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

The balanced equations are: (a) Mg₃N₂(s) + 6H₂O(l) → 3Mg(OH)₂(s) + 2NH₃(g) (redox) (b) 2NO(g) + O₂(g) → 2NO₂(g) (redox) (c) N₂O₅(g) + H₂O(l) → 2HNO₃(aq) (redox) (d) NH₃(aq) + H⁺(aq) → NH₄⁺(aq) (not redox) (e) N₂H₄(l) + 2O₂(g) → N₂(g) + 4H₂O(l) (redox) Redox reactions: (a), (b), (c), and (e).

Step-by-step solution

1. Complete and balance Reaction (a)

We need to balance the given equation: \[ \mathrm{Mg}_{3} \mathrm{N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \] First, complete the reaction by writing the expected product, which is magnesium hydroxide and ammonia gas: \[ \mathrm{Mg}_{3} \mathrm{N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s)+\mathrm{NH}_{3}(g) \] Now, let's balance it by adjusting the number of moles of reactants and products: \[ \mathrm{Mg}_{3} \mathrm{N}_{2}(s)+6 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 3 \mathrm{Mg}(\mathrm{OH})_{2}(s)+2 \mathrm{NH}_{3}(g) \]

2. Complete and balance Reaction (b)

We need to balance the given equation: \[ \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow \] First, complete the reaction by providing the expected product, which is nitrogen dioxide: \[ \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{NO}_{2}(g) \] Now, let's balance the equation by adjusting the number of moles of reactants and products: \[ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \]

3. Complete and balance Reaction (c)

We need to balance the given equation: \[ \mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \] First, complete the reaction by providing the expected product, which is nitric acid: \[ \mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(aq) \] Now, let's balance the equation by adjusting the number of moles of reactants and products: \[ \mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(aq) \]

4. Complete and balance Reaction (d)

We need to balance the given equation: \[ \mathrm{NH}_{3}(a q)+\mathrm{H}^{+}(a q) \longrightarrow \] First, complete the reaction by providing the expected product, which is ammonium ion: \[ \mathrm{NH}_{3}(a q)+\mathrm{H}^{+}(a q) \longrightarrow \mathrm{NH}_{4}^{+}(aq) \] The equation is already balanced.

5. Complete and balance Reaction (e)

We need to balance the given equation: \[ \mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow \] First, complete the reaction by providing the expected product, which is nitrogen gas and water: \[ \mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{N}_{2}(g)+\mathrm{H}_{2}\mathrm{O}(l) \] Now, let's balance the equation by adjusting the number of moles of reactants and products: \[ \mathrm{N}_{2} \mathrm{H}_{4}(l)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{N}_{2}(g)+4 \mathrm{H}_{2}\mathrm{O}(l) \]

6. Identify redox reactions

To find out which of these reactions are redox reactions, we need to check if there are any changes in the oxidation numbers of the elements in the reactants and products. (a) Mg₃N₂(s) + 6H₂O(l) → 3Mg(OH)₂(s) + 2NH₃(g) Oxidation numbers change, so it is a redox reaction. (b) 2NO(g) + O₂(g) → 2NO₂(g) Oxidation numbers change, so it is a redox reaction. (c) N₂O₅(g) + H₂O(l) → 2HNO₃(aq) Oxidation numbers change, so it is a redox reaction. (d) NH₃(aq) + H⁺(aq) → NH₄⁺(aq) No changes in oxidation numbers, so it is not a redox reaction. (e) N₂H₄(l) + 2O₂(g) → N₂(g) + 4H₂O(l) Oxidation numbers change, so it is a redox reaction. Hence, the redox reactions are (a), (b), (c), and (e).

Key concepts

Balancing Equations

When you come across a chemical equation, it is crucial to ensure that it is balanced. A balanced equation has the same number of atoms for each element on both the reactant side and the product side. This follows the law of conservation of mass, which states that matter cannot be created or destroyed in a chemical reaction. To balance an equation, follow these steps:

• Write the unbalanced equation with all reactants and products.
• Count the number of atoms for each element in both reactants and products.
• Use coefficients to multiply the number of molecules of each substance to balance the atoms on both sides.
• Check and ensure that all coefficients are in the simplest ratio.

Remember that only coefficients are changed while leaving the subscripts in the chemical formulas unaltered. Balancing equations accurately helps in understanding the stoichiometry of the reactions and ensures that all the mass relationships are maintained.

Redox Reactions

Redox, or reduction-oxidation reactions, involve the transfer of electrons between two substances. These reactions are characterized by changes in the oxidation states of the elements involved. The substance that gives up electrons is said to be oxidized, while the one that gains electrons is reduced.
Some characteristics of redox reactions include:

• Oxidation - Loss of electrons.
• Reduction - Gain of electrons.
• Oxidizing agent - Accepts electrons and gets reduced.
• Reducing agent - Donates electrons and gets oxidized.

Redox reactions play a critical role in various processes, such as metabolism in living organisms, combustion, and corrosion. In some cases, redox reactions can be identified by changes in color or heat release. Identifying these reactions involves assessing the changes in oxidation numbers across the two sides of the chemical equation.

Oxidation Numbers

Oxidation numbers are a theoretical charge assigned to different elements in a molecule, enabling us to track electron transfer in redox reactions. These numbers help identify which elements are oxidized and reduced during the reaction.
Some simple rules for determining oxidation numbers include:

• The oxidation number of an atom in a free element is always zero.
• For simple ions, the oxidation number equals the charge of the ion.
• In compounds, oxygen usually has an oxidation number of -2, and hydrogen is generally +1.
• The sum of oxidation numbers in a neutral compound equals zero, while in a polyatomic ion, it equals the charge of the ion.

By following these rules, you can identify the changes in oxidation numbers to determine if a reaction is a redox reaction. Understanding oxidation numbers is vital for mastering concepts in electrochemistry and for analyzing redox reactions in chemistry.