Question

Each of the following elements is capable of forming an ion in chemical reactions. By referring to the periodic table, predict (b) \(\mathrm{Al},(\mathbf{c}) \mathrm{K}\) the charge of the most stable ion of each: (a) \(\mathrm{Mg}\), (d) S, (e) \(\mathrm{F}\).

Short answer

The most stable ions for the given elements are: \(\mathrm{Mg^{2+}}\), \(\mathrm{Al^{3+}}\), \(\mathrm{K^{+}}\), \(\mathrm{S^{2-}}\), and \(\mathrm{F^{-}}\).

Step-by-step solution

Examine Mg (Magnesium)

Magnesium is in group 2 of the periodic table. Elements from group 2 have 2 valence electrons, and to achieve a full outer shell, they must lose those 2 electrons. Therefore, Mg forms a +2 ion: \(\mathrm{Mg^{2+}}\).

Examine Al (Aluminum)

Aluminum is in group 13 (3A) of the periodic table. Elements in group 13 have 3 valence electrons, and to achieve a full outer shell, they must lose those 3 electrons. Thus, Al forms a +3 ion: \(\mathrm{Al^{3+}}\).

Examine K (Potassium)

Potassium is in group 1 of the periodic table. Elements in group 1 have 1 valence electron, and to achieve a full outer shell, they must lose that one electron. Hence, K forms a +1 ion: \(\mathrm{K^{+}}\).

Examine S (Sulfur)

Sulfur is in group 16 (6A) of the periodic table. Elements of this group have 6 valence electrons. To achieve a full outer shell with 8 electrons, they must gain two electrons. In this case, S will form a -2 ion: \(\mathrm{S^{2-}}\).

Examine F (Fluorine)

Fluorine is in group 17 (7A) of the periodic table. Elements of group 17 have 7 valence electrons. To achieve a full outer shell with 8 electrons, they must gain one electron. Therefore, F forms a -1 ion: \(\mathrm{F^{-}}\). Based on the information from the periodic table, we can predict the most stable ions for the given elements as follows: Mg forms \(\mathrm{Mg^{2+}}\), Al forms \(\mathrm{Al^{3+}}\), K forms \(\mathrm{K^{+}}\), S forms \(\mathrm{S^{2-}}\), and F forms \(\mathrm{F^{-}}\).

Key concepts

Group 1 elements

Group 1 elements, known as the alkali metals, are found in the first column of the periodic table. These elements include

• Lithium (Li)
• Sodium (Na)
• Potassium (K)
• Rubidium (Rb)
• Cesium (Cs)
• Francium (Fr)

One of the distinctive features of group 1 elements is that they all have one valence electron. This lone electron is in their outermost shell, making them highly reactive.

To achieve stability, alkali metals often form ions by losing this single electron, resulting in a "+1" charge. For example, when Potassium (K) loses its one valence electron, it forms a potassium ion \(\mathrm{K^{+}}\).This ionization process provides them a full outer electron shell, typically akin to the nearest noble gas configuration.

Group 2 elements

The Group 2 elements, called alkaline earth metals, are positioned in the second column of the periodic table. Some common elements in this group are:

• Beryllium (Be)
• Magnesium (Mg)
• Calcium (Ca)
• Strontium (Sr)
• Barium (Ba)
• Radium (Ra)

Each of these elements has two valence electrons. These electrons reside in their outermost shell, rendering them reactive but less so than the alkali metals of Group 1.

Group 2 elements tend to form ions by losing both of their valence electrons. This results in an ion with a "+2" charge. For instance, Magnesium (Mg) loses its two valence electrons to become \(\mathrm{Mg^{2+}}\).This ionic form allows them to achieve a stable electronic configuration, similar to the nearest noble gas.

Valence electrons

Valence electrons are the electrons located in the outermost electron shell of an atom. These electrons play a crucial role in chemical bonding because they are involved in the formation of chemical bonds.

The number of valence electrons determines an element's reactivity and the types of ions or bonds it can form. For instance:

• Elements in Group 1 have one valence electron
  
• Elements in Group 2 have two valence electrons
• Elements in Group 17 have seven valence electrons

Understanding valence electrons is essential because:

• They dictate how elements interact with each other
• They help predict the charge of ions
• They determine the number and type of bonds an element can form

This knowledge is fundamental in predicting chemical reactions and creating new compounds.

Chemical ion formation

Chemical ion formation is a process where atoms gain or lose electrons to attain stable electron configurations, similar to those of noble gases. This process significantly influences chemical reactions and compound formation.

Atoms form ions to achieve an outer electron shell that emulates the nearest noble gas. The noble gases are stable because they have full valence shells. By gaining or losing electrons, other atoms achieve this stability:

• Metals, like those in Groups 1 and 2, lose electrons to form positive ions (cations).
• Non-metals, like those in Groups 16 and 17, gain electrons to form negative ions (anions).

For example, a sulfur atom (S) from Group 16 gains two electrons to form a sulfide ion:
\(\mathrm{S^{2-}}\).

Understanding ion formation is crucial for predicting how elements will interact in chemical reactions and determining the properties of the resulting compounds.