Question

Write the electron configuration for each ion. What do all of the electron configurations have in common? (a) \(\mathrm{Ca}^{2+}\) (b) \(\mathrm{K}^{+}\) (c) \(\mathrm{S}^{2-}\) (d) \(\mathrm{Br}^{-}\)

Short answer

The electron configurations are: (a) Ca^{2+}: 1s2 2s2 2p6 3s2 3p6, (b) K^{+}: 1s2 2s2 2p6 3s2 3p6, (c) S^{2-}: 1s2 2s2 2p6 3s2 3p6, (d) Br^{-}: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6. All ions have electron configurations that are isoelectronic with the nearest noble gas, with fully filled s and p subshells in the highest energy level.

Step-by-step solution

Determine the electron configurations for neutral atoms

Find the electron configurations of the neutral atoms before they become ions. Use the periodic table to determine the number of electrons for calcium (Ca), potassium (K), sulfur (S), and bromine (Br). Write down their electron configurations: - Calcium (Ca, atomic number 20): 1s2 2s2 2p6 3s2 3p6 4s2 - Potassium (K, atomic number 19): 1s2 2s2 2p6 3s2 3p6 4s1 - Sulfur (S, atomic number 16): 1s2 2s2 2p6 3s2 3p4 - Bromine (Br, atomic number 35): 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5

Adjust electron configurations for ions

Adjust the electron configurations for each atom to match the ions' charges. Ions are formed by the loss or gain of electrons: - For calcium ion (Ca^{2+}): Remove two electrons (from 4s). Electron configuration: 1s2 2s2 2p6 3s2 3p6 - For potassium ion (K^{+}): Remove one electron (from 4s). Electron configuration: 1s2 2s2 2p6 3s2 3p6 - For sulfur ion (S^{2-}): Add two electrons (to 3p). Electron configuration: 1s2 2s2 2p6 3s2 3p6 - For bromine ion (Br^{-}): Add one electron (to 4p). Electron configuration: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6

Identify commonalities in electron configurations

Compare the electron configurations of all four ions to find a common pattern. All the configurations end with a fully filled s and p subshells in their respective highest energy level, making them isoelectronic with the nearest noble gas.

Key concepts

Periodic Table

The periodic table is a comprehensive chart that organizes all known chemical elements in an informative and systematic array. Elements are arranged in order of increasing atomic number — the number of protons in an atom's nucleus. The layout of the table reflects the periodic law which states that certain properties of elements recur periodically as their atomic numbers increase. This is seen in the table's rows (periods) and columns (groups).

• Groups consist of elements with similar properties and trends in chemical reactivity.
• Periods are horizontal sequences where outer-shell electron configurations progress from one noble gas to the next.

To locate an element on the periodic table, one should know its atomic number. Understanding the periodic table allows us to predict how different elements will react and informs us about their properties, such as electron configurations which are crucial when discussing ions.

Ions and Electron Gain/Loss

Ions are atoms or molecules that have lost or gained one or more of their valence electrons, giving them a positive or negative electrical charge. The process of losing electrons results in a positively charged ion or cation, whereas gaining electrons results in a negatively charged ion or anion. The electron configurations of ions are thus different from their neutral atoms as they achieve stability by either filling or emptying their outermost electron shell.

• Cations are formed when atoms lose electrons, resulting in a lower energy level becoming the outermost shell.
• Anions are formed when atoms gain electrons, filling up their outer shell closer to the configuration of a noble gas.

By examining the periodic table, we can predict the most likely ions an element will form based on its group and period. For example, elements in group 1 are most likely to form +1 cations, while group 17 elements often form -1 anions.

Electron Configuration Notation

Electron configuration notation is a shorthand representation of the arrangement of electrons in an atom or ion based on the increasing energy levels and subshells. These notations reflect the quantum mechanical model of the atom and are essential for understanding chemical bonding and reactivity. There are a few key principles:

• Each electron is described by four quantum numbers that determine its energy level, the type of subshell, and its orientation and spin within that subshell.
• The notation uses a series of numbers and letters where the number indicates the energy level (principal quantum number) and the letter represents the subshell type (s,p,d,f).
• The superscript indicates how many electrons are in that subshell.

For example, the notation 1s2 means an energy level 1 subshell 's' with 2 electrons. By following the rules of electron configurations such as the Aufbau principle, the Pauli exclusion principle, and Hund's rule, one can determine the electron configuration for a given element.

Isoelectronic Noble Gases

Ions often gain or lose electrons to become isoelectronic with noble gases — meaning they have the same number of electrons and hence an identical electron configuration. Noble gases are elements such as helium (He), neon (Ne), argon (Ar), etc., that have their outermost electron shells fully occupied, which makes them exceptionally stable.

• For example, a Ca^{2+} ion is isoelectronic with the noble gas argon (Ar) as it has 18 electrons, which is the same number as argon.
• Similarly, a K^{+} ion has an electron configuration corresponding to the noble gas argon due to the loss of one electron.
• Anions add electrons to reach the electron configuration of a noble gas, as seen with S^{2-} and Br^{-}, which become isoelectronic with argon and krypton (Kr) respectively.

This characteristic of sharing similar electron configurations with noble gases is what provides ions with chemical stability. During chemical reactions, atoms will often gain or lose electrons until they achieve a noble gas configuration.