Question

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

Short answer

The balanced reactions are: (a) \(\mathrm{Mg}_3\mathrm{N}_{2}(s) + 6\,\mathrm{H}_{2}\mathrm{O}(l) \longrightarrow 3\,\mathrm{Mg}(\mathrm{OH})_2(aq) + 2\,\mathrm{NH}_3(g)\) (b) \(2\,\mathrm{NO}(g) + \mathrm{O}_{2}(g) \longrightarrow 2\,\mathrm{NO}_{2}(g)\) (c) \(\mathrm{N}_{2}\mathrm{O}_{5}(g) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow 2\,\mathrm{HNO}_{3}(aq)\) (d) \(\mathrm{NH}_{3}(aq) + \mathrm{H}^{+}(aq) \longrightarrow \mathrm{NH}_{4}^{+}(aq)\) (e) \(2\,\mathrm{N}_{2}\mathrm{H}_{4}(l) + \mathrm{O}_{2}(g) \longrightarrow 4\,\mathrm{H}_{2}\mathrm{O}(l) + \mathrm{N}_{2}(g)\) Redox reactions are (a), (b), and (e).

Step-by-step solution

Complete and balance the given reaction (a)

Given reaction: \(\mathrm{Mg}_3\mathrm{N}_{2}(s) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow\) The magnesium nitride (\(\mathrm{Mg}_3\mathrm{N}_{2}\)) reacts with water (\(\mathrm{H}_{2}\mathrm{O}\)) to form magnesium hydroxide (\(\mathrm{Mg}(\mathrm{OH})_2\)) and ammonia (\(\mathrm{NH}_3\)). Balanced reaction: \(\mathrm{Mg}_3\mathrm{N}_{2}(s) + 6\,\mathrm{H}_{2}\mathrm{O}(l) \longrightarrow 3\,\mathrm{Mg}(\mathrm{OH})_2(aq) + 2\,\mathrm{NH}_3(g)\)

Complete and balance the given reaction (b)

Given reaction: \(\mathrm{NO}(g) + \mathrm{O}_{2}(g) \longrightarrow\) Nitric oxide (\(\mathrm{NO}\)) reacts with oxygen gas (\(\mathrm{O}_{2}\)) to form nitrogen dioxide (\(\mathrm{NO}_{2}\)). Balanced reaction: \(2\,\mathrm{NO}(g) + \mathrm{O}_{2}(g) \longrightarrow 2\,\mathrm{NO}_{2}(g)\)

Complete and balance the given reaction (c)

Given reaction: \(\mathrm{N}_{2}\mathrm{O}_{5}(g) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow\) Dinitrogen pentoxide (\(\mathrm{N}_{2}\mathrm{O}_{5}\)) reacts with water (\(\mathrm{H}_{2}\mathrm{O}\)) to form nitric acid (\(\mathrm{HNO}_{3}\)). Balanced reaction: \(\mathrm{N}_{2}\mathrm{O}_{5}(g) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow 2\,\mathrm{HNO}_{3}(aq)\)

Complete and balance the given reaction (d)

Given reaction: \(\mathrm{NH}_{3}(aq) + \mathrm{H}^{+}(aq) \longrightarrow\) Ammonia (\(\mathrm{NH}_{3}\)) reacts with hydrogen ions (\(\mathrm{H}^{+}\)) to form the ammonium ion (\(\mathrm{NH}_{4}^{+}\)). Balanced reaction: \(\mathrm{NH}_{3}(aq) + \mathrm{H}^{+}(aq) \longrightarrow \mathrm{NH}_{4}^{+}(aq)\)

Complete and balance the given reaction (e)

Given reaction: \(\mathrm{N}_{2}\mathrm{H}_{4}(l) + \mathrm{O}_{2}(g) \longrightarrow\) Hydrazine (\(\mathrm{N}_{2}\mathrm{H}_{4}\)) reacts with oxygen gas (\(\mathrm{O}_{2}\)) to form water (\(\mathrm{H}_{2}\mathrm{O}\)) and nitrogen gas (\(\mathrm{N}_{2}\)). Balanced reaction: \(2\,\mathrm{N}_{2}\mathrm{H}_{4}(l) + \mathrm{O}_{2}(g) \longrightarrow 4\,\mathrm{H}_{2}\mathrm{O}(l) + \mathrm{N}_{2}(g)\)

