Question

Which element is oxidized and which is reduced in the follow- ing reactions? $$ \begin{array}{l} \text { (a) } \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g) \\ \text { (b) } 3 \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Al}(s) \longrightarrow \\ \text { (c) } \mathrm{Cl}_{2}(a q)+2 \mathrm{NaI}(a q) \longrightarrow \mathrm{I}_{2}(a q)+2 \mathrm{NaCl}(a q) \\ \text { (d) } \mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l) \end{array} $$

Short answer

In the given reactions: (a) Nitrogen is reduced (0 to -3), and Hydrogen is oxidized (+1 to +1, no change, but participates). (b) Iron is oxidized (+2 to a higher state), and Aluminum is reduced (0 to a lower state). (c) Chlorine is reduced (0 to -1), and Iodine is oxidized (-1 to 0). (d) Sulfur is oxidized (-2 to +6), and Oxygen in H_2O_2 is reduced (-1 to -2). Lead is not oxidized/reduced.

Step-by-step solution

Determine the oxidation states of each element

For each of the given reactions, we need to assign oxidation states to the elements on both the reactant and product sides: Reaction (a): \(N_2(g) + 3H_2(g) \rightarrow 2NH_3(g)\) Oxidation states: N: 0, H: +1, N: -3, H: +1 Reaction (b): \(3Fe(NO_3)_2(aq) + 2Al(s) \rightarrow \) Oxidation states: Fe: +2, N: +5, O: -2, Al: 0 Reaction (c): \(Cl_2(aq) + 2NaI(aq) \rightarrow I_2(aq) + 2NaCl(aq)\) Oxidation states: Cl: 0, Na: +1, I: -1, I: 0, Na: +1, Cl: -1 Reaction (d): \(PbS(s) + 4H_2O_2(aq) \rightarrow PbSO_4(s) + 4H_2O(l)\) Oxidation states: Pb: +2, S: -2, H: +1, O: -1, Pb: +2, S: +6, O: -2, H: +1, O: -2

Identify elements oxidized and reduced in each reaction

Now that we have the oxidation states for each reaction, we can identify the elements that are oxidized and reduced: Reaction (a): Nitrogen is reduced (oxidation state changes from 0 to -3) and Hydrogen is oxidized (oxidation state changes from +1 to +1, no change in oxidation state, but still participating in the reaction). Reaction (b): Iron is oxidized (oxidation state changes from +2 to an unknown higher state) and Aluminum is reduced (oxidation state changes from 0 to an unknown lower state). Reaction (c): Chlorine is reduced (oxidation state changes from 0 to -1) and Iodine is oxidized (oxidation state changes from -1 to 0). Reaction (d): Lead is neither oxidized nor reduced (oxidation state remains +2), while Sulfur is oxidized (oxidation state changes from -2 to +6), and Oxygen in hydrogen peroxide (H_2O_2) is reduced (oxidation state changes from -1 to -2).

Key concepts

Oxidation States Determination

Understanding oxidation states is crucial for analyzing chemical reactions, particularly redox reactions. Oxidation states (also known as oxidation numbers) can be considered as the 'bookkeeping' tool that allows chemists to keep track of electron transfer between atoms in a molecule or compound.

Determining oxidation states involves a set of rules: free elements have an oxidation state of zero; the oxidation state of an atom in a monatomic ion is equal to the charge of the ion; and certain elements have specific common oxidation states (e.g., oxygen is usually -2, hydrogen is usually +1). In covalent compounds, the less electronegative atom is assigned positive oxidation state, and the more electronegative atom is assigned a negative oxidation state.

For example, in the reaction \( NH_{3}(g) \) nitrogen moves from an oxidation state of 0 in \({N_{2}}\) to -3 in \({NH_{3}}\), indicating it gains electrons, hence is reduced. Applying these rules helps us identify which elements are oxidized or reduced, paving the way for a deeper understanding of the reaction's mechanism.

Redox Reactions

Redox reactions are types of chemical reactions that involve a change in oxidation states of the involved atoms through the transfer of electrons. The term 'redox' is a shorthand for reduction-oxidation reactions. In these processes, one species undergoes oxidation (loss of electrons) while another undergoes reduction (gain of electrons).

To properly identify the changes, we look at the oxidation states before and after the reaction. For instance, in the given exercise, chlorine goes from an oxidation state of 0 in \(Cl_{2}\) to -1 in \({NaCl}\), thus it is reduced. Meanwhile, iodine is oxidized as its oxidation state increases from -1 in \(NaI\) to 0 in \({I_{2}}\).

Recognizing the elements that are oxidized or reduced allows us to predict the movement of electrons and provide insight into the energetics and directionality of the reaction. Redox reactions are fundamental to numerous natural processes and industrial applications, such as energy production in batteries and the metabolism of food in biological systems.

Chemical Reactions Analysis

The analysis of chemical reactions extends beyond the identification of substances that are oxidized or reduced; it encompasses the overall transformation of reactants into products, the stoichiometry, and the reaction conditions. This systematic study lies at the heart of understanding and controlling chemical processes.

Reaction analysis can involve determining reaction rates, equilibrium constants, and energetics (thermodynamics) as well as understanding the reaction mechanism or the step-by-step sequence of elementary steps that lead to product formation. For instance, in reaction \(d\), \(PbS\) reacts with hydrogen peroxide to form \(PbSO_{4}\) and water. This reaction indicates that sulfur is oxidized from -2 to +6, while the oxygen in hydrogen peroxide is reduced, showcasing the conservation of charge in isolated systems.

By meticulously analyzing each aspect of the reaction, chemists can predict outcomes, design new reactions for synthesis, and optimize conditions for large-scale industrial processes.