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Chapter 18: Problem 2

When balancing reactions in Chapter 3, we did not mention that reactions must be charge balanced as well as mass balanced. What do charge balance and mass balance mean? What happens in a redox reaction that requires charge to be balanced?

Short Answer

Expert verified
Mass balance refers to the conservation of mass in a chemical reaction, ensuring the number and type of atoms are equal on both sides of the equation. Charge balance refers to the conservation of charge in a reaction, ensuring the total charge of reactants is equal to the total charge of products. In redox reactions, where electrons are transferred between atoms, maintaining charge balance is crucial as every electron lost in oxidation must be gained in reduction to keep the system charge-neutral.

Step by step solution

01

Definition of Mass Balance

Mass balance refers to the principle that the total mass of reactants in a chemical reaction must be equal to the total mass of products, indicating that matter cannot be created or destroyed during the reaction. This principle is based on the conservation of mass. In other words, the number and type of atoms on both sides of a chemical reaction must be identical.
02

Definition of Charge Balance

Charge balance refers to the principle that the total charge of reactants in a chemical reaction must be equal to the total charge of products. This is because charge is also conserved during processes such as chemical reactions. In other words, the sum of charges on both sides of a chemical equation must be equal to maintain charge neutrality.
03

Redox Reactions

Redox reactions are chemical reactions in which the oxidation states of some atoms change due to transfer of electrons. They comprise two half-reactions: reduction and oxidation. During reduction, an atom gains electrons and its oxidation state decreases. In the oxidation half-reaction, an atom loses electrons, which causes its oxidation state to increase.
04

Charge Balance in Redox Reactions

In redox reactions, maintaining charge balance is crucial as electrons are transferred between species. This means that, for every electron lost by one species (oxidation half-reaction) during the reaction, there must be an equal number of electrons gained by another species (reduction half-reaction). This ensures that the overall charge of the reaction remains balanced and the system remains charge-neutral. To summarize, charge balance and mass balance are two essential principles that must be followed when balancing chemical reaction equations. In redox reactions, the charge balance plays a critical role in ensuring the conservation of both mass and charge.

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Most popular questions from this chapter

A galvanic cell is based on the following half-reactions at $25^{\circ} \mathrm{C} :$ $$\begin{array}{c}{\mathrm{Ag}^{+}+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}} \\\ {\mathrm{H}_{2} \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}}\end{array}$$ Predict whether \(\mathscr{E}_{\text{cell}}\) is larger or smaller than \(\mathscr{E}^{\circ}_{\text{cell}}\) for the following cases. a. [Ag1] 5 1.0 a. $\left[\mathrm{Ag}^{+}\right]=1.0 M,\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]=2.0 M,\left[\mathrm{H}^{+}\right]=2.0 \mathrm{M}$ b. $\left[\mathrm{Ag}^{+}\right]=2.0 \mathrm{M},\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]=1.0 M,\left[\mathrm{H}^{+}\right]=1.0 \times 10^{-7} \mathrm{M}$

Calculate \(\mathscr{E}^{\circ}\) values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations. Standard reduction potentials are found in Table 18.1. a. $\mathrm{MnO}_{4}^{-(a q)}+\mathrm{I}^{-}(a q) \longrightarrow \mathrm{I}_{2}(a q)+\mathrm{Mn}^{2+}(a q)$ b. $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{F}^{-}(a q) \longrightarrow \mathrm{F}_{2}(g)+\mathrm{Mn}^{2+}(a q)$

What volumes of \(\mathrm{H}_{2}(g)\) and \(\mathrm{O}_{2}(g)\) at STP are produced from the electrolysis of water by a current of 2.50 \(\mathrm{A}\) in 15.0 \(\mathrm{min}\) ?

Specify which of the following equations represent oxidation–reduction reactions, and indicate the oxidizing agent, the reducing agent, the species being oxidized, and the species being reduced. a. $\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)$ b. $2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)$ c. $\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)$ d. $2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

Sketch the galvanic cells based on the following half- reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine \(\mathscr{E}^{\circ}\) for the galvanic cells. Assume that all concentrations are 1.0 \(M\) and that all partial pressures are 1.0 atm. a. $\mathrm{Cl}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cl}^{-} \quad \mathscr{E}^{\circ}=1.36 \mathrm{V}$ $\mathrm{Br}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br}^{-} \quad \mathscr{E}^{\circ}=1.09 \mathrm{V}$ b. $\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.51 \mathrm{V}$ $\mathrm{IO}_{4}^{-}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{IO}_{3}^{-}+\mathrm{H}_{2} \mathrm{O} \quad \quad \mathscr{E}^{\circ}=1.60 \mathrm{V}$
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