Question

Use the periodic table, but not Table \(2.2,\) to write the (b) Br, (c) \(\mathrm{Mg},\) (d) \(\mathrm{S},\) and symbols for the ions of (a) \(\mathrm{K}\) (e) Al.

Short answer

The symbols for the ions are (a) \text{K}^+, (b) \text{Br}^-, (c) \text{Mg}^{2+}, (d) \text{S}^{2-}, (e) \text{Al}^{3+}.

Step-by-step solution

- Identify the Charge of Potassium (K)

Using the periodic table, locate Potassium (K), which is in group 1. Atoms in this group typically have a charge of +1 when they form ions, as they tend to lose one electron to achieve a full outer shell.

- Identify the Charge of Bromine (Br)

Locate Bromine (Br) on the periodic table, which is in group 17. Atoms in this group typically have a charge of -1 when they form ions, as they tend to gain one electron to complete their valence shell.

- Identify the Charge of Magnesium (Mg)

Find Magnesium (Mg) in group 2 of the periodic table. Atoms in this group usually have a charge of +2 as ions because they lose two electrons to achieve a stable electron configuration.

- Identify the Charge of Sulfur (S)

Locate Sulfur (S) on the periodic table, which can be found in group 16. Atoms in this group often have a charge of -2 when they form ions, as they tend to gain two electrons to fill their outer electron shell.

- Identify the Charge of Aluminum (Al)

On the periodic table, find Aluminum (Al), which is in group 13. Atoms in this group typically form ions with a charge of +3, due to the loss of three electrons for stability.

- Write the Symbols for the Resulting Ions

Write the symbols for the ions of each element with their respective charges. For positive ions (cations), write the element symbol followed by the charge; for negative ions (anions), do the same.

Key concepts

Understanding the Periodic Table

The periodic table is an essential tool in chemistry, allowing scientists and students alike to quickly understand the properties of the elements and predict how they will react with one another. It is organized into rows called periods and columns called groups. The elements are arranged based on increasing atomic number — the number of protons in an atom's nucleus.

As one moves across a period, the properties of elements gradually change due to the increasing number of protons and electrons. In comparison, moving down a group illustrates a consistency in chemical properties. This regularity arises because the elements in a group have the same number of electrons in their outermost shell, which governs their bonding behavior.

For instance, in the exercise, potassium (K) in group 1 is known for losing one electron to form a +1 ion, while bromine (Br) in group 17 gains an electron resulting in a -1 ion. Understanding both the layout and the common properties of the elements in each group is crucial for predicting ion formation and reactiveness.

The Charge of Ions

Ions are atoms or molecules that have gained or lost electrons, giving them a net positive or negative charge. The formation of ions is driven by an element's desire to achieve a stable electron configuration, often resembling the nearest noble gas. This involves either losing electrons to reduce the energy level of the valence shell or gaining electrons to fill it.

For example, as seen in the exercise, magnesium (Mg), located in group 2, tends to lose two electrons to achieve the electron configuration of neon, resulting in a +2 charge. On the other hand, sulfur (S), in group 16, tends to gain two electrons to achieve the electron configuration of argon, ending with a -2 charge. The ability to predict the charge of ions helps in understanding the nature of the bonds they will form and the compounds they could produce.

Electron Configuration and Ion Formation

Electron configuration describes the distribution of electrons of an atom or molecule in atomic or molecular orbitals. The most stable state of an atom is when it has a full outer shell of electrons, similar to the noble gases. Elements achieve this stable state through the loss or gain of electrons which results in ion formation.

Groups in the periodic table often share similar electron configurations, especially in their valence shells, which determines their ion charge. For instance, aluminum (Al) which is in group 13, has three electrons in its outer shell and loses them to form an Al³⁺ cation as demonstrated in the exercise. Visualizing and understanding electron configurations are essential in predicting how an element will interact chemically, including the specific ions they will form and their resulting charge.