Electrochem Practice Flashcards
A voltaic cell is constructed using a Cu(s) / Cu2+(aq) half–cell connected to a C(s) / Cr2O72–(aq), H+(aq) half–cell.
The correct net reaction of the cell is
A) 3Cu2+(aq) + Cr2O72–(aq) + 14H+(aq) 3Cu(s) + 2Cr3+(aq) + 7H2O(l)
B) 3Cu(s) + Cr2O72–(aq) + 14H+(aq) 3Cu2+(aq) + 2Cr3+(aq) + 7H2O(l)
C) Cu2+(aq) + Cr2O72–(aq) + 14H+(aq) Cu(s) + 2Cr3+(aq) + 7H2O(l)
D) Cu(s) + Cr2O72–(aq) + 14H+(aq) Cu2+(aq) + 2Cr3+(aq) + 7H2O(l)
Explanation: List all possible redox participants:
Cu(s), Cu2+(aq), Cr2O72–(aq), H+(aq)
note: C(s) is an inert electrode and will not be involved.
Next identify the OAs and RAs and chose half reactions for oxidation and reduction. Refer to the Table of Relative Strengths of Oxidizing and Reducing Agents.
cathode:
Cr2O72–(aq) + 14 H+(aq) + 6e– 2Cr3+ +7H2O(l)
anode:
Cu(s) Cu2+(aq) + 2e–
The net reaction requires multiplying the copper half reaction by 3 to balance the electrons. Choice B is the correct equation.
In a breathalyzer, ethanol from the suspect’s breath is oxidized by an acidic dichromate solution in the reaction ampule as in the following reaction:
2Cr2O72–(aq) + 16H+(aq) + 3C2H5OH(aq) 4Cr3+(aq) + 11H2O(l) + 3CH3COOH(aq)
The ampule which contains the acidic dichromate testing solution would best be made of:
A) chromium.
B) zinc.
C) plastic.
D) copper.
Explanation:
Metals like Cr(s), Zn(s), and Cu(s), are all reducing agents and might interfere with the redox reaction between chromium and ethanol. This could produce falsely low or erroneously high results.
An ampule made of a non-reactive material, like plastic, would provide a more accurate test result.
PbO2 + Cl– PbCl2 + Cl2
When the following reaction that occurs in acid solution is balanced, the coefficient in front of Cl2 is
A) 1
B) 2
C) 3
D) 4
Explanation:
reduction half reaction: Pb02 –> PbCl2
balance Cl atoms : PbO2 + Cl– –> PbCl2
balance O atoms using H2O: PbCl2 + 2Cl– –> PbCl2 + 2H2O
balance H atoms using H+: PbO2 + 2Cl– + 4H+ –> PbCl2 + 2H2O
balance charge by adding e–: PbP2 + 2Cl– 4H+ + 2e– –> PbCl + 2H2O (1)
oxidation half reaction: Cl– –> Cl2
balance the Cl atoms: 2Cl– –> 2Cl2
balance charge using e–: 2Cl– –> Cl2 + 2e– (2)
Gain of e– in reduction must equal the loss of e– in oxidation
Add equations (1) and (2). PbO2 + 4Cl– + 4H+ PbCl2 + 2H2O + CL2
Which of the following generalizations is true of the Table of Selected Standard Electrode Potentials?
A) Metal ions and non-metallic elements are generally oxidizing agents.
B) Metal ions and non-metallic ions are generally oxidized.
C) Metals and non-metal elements are generally oxidized.
D) Metals and non-metal ions are generally reduced.
Explanation: Metals in elemental form behave as RAs (reducing agents, i.e. they are oxidized, they rust). When the reaction goes in the other direction, the metal ions are OAs (oxidizing agents). Non-metal ions are usually RAs and non-metal elements are OAs. Remember that OAs reduce and RAs oxidize. Choices B, C and D are all false and only choice A is a true statement.
Galvanizing is a process used to prevent corrosion. The process of galvanization involves coating iron or steel with a thin layer of zinc metal.
In an electrolytic cell, 100.0 g of Zn(s) was plated in 2.50 h. The current that would have to be supplied (to the nearest tenth) is ____________ A.
Explanation:
Zn2+(aq) + 2e– Zn(s)
Time = 2.5 h × 60 min/h × 60 s/min = 9000 s
moles of Zn(s) = nZn = g of Zn / molar mass of Zn = 100.0 g /65.4 g/mol = 1.53 mol
moles of electrons = ne– = nZn × 2 = 1.53 × 2 = 3.06 mol
Faraday constant = F = 9.65 × 104 C/mol = 9.65 × 104 A/(s × mol)
ne– = It / F I = ne– × F / t = 3.06 mol × 9.65 × 104 A/(s × mol) / 9000 s = 32.8 A
One way in which electrolytic cells differ from voltaic cells is that
A) electrons transfer from the anode in one, but from the cathode in the other
B) one cell requires energy while the other releases energy
C) electrolytic cells produce energy but voltaic consume it
D) oxidation occurs at the anode at one, but at the cathode in the other
Explanation: Electrolytic cells force a reaction which requires an input of energy to do the work. In voltaic cells the redox reaction is spontaneous and results in a release of energy. Choice C declares the opposite and is therefore false. In both electrolytic and voltaic cells, the movement of electrons is from the anode, the site of oxidation, to the cathode, the site of reduction, so choices A and D are also false.
Hydrochloric acid may be safely stored in a container made of:
A) copper.
B) cadmium.
C) iron.
D) aluminum.
Explanation:
Refer to a Table of Relative Strengths of Oxidizing and Reducing Agents.
