Question

State the number of electrons lost or gained when the following elements form ions: a. O b. Group \(2 \mathrm{~A}(2)\) c. \(\mathrm{F}\) d. K e. \(\mathrm{Rb}\)

Short answer

a. Gains 2 electronsb. Loses 2 electronsc. Gains 1 electrond. Loses 1 electrone. Loses 1 electron.

Step-by-step solution

Understand the Loss or Gain of Electrons

Elements form ions by either losing or gaining electrons to achieve a stable electron configuration, often aiming for a full valence shell similar to the nearest noble gas.

Determine Electrons for Oxygen (O)

Oxygen is in Group 16. It needs to gain 2 electrons to achieve the electron configuration of neon. Therefore, oxygen gains 2 electrons.

Determine Electrons for Group 2A Elements

Elements in Group 2A (like Mg, Ca) have 2 electrons in their outer shell. They achieve stability by losing these 2 electrons. Therefore, Group 2A elements lose 2 electrons.

Determine Electrons for Fluorine (F)

Fluorine is in Group 17. It needs to gain 1 electron to achieve the electron configuration of neon. Therefore, fluorine gains 1 electron.

Determine Electrons for Potassium (K)

Potassium is in Group 1. It needs to lose 1 electron to achieve the electron configuration of argon. Therefore, potassium loses 1 electron.

Determine Electrons for Rubidium (Rb)

Rubidium is also in Group 1. Similar to potassium, it needs to lose 1 electron to achieve the electron configuration of krypton. Therefore, rubidium loses 1 electron.

Key concepts

Electron Configuration

Electron configuration refers to the arrangement of electrons in an atom's orbitals. It tells us where each electron is most likely to be found and helps predict how atoms will bond and react.
Each element has a unique electron configuration that can be determined using the periodic table. The electrons are arranged in shells and subshells around the nucleus and follow the Pauli exclusion principle, Hund's rule, and the Aufbau principle.
For example, oxygen has an atomic number of 8, meaning it has 8 electrons. Its electron configuration is \(1s^2 2s^2 2p^4\). This shows there are 2 electrons in the first shell and 6 in the second shell.

Ion Formation

Ion formation occurs when atoms gain or lose electrons to achieve a stable electron configuration, usually similar to the nearest noble gas.
Atoms can form either cations (positively charged ions) by losing electrons or anions (negatively charged ions) by gaining electrons.

• For example, oxygen (O) from the exercise gains 2 electrons to become \(\text{O}^{2-}\).
• Potassium (K) loses 1 electron to become \(\text{K}^+\).

In both cases, the elements achieve a more stable, noble gas-like configuration.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. These electrons play a crucial role in chemical bonding and reactions.
Elements in the same group on the periodic table have the same number of valence electrons, which is why they exhibit similar chemical properties.
For example, elements in Group 2A, such as magnesium (Mg) and calcium (Ca), each have 2 valence electrons. They tend to lose these 2 electrons to form cations with a \(2+\) charge, achieving a stable electron configuration.

Noble Gas Configuration

A noble gas configuration is when an atom has a full valence shell similar to that of the noble gases (Group 18 elements). This is associated with maximum stability.
Noble gases have complete electron shells and are very stable. Other elements often lose or gain electrons to achieve this stable arrangement.
For instance, fluorine (F) has 7 valence electrons. By gaining 1 electron, it achieves the electron configuration of neon \(\text{Ne}\), making it stable.

Periodic Table Groups

The periodic table groups (or families) vertically organize elements with similar chemical properties. These groups are numbered from 1 to 18.
Each group has a characteristic number of valence electrons. For example:

• Group 1A (e.g., Li, Na, K) has 1 valence electron and tends to lose it, forming a \(1+\) ion.
• Group 17 (e.g., F, Cl), has 7 valence electrons and typically gains 1 electron to form a \(1-\) ion.

Understanding the properties of different groups helps predict the ion formation and reactivity of elements.