Chapter 18 – Review
18.1 – Redox Reactions and Oxidation Numbers
- Determine the oxidation states of the elements in the following compounds:
- NaI
- GdCl3
- LiNO3
- H2Se
- Mg2Si
- RbO2, rubidium superoxide
- HF
- Determine the oxidation states of the elements in the compounds listed. None of the oxygen-containing compounds are peroxides or superoxides.
- H3PO4
- Al(OH)3
- SeO2
- KNO2
- In2S3
- P4O6
Check Answer: [1]
- Determine the oxidation states of the elements in the compounds listed. None of the oxygen-containing compounds are peroxides or superoxides.
- H2SO4
- Ca(OH)2
- BrOH
- ClNO2
- TiCl4
- NaH
- Determine which elements are oxidized and which are reduced in the following reactions:
- 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g)
- Mg(s) + Cl2(g) → MgCl2(aq)
- K3P(s) + 2O2(g) → K3PO4(s)
Check Answer: [2]
- Identify the atoms that are oxidized and reduced, the change in oxidation state for each, and the oxidizing and reducing agents in each of the following equations:
- Mg(s) + NiCl2(aq) → MgCl2(aq) + Ni(s)
- PCl3(l) + Cl2(g) → PCl5(s)
- C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(g)
- Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
- 2K2S2O3(s) + I2(s) → 2K2S4O6(s) + 2KI(s)
18.2 – Balancing Redox Reactions
- Complete and balance each of the following half-reactions (steps 2–5 in half-reaction method):
- Sn4+(aq) → Sn2+(aq)
- [Ag(NH3)2]+(aq) → Ag(s) + NH3(aq)
- Hg2Cl2(s) → Hg(l) + Cl−(aq)
- H2O(l) → O2(g) (in acidic solution)
- IO3−(aq) → I2(s)
- SO32−(aq) → SO42−(aq) (in acidic solution)
- MnO4−(aq) → Mn2+(aq) (in acidic solution)
- Cl−(aq) → ClO3−(aq) (in basic solution)Check Answer: [3]
- Complete and balance each of the following half-reactions (steps 2–5 in half-reaction method):
- Cr2+(aq) → Cr3+(aq)
- Hg(l) + Br−(aq) → HgBr42−(aq)
- ZnS(s) → Zn(s) + S2−(aq)
- H2(g) → H2O(l) (in basic solution)
- H2(g) → H3O+(aq) (in acidic solution)
- NO3−(aq) → HNO2(aq) (in acidic solution)
- MnO2(s) → MnO4−(aq) (in basic solution)
- Cl−(aq) → ClO3−(aq) (in acidic solution)
- Balance each of the following equations according to the half-reaction method:
- Sn2+(aq) + Cu2+(aq) → Sn4+(aq) + Cu+(aq)
- H2S(g) + Hg22+(aq) → Hg(l) + S(s) (in acid)
- CN−(aq) + ClO2(aq) → CNO−(aq) + Cl−(aq) (in acid)
- Fe2+(aq) + Ce4+(aq) → Fe3+(aq) + Ce3+(aq)
- HBrO(aq) → Br−(aq) + O2(g) (in acid)
Check Answer: [4]
- Balance each of the following equations according to the half-reaction method:
- Zn(s) + NO3−(aq) → Zn2+(aq) + N2(g) (in acid)
- Zn(s) + NO3−(aq) → Zn2+(aq) + NH3(aq) (in base)
- CuS(s) + NO3−(aq) → Cu2+ + S(s) + NO(g) (in acid)
- NH3(aq) + O2(g) → NO2(g) (gas phase)
- Cl2(g) + OH−(aq) → Cl−(aq) + ClO3−(aq) (in base)
- H2O2(aq) + MnO4−(aq) → Mn2+(aq) + O2(g) (in acid)
- NO2(g) → NO3−(aq) + NO2−(aq) (in base)
- Fe3+(aq) + I−(aq) → Fe2+(aq) + I2(aq)
- Balance each of the following equations according to the half-reaction method:
- MnO4−(aq) + NO2−(aq) → MnO2(s) + NO3−(aq) (in base)
- MnO42−(aq) → MnO4−(aq) + MnO2(s) (in base)
- Br2(l) + SO2(g) → Br−(aq) + SO42−(aq) (in acid)
Check Answer: [5]
18.3 – Galvanic Cells
- Write the following balanced reactions using cell notation. Use platinum as an inert electrode, if needed.
