Question

Why is the number of common oxidation states for the elements at the beginning and those at the end of the first transition series less than for elements in the middle of the series?

Short answer

The elements at the beginning and end of a series have fewer oxidation states because their 'd' orbitals are either empty or filled, causing 'd' electrons to be tightly bound and not readily participate in bonding. This is different from elements in the middle of the series which have more electrons in their 'd' orbitals, hence, more available electrons for bonding resulting in more common oxidation states.

Step-by-step solution

Understanding Transition Elements

Transition elements are the elements that are located in the d-block of the periodic table. These elements are characterized by their variable oxidation states.

Defining Oxidation States

Oxidation state is a concept in chemistry that allows the counting of electrons transferred to or from a substance in a redox reaction. The oxidation state can be a positive, negative, or zero number.

Explaining Oxidation States in Transition Elements

In a transition series, the number of oxidation states corresponds to the number of electrons in the outer 'd' orbitals that can participate in bond formation. This is why elements in the middle of a transition series typically have more common oxidation states, because they have more 'd' electrons that can be involved in forming bonds.

Explaining Lower Oxidation States at the Beginning and End

Elements at the beginning and end of a transition series have fewer common oxidation states as the 'd' orbital is empty or filled at the start, and at the end of the series, the 'd' electrons are tightly bound and do not readily participate in bonding.

Key concepts

Oxidation States

In chemistry, oxidation states are crucial for understanding how elements interact and bind with one another. The oxidation state of an element can be thought of as the theoretical charge an atom would have if all bonds to it were purely ionic. This concept helps chemists to follow electron movement between atoms during reactions, especially in redox (oxidation-reduction) reactions.

For example, in a compound like \(\text{Fe}_2\text{O}_3\), iron can have an oxidation state of +3, whereas oxygen typically has an oxidation state of -2. The sum of these oxidation states must match the overall charge of the compound, which in neutrals is zero. Transition elements, which we will explore further, can have multiple oxidation states due to the participation of their 'd' electrons in bonding, allowing for a greater variety of chemical compounds.

Periodic Table

The periodic table is not just an arrangement of elements; it is an invaluable tool that reflects an element's chemical behavior based on its position. Transition elements occupy the center of the periodic table in the 'd-block', stretching from groups 3 through 12. These elements bridge the section between the s-block and p-block elements, often showing properties characteristic of metals, such as high melting points and electrical conductivity.

• Location in d-block: They are situated in the middle of the periodic table.
• Variety in oxidation states: The transition metals are notable for exhibiting more than one oxidation state, contrasting with s and p block elements.

As you move across a transition series, each element adds an additional electron to the 'd' orbital. This allows them to form a range of oxidation states and complex ions. The central location on the table correlates with a higher number of oxidation states because the mid-series elements have partially filled 'd' sublevels, providing more electrons available for bonding.

Redox Reactions

Redox reactions are chemical reactions that involve the transfer of electrons between two substances. These reactions can be identified by changes in oxidation states. A simple mnemonic to remember this is "OIL RIG" – Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

In redox reactions involving transition elements, oxidation states change as electrons are transferred. For instance, when \(\text{Fe}^{2+}\) is oxidized to \(\text{Fe}^{3+}\), iron loses an electron. Conversely, a reduction occurs if it gains an electron back.

• Redox reactions are crucial in metabolic processes, batteries, and corrosion.
• They involve reduction (gaining electrons) and oxidation (losing electrons).

Transition elements are particularly significant in redox chemistry due to their variable oxidation states, which allow them to participate in multiple reaction pathways. Their ability to cycle between different states makes them indispensable in numerous industrial processes, such as catalysis.