Question

Elements in group \(7 \mathrm{~A}\) in the periodic table are the halogens; elements in group \(6 \mathrm{~A}\) are called the chalcogens. (a) What is the most common oxidation state of the chalcogens compared to the halogens? Can you suggest an explanation for the difference? (b) For each of the following periodic properties, state whether the halogens or the chalcogens have larger values: atomic radii; ionic radii of the most common oxidation state; first ionization energy; second ionization energy.

Short answer

(a) The most common oxidation state of chalcogens is -2, while for halogens, it is -1. The difference is due to the number of valence electrons in their outer shell, with chalcogens needing to gain two electrons and halogens needing one to achieve a stable noble gas electron configuration. (b) Chalcogens have larger atomic radii and ionic radii in their most common oxidation state (-2) compared to halogens. Halogens have larger first and second ionization energies than chalcogens due to the increasing effective nuclear charge across a period.

Step-by-step solution

Part (a): Identifying the most common oxidation states

Chalcogens are in group 6A, which means they have six valence electrons. To achieve a stable noble gas electron configuration, they can gain two electrons, giving them a common oxidation state of -2. Halogens are in group 7A, so they have seven valence electrons. To achieve a stable noble gas electron configuration, they just need to gain one electron, which results in a common oxidation state of -1.

Part (a): Explanation for the difference in oxidation states

The difference in the most common oxidation state between chalcogens and halogens boils down to the number of valence electrons in their outer shell. Chalcogens have six valence electrons, so they gain two electrons to have a full octet, achieving an oxidation state of -2. Halogens, on the other hand, have seven valence electrons and need only one more to reach the full octet, resulting in an oxidation state of -1.

Part (b): Comparing atomic radii

In general, atomic radii increase going down a group and decrease going across a period. Since chalcogens are to the left of halogens in the periodic table, they will generally have larger atomic radii compared to halogens.

Part (b): Comparing ionic radii of the most common oxidation state

When an atom gains electrons, it forms a negatively charged ion. In their common oxidation states, both chalcogens and halogens form negative ions, with chalcogens forming -2 ions and halogens forming -1 ions. In general, ions with more negative charges will have larger ionic radii due to increased electron repulsion. Therefore, the ionic radii of chalcogens in their most common oxidation state will be larger compared to that of halogens.

Part (b): Comparing first ionization energies

Ionization energy is the energy required to remove an electron from an atom. As you move across a period from left to right, ionization energy generally increases due to an increasing effective nuclear charge. Therefore, halogens will typically have larger first ionization energies compared to chalcogens.

Part (b): Comparing second ionization energies

Second ionization energy is the energy required to remove a second electron from an atom. Similar to the first ionization energy, the second ionization energy will be higher for halogens than chalcogens due to the increasing effective nuclear charge across a period.

Key concepts

Oxidation States

The oxidation state is an indication of the degree of oxidation (loss of electrons) of an atom in a chemical compound. For chalcogens (Group 6A elements), the most common oxidation state is -2. This is because chalcogens have six valence electrons and need two more to complete their outer electron shell, achieving a stable, noble gas configuration.
Halogens (Group 7A elements), on the other hand, have seven valence electrons. They readily gain one electron to attain the full octet, resulting in their common oxidation state of -1. The difference in oxidation states between these groups arises from their electron configuration and the number of electrons required to fill their outer shells.

Chalcogens

Chalcogens are the elements found in Group 6A of the periodic table. This group includes Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), and Polonium (Po).

• They have six valence electrons, needing two additional ones to reach a stable state.
• Chalcogens form compounds with many metals and nonmetals.
• Their most common oxidation state is -2, as they gain electrons to achieve a full octet.

Chalcogens possess varying properties such as increasing atomic number and mass, with each having unique applications and roles in chemistry and biology.

Halogens

Halogens are the elements in Group 7A of the periodic table, including Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At).

• With seven valence electrons, they need only one more to complete their octet.
• Halogens typically form diatomic molecules, such as \( ext{Cl}_2 \) and \( ext{I}_2 \).
• They are highly reactive, particularly with alkali metals and alkaline earth metals, to form salts.

Their familiar usage spans industries and household products, including disinfectants and bleach, due to their potent reactivity.

Periodic Trends

Periodic trends refer to patterns within the periodic table that highlight properties of elements. These trends help us understand the behavior and characteristics of elements more intuitively.

• **Atomic Radii**: As you move down a group, atomic radii increase. Across a period from left to right, atomic radii decrease. Therefore, chalcogens, being to the left of halogens, generally have larger atomic radii.
• **Ionic Radii**: For the most common oxidation states, chalcogens form -2 ions while halogens form -1 ions, meaning chalcogen ions are larger due to their increased electron count and repulsion.
• **Ionization Energy**: Generally increasing across a period, halogens feature higher first and second ionization energies than chalcogens.

Understanding these trends enhances our ability to predict the properties and reactivities of elements, aiding significantly in the study of chemical interactions.