(a) [latex]\text{Mg}(s)\;+\;\text{Ni}^{2+}(aq)\;{\longrightarrow}\;\text{Mg}^{2+}(aq)\;+\;\text{Ni}(s)[/latex]
(b) [latex]2\text{Ag}^{+}(aq)\;+\;\text{Cu}(s)\;{\longrightarrow}\;\text{Cu}^{2+}(aq)\;+\;2\text{Ag}(s)[/latex]
(c) [latex]\text{Mn}(s)\;+\;\text{Sn(NO}_3)_2(aq)\;{\longrightarrow}\;\text{Mn(NO}_3)_2(aq)\;+\;\text{Au}(s)[/latex]
(d) [latex]3\text{CuNO}_3(aq)\;+\;\text{Au(NO}_3)_3(aq)\;{\longrightarrow}\;3\text{Cu(NO}_3)_2(aq)\;+\;\text{Au}(s)[/latex]
Check Answer: [6] - Given the following cell notations, determine the species oxidized, species reduced, and the oxidizing agent and reducing agent, without writing the balanced reactions.(a) [latex]\text{Mg}(s){\mid}\text{Mg}^{2+}(aq){\parallel}\text{Cu}^{2+}(aq){\mid}\text{Cu}(s)[/latex](b) [latex]\text{Ni}(s){\mid}\text{Ni}^{2+}(aq){\parallel}\text{Ag}^{+}(aq){\mid}\text{Ag}(s)[/latex]
- For the cell notations in the previous problem, write the corresponding balanced reactions.
Check Answer: [7] - Balance the following reactions and write the reactions using cell notation. Ignore any inert electrodes, as they are never part of the half-reactions.
(a) [latex]\text{Al}(s)\;+\;\text{Zr}^{4+}(aq)\;{\longrightarrow}\;\text{Al}^{3+}(aq)\;+\;\text{Zr}(s)[/latex]
(b) [latex]\text{Ag}^{+}(aq)\;+\;\text{NO}(g)\;{\longrightarrow}\;\text{Ag}(s)\;+\;\text{NO}_3^{\;\;-}(aq)\;\;\;\;\;\;\;\text{(acidic solution)}[/latex]
(c) [latex]\text{SiO}_3^{\;\;2-}(aq)\;+\;\text{Mg}(s)\;{\longrightarrow}\;\text{Si}(s)\;+\;\text{Mg(OH)}_2(s)\;\;\;\;\;\;\;\text{(basic solution)}[/latex]
(d) [latex]\text{ClO}_3^{\;\;-}(aq)\;+\;\text{MnO}_2(s)\;{\longrightarrow}\;\text{Cl}^{\;\;-}(aq)\;+\;\text{MnO}_4^{\;\;-}(aq)\;\;\;\;\;\;\;\text{(basic solution)}[/latex]
- Identify the species oxidized, species reduced, and the oxidizing agent and reducing agent for all the reactions in the previous problem.
Check Answer: [8] - From the information provided, use cell notation to describe the following systems:(a) In one half-cell, a solution of Pt(NO3)2 forms Pt metal, while in the other half-cell, Cu metal goes into a Cu(NO3)2 solution with all solute concentrations 1 M.(b) The cathode consists of a gold electrode in a 0.55 M Au(NO3)3 solution and the anode is a magnesium electrode in 0.75 M Mg(NO3)2 solution.(c) One half-cell consists of a silver electrode in a 1 M AgNO3 solution, and in the other half-cell, a copper electrode in 1 M Cu(NO3)2 is oxidized.
- Why is a salt bridge necessary in galvanic cells like the one in Figure 18.3b?
Check Answer: [9] - An active (metal) electrode was found to gain mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode part of the anode or cathode? Explain.
- An active (metal) electrode was found to lose mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode part of the anode or cathode? Explain. Check Answer: [10]
- The mass of three different metal electrodes, each from a different galvanic cell, were determined before and after the current generated by the oxidation-reduction reaction in each cell was allowed to flow for a few minutes. The first metal electrode, given the label A, was found to have increased in mass; the second metal electrode, given the label B, did not change in mass; and the third metal electrode, given the label C, was found to have lost mass. Make an educated guess as to which electrodes were active and which were inert electrodes, and which were anode(s) and which were the cathode(s).
18.4 – Electrode and Cell Potentials
- For each reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard conditions.
