Problem 19 - Chemistry The Central Science In Si Units

Question

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number. $$ \begin{array}{l} \text { (a) } 2 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow 2 \mathrm{HNO}_{3}(a q) \\ \text { (b) } \mathrm{FeS}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{FeCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{~S}(g) \\ \text { (c) } \mathrm{Fe}(s)+2 \mathrm{HNO}_{3}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+ \end{array} $$ \(2 \mathrm{NO}_{2}(g)+\mathrm{FeO}(s)\)

Step-by-Step Solution

Answer

Equations (a) and (c) are redox reactions. In (a), \(N\) is oxidized and \(O\) is reduced; in (c), \(Fe\) is oxidized and \(N\) is reduced.

Step 1: Define Oxidation and Reduction

Oxidation involves the increase in oxidation number, while reduction involves the decrease in oxidation number of elements in a chemical reaction. An oxidation-reduction (redox) reaction occurs when the oxidation states of atoms are changed.

Step 2: Analyze Equation (a)

Examine each element in the equation: \(2 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow 2 \mathrm{HNO}_{3}(a q)\). The oxidation states of \(\mathrm{N}\) in \(\mathrm{NO}_{2}\) is \(+4\), and in \(\mathrm{HNO}_{3}\) it is \(+5\). \(\mathrm{N}\) is oxidized. The oxidation state of \(\mathrm{O}\) in \(\mathrm{H}_{2} \mathrm{O}_{2}\) is \(-1\), and in \(\mathrm{HNO}_{3}\) it is \(-2\). \(\mathrm{O}\) is reduced.

Step 3: Analyze Equation (b)

Consider the equation: \(\mathrm{FeS}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{FeCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{S}(g)\). Here, the oxidation state of \(\mathrm{Fe}\) remains \(+2\) and that of \(\mathrm{S}\) remains \(-2\). No changes in oxidation states are observed, meaning it is not a redox reaction.

Step 4: Analyze Equation (c)

Look at the equation: \(\mathrm{Fe}(s)+2 \mathrm{HNO}_{3}(a q) \rightarrow \mathrm{H}_{2} \mathrm{O}(l)+ 2 \mathrm{NO}_{2}(g)+\mathrm{FeO}(s)\). \(\mathrm{Fe}\) changes from 0 in \(\mathrm{Fe}(s)\) to \(+2\) in \(\mathrm{FeO(s)}\), it is oxidized. \(\mathrm{N}\) changes from \(+5\) in \(\mathrm{HNO}_{3}\) to \(+4\) in \(\mathrm{NO}_{2}\), so it's reduced.

Key Concepts

Oxidation States

Understanding oxidation states is crucial for recognizing oxidation-reduction reactions. An oxidation state, often called an oxidation number, is a number assigned to an element in a chemical compound. It represents the number of electrons an atom either gains or loses or seems to use. The rules for assigning oxidation states are:

• The oxidation state of an element in its natural form is zero. For instance, in \(\mathrm{Fe}(s)\) or \(\mathrm{O}_{2}(g)\), the oxidation state is 0.
• For a simple ion, the oxidation state is equal to the charge on the ion. For example, the oxidation state of \(\mathrm{Na}^+\) is +1.
• In compounds, hydrogen has an oxidation state of +1, and oxygen has an oxidation state of -2.
• The sum of oxidation states in a neutral compound is zero, while in a polyatomic ion, it equals the ion charge.

By tracking oxidation states, we can identify whether elements undergo oxidation (an increase in oxidation state) or reduction (a decrease in oxidation state).

Balanced Chemical Equations

Balanced chemical equations are vital to accurately represent a chemical reaction. They demonstrate the conservation of mass by having the same number of atoms for each element in both reactants and products. A balanced equation ensures that matter is neither created nor destroyed during the reaction.
Here's how to balance an equation:

• Write the unbalanced equation.
• Count the number of atoms for each element in reactants and products. These must balance.
• Add coefficients, the numbers before compounds, to balance the atoms of each element. Only adjust coefficients, never the subscripts of a compound.
• Check the equation to ensure all atoms balance and all coefficients are in the lowest possible ratio.

Balancing equations is essential for quantitative chemical analyses and for understanding the ratios of substances that react and are produced. This foundation is significantly beneficial when tackling complex redox reactions.

Redox Reactions

Redox reactions are a type of chemical reaction where oxidation and reduction occur simultaneously. A key aspect of redox reactions is electron transfer between the reacting species, significantly altering their oxidation states.
In every redox reaction:

• An element that undergoes oxidation loses electrons, hence its oxidation state increases. An example is \(\mathrm{Fe}(s)\) in the exercise, changing from 0 to +2, indicating it lost electrons.
• An element that undergoes reduction gains electrons, resulting in a decrease in its oxidation state. For example, \(\mathrm{N}\) in \(\mathrm{HNO}_3\) reducing from +5 to +4 in \(\mathrm{NO}_{2}\).

To identify redox reactions in equations:

• Examine the oxidation states of all elements in the reactants and the products.
• If there's a change in the oxidation state of elements, a redox reaction occurs.

Practicing the identification of redox reactions helps in understanding various chemical processes like combustion, cellular respiration, and corrosion.