Look for the placement of the half reactions of the metals and H+(aq). Only Cu is above H+, indicating that the redox reaction: 2H+(aq) + Cu(s) H2(g) + Cu2+(aq) would not be spontaneous, and the acid would be safe in a copper container.
All the other metals are below the 2H+(aq) + 2e– H2 reaction, so would all be stronger oxidizing agents than H+. The HCL would react spontaneously with Ca, Fe, and Al, and the breakdown of the metal containers would make the storage of acid in these unsafe.
In a reaction, Fe2+(aq):
A) will undergo reduction when combined with Pb(s).
B) will act as an oxidizing agent when combined with Sn(s).
C) will always act as an oxidizing agent.
D) will act as an reducing agent when combined with Ag+(aq).
Explanation:
Refer to the Table of Relative Strengths of Oxidizing and Reducing Agents.
Fe2+(aq) appears twice in the table; once as a RA and further down as an OA, so choice C) is untrue.
To have a spontaneous reaction as a RA, Fe2+(aq) must appear below and in the right hand column relative to the OA, as is the case with Ag+(aq). Choice D) is correct. Fe2+(aq) as OA is below both Pb(s) and Sn(s) as RA so no reactions will occur.
One of the main reasons that copper pipes are used in household plumbing (rather than iron pipes) is that
A) copper is a better conductor of heat
B) iron will react with dissolved hard water minerals such as calcium ions
C) iron has a greater tendency to act as a reducing agent than copper
D) drain cleaners and soaps containing sodium hydroxide will react with iron
Explanation: Pipes will last longer and not contaminate the fluid they carry if they are inert and non-reactive. The fact that Cu(s) is above Fe(s) on the Table of relative Strengths of Oxidizing and Reducing Agents indicates that copper metal is a weaker reducing agent than iron and so will be less reactive. Although the other choices are not totally incorrect, Choice C is a better reason to choose copper over iron.
During photosynthesis, CO2(g) + H2O(g) + energy C6H12O6(aq) + O2(g),
A) carbon in carbon dioxide is reduced.
B) hydrogen in water is reduced.
C) oxygen in carbon dioxide and/or water is reduced.
D) hydrogen in glucose is oxidized.
Photosynthesis is a series of redox reactions. In the above summary reaction, the oxidation state of O changes from 2– in water to 0 in oxygen gas indicating a loss of electrons, or oxidation. The reduction occurs on the C in carbon dioxide (oxidation number 4+) when it becomes part of the glucose molecule. The oxidation state of a hydrogen ion is always 1+.
Under standard conditions, hydrogen gas reacts with Hg2+(aq) ions to produce Hg(s). The net cell potential for this reaction (to the nearest hundredth) is ± ____________ V.
Explanation: See Table of Relative Strengths of Oxidizing and Reducing Agents.
Cathode: Hg2+(aq) + 2e– Hg(s) E° = +0.85 V
Anode: H2(g) 2H+(aq) + 2e– E° = 0.00 V
net cell potential = E° cathode – E° anode
= + 0.85 V – 0.00 V = 0.85 V
When nickel metal is placed into aqueous lead (II) nitrate a spontaneous reaction is observed. This reaction is best explained by theorizing that:
A) nickel (II) ions have a greater attraction for electrons than do lead (II) ions.
B) nickel (II) ions have a lesser attraction for electrons than do nickel atoms.
C) lead (II) ions have a greater attraction for electrons than do nickel (II) ions.
D) lead (II) ions have a lesser attraction for electrons than do lead atoms.
Explanation:
Since the reaction is spontaneous, lead(II), as the RA(reducing agent), must be below metallic nickel atoms, the OA (oxidizing agent), on the Table of Relative Strengths of Oxidizing and Reducing Agents.
The half reactions are:
Ni(s) Ni2+(aq) + 2e–
Pb2+(aq) + 2e– Pb(s)
Nickel atoms are reduced and lead (II) ions are oxidized. Lead(II) ions have a stronger attraction for electrons than nickel(II) ions and than lead atoms.
Properties:
1) Reacts spontaneously with Zn2+
2) Reacts spontaneously with Ag
3) Is an oxidizing and a deducting Agent
4) Is reduced by Au
5) Reacts spontaneously with K+
6) Reacts spontaneously with Al3+
Which property in the list is most appropriate for Hg2+(aq)?
Explanation:
Refer to the Table of Relative Strengths of Oxidizing and Reducing Agents and apply the spontaneity rule by determining the relative position of the OA and RA.
Hg2+, the OA is higher on the table than Ag(s), the RA, so the reaction would be spontaneous.
Zn2+, K+ and Al3+ as RAs are all above Hg2+ as OA so no spontaneous reaction would occur, ruling out choices 1, 5 & 6. The half-cell reduction potentials show that Au cannot react spontaneously as a RA with Hg2+.
Under standard conditions, hydrogen gas reacts with Ag+(aq) ions to produce Ag(s). The net cell potential for this reaction (to the nearest hundredth) is ± ____________ V.
Explanation:
Cathode: Ag+(aq) + e– Ag(s) (E° = +0.80 V)
Anode: H2(g) 2H+(aq) + 2e– (E° = 0.00 V)
Values for standard cell potentials are found on the Table of Relative Strengths of Oxidizing and Reducing Agents.
net cell potential = E° cathode – E° anode
= + 0.80 V – 0.00 V = 0.80 V
A reducing agent can be described as a chemical that:
A) loses electrons and becomes oxidized.
B) loses electrons and causes oxidation.
C) gains electrons and causes reduction.
D) gains electrons and becomes reduced.
Explanation:
A reducing agent causes reduction. A reducing agent undergoes oxidation, which means it loses electrons and its oxidation number increases.
