Electron Configuration Practice Worksheet

Struggling to grasp the intricacies of electron configurations? You’re not alone! Electron configuration, the arrangement of electrons within the energy levels and sublevels of an atom, can seem daunting at first. But mastering it is crucial for understanding chemical bonding, predicting reactivity, and unraveling the properties of elements. That’s why practice is key! This post provides a valuable resource: an Electron Configuration Practice Worksheet designed to help you hone your skills. We’ll walk you through the basics and then provide the answers so you can check your work and identify areas where you might need more practice. So, grab a periodic table, a pencil, and get ready to dive into the exciting world of electron configurations!

Understanding Electron Configurations

Before we get to the worksheet answers, let’s recap the fundamental principles behind electron configurations. Remember these key concepts:

* **Energy Levels (n):** Electrons occupy different energy levels, numbered 1, 2, 3, and so on (n=1, n=2, n=3…). Higher numbers represent higher energy levels.
* **Sublevels (l):** Within each energy level, there are sublevels denoted by letters: *s*, *p*, *d*, and *f*.
* *s* sublevels can hold a maximum of 2 electrons.
* *p* sublevels can hold a maximum of 6 electrons.
* *d* sublevels can hold a maximum of 10 electrons.
* *f* sublevels can hold a maximum of 14 electrons.
* **Orbitals:** Each sublevel is further divided into orbitals, each of which can hold a maximum of two electrons. The *s* sublevel has one orbital, the *p* sublevel has three orbitals, the *d* sublevel has five orbitals, and the *f* sublevel has seven orbitals.
* **Aufbau Principle:** Electrons fill the lowest energy levels and sublevels first. This means you generally fill the levels in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
* **Hund’s Rule:** Within a sublevel, electrons will individually occupy each orbital before doubling up in any one orbital. All of the unpaired electrons will have the same spin (either spin up or spin down)
* **Pauli Exclusion Principle:** No two electrons in an atom can have the same set of four quantum numbers. This means each orbital can hold a maximum of two electrons with opposite spins.

Remembering these rules will make writing electron configurations much easier. The periodic table itself is a fantastic tool for predicting the filling order! The s-block elements (groups 1 and 2) are filling the *s* sublevels, the p-block elements (groups 13-18) are filling the *p* sublevels, the d-block elements (groups 3-12) are filling the *d* sublevels, and the f-block elements (lanthanides and actinides) are filling the *f* sublevels.

Shortcut Notation (Noble Gas Configuration)

For larger atoms, writing out the full electron configuration can be tedious. That’s where noble gas configuration comes in handy. Identify the noble gas that comes before the element in question on the periodic table. Write its symbol in brackets, then continue the electron configuration from that point onward. For example, instead of writing the full electron configuration for Sodium (Na), you can use [Ne]3s¹. [Ne] represents the electron configuration of Neon.

Electron Configuration Practice Worksheet – Answer Key

Below you’ll find the answers to a typical Electron Configuration Practice Worksheet. Use this to check your own work. Remember to pay attention to the correct superscripts, which indicate the number of electrons in each sublevel.

• Hydrogen (H): 1s¹
• Helium (He): 1s²
• Lithium (Li): 1s²2s¹
• Beryllium (Be): 1s²2s²
• Boron (B): 1s²2s²2p¹
• Carbon (C): 1s²2s²2p²
• Nitrogen (N): 1s²2s²2p³
• Oxygen (O): 1s²2s²2p⁴
• Fluorine (F): 1s²2s²2p⁵
• Neon (Ne): 1s²2s²2p⁶
• Sodium (Na): 1s²2s²2p⁶3s¹ or [Ne]3s¹
• Magnesium (Mg): 1s²2s²2p⁶3s² or [Ne]3s²
• Aluminum (Al): 1s²2s²2p⁶3s²3p¹ or [Ne]3s²3p¹
• Silicon (Si): 1s²2s²2p⁶3s²3p² or [Ne]3s²3p²
• Phosphorus (P): 1s²2s²2p⁶3s²3p³ or [Ne]3s²3p³
• Sulfur (S): 1s²2s²2p⁶3s²3p⁴ or [Ne]3s²3p⁴
• Chlorine (Cl): 1s²2s²2p⁶3s²3p⁵ or [Ne]3s²3p⁵
• Argon (Ar): 1s²2s²2p⁶3s²3p⁶
• Potassium (K): 1s²2s²2p⁶3s²3p⁶4s¹ or [Ar]4s¹
• Calcium (Ca): 1s²2s²2p⁶3s²3p⁶4s² or [Ar]4s²
• Scandium (Sc): 1s²2s²2p⁶3s²3p⁶4s²3d¹ or [Ar]4s²3d¹
• Titanium (Ti): 1s²2s²2p⁶3s²3p⁶4s²3d² or [Ar]4s²3d²
• Vanadium (V): 1s²2s²2p⁶3s²3p⁶4s²3d³ or [Ar]4s²3d³
• Chromium (Cr): 1s²2s²2p⁶3s²3p⁶4s¹3d⁵ or [Ar]4s¹3d⁵ (Note the exception!)
• Manganese (Mn): 1s²2s²2p⁶3s²3p⁶4s²3d⁵ or [Ar]4s²3d⁵
• Iron (Fe): 1s²2s²2p⁶3s²3p⁶4s²3d⁶ or [Ar]4s²3d⁶
• Cobalt (Co): 1s²2s²2p⁶3s²3p⁶4s²3d⁷ or [Ar]4s²3d⁷
• Nickel (Ni): 1s²2s²2p⁶3s²3p⁶4s²3d⁸ or [Ar]4s²3d⁸
• Copper (Cu): 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰ or [Ar]4s¹3d¹⁰ (Note the exception!)
• Zinc (Zn): 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰ or [Ar]4s²3d¹⁰

Important Notes: Chromium (Cr) and Copper (Cu) are exceptions to the Aufbau principle. They are more stable with a half-filled (d⁵) and completely filled (d¹⁰) d-sublevel, respectively. This results in one electron being promoted from the 4s orbital to the 3d orbital.

Practice makes perfect! Keep working through electron configuration problems, and you’ll become more comfortable with the rules and patterns. Good luck!

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