Question

What element is oxidized in the following equation and what element is reduced? $$ \mathrm{Sn}^{2+}+2 \mathrm{Ag} \longrightarrow \mathrm{Sn}+2 \mathrm{Ag}^{+} $$

Short answer

Ag is oxidized, Sn^{2+} is reduced.

Step-by-step solution

Identify the Oxidation States

Determine the oxidation states of each element in the reaction. \( \mathrm{Sn}^{2+} \) has an oxidation state of +2, \( \mathrm{Ag} \) has an oxidation state of 0, \( \mathrm{Sn} \) is 0, and \( \mathrm{Ag}^{+} \) is +1.

Determine Oxidation Changes

Compare the oxidation states on both sides of the equation. \( \mathrm{Sn}^{2+} \) goes from +2 to 0, indicating a gain of electrons (reduction). \( \mathrm{Ag} \) goes from 0 to +1, indicating a loss of electrons (oxidation).

Identify Oxidized Element

Since \( \mathrm{Ag} \) transitions from an oxidation state of 0 to +1, it loses electrons. Therefore, \( \mathrm{Ag} \) is the element that is oxidized.

Identify Reduced Element

\( \mathrm{Sn}^{2+} \) transitions from an oxidation state of +2 to 0, gaining electrons. This means that \( \mathrm{Sn}^{2+} \) is the element that is reduced.

Key concepts

Oxidation States

In chemistry, the concept of oxidation states allows us to understand how electrons are distributed among atoms in a chemical reaction. Each element in a chemical reaction can gain or lose electrons, which changes its oxidation state. This change provides a vital clue in identifying oxidation-reduction reactions.

In the given reaction \[\text{Sn}^{2+} + 2 \text{Ag} \rightarrow \text{Sn} + 2 \text{Ag}^{+}\]we need to determine the oxidation states of the participating elements. Initially, the oxidation state of \( \text{Sn}^{2+} \) is +2, while the oxidation state of \( \text{Ag} \) is 0. After the reaction, the oxidation state of \( \text{Sn} \) becomes 0, and \( \text{Ag}^{+} \) becomes +1.

• \( \text{Sn}^{2+} \): transitions from +2 to 0
• \( \text{Ag} \): transitions from 0 to +1

These changes in oxidation states signify a transfer of electrons and form the basis for analyzing which elements are oxidized and reduced.

Electron Transfer

Electron transfer is at the heart of oxidation-reduction reactions, often abbreviated as redox reactions. During a redox reaction, certain elements will gain electrons, while others will lose them. This electron transfer results in the change in oxidation states.

In the example reaction: \[\text{Sn}^{2+} + 2 \text{Ag} \rightarrow \text{Sn} + 2 \text{Ag}^{+}\]electrons are shuffled between the elements. Each \(\text{Sn}^{2+}\) ion gains two electrons, reducing its oxidation state from +2 to 0. This change signifies that \(\text{Sn}^{2+}\) is reduced. Conversely, each \(\text{Ag}\) atom loses one electron, resulting in its oxidation state change from 0 to +1, which means \(\text{Ag}\) is oxidized.

• Reduction: gain of electron(s), decrease in oxidation state
• Oxidation: loss of electron(s), increase in oxidation state

Redox Processes

Redox processes involve tandem reactions where two essential parts work simultaneously: oxidation and reduction. One element's electron loss (oxidation) coincides with another element's electron gain (reduction). This dual process maintains the balance in a chemical reaction and is why these are called oxidation-reduction reactions, or redox for short.

In the equation \[\text{Sn}^{2+} + 2 \text{Ag} \rightarrow \text{Sn} + 2 \text{Ag}^{+}\]the silver \(\text{Ag}\) atoms are oxidized as they lose electrons and increase in oxidation state from 0 to +1. Meanwhile, \(\text{Sn}^{2+}\) ions are reduced, as they gain electrons and see their oxidation state fall from +2 to 0. It's this electron exchange that drives the redox process, acting as a crucial method in understanding many chemical and biological systems.

To sum up, remember the key components:

• Oxidation: increase in oxidation state (loss of electrons)
• Reduction: decrease in oxidation state (gain of electrons)
• Redox reactions require both processes to occur simultaneously