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PROBLEM

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1. Dichromate ion (Cr2O72)\left(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\right) in acidic solution is a good oxidizing agent. Which of the following oxidations can be accomplished with dichromate ion in acidic solution?
Explain
a. Sn2+(aq)\mathrm{Sn}^{2+}(\mathrm{aq}) to Sn4+(aq)\mathrm{Sn}^{4+}(\mathrm{aq})
b. Ag(s)\mathrm{Ag}(\mathrm{s}) to Ag+(aq)\mathrm{Ag}^{+}(\mathrm{aq})
2. For each of the following cell diagrams
a. Pt(s)Fe2+(aq),Fe3+(aq)Ag+(aq)Ag(s)\operatorname{Pt}(\mathrm{s})\left|\mathrm{Fe}^{2+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq}) \|\left|\mathrm{Ag}^{+}(\mathrm{aq})\right| \mathrm{Ag}(\mathrm{s})\right.
b. Pt(s)Mn2+(aq),MnO4(aq)Cu2+(aq)Cu(s)\mathrm{Pt}(\mathrm{s})\left|\mathrm{Mn}^{2+}(\mathrm{aq}), \mathrm{MnO}_{4}^{-}(\mathrm{aq})\right|\left|\mathrm{Cu}^{2+}(\mathrm{aq})\right| \mathrm{Cu}(\mathrm{s})
i. calculate the Ecell oE_{\text {cell }}^{o}
ii. what is the change in Gibbs free energy?
iii. is the cell reaction spontaneous or nonspontaneous? Explain.
3. Will the following reaction
Cr2O72(aq)+14H++6Ag( s)2Cr3+(aq)+6Ag+(aq)+7H2O(l)Ecell =0.023 V\mathrm{Cr}_{2} \mathrm{O}_{7}^{2 \cdot}(\mathrm{aq})+14 \mathrm{H}^{+}+6 \mathrm{Ag}(\mathrm{~s}) \longrightarrow 2 \mathrm{Cr}^{3+}(\mathrm{aq})+6 \mathrm{Ag}^{+}(\mathrm{aq})+7 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) E_{\text {cell }}=-0.023 \mathrm{~V} be spontaneous if [Cr2O72]=[Ag+]=0.675M,[Cr3+]=0.6M\left[\mathrm{Cr}_{2} \mathrm{O}_{7}{ }^{2-}\right]=[\mathrm{Ag}+]=0.675 \mathrm{M},\left[\mathrm{Cr}^{3+}\right]=0.6 \mathrm{M} and pH=2\mathrm{pH}=2 ? Explain.
4. Define the following terms
a) Oxidation
b) Oxidizing agent
c) Salt bridge
d) Liquid junction
e) Standard electrode potential

STEP 1

1. The dichromate ion Cr2O72\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} is a strong oxidizing agent in acidic solutions.
2. Standard electrode potentials (EE^\circ) are available for the relevant half-reactions.
3. Gibbs free energy change (ΔG\Delta G) can be calculated using the formula ΔG=nFEcello\Delta G = -nFE_{\text{cell}}^{o}.
4. The Nernst equation can be used to determine cell potential under non-standard conditions.
5. Definitions are based on standard chemistry terminology.

STEP 2

1. Analyze the oxidation reactions with dichromate ion.
2. Calculate EcelloE_{\text{cell}}^{o} for given cell diagrams.
3. Determine Gibbs free energy change and spontaneity.
4. Evaluate the spontaneity of the given reaction under specified conditions.
5. Define the given chemistry terms.

STEP 3

Analyze the oxidation reactions with dichromate ion:
a. Check if Sn2+(aq)\mathrm{Sn}^{2+}(\mathrm{aq}) can be oxidized to Sn4+(aq)\mathrm{Sn}^{4+}(\mathrm{aq}) using dichromate ion.
- The standard reduction potential for Sn4+/Sn2+\mathrm{Sn}^{4+}/\mathrm{Sn}^{2+} is E=+0.15VE^\circ = +0.15 \, \text{V}.
- The standard reduction potential for Cr2O72/Cr3+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}/\mathrm{Cr}^{3+} is E=+1.33VE^\circ = +1.33 \, \text{V}.
- Since EE^\circ for dichromate is higher, it can oxidize Sn2+\mathrm{Sn}^{2+}.
b. Check if Ag(s)\mathrm{Ag}(\mathrm{s}) can be oxidized to Ag+(aq)\mathrm{Ag}^{+}(\mathrm{aq}).
- The standard reduction potential for Ag+/Ag\mathrm{Ag}^{+}/\mathrm{Ag} is E=+0.80VE^\circ = +0.80 \, \text{V}.
- Since EE^\circ for dichromate is higher, it can oxidize Ag\mathrm{Ag}.

