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Chapter 19: Problem 5

Identify the following reactions as redox or nonredox: a. \(2 \mathrm{NH}_{4} \mathrm{Cl}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \longrightarrow\) \(2 \mathrm{NH}_{3}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{CaCl}_{2}(a q)\) b. \(2 \mathrm{HNO}_{2}(a q)+3 \mathrm{H}_{2} \mathrm{S}(\mathrm{g})\) \(2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l)+3 \mathrm{S}(s)\) c. \(\left[\operatorname{Be}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2}+(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) \(\mathrm{H}_{3} \mathrm{O}+(a q)+\left[\mathrm{Be}\left(\mathrm{H}_{2} \mathrm{O}\right)_{3} \mathrm{OH}\right]+(a q)\)

Short Answer

Expert verified
a. Nonredox, b. Redox, c. Nonredox

Step by step solution

01

Understand the given reactions

Review the given reactions and note the reactants and products for each case. a. \(2 \mathrm{NH}_{4} \mathrm{Cl}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \longrightarrow 2 \mathrm{NH}_{3}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{CaCl}_{2}(a q)\) b. \(2 \mathrm{HNO}_{2}(a q)+3 \mathrm{H}_{2} \mathrm{S}(\mathrm{g}) \longrightarrow 2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l)+3 \mathrm{S}(s)\) c. \(\left[ \operatorname{Be} \left(\mathrm{H}_2 \mathrm{O}\right)_{4} \right]^{2}+(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}+(a q)+\left[ \mathrm{Be} \left(\mathrm{H}_{2} \mathrm{O}\right)_{3} \mathrm{OH}\right]+(a q)\)
02

Check for oxidation and reduction in each reaction

For each reaction, determine if any of the reactants are undergoing a change in oxidation state (Oxidation: increase in oxidation state, Reduction: decrease in oxidation state).
03

Identify redox or nonredox for Reaction a

In reaction (a), evaluate the oxidation states of nitrogen, hydrogen, calcium, chlorine, and oxygen. None of the elements involved undergo a change in oxidation state. This indicates that reaction (a) is not a redox reaction.
04

Identify redox or nonredox for Reaction b

In reaction (b), observe the oxidation states: - Nitrogen: - Reactant \( \mathrm{HNO}_{2}\) \( (+3)\) to Product \(\mathrm{NO}\) \( (+2)\) (reduction) - Sulfur: - Reactant \(\mathrm{H}_{2} \mathrm{S}\) \( (-2)\) to Product \(\mathrm{S}\) \( 0\) (oxidation) This indicates that reaction (b) is a redox reaction.
05

Identify redox or nonredox for Reaction c

In reaction (c), no element changes its oxidation state since Be remains \(2+\), H remains \(1+\), and O remains \(2-\). This indicates that reaction (c) is not a redox reaction.

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Key Concepts

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

Oxidation
Oxidation is a fundamental concept in redox chemistry. It occurs when an atom, ion, or molecule loses electrons. This loss causes an increase in the oxidation state. For example, consider the reaction where magnesium reacts with oxygen: \[2 \text{Mg (s)} + \text{O}_2 \text{(g)} \rightarrow 2 \text{MgO (s)}\]. Magnesium loses two electrons and becomes Mg2+, while oxygen gains those electrons and becomes O2-. In this case, magnesium undergoes oxidation.
  • Oxidation involves the loss of electrons.
  • It results in an increase in the oxidation state.
  • This process often involves the formation of positive ions, or cations.
Reduction
Reduction is the other half of a redox reaction, occurring simultaneously with oxidation. It involves the gain of electrons, which leads to a decrease in the oxidation state. For instance, in the reaction \[ \text{Cu (s)} + \text{2} \text{Ag}^+\text{ (aq)} \rightarrow \text{Cu}^{2+} \text{(aq)} + \text{2} \text{Ag (s)}\], copper loses electrons and is oxidized, while silver ions gain electrons and are reduced to metallic silver.
  • Reduction is defined by the gain of electrons.
  • It leads to a decrease in the oxidation state.
  • This process often results in the formation of negative ions, or anions, or neutral atoms.
Oxidation States
Oxidation states (or oxidation numbers) are values assigned to individual atoms that help in understanding redox reactions. They represent the hypothetical electric charge that an atom would have if all bonds were 100% ionic.
For example, in water (\( \text{H}_2\text{O} \)), hydrogen has an oxidation state of \(+1\) and oxygen has an oxidation state of \(-2\). The sum of the oxidation states in a neutral molecule always equals zero.
  • Oxidation state increases with oxidation.
  • Oxidation state decreases with reduction.
  • The sum of oxidation states in a molecule or ion reflects its overall charge.
Redox Reactions
Redox reactions are chemical reactions that involve the transfer of electrons between two species, encompassing both oxidation and reduction processes. These reactions are crucial for many biological and industrial processes, such as cellular respiration, photosynthesis, and combustion.
In the redox reaction \(2\text { H } _2\text { O } \text { ( l ) } +2\text { Na } \text { ( s ) } \rightarrow 2 \text { NaOH } (\text { aq } ) + \text { H } _2 ( \text { g } )\), sodium is oxidized and water is reduced. Identifying redox reactions involves checking for changes in oxidation states of the involved species.
  • Both oxidation and reduction occur simultaneously in a redox reaction.
  • They are integral to energy conversion processes.
  • Transferred electrons can be observed through changes in oxidation states.

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Most popular questions from this chapter

Balance the following equations in basic solution: a) \(\mathrm{PbO}_{2}+\mathrm{KCl} \longrightarrow \mathrm{KClO}+\mathrm{KPb}(\mathrm{OH})_{3}\) b) \(\mathrm{KMnO}_{4}+\mathrm{KIO}_{3} \longrightarrow \mathrm{MnO}_{2}+\mathrm{KIO}_{4}\) c) \(\mathrm{K}_{2} \mathrm{MnO}_{4} \longrightarrow \mathrm{MnO}_{2}+\mathrm{KMnO}_{4}\)

Oxidizing agents are used in the cleaning industry. Research three different oxidizing agents used in this area, and write a report on the advantages and disadvantages of these compounds.

Balance the equation for the reaction in which hot, concentrated sulfuric acid reacts with zinc to form zinc sulfate, hydrogen sulfide, and water.

Determine the oxidation number of each atom indicated in the following: a. \({H}_{2}\) f. \({HNO}_{3}\) b. \({H}_{2} {O}\) g. \({H}_{2} {SO}_{4}\) c. Al h. \({Ca}({OH})_{2}\) d. \({MgO}\) i. \({Fe}({NO}_{3})_{2}\) e. \({Al}_{2} {S}_{3}\) j. \({O}_{2}\)

For each requested step, use the half-reaction method to balance the oxidation-reduction equation below. (Hint: See Sample Problem A.) \({K}+{H}_{2} {O} \longrightarrow {KOH}+{H}_{2}\) a. Write the ionic equation, and assign oxidation numbers to all atoms to determine what is oxidized and what is reduced. b. Write the equation for the reduction, and balance it for both atoms and charge. c. Write the equation for the oxidation, and balance it for both atoms and charge. d. Multiply the coefficients of the oxidation and reduction equations so that the number of electrons lost equals the number of electrons gained. Add the two equations. e. Add species as necessary to balance the overall formula equation.
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