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Chapter 18: Problem 16

Assign oxidation numbers to all the atoms in each of the following. a. \(\mathrm{HNO}_{3}\) e. \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}\) i. \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) b. \(\mathrm{CuCl}_{2}\) f. Ag j. \(\mathrm{CO}_{2}\) c. \(\mathrm{O}_{2}\) g. \(\mathrm{PbSO}_{4}\) k. \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Ce}\left(\mathrm{SO}_{4}\right)_{3}\) d. \(\mathrm{H}_{2} \mathrm{O}_{2}\) h. \(\mathrm{PbO}_{2}\) 1\. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)

Short Answer

Expert verified
a. H: +1, N: +5, O: -2 e. C: 0, H: +1, O: -2 i. Na: +1, C: +3, O: -2 b. Cu: +2, Cl: -1 f. Ag: 0 j. C: +4, O: -2 c. O: 0 g. Pb: +2, S: +6, O: -2 k. N: -3, H: +1, Ce: +3, S: +6, O: -2 d. H: +1, O: -1 h. Pb: +4, O: -2 1. Cr: +3, O: -2

Step by step solution

01

a. HNO3

(+1) + (x) + 3(-2) = 0 (Since the total charge of the compound is 0) x = +5 Oxidation numbers: H: +1, N: +5, O: -2
02

e. C6H12O6

6(x) + 12(+1) + 6(-2) = 0 x = 0 Oxidation numbers: C: 0, H: +1, O: -2
03

i. Na2C2O4

2(+1) + 2(x) + 4(-2) = 0 x = +3 Oxidation numbers: Na: +1, C: +3, O: -2
04

b. CuCl2

(x) + 2(-1) = 0 x = +2 Oxidation numbers: Cu: +2, Cl: -1
05

f. Ag

This is a free element, so its oxidation number is 0.
06

j. CO2

(x) + 2(-2) = 0 x = +4 Oxidation numbers: C: +4, O: -2
07

c. O2

This is a free element, so its oxidation number is 0.
08

g. PbSO4

(x) + (+6) + 4(-2) = 0 x = +2 Oxidation numbers: Pb: +2, S: +6, O: -2
09

k. (NH4)2Ce(SO4)3

2(+1) * 2 + (x) + 3(+6) + 12(-2) = 0 x = +3 Oxidation numbers: N: -3, H: +1, Ce: +3, S: +6, O: -2
10

d. H2O2

2(+1) + 2(-1) = 0 Oxidation numbers: H: +1, O: -1
11

h. PbO2

(x) + 2(-2) = 0 x = +4 Oxidation numbers: Pb: +4, O: -2
12

1. Cr2O3

2(x) + 3(-2) = 0 x = +3 Oxidation numbers: Cr: +3, O: -2

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Most popular questions from this chapter

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 \(1.0 \mathrm{~atm} .\) \(\begin{array}{ll}\text { a. } \mathrm{H}_{2} \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O} & \mathscr{6}^{\circ}=1.78 \mathrm{~V} \\ \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2} \mathrm{O}_{2} & \mathscr{6}^{\circ}=0.68 \mathrm{~V}\end{array}\) b. \(\mathrm{Mn}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Mn}\) \(\mathscr{6}^{\circ}=-1.18 \mathrm{~V}\) \(\mathrm{Fe}^{3+}+3 \mathrm{e}^{-} \rightarrow \mathrm{Fe} \quad \mathscr{6}^{\circ}=-0.036 \mathrm{~V}\)

The following standard reduction potentials have been determined for the aqueous chemistry of indium: $$\begin{array}{ll}\mathrm{In}^{3+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{In}^{+}(a q) & \mathscr{E}^{\circ}=-0.444 \mathrm{~V} \\ \mathrm{In}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{In}(s) & \mathscr{E}^{\circ}=-0.126 \mathrm{~V}\end{array}$$ a. What is the equilibrium constant for the disproportionation reaction, where a species is both oxidized and reduced, shown below? $$3 \operatorname{In}^{+}(a q) \longrightarrow 2 \operatorname{In}(s)+\operatorname{In}^{3+}(a q)$$ b. What is \(\Delta G_{\mathrm{f}}^{\circ}\) for \(\mathrm{In}^{+}(a q)\) if \(\Delta G_{\mathrm{f}}^{\circ}=-97.9 \mathrm{~kJ} / \mathrm{mol}\) for \(\mathrm{In}^{3+}(a q)\) ?

The amount of manganese in steel is determined by changing it to permanganate ion. The steel is first dissolved in nitric acid, producing \(\mathrm{Mn}^{2+}\) ions. These ions are then oxidized to the deeply colored \(\mathrm{MnO}_{4}^{-}\) ions by periodate ion \(\left(\mathrm{IO}_{4}^{-}\right)\) in acid solution. a. Complete and balance an equation describing each of the above reactions. b. Calculate \(\mathscr{C}^{\circ}\) and \(\Delta G^{\circ}\) at \(25^{\circ} \mathrm{C}\) for each reaction.

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 \(1.0 \mathrm{~atm}\). a. \(\mathrm{Cl}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cl}^{-} \quad \mathscr{E}^{\circ}=1.36 \mathrm{~V}\) \(\mathrm{Br}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br}^{-} \quad 8^{\circ}=1.09 \mathrm{~V}\) b. \(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}\) \(\mathscr{b}^{\circ}=1.51 \mathrm{~V}\) \(\mathrm{IO}_{4}^{-}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{IO}_{3}^{-}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E ^ { \circ }}=1.60 \mathrm{~V}\)

Is the following statement true or false? Concentration cells work because standard reduction potentials are dependent on concentration. Explain.
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