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Chapter 4: Problem 92

Identify the oxidizing and reducing agents in the following reactions: $$ \begin{array}{l} \text { (a) } 5 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(a q)+2 \mathrm{MnO}_{4}^{-}(a q)+6 \mathrm{H}^{*}(a q) \longrightarrow \\ 2 \mathrm{Mn}^{2+}(a q)+10 \mathrm{CO}_{2}(g)+8 \mathrm{H}_{2} \mathrm{O}(l) \end{array} $$ (b) \(3 \mathrm{Cu}(s)+8 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q) \longrightarrow\) $$ 3 \mathrm{Cu}^{2+}(a q)+2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l) $$

Short Answer

Expert verified
a) Oxidizing agent: \(\mathrm{MnO}_{4}^{-}\), reducing agent: \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\). b) Oxidizing agent: \(\mathrm{NO}_{3}^{-}\), reducing agent: \(\mathrm{Cu}\).

Step by step solution

01

Identify Oxidation States - Part (a)

Determine the oxidation states of each element in the reaction (a). For example, in \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), hydrogen has an oxidation state of +1, oxygen has -2, and carbon balanced to +3 (since two carboxyl groups are present). For \(\mathrm{MnO}_{4}^{-}\), manganese has an oxidation state of +7 and oxygen is -2.
02

Identify Changes in Oxidation States - Part (a)

In the product side of the reaction (a), \(\mathrm{Mn}^{2+}\) indicates manganese has an oxidation state of +2. Compare this change: Manganese goes from +7 to +2 (reduction), and carbon goes from +3 in \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) to +4 in \(\mathrm{CO}_{2}\) (oxidation).
03

Identify Agents - Part (a)

The substance that gets reduced is the oxidizing agent. Hence, \(\mathrm{MnO}_{4}^{-}\) is the oxidizing agent, and \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) is the reducing agent because it is oxidized.
04

Identify Oxidation States - Part (b)

Determine the oxidation states of each element in the reaction (b). For \(\mathrm{Cu}\), copper in elemental form is 0, and in \(\mathrm{Cu}^{2+}\), it is +2. For \(\mathrm{NO}_{3}^{-}\), nitrogen is +5.
05

Identify Changes in Oxidation States - Part (b)

In the product side of reaction (b), nitrogen in \(\mathrm{NO}\) is +2. Compare this change: Copper goes from 0 in \(\mathrm{Cu}\) to +2 in \(\mathrm{Cu}^{2+}\) (oxidation), and nitrogen goes from +5 in \(\mathrm{NO}_{3}^{-}\) to +2 in \tt{NO}\ (reduction).
06

Identify Agents - Part (b)

The substance that gets reduced is the oxidizing agent. Hence, \(\mathrm{NO}_{3}^{-}\) is the oxidizing agent, and \(\mathrm{Cu}\) is the reducing agent because it is oxidized.

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Key Concepts

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

Oxidation States
Understanding oxidation states is crucial for identifying oxidizing and reducing agents in redox reactions.
An oxidation state, also called the oxidation number, is an indicator of the degree of oxidation of an atom in a chemical compound.
It shows the hypothetical charge an atom would have if all bonds were 100% ionic.
Oxidation numbers are assigned according to a set of rules, such as:
  • Elements in their elemental form have an oxidation state of 0. For example, O鈧� and H鈧� have oxidation states of 0.
  • The oxidation state of a monoatomic ion equals the charge of the ion. For instance, Na鈦� has an oxidation state of +1.
  • In compounds, hydrogen usually has an oxidation state of +1, while oxygen usually has -2.
  • The sum of oxidation states in a neutral molecule must be zero, and in an ion, it must be equal to the ion's charge.
For example, in H2C2O4, hydrogen has an oxidation state of +1, oxygen has -2, and carbon's state is balanced to +3.
Redox Reactions
Redox reactions, also known as oxidation-reduction reactions, are processes where there is a transfer of electrons between two species.
These reactions are characterized by the change in oxidation states of the involved chemical elements.

In a redox reaction, one substance undergoes oxidation (loses electrons), and another undergoes reduction (gains electrons).
  • Oxidation: An increase in oxidation state.
  • Reduction: A decrease in oxidation state.
For instance, in the reaction 5 H2C2O4 + 2 MnO4- + 6 H+ 鈫� 2 Mn2+ + 10 CO2 + 8 H2O, oxalic acid (H2C2O4) gets oxidized, whereas permanganate (MnO4-) gets reduced.
Oxidizing Agent
An oxidizing agent, or oxidant, is a substance that gains electrons in a redox reaction and gets reduced.
It effectively causes the oxidation of another substance.
Identifying the oxidizing agent involves finding which substance undergoes a decrease in oxidation state.