(a) [latex]\text{Mg}(s)\;+\;\text{Ni}^{2+}(aq)\;{\longrightarrow}\;\text{Mg}^{2+}(aq)\;+\;\text{Ni}(s)[/latex]
(b) [latex]2\text{Ag}^{+}(aq)\;+\;\text{Cu}(s)\;{\longrightarrow}\;\text{Cu}^{2+}(aq)\;+\;2\text{Ag}(s)[/latex]
(c) [latex]\text{Mn}(s)\;+\;\text{Sn(NO}_3)_2(aq)\;{\longrightarrow}\;\text{Mn(NO}_3)_2(aq)\;+\;\text{Sn}(s)[/latex]
(d) [latex]3\text{Fe(NO}_3)_2(aq)\;+\;\text{Au(NO}_3)_3(aq)\;{\longrightarrow}\;3\text{Fe(NO}_3)_3(aq)\;+\;\text{Au}(s)[/latex]
Check Answer: [11] - For each reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard conditions.
(a) [latex]\text{Mn}(s)\;+\;\text{Ni}^{2+}(aq)\;{\longrightarrow}\;\text{Mn}^{2+}(aq)\;+\;\text{Ni}(s)[/latex]
(b) [latex]3\text{Cu}^{2+}(aq)\;+\;2\text{Al}(s)\;{\longrightarrow}\;2\text{Al}^{3+}(aq)\;+\;2\text{Cu}(s)[/latex]
(c) [latex]\text{Na}(s)\;+\;\text{LiNO}_3(aq)\;{\longrightarrow}\;\text{NaNO}_3(aq)\;+\;\text{Li}(s)[/latex]
(d) [latex]\text{Ca(NO}_3)_2(aq)\;+\;\text{Ba}(s)\;{\longrightarrow}\;\text{Ba(NO}_3)_2(aq)\;+\;\text{Ca}(s)[/latex]
- Determine the overall reaction and its standard cell potential at 25 °C for this reaction. Is the reaction spontaneous at standard conditions?
[latex]\text{Cu}(s){\mid}\text{Cu}^{2+}(aq){\parallel}\text{Au}^{3+}(aq){\mid}\text{Au}(s)[/latex]
Check Answer: [12] - Determine the overall reaction and its standard cell potential at 25 °C for the reaction involving the galvanic cell made from a half-cell consisting of a silver electrode in 1 M silver nitrate solution and a half-cell consisting of a zinc electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?
- Determine the overall reaction and its standard cell potential at 25 °C for the reaction involving the galvanic cell in which cadmium metal is oxidized to 1 M cadmium(II) ion and a half-cell consisting of an aluminum electrode in 1 M aluminum nitrate solution. Is the reaction spontaneous at standard conditions?
Check Answer: [13] - Determine the overall reaction and its standard cell potential at 25 °C for these reactions. Is the reaction spontaneous at standard conditions? Assume the standard reduction for Br2(l) is the same as for Br2(aq).
[latex]\text{Pt}(s){\mid}\text{H}_2(g){\mid}\text{H}^{+}(aq){\parallel}\text{Br}_2(aq){\mid}\text{Br}^{-}(aq){\mid}\text{Pt}(s)[/latex]
18.5 – Batteries and Fuel Cells
- What are the desirable qualities of an electric battery?
- List some things that are typically considered when selecting a battery for a new application.
Check Answer: [14] - Consider a battery made from one half-cell that consists of a copper electrode in 1 M CuSO4 solution and another half-cell that consists of a lead electrode in 1 M Pb(NO3)2 solution.(a) What are the reactions at the anode, cathode, and the overall reaction?(b) What is the standard cell potential for the battery?(c) Most devices designed to use dry-cell batteries can operate between 1.0 and 1.5 V. Could this cell be used to make a battery that could replace a dry-cell battery? Why or why not.(d) Suppose sulfuric acid is added to the half-cell with the lead electrode and some PbSO4(s) forms. Would the cell potential increase, decrease, or remain the same?
- Consider a battery with the overall reaction:
[latex]\text{Cu}(s)\;+\;2\text{Ag}^{+}(aq)\;{\longrightarrow}\;2\text{Ag}(s)\;+\;\text{Cu}^{2+}(aq)[/latex].
(a) What is the reaction at the anode and cathode?
(b) A battery is “dead” when it has no cell potential. What is the value of Q when this battery is dead?
(c) If a particular dead battery was found to have [Cu2+] = 0.11 M, what was the concentration of silver ion?
