Question

For each of the following oxidation-reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced. a. \(3 \mathrm{Zn}(s)+\mathrm{N}_{2}(g) \rightarrow \mathrm{Zn}_{3} \mathrm{N}_{2}(s)\) b. \(\operatorname{Co}(s)+S(s) \rightarrow \cos (s)\) c. \(4 \mathrm{K}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{K}_{2} \mathrm{O}(s)\) d. \(4 \mathrm{Ag}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Ag}_{2} \mathrm{O}(s)\)

Short answer

a. Zn is oxidized, N is reduced. b. Co is oxidized, S is reduced. c. K is oxidized, O is reduced. d. Ag is oxidized, O is reduced.

Step-by-step solution

Identify the oxidation states of each element in the reactions

To do this, we'll look at the initial and final states of each element in the reactions. a. Zn(s) has an oxidation state of 0, while N2(g) has an oxidation state of 0 for each N atom. In Zn3N2(s), Zn has an oxidation state of +2 and N has an oxidation state of -3. b. Co(s) has an oxidation state of 0, while S(s) has an oxidation state of 0. In CoS(s), Co has an oxidation state of +2 and S has an oxidation state of -2. c. K(s) has an oxidation state of 0, while O2(g) has an oxidation state of 0 for each O atom. In K2O(s), K has an oxidation state of +1 and O has an oxidation state of -2. d. Ag(s) has an oxidation state of 0, while O2(g) has an oxidation state of 0 for each O atom. In Ag2O(s), Ag has an oxidation state of +1 and O has an oxidation state of -2.

Determine which elements are oxidized and reduced

Now, we'll compare the oxidation states of the elements before and after the reactions. a. Zn's oxidation state changes from 0 to +2, meaning it loses electrons - so Zn is oxidized. N's oxidation state changes from 0 to -3, meaning it gains electrons - so N is reduced. b. Co's oxidation state changes from 0 to +2, meaning it loses electrons - so Co is oxidized. S's oxidation state changes from 0 to -2, meaning it gains electrons - so S is reduced. c. K's oxidation state changes from 0 to +1, meaning it loses electrons - so K is oxidized. O's oxidation state changes from 0 to -2, meaning it gains electrons - so O is reduced. d. Ag's oxidation state changes from 0 to +1, meaning it loses electrons - so Ag is oxidized. O's oxidation state changes from 0 to -2, meaning it gains electrons - so O is reduced.

Results

By analyzing the changes in oxidation states for each reaction, we obtained the following results: a. Zn is oxidized, and N is reduced. b. Co is oxidized, and S is reduced. c. K is oxidized, and O is reduced. d. Ag is oxidized, and O is reduced.

Key concepts

Oxidation and Reduction

Understanding redox reactions, which stands for reduction-oxidation reactions, is key to grasping the fundamental processes in chemistry. These reactions involve a transfer of electrons between substances. Importantly, oxidation refers to the loss of electrons, and as a result, an increase in the oxidation state of an element. Conversely, reduction is the gain of electrons, leading to a decrease in the oxidation state.

For example, when zinc (Zn) reacts with nitrogen (N₂) to form zinc nitride (Zn₃N₂), zinc atoms lose electrons and are oxidized. On the flip side, nitrogen molecules gain electrons and are reduced.

To make the redox process clearer, consider this handy tip: think of oxidation as 'loss of electrons' or 'gain of oxygen' and reduction as 'gain of electrons' or 'loss of oxygen'. This can be remembered through the mnemonic 'OIL RIG' - Oxidation Is Loss, Reduction Is Gain. Redox reactions are fundamental in many natural and industrial processes, including cellular respiration, metal corrosion, and the operation of batteries.

Oxidation States

The oxidation state, or oxidation number, is a concept that provides insights into the hypothetical charge that an atom would have if all bonds to atoms of different elements were completely ionic. Oxidation states play a crucial role in redox reactions as they help determine which substances are oxidized and which are reduced.

For an atom in its elemental form, like zinc (Zn) or oxygen (O₂), the oxidation state is always zero. However, when these atoms form compounds, their oxidation states change due to electron transfer. For instance, in the reaction of potassium (K) with oxygen, potassium has an initial oxidation state of zero and changes to +1 in the compound potassium oxide (K₂O).

Here's a simple approach to determining oxidation states: use known oxidation states of elements, like hydrogen (+1) and oxygen (-2), as references. For neutral compounds, the sum of oxidation states equals zero, and for polyatomic ions, the sum equals the ion's charge. By tracking these changes, we can decipher which elements undergo oxidation or reduction in chemical reactions.

Chemical Reactions

Chemical reactions are processes that result in the transformation of one set of chemical substances to another. Classified into various types such as synthesis, decomposition, single displacement, and double displacement; redox reactions form a category where electron transfer occurs, leading to changes in oxidation states.

In redox reactions, like the synthesis of silver oxide from silver and oxygen, there's a neat interplay of the oxidation and reduction processes. The elemental form of silver (Ag) and oxygen (O₂) react to produce silver oxide (Ag₂O), signifying a redox reaction where silver is oxidized and oxygen is reduced.

To consolidate understanding of these reactions, it helps to identify the reactants (starting materials) and the products (substances formed). Balancing redox reactions can be complex and often requires breaking them down into half-reactions - one for oxidation and one for reduction. This technique helps in visualizing the electron transfer and ensuring that mass and charge are conserved throughout the reaction.