Question

Predict the ion formed by each element. (a) \(\mathrm{Rb}\) (b) \(\mathrm{K}\) (c) \(\mathrm{Al}\) (d) \(\mathrm{O}\)

Short answer

The predicted ions are: (a) \text{Rb}^+, (b) \text{K}^+, (c) \text{Al}^{3+}, (d) \text{O}^{2-}.

Step-by-step solution

Identify Group on Periodic Table

Locate each element on the periodic table to determine which group it belongs to. The group number will help predict the charge of the ion it forms.

Predict the Charge Based on Group

For elements in Group 1 (alkali metals), such as Rb (rubidium) and K (potassium), they typically lose one electron to form a +1 charge. For elements in Group 13, like Al (aluminum), they tend to lose three electrons to form a +3 charge. Nonmetals like oxygen, which is in Group 16, tend to gain two electrons to reach a noble gas electron configuration, resulting in a -2 charge.

Write Down the Predicted Ions

Based on the expected charges from Step 2, write the symbol for the element followed by its ionic charge as a superscript. Rb will form Rb^+, K will form K^+, Al will form Al^3+, and O will form O^2-.

Key concepts

Understanding Periodic Table Groups

When it comes to predicting ion charges, the layout of the periodic table is incredibly insightful. Elements are arranged in columns known as 'groups' which reflect how they react chemically. Generally, elements within the same group form ions with the same charge due to their similar electron configurations.

Since elements tend to adopt the stable electron configuration of the nearest noble gas, the group number is a handy predictor of ion charge. Groups are numbered 1-18, and as we move across the table from left to right, the elements' inclination to lose electrons decreases while their tendency to gain electrons increases. In simple terms, metals located in groups 1-3 typically lose electrons and form positive ions, while non-metals in groups 15-17 usually gain electrons and become negatively charged ions.

Alkali Metals & Their Ion Formation

The alkali metals, which occupy Group 1 of the periodic table, are renowned for being highly reactive and have a specific trait: they all have one electron in their outermost shell. This singular electron is relatively loosely bound, making alkali metals eager to lose it in chemical reactions, which results in the formation of ions with a +1 charge.

For instance, rubidium (Rb) and potassium (K) are both alkali metals. Following the pattern, they tend to lose their one valence electron to achieve a more stable configuration, thus forming Rb^+ and K^+, respectively. An important note for students to remember is that despite their reactivity, the predictability of the ion formation for alkali metals is straightforward owing to their consistent +1 ionic charge across the group.

Electron Configuration and Ion Charges

Electron configuration describes the distribution of electrons in an atom's shells and subshells. This arrangement is pivotal in determining how an element behaves in chemical reactions, particularly in the gain and loss of electrons to form ions.

The aim for most atoms is to achieve a noble gas electron configuration, which is a stable and energetically favorable state. For example, aluminium (Al), from Group 13, has three electrons in its outer shell. It tends to lose all three to emulate the noble gas preceding it, resulting in a +3 charge (Al^3+). In contrast, oxygen (O) from Group 16 has six valence electrons and needs two more to fill its outer shell, as a consequence, it often gains two electrons, forming an O^2- ion.

An easy tip for students: The electron configuration not only helps to predict the charge of the ions but also lends insight into the chemical properties and reactivity of an element.