Check Answer: [15] - An inventor proposes using a SHE (standard hydrogen electrode) in a new battery for smartphones that also removes toxic carbon monoxide from the air:
[latex]\begin{array}{lr @{{}\longrightarrow{}} ll} \text{Anode:} & \text{CO}(g)\;+\;\text{H}_2\text{O}(l) & \text{CO}_2(g)\;+\;2\text{H}^{+}(aq)\;+\;2\text{e}^{-} & E_{\text{anode}}^{\circ} = -0.53\;\text{V} \\[0.5em] \text{Cathode:} & 2\text{H}^{+}(aq)\;+\;2\text{e}^{-} & \text{H}_2(g) & E_{\text{cathode}}^{\circ} = 0\;\text{V} \\[0.5em] \hline \\[-0.25em] \text{Overall:} & \text{CO}(g)\;+\;\text{H}_2\text{O}(l) & \text{CO}_2(g)\;+\;\text{H}_2(g) & E_{\text{cell}}^{\circ} = +0.53\;\text{V} \end{array}[/latex] Would this make a good battery for smartphones? Why or why not? - Why do batteries go dead, but fuel cells do not? Check Answer: [16]
- Explain what happens to battery voltage as a battery is used, in terms of the Nernst equation.
- Using the information thus far in this chapter, explain why battery-powered electronics perform poorly in low temperatures. Check Answer: [17]
18.6 – Corrosion
- Which member of each pair of metals is more likely to corrode (oxidize)?(a) Mg or Ca(b) Au or Hg(c) Fe or Zn(d) Ag or Pt
- Consider the following metals: Ag, Au, Mg, Ni, and Zn. Which of these metals could be used as a sacrificial anode in the cathodic protection of an underground steel storage tank? Steel is mostly iron, so use −0.447 V as the standard reduction potential for steel.
Check Answer: [18] - Aluminum [latex](E_{\text{Al}^{3+}/\text{Al}}^{\circ} = -2.07\;\text{V})[/latex] is more easily oxidized than iron [latex](E_{\text{Fe}^{3+}/\text{Fe}}^{\circ} = -0.477\;\text{V})[/latex], and yet when both are exposed to the environment, untreated aluminum has very good corrosion resistance while the corrosion resistance of untreated iron is poor. Explain this observation.
- If a sample of iron and a sample of zinc come into contact, the zinc corrodes but the iron does not. If a sample of iron comes into contact with a sample of copper, the iron corrodes but the copper does not. Explain this phenomenon.
Check Answer: [19] - Suppose you have three different metals, A, B, and C. When metals A and B come into contact, B corrodes and A does not corrode. When metals A and C come into contact, A corrodes and C does not corrode. Based on this information, which metal corrodes and which metal does not corrode when B and C come into contact?
- Why would a sacrificial anode made of lithium metal be a bad choice despite its [latex]E_{\text{Li}^{+}/\text{Li}}^{\circ} = -3.04\;\text{V}[/latex], which appears to be able to protect all the other metals listed in the standard reduction potential table?
Check Answer: [20]
18.7 – Electrolysis
- Identify the reaction at the anode, reaction at the cathode, the overall reaction, and the approximate potential required for the electrolysis of the following molten salts. Assume standard states and that the standard reduction potentials in Appendix M are the same as those at each of the melting points. Assume the efficiency is 100%.
(a) CaCl2
(b) LiH
(c) AlCl3
(d) CrBr3 - What mass of each product is produced in each of the electrolytic cells of the previous problem if a total charge of 3.33 × 105 C passes through each cell? Assume the voltage is sufficient to perform the reduction.
Check Answer: [21] - How long would it take to reduce 1 mole of each of the following ions using the current indicated? Assume the voltage is sufficient to perform the reduction.(a) Al3+, 1.234 A(b) Ca2+, 22.2 A(c) Cr5+, 37.45 A(d) Au3+, 3.57 A
- A current of 2.345 A passes through the cell shown in Figure 18.7b for 45 minutes. What is the volume of the hydrogen collected at room temperature if the pressure is exactly 1 atm? Assume the voltage is sufficient to perform the reduction. (Hint: Is hydrogen the only gas present above the water?)
Check Answer: [22] - An irregularly shaped metal part made from a particular alloy was galvanized with zinc using a Zn(NO3)2 solution. When a current of 2.599 A was used, it took exactly 1 hour to deposit a 0.01123-mm layer of zinc on the part. What was the total surface area of the part? The density of zinc is 7.140 g/cm3. Assume the efficiency is 100%.
