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Chapter 20: Problem 19

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

Short Answer

Expert verified
(a) Equation (a) involves an oxidation-reduction reaction. Nitrogen undergoes oxidation (oxidation number increases from +4 to +5), and oxygen undergoes reduction (oxidation number decreases from -1 to -2). (b) Equation (b) does not involve an oxidation-reduction reaction, as there is no change in oxidation numbers of any element. (c) Equation (c) involves an oxidation-reduction reaction. Iron undergoes oxidation (oxidation number increases from 0 to +2), and nitrogen undergoes reduction (oxidation number decreases from +5 to +4).

Step by step solution

01

(a) Check if equation (a) involves oxidation-reduction

For equation (a), first assign the oxidation numbers to each element in the reactants and products: \(NO_2\): N = +4, O = -2 \(H_2O_2\): H = +1, O = -1 \(HNO_3\): H = +1, N = +5, O = -2 Now compare the oxidation numbers before and after the reaction: - Nitrogen (N) has its oxidation number increased from +4 to +5 - Oxygen (O) has its oxidation number decreased from -1 to -2 Since there is a change in oxidation numbers, equation (a) involves an oxidation-reduction reaction. Nitrogen undergoes oxidation, and oxygen undergoes reduction.
02

(b) Check if equation (b) involves oxidation-reduction

For equation (b), assign oxidation numbers to each element in the reactants and products: \(FeS\): Fe = +2, S = -2 \(HCl\): H = +1, Cl = -1 \(FeCl_2\): Fe = +2, Cl = -1 \(H_2S\): H = +1, S = -2 Now compare the oxidation numbers before and after the reaction: - There is no change in the oxidation numbers of any element. Since there is no change in oxidation numbers, equation (b) does not involve an oxidation-reduction reaction.
03

(c) Check if equation (c) involves oxidation-reduction

For equation (c), assign oxidation numbers to each element in the reactants and products: \(Fe\): Fe = 0 \(HNO_3\): H = +1, N = +5, O = -2 \(H_2O\): H = +1, O = -2 \(NO_2\): N = +4, O = -2 \(FeO\): Fe = +2, O = -2 Now compare the oxidation numbers before and after the reaction: - Iron (Fe) has its oxidation number increased from 0 to +2 - Nitrogen (N) has its oxidation number decreased from +5 to +4 Since there is a change in oxidation numbers, equation (c) involves an oxidation-reduction reaction. Iron undergoes oxidation, and nitrogen undergoes reduction.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidation Numbers
Oxidation numbers are numerical values assigned to atoms to help determine how many electrons are gained or lost in a chemical reaction. They provide a way to track electron transfers, which is crucial for identifying oxidation-reduction reactions. Understanding how to assign these numbers is a key step in spotting changes during reactions.
  • Each element in its natural state has an oxidation number of zero. For example, O2 or H2 have oxidation numbers of 0.
  • For simple ions, the oxidation number is the charge of the ion. For example, Na+ is +1, and Cl鈭� is -1.
  • In compounds, hydrogen is typically +1, and oxygen is typically -2.
  • The sum of oxidation numbers in a neutral compound is 0. For polyatomic ions, the sum is the charge of the ion.
Assigning oxidation numbers consistently can help reveal whether a reaction involves exchange of electrons, indicating oxidation and reduction processes.
Reduction
Reduction is a critical component of redox (reduction-oxidation) reactions. It involves the gain of electrons by a molecule, atom, or ion. In terms of oxidation numbers, reduction leads to a decrease in the oxidation number. Key points about reduction:
  • It often involves the transfer of oxygen atoms away from the substance.
  • A good example is when oxygen gains electrons reducing its oxidation state.
  • Watch for a decrease in oxidation numbers to identify which species is being reduced.
For example, in the reaction: Oxygen in hydrogen peroxide ( H_2O_2 ) is reduced, as it decreases from an oxidation number of -1 to -2 in water ( H_2O ). Recognizing the species being reduced helps to identify the half of the process tied with electron acceptance.
Oxidation
Oxidation is the counterpart to reduction and involves the loss of electrons. When a substance undergoes oxidation, its oxidation number increases. This is an integral part of redox reactions. A few important notes about oxidation:
  • Often involves the addition of oxygen or removal of hydrogen.
  • You can spot oxidation by observing increases in oxidation numbers.
  • Common in metals when they react with oxygen to form oxides.
In the biochemical breakdown, for instance, where iron ( Fe ) goes from an oxidation number of 0 in its elemental form to +2 in FeO , iron undergoes oxidation. Identifying oxidation steps helps clarify which molecules donate electrons in reactions.
Chemical Equations
Chemical equations are symbolic representations of chemical reactions. They show the reactants turning into products and often indicate the states of matter. Criteria to check in chemical equations include balance and whether there's a change in oxidation numbers. Things to remember when examining chemical equations:
  • Both sides must have the same number of each type of atom, reflecting the conservation of mass.
  • Check the oxidation numbers for each element; a change indicates an oxidation-reduction process.
  • Look for products and reactants that signify common redox reactions, such as combustion or single replacement reactions.
Taking Equation (a) as an example, where NO_2 reacts with H_2O_2 and leads to a change in oxidation numbers of N and O , this equation involves a redox reaction. Observing oxidation numbers within these equations helps you understand the electron transfer dynamics.

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Most popular questions from this chapter

Magnesium is obtained by electrolysis of molten \(\mathrm{MgCl}_{2}\). (a) Why is an aqueous solution of \(\mathrm{MgCl}_{2}\) not used in the electrolysis? (b) Several cells are connected in parallel by very large copper bars that convey current to the cells. Assuming that the cells are \(96 \%\) efficient in producing the desired products in electrolysis, what mass of \(\mathrm{Mg}\) is formed by passing a current of 97,000 A for a period of 24 h?

(a) What conditions must be met for a reduction potential to be a standard reduction potential? (b) What is the standard reduction potential of a standard hydrogen electrode? (c) Why is it impossible to measure the standard reduction potential of a single half-reaction?

In the Br酶nsted-Lowry concept of acids and bases, acidbase reactions are viewed as proton-transfer reactions. The stronger the acid, the weaker is its conjugate base. If we were to think of redox reactions in a similar way, what particle would be analogous to the proton? Would strong oxidizing agents be analogous to strong acids or strong bases?

Complete and balance the following half-reactions. In each case indicate whether the half-reaction is an oxidation or a reduction. (a) \(\mathrm{Mo}^{3+}(a q) \longrightarrow \mathrm{Mo}(s)\) (acidic solution) (b) \(\mathrm{H}_{2} \mathrm{SO}_{3}(a q) \longrightarrow \mathrm{SO}_{4}^{2-}(a q)\) (acidic solution) (c) \(\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{NO}(g)\) (acidic solution) (d) \(\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)\) (acidic solution) (e) \(\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)\) (basic solution) (f) \(\mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{2}(s)\) (basic solution) (g) \(\mathrm{Cr}(\mathrm{OH})_{3}(s) \longrightarrow \mathrm{CrO}_{4}^{2-}(a q)\) (basic solution)

From each of the following pairs of substances, use data in Appendix \(\mathrm{E}\) to choose the one that is the stronger reducing agent: (a) \(\mathrm{Al}(s)\) or \(\mathrm{Mg}(s)\) (b) \(\mathrm{Fe}(s)\) or \(\mathrm{Ni}(s)\) (c) \(\mathrm{H}_{2}(g\), acidic solution) or \(\operatorname{Sn}(s)\) (d) \(\mathrm{I}^{-}(a q)\) or \(\mathrm{Br}^{-}(a q)\)
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