STEP 4

Calculate EcelloE_{\text{cell}}^{o} for given cell diagrams:
a. For Pt(s)Fe2+(aq),Fe3+(aq)Ag+(aq)Ag(s)\operatorname{Pt}(\mathrm{s})\left|\mathrm{Fe}^{2+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq}) \|\left|\mathrm{Ag}^{+}(\mathrm{aq})\right| \mathrm{Ag}(\mathrm{s})\right.:
- Ered(Fe3+/Fe2+)=+0.77VE^\circ_{\text{red}}(\mathrm{Fe}^{3+}/\mathrm{Fe}^{2+}) = +0.77 \, \text{V}
- Ered(Ag+/Ag)=+0.80VE^\circ_{\text{red}}(\mathrm{Ag}^{+}/\mathrm{Ag}) = +0.80 \, \text{V}
- Ecello=EcathodeEanode=0.80V0.77V=0.03VE_{\text{cell}}^{o} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0.80 \, \text{V} - 0.77 \, \text{V} = 0.03 \, \text{V}
b. For Pt(s)Mn2+(aq),MnO4(aq)Cu2+(aq)Cu(s)\mathrm{Pt}(\mathrm{s})\left|\mathrm{Mn}^{2+}(\mathrm{aq}), \mathrm{MnO}_{4}^{-}(\mathrm{aq})\right|\left|\mathrm{Cu}^{2+}(\mathrm{aq})\right| \mathrm{Cu}(\mathrm{s}):
- Ered(MnO4/Mn2+)=+1.51VE^\circ_{\text{red}}(\mathrm{MnO}_{4}^{-}/\mathrm{Mn}^{2+}) = +1.51 \, \text{V}
- Ered(Cu2+/Cu)=+0.34VE^\circ_{\text{red}}(\mathrm{Cu}^{2+}/\mathrm{Cu}) = +0.34 \, \text{V}
- Ecello=0.34V1.51V=1.17VE_{\text{cell}}^{o} = 0.34 \, \text{V} - 1.51 \, \text{V} = -1.17 \, \text{V}

STEP 5

Determine Gibbs free energy change and spontaneity:
a. For cell diagram (a):
- ΔG=nFEcello\Delta G = -nFE_{\text{cell}}^{o}
- n=1n = 1 (for Fe3+/Fe2+\mathrm{Fe}^{3+}/\mathrm{Fe}^{2+} and Ag+/Ag\mathrm{Ag}^{+}/\mathrm{Ag})
- ΔG=(1)(96485C/mol)(0.03V)=2894.55J/mol\Delta G = -(1)(96485 \, \text{C/mol})(0.03 \, \text{V}) = -2894.55 \, \text{J/mol}
- Reaction is spontaneous as ΔG<0\Delta G < 0.
b. For cell diagram (b):
- ΔG=(2)(96485C/mol)(1.17V)=226,902.3J/mol\Delta G = -(2)(96485 \, \text{C/mol})(-1.17 \, \text{V}) = 226,902.3 \, \text{J/mol}
- Reaction is nonspontaneous as ΔG>0\Delta G > 0.

STEP 6

Evaluate the spontaneity of the given reaction under specified conditions:
- Use the Nernst equation to find EcellE_{\text{cell}} under non-standard conditions.
- Ecell=EcellRTnFlnQE_{\text{cell}} = E^\circ_{\text{cell}} - \frac{RT}{nF} \ln Q
- Q=[Cr3+]2[Ag+]6[Cr2O72][H+]14Q = \frac{[\mathrm{Cr}^{3+}]^2 [\mathrm{Ag}^{+}]^6}{[\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}][\mathrm{H}^{+}]^{14}}
- Substitute values: Q=(0.6)2(0.675)6(0.675)(102)14Q = \frac{(0.6)^2 (0.675)^6}{(0.675)(10^{-2})^{14}}
- Calculate EcellE_{\text{cell}} and determine if the reaction is spontaneous (Ecell>0E_{\text{cell}} > 0).

SOLUTION

Define the given chemistry terms:
a) Oxidation: The process of losing electrons in a chemical reaction.
b) Oxidizing agent: A substance that gains electrons and is reduced in a chemical reaction.
c) Salt bridge: A device used in an electrochemical cell to maintain electrical neutrality by allowing the flow of ions.
d) Liquid junction: A boundary between two different electrolyte solutions in an electrochemical cell.
e) Standard electrode potential: The measure of the individual potential of a reversible electrode at standard state, which is 1M1 \, \text{M} concentration, 1atm1 \, \text{atm} pressure, and 25C25^\circ \text{C}.

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