Attribution & References
Except where otherwise noted, this page is adapted by David Wegman from “7.2 Classifying Chemical Reactions”, “17.2 Galvanic Cells”, “17.3 Standard Reduction Potentials”, “17.5 Batteries and Fuel Cells”, “17.6 Corrosion”, “17.7 Electrolysis” In General Chemistry 1 & 2 by Rice University, a derivative of Chemistry (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson and is licensed under CC BY 4.0. Access for free at Chemistry (OpenStax). / Extracted end of chapter exercises from 7.2, 17.2, 17.3, 17.5, 17.6, 17.7 for use on this page.
- (a) H +1, P +5, O −2; (b) Al +3, H +1, O −2; (c) Se +4, O −2; (d) K +1, N +3, O −2; (e) In +3, S −2; (f) P +3, O −2 ↵
- (a) Na is oxidized, H+ is reduced; (b) Mg is oxidized, Cl2 is reduced; (c) P3− is oxidized, O2 is reduced; ↵
- (a)Sn4+(aq) + 2e− → Sn2+(aq) (b)[Ag(NH3)2]+(aq) + e− → Ag(s) + 2NH3(aq) (c) Hg2Cl2(s) + 2e− → 2Hg(l) + 2Cl−(aq) (d) 2H2O(l) → O2 + 4H+(aq) + 4e− (e) 6H2O(l) + 2IO3−(aq) + 10e− → I2(s) + 12OH−(aq) (f) H2O(l) + SO32−(aq) → SO42−(aq) + 2H+(aq) + 2e− (g) 8H+(aq) + MnO4−(aq) + 5e− → Mn2+(aq) + 4H2O(l) (a) Cl−(aq) + 6OH−(aq) → ClO3−(aq) + 3H2O(l) + 6e− ↵
- (a) Sn2+(aq) + 2Cu2+(aq) → Sn4+(aq) + 2Cu+(aq) (b) H2S(g) + Hg22+(aq) + 2H2O(l) → 2Hg(l) + S(s) + 2H3O+(aq) (c) 5CN−(aq) + 2ClO2(aq) + 3H2O(l) → 5CNO−(aq) + 2Cl−(aq) + 2H3O+(aq) (d)Fe2+(aq) + Ce4+(aq) → Fe3+(aq) + Ce3+(aq) (e) 2HBrO(aq) + 2H2O(l) → 2H3O(aq) + 2Br−(aq) + O2(g) ↵
- (a) 2MnO4−(aq) + 3NO2−(aq) + H2O(l) → 2MnO2(s) + 3NO3−(aq) + 2OH−(aq) (b) 3MnO42−(aq) + 2H2O(l) → 2MnO4−(aq) + 4OH−(aq) + MnO2(s) (in base) (c) Br2(l) + SO2(g) + 2H2O(l) → 4H+(aq) + 2Br−(aq) + SO42−(aq) ↵
- (a) [latex]\text{Mg}(s){\mid}\text{Mg}^{2+}(aq){\parallel}\text{Ni}^{2+}(aq){\mid}\text{Ni}(s)[/latex]; (b) [latex]\text{Cu}(s){\mid}\text{Cu}^{2+}(aq){\parallel}\text{Ag}^{+}(aq){\mid}\text{Ag}(s)[/latex]; (c) [latex]\text{Mn}(s){\mid}\text{Mn}^{2+}(aq){\parallel}\text{Sn}^{2+}(aq){\mid}\text{Sn}(s)[/latex]; (d) [latex]\text{Pt}(s){\mid}\text{Cu}^{+}(aq)\text{,\;Cu}^{2+}(aq){\parallel}\text{Au}^{3+}(aq){\mid}\text{Au}(s)[/latex] ↵
- (a) [latex]\text{Mg}(s)\;+\;\text{Cu}^{2+}(aq)\;{\longrightarrow}\;\text{Mg}^{2+}(aq)\;+\;\text{Cu}(s)[/latex]; (b) [latex]2\text{Ag}^{+}(aq)\;+\;\text{Ni}(s)\;{\longrightarrow}\;\text{Ni}^{2+}(aq)\;+\;2\text{Ag}(s)[/latex] ↵
- Species oxidized = reducing agent: (a) Al(s); (b) NO(g); (c) Mg(s); and (d) MnO2(s); Species reduced = oxidizing agent: (a) Zr4+(aq); (b) Ag+(aq); (c) [latex]\text{SiO}_3^{\;\;2-}(aq)[/latex]; and (d) [latex]\text{ClO}_3^{\;\;-}(aq)[/latex] ↵
- Without the salt bridge, the circuit would be open (or broken) and no current could flow. With a salt bridge, each half-cell remains electrically neutral and current can flow through the circuit. ↵
- Active electrodes participate in the oxidation-reduction reaction. Since metals form cations, the electrode would lose mass if metal atoms in the electrode were to oxidize and go into solution. Oxidation occurs at the anode. ↵
- (a) +2.115 V (spontaneous); (b) +0.4626 V (spontaneous); (c) +1.0589 V (spontaneous); (d) +0.727 V (spontaneous) ↵
