Like copper and silver, gold also exhibits an exception to the Aufbau principle to achieve a more stable, completely filled d subshell.
• One electron from the 6s orbital is promoted to the 5d orbital.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s¹ 4f¹⁴ 5d¹⁰

12. Uranium (U):

• Atomic number: 92 (92 protons and 92 electrons)
• Uranium is a more complex element, requiring careful adherence to the Aufbau principle and Hund's rule.
• Full electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁶ 7s² 5f³ 6d¹

Exceptions to the Aufbau Principle

As seen with copper, silver, and gold, there are exceptions to the Aufbau principle. These exceptions arise because completely filled and half-filled d and f subshells have extra stability. This stability is due to the symmetrical distribution of electron density, which minimizes electron-electron repulsion.

• Chromium (Cr): Predicted: [Ar] 4s² 3d⁴. Actual: [Ar] 4s¹ 3d⁵ (half-filled 4s and half-filled 3d)
• Molybdenum (Mo): Predicted: [Kr] 5s² 4d⁴. Actual: [Kr] 5s¹ 4d⁵ (half-filled 5s and half-filled 4d)

Tips for Success

• Master the Aufbau Principle and Hund's Rule: These are your primary tools.
• Use the Diagonal Rule (Madelung Rule): This visual aid helps you remember the filling order of orbitals.
• Know the Exceptions: Be aware of elements that deviate from the Aufbau principle due to the stability of half-filled and completely filled d and f subshells.
• Practice, Practice, Practice: The more you practice, the more comfortable you'll become with writing electron configurations.
• Double-Check Your Work: Ensure that the sum of the superscripts (number of electrons in each orbital) equals the atomic number of the element.

The Relationship to the Periodic Table

The periodic table is a visual representation of the electronic structure of elements. The arrangement of elements in the periodic table directly reflects the filling of electron orbitals.

• Groups (Vertical Columns): Elements within the same group have similar valence electron configurations (electrons in the outermost shell), leading to similar chemical properties.
• Periods (Horizontal Rows): Elements within the same period have the same highest principal quantum number (n) for their valence electrons.
• Blocks:
  • s-block: Groups 1 and 2 (alkali and alkaline earth metals). The last electron enters an s orbital.
  • p-block: Groups 13-18. The last electron enters a p orbital.
  • d-block: Transition metals. The last electron enters a d orbital.
  • f-block: Lanthanides and actinides. The last electron enters an f orbital.

By understanding the relationship between electron configurations and the periodic table, you can predict the electron configuration of an element based on its position in the table.

Full vs. Condensed Electron Configurations

The examples above showed full electron configurations, writing out every orbital and its electron occupancy. A condensed (or noble gas) electron configuration is a shorthand notation that uses the preceding noble gas symbol in brackets to represent the core electrons.

For example, Sodium (Na):

• Full electron configuration: 1s² 2s² 2p⁶ 3s¹
• Condensed electron configuration: [Ne] 3s¹ (Neon (Ne) has the configuration 1s² 2s² 2p⁶)

Condensed electron configurations are particularly useful for larger elements, as they save space and focus on the valence electrons, which are most important for chemical bonding.

Applications of Electron Configurations

Understanding electron configurations has broad applications in chemistry and related fields:

• Predicting Chemical Properties: The valence electron configuration determines how an element will interact with other elements, forming chemical bonds.
• Understanding Bonding: Electron configurations explain the types of bonds (ionic, covalent, metallic) that an element can form.
• Explaining Spectra: The electronic transitions between energy levels within an atom result in the absorption or emission of light at specific wavelengths, creating atomic spectra. Electron configurations help explain these spectra.
• Predicting Magnetic Properties: Electron configurations determine whether an element or ion is paramagnetic (attracted to a magnetic field due to unpaired electrons) or diamagnetic (repelled by a magnetic field due to all paired electrons).
• Materials Science: Electron configurations are used to design and understand the properties of new materials.

Conclusion

Writing full electron configurations is a crucial skill for understanding the behavior of atoms and molecules. By mastering the Aufbau principle, Hund's rule, and recognizing exceptions, you can accurately determine the electronic structure of any element. This knowledge provides a foundation for understanding chemical properties, bonding, and a wide range of phenomena in chemistry and related fields. Practice consistently, and you'll find yourself confidently navigating the world of electron configurations.