Identify redox reactions

Now, we will identify which reactions are redox reactions by determining if there is any change in the oxidation state of the involved chemical species. Reaction (a) is a redox reaction: \(3\,\mathrm{MG}(+2) + 2\,\mathrm{N}(-3) \longrightarrow 3\,\mathrm{Mg}^{2+} + 6\,\mathrm{OH}^{-} + 2\,\mathrm{NH}_{3}\) Reaction (b) is a redox reaction: \(2\,\mathrm{NO} \longrightarrow 2\,\mathrm{NO}_{2}\) Reaction (c) is not a redox reaction: \(\mathrm{N}_{2}\mathrm{O}_{5} + \mathrm{H}_{2}\mathrm{O} \longrightarrow 2\,\mathrm{HNO}_{3}\) Reaction (d) is not a redox reaction: \(\mathrm{NH}_{3} + \mathrm{H}^{+} \longrightarrow \mathrm{NH}_{4}^{+}\) Reaction (e) is a redox reaction: \(2\,\mathrm{N}_{2}\mathrm{H}_{4} + \mathrm{O}_{2} \longrightarrow 4\,\mathrm{H}_{2}\mathrm{O} + \mathrm{N}_{2}\) Redox reactions are (a), (b), and (e).

Key concepts

Chemical Equations Balancing

Understanding and balancing chemical equations lies at the core of mastering chemistry. Balancing an equation is a process of ensuring that the number of atoms of each element is the same on both sides of the equation, thus obeying the Law of Conservation of Mass. When balancing a chemical equation, it's essential to account for each atom. In the provided exercise,
Mg reacts with H to form Mg(OH)₂ and NH₃ as products. Balancing requires making sure that the number of Mg, N, H, and O atoms in the reactants matches those in the products.

Starting Tips for Balancing Equations

• Write down the number of atoms of each element present in the unbalanced equation.
• Start by balancing the atom that appears in the least number of compounds on both sides of the equation.
• Adjust the coefficients - the numbers placed before compounds - to get the same number of atoms of an element on both sides.
• Iterate the process for each element until the equation is balanced.
• Check your work by counting the atoms of each element one last time.

This step-by-step balancing act not only makes the equation representative of the actual reaction but also sets the stage for stoichiometric calculations, a vital skill in chemistry.

Oxidation State Changes

Oxidation state changes are indicative of redox reactions, a type of chemical reaction where the oxidation state (or oxidation number) of the elements involved changes due to the transfer of electrons. Redox reactions are all around us, from the rusting of iron to the batteries powering devices.

Recognizing Oxidation State Changes

• An increase in the oxidation state of an element during a reaction implies it has lost electrons, known as oxidation.
• A decrease in the oxidation state signifies electron gain, termed as reduction.
• In any redox reaction, one species will be oxidized, and another will be reduced, summing up to a zero net charge change.

In the solutions provided, reactions (a), (b), and (e) are redox reactions, as exemplified by the change in the oxidation state of nitrogen from -3 in Mg₃N₂ to 0 in NH₃ in reaction (a). The understanding of oxidation states is crucial to identifying redox reactions and predicting the outcome of chemical processes.

Chemical Reactions Identification

The ability to identify chemical reactions is a basic skill required in chemistry. Chemical reactions can be categorized into different types like synthesis, decomposition, single replacement, double replacement, combustion, and redox, which often involve changes in energy and the formation of new substances from the reactants.

Distinguishing Between Reaction Types

• Synthesis reactions involve the combination of elements or simpler compounds to form a more complex compound.
• Decomposition reactions see a single compound breaking down into two or more simpler substances.
• Single and double replacement reactions involve the exchange of one or more parts between two reactant chemicals.
• Combustion usually involves a hydrocarbon reacting with oxygen to produce carbon dioxide and water.
• Redox reactions involve the transfer of electrons between species, altering their oxidation states.

In the exercise provided, reactions (a), (b), and (e) were identified as redox reactions due to changes in oxidation states, providing a clear example of identifying reaction types. Recognizing the type of chemical reaction is vital for understanding the behavior of different substances in various chemical processes.