- [latex]3\text{Cu}(s)\;+\;2\text{Au}^{3+}(aq)\;{\longrightarrow}\;3\text{Cu}^{2+}(aq)\;+\;2\text{Au}(s)[/latex]; +1.16 V; spontaneous ↵
- [latex]3\text{Cd}(s)\;+\;2\text{Al}^{3+}(aq)\;{\longrightarrow}\;3\text{Cd}^{2+}(aq)\;+\;2\text{Al}(s)[/latex]; −1.259 V; nonspontaneous ↵
- Considerations include: cost of the materials used in the battery, toxicity of the various components (what constitutes proper disposal), should it be a primary or secondary battery, energy requirements (the “size” of the battery/how long should it last), will a particular battery leak when the new device is used according to directions, and its mass (the total mass of the new device). ↵
- (a) [latex]\begin{array}{lr @{{}\longrightarrow{}} ll} \text{anode:} & \text{Cu}(s) & \longrightarrow \text{Cu}^{2+}(aq)\;+\;2\text{e}^{-} & E_{\text{anode}}^{\circ} = 0.34\;\text{V} \\[0.5em] \text{cathode:} & 2\;\times\;(\text{Ag}^{+}(aq)\;+\;\text{e}^{-} & \longrightarrow \text{Ag}(s)) & E_{\text{cathode}}^{\circ} = 0.7996\;\text{V} \end{array}[/latex] ↵
- Batteries are self-contained and have a limited supply of reagents to expend before going dead. Alternatively, battery reaction byproducts accumulate and interfere with the reaction. Because a fuel cell is constantly resupplied with reactants and products are expelled, it can continue to function as long as reagents are supplied. ↵
- Ecell, as described in the Nernst equation, has a term that is directly proportional to temperature. At low temperatures, this term is decreased, resulting in a lower cell voltage provided by the battery to the device—the same effect as a battery running dead. ↵
- Mg and Zn ↵
- Both examples involve cathodic protection. The (sacrificial) anode is the metal that corrodes (oxidizes or reacts). In the case of iron (−0.447 V) and zinc (−0.7618 V), zinc has a more negative standard reduction potential and so serves as the anode. In the case of iron and copper (0.34 V), iron has the smaller standard reduction potential and so corrodes (serves as the anode). ↵
- While the reduction potential of lithium would make it capable of protecting the other metals, this high potential is also indicative of how reactive lithium is; it would have a spontaneous reaction with most substances. This means that the lithium would react quickly with other substances, even those that would not oxidize the metal it is attempting to protect. Reactivity like this means the sacrificial anode would be depleted rapidly and need to be replaced frequently. (Optional additional reason: fire hazard in the presence of water.) ↵
- (a) [latex]\begin{array}{r @{{}={}} l} \text{mass Ca} & 69.1\;\text{g} \\[0.5em] \text{mass Cl}_2 & 122\;\text{g} \end{array}[/latex] (b) [latex]\begin{array}{r @{{}={}} l} \text{mass Li} & 23.9\;\text{g} \\[0.5em] \text{mass H}_2 & 3.48\;\text{g} \end{array}[/latex] (c) [latex]\begin{array}{r @{{}={}} l} \text{mass Al} & 31.0\;\text{g} \\[0.5em] \text{mass Cl}_2 & 122 \;\text{g} \end{array}[/latex] (d) [latex]\begin{array}{r @{{}={}} l} \text{mass Cr} & 59.8\;\text{g} \\[0.5em] \text{mass Br}_2 & 276\;\text{g} \end{array}[/latex] ↵
- 0.79 L ↵