For example, in the reaction 5 H2C2O4 + 2 MnO4- + 6 H+ 鈫� 2 Mn2+ + 10 CO2 + 8 H2O, manganese in MnO4- goes from +7 to +2.
Here, MnO4- acts as the oxidizing agent.
Reducing Agent
A reducing agent, or reductant, is a substance that loses electrons in a redox reaction and gets oxidized.
It causes the reduction of another substance.
Identifying the reducing agent involves finding which substance undergoes an increase in oxidation state.

For example, in the reaction 5 H2C2O4 + 2 MnO4- + 6 H+ 鈫� 2 Mn2+ + 10 CO2 + 8 H2O, carbon in H2C2O4 goes from +3 to +4.
Here, H2C2O4 acts as the reducing agent.
Chemical Reactions
Chemical reactions involve the transformation of reactants into products through the breaking and forming of bonds.
Redox reactions are a specific type of chemical reaction where oxidation and reduction occur simultaneously.
The understanding of redox reactions is essential because they are involved in various processes such as cellular respiration, photosynthesis, combustion, and industrial applications.

Redox reactions can often be identified by observing changes in oxidation states of the involved substances, signaling the transfer of electrons.
  • Important real-world applications include energy production in batteries, metal corrosion and prevention, and biochemical pathways like respiration.
Hence, learning to balance redox reactions and identifying oxidation and reduction processes is vital for chemistry students.

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Most popular questions from this chapter

Why must every redox reaction involve an oxidizing agent and a reducing agent?

A student forgets to weigh a mixture of sodium bromide dihydrate and magnesium bromide hexahydrate. Upon strong heating, the sample loses \(252.1 \mathrm{mg}\) of water. The mixture of anhydrous salts reacts with excess AgNO \(_{3}\) solution to form \(6.00 \times 10^{-3} \mathrm{~mol}\) of solid AgBr. Find the mass \(\%\) of each compound in the original mixture.

In which of the following equations does sulfuric acid act as an oxidizing agent? In which does it act as an acid? Explain. $$ \begin{array}{l} \text { (a) } 4 \mathrm{H}^{+}(a q)+\mathrm{SO}_{4}^{2-}(a q)+2 \mathrm{NaI}(s) \longrightarrow \\ \quad 2 \mathrm{Na}^{+}(a q)+\mathrm{I}_{2}(s)+\mathrm{SO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(I) \\ \text { (b) } \mathrm{BaF}_{2}(s)+2 \mathrm{H}^{+}(a q)+\mathrm{SO}_{4}^{2-}(a q) \longrightarrow 2 \mathrm{HF}(a q)+\mathrm{BaSO}_{4}(s) \end{array} $$

What occurs on the molecular level when an ionic compound dissolves in water?

Over time, as their free fatty acid (FFA) content increases, edible fats and oils become rancid. To measure rancidity, the fat or oil is dissolved in ethanol, and any FFA present is titrated with KOH dissolved in ethanol. In a series of tests on olive oil, a stock solution of \(0.050 \mathrm{M}\) ethanolic \(\mathrm{KOH}\) was prepared at \(25^{\circ} \mathrm{C},\) stored at \(0^{\circ} \mathrm{C},\) and then placed in a \(100-\mathrm{mL}\) buret to titrate oleic acid [an FFA with formula \(\left.\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{7} \mathrm{CH}=\mathrm{CH}\left(\mathrm{CH}_{2}\right)_{7} \mathrm{COOH}\right]\) in the oil. Each of four \(10.00-\mathrm{g}\) samples of oil took several minutes to titrate: the first required \(19.60 \mathrm{~mL}\), the second \(19.80 \mathrm{~mL},\) and the third and fourth \(20.00 \mathrm{~mL}\) of the ethanolic \(\mathrm{KOH}\). (a) What is the apparent acidity of each sample, in terms of mass \(\%\) of oleic acid? (Note: As the ethanolic KOH warms in the buret, its volume increases by a factor of \(0.00104 /{ }^{\circ} \mathrm{C}\).) (b) Is the variation in acidity a random or systematic error? Explain. (c) What is the actual acidity? How would you demonstrate this?
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