Question

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (d) \(\mathrm{NH}_{3}(a q)+\mathrm{H}^{+}(a q) \longrightarrow\) (e) \(\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow\) Which ones of these are redox reactions?

Short answer

Balanced equations: (a) \(\mathrm{Mg_3N_2} + 6\mathrm{H_2O} \longrightarrow 3\mathrm{Mg(OH)_2} + 2\mathrm{NH_3}\) - Redox (b) \(\mathrm{2NO} + \mathrm{O_2} \longrightarrow 2\mathrm{NO_2}\) - Redox (c) \(\mathrm{N_2O_5} + \mathrm{H_2O} \longrightarrow 2\mathrm{HNO_3}\) - Not redox (d) \(\mathrm{NH_3} + \mathrm{H^+} \longrightarrow \mathrm{NH_4^+}\) - Redox (e) \(\mathrm{2N_2H_4} + \mathrm{O_2} \longrightarrow 4\mathrm{N_2} + 8\mathrm{H_2O}\) - Redox

Step-by-step solution

Equation (a) Balancing

First, we identify that the products are magnesium hydroxide (\(\mathrm{Mg(OH)_2}\)) and ammonia (\(\mathrm{NH_3}\)). Now we balance the equation: \[\mathrm{Mg_3N_2} + 6\mathrm{H_2O} \longrightarrow 3\mathrm{Mg(OH)_2} + 2\mathrm{NH_3}\]

Equation (b) Balancing

We see that the product is nitrogen dioxide (\(\mathrm{NO_2}\)). Balancing the equation, we get: \[\mathrm{2NO} + \mathrm{O_2} \longrightarrow 2\mathrm{NO_2}\]

Equation (c) Balancing

The products are nitric acid (\(\mathrm{HNO_3}\)) and water. Balancing the equation: \[\mathrm{N_2O_5} + \mathrm{H_2O} \longrightarrow 2\mathrm{HNO_3}\]

Equation (d) Balancing

The product is the ammonium ion (\(\mathrm{NH_4^+}\)). Balancing the equation: \[\mathrm{NH_3} + \mathrm{H^+} \longrightarrow \mathrm{NH_4^+}\]

Equation (e) Balancing

The products are nitrogen gas (\(\mathrm{N_2}\)) and water. Balancing the equation: \[\mathrm{2N_2H_4} + \mathrm{O_2} \longrightarrow 4\mathrm{N_2} + 8\mathrm{H_2O}\]

Identifying Redox Reactions

(a) \(\mathrm{Mg^{2+}}\) and \(\mathrm{N^{3-}}\) are reduced to \(\mathrm{Mg}\) and \(\mathrm{N}\). (b) \(\mathrm{N}\) is oxidized (from \(\mathrm{+2}\) to \(\mathrm{+4}\)) and \(\mathrm{O}\) is reduced (from \(\mathrm{0}\) to \(\mathrm{-2}\)). (c) \(\mathrm{N}\) has the same oxidation state in both reactants and products (\(\mathrm{+5}\)). (d) \(\mathrm{N}\) is reduced (from \(\mathrm{-3}\) to \(\mathrm{-2}\)). (e) Both \(\mathrm{N}\) and \(\mathrm{O}\) undergo changes in their oxidation states.

Results

After balancing and identifying redox reactions: (a) \(\mathrm{Mg_3N_2} + 6\mathrm{H_2O} \longrightarrow 3\mathrm{Mg(OH)_2} + 2\mathrm{NH_3}\) - Redox (b) \(\mathrm{2NO} + \mathrm{O_2} \longrightarrow 2\mathrm{NO_2}\) - Redox (c) \(\mathrm{N_2O_5} + \mathrm{H_2O} \longrightarrow 2\mathrm{HNO_3}\) - Not redox (d) \(\mathrm{NH_3} + \mathrm{H^+} \longrightarrow \mathrm{NH_4^+}\) - Redox (e) \(\mathrm{2N_2H_4} + \mathrm{O_2} \longrightarrow 4\mathrm{N_2} + 8\mathrm{H_2O}\) - Redox

Key concepts

Redox Reactions

Redox reactions, or reduction-oxidation reactions, are processes where the oxidation states of atoms are changed due to the transfer of electrons. In these reactions, one species undergoes oxidation (loses electrons) while another undergoes reduction (gains electrons). Understanding redox reactions is crucial because they are prominent in numerous chemical processes, including combustion, corrosion, and in biological systems. For instance, equation (b) in our exercise, \(\mathrm{2NO} + \mathrm{O_2} \longrightarrow 2\mathrm{NO_2}\), is a classic example of a redox reaction where nitrogen in \(\mathrm{NO}\) is oxidized from a +2 to a +4 oxidation state, whereas oxygen is reduced from 0 to -2.

Identifying redox reactions requires knowledge of oxidation states and understanding that total electron loss must equal total electron gain. For example, equation (a) demonstrates reduction of the magnesium (\(\mathrm{Mg^{2+}}\)) and nitrogen (\(\mathrm{N^{3-}}\)) ions to their elemental forms. Similarly, equation (e) shows a more complex redox reaction where both nitrogen and oxygen experience changes in their oxidation states.

Chemical Reaction Balancing

Balancing chemical equations is a fundamental skill in chemistry which ensures that the Law of Conservation of Mass is respected, implying that matter is neither created nor destroyed in a chemical reaction. To achieve a balanced equation, one must ensure that the same number of each type of atom appears on both sides of the equation. For instance, in equation (a) from our exercise, \(\mathrm{Mg_3N_2} + 6\mathrm{H_2O} \longrightarrow 3\mathrm{Mg(OH)_2} + 2\mathrm{NH_3}\), the equation is balanced by adjusting coefficients to ensure there are equal numbers of magnesium, nitrogen, and hydrogen atoms on both sides. Effective balancing often begins with balancing atoms of elements that are not oxygen or hydrogen and saving these two elements for last. This technique was applied in all the given equations to successfully balance them.

When students approach balancing equations, it's advisable to systematically change one coefficient at a time and check the balance of each element before proceeding to the next. This step-by-step approach minimizes confusion and helps avoid common errors.

Oxidation States

Understanding oxidation states is essential when dealing with redox reactions, as they provide insight into the electron flow during a reaction. An oxidation state is a number that represents the total number of electrons an atom either gains or loses to become a stable or chemically bonded state. It's often indicated as a positive or negative value based on conventional rules.

In our step-by-step solutions, we identified changes in oxidation states as the driving force behind redox reactions. Take equation (d), where the nitrogen in ammonia (\(\mathrm{NH_3}\)) undergoes reduction, its oxidation state changes from -3 in the reactants to -2 in the ammonium ion (\(\mathrm{NH_4^+}\)). Each element has its own set of common oxidation states, which can be predicted using periodic table trends and known compound structures.

To identify oxidation states, students should become familiar with these trends and rules, such as elements in their elemental form having an oxidation state of zero and the oxidation state of oxygen typically being -2 (with some exceptions). By grasping the concept of oxidation states, students can not only determine whether a chemical equation is a redox or non-redox reaction but can also predict which species in a reaction are likely to get oxidized or reduced.