Electrochemistry Essentials: Integrated Practice—Reading Tables, Balancing Reactions, and Explaining Observations

How to Use This Integrated Practice Chapter

This chapter is a skills-integration workout: each problem forces you to connect (1) what is oxidized/reduced, (2) a correctly balanced net ionic equation, (3) a prediction from an E° table, (4) cell notation/diagram meaning, and (5) what you would actually observe (mass changes, gas, color change, plating, corrosion patterns). You are expected to use your prior notes for reference; the focus here is on workflow and checking.

Standard workflow (use on every multi-step problem)

• Step A: Identify the redox pair (species changing oxidation state) and write the two half-reactions you expect.
• Step B: Balance each half-reaction (atoms, then charge with e⁻; add H₂O/H⁺/OH⁻ as needed for aqueous conditions).
• Step C: Electron audit: multiply half-reactions so electrons cancel; add to get the net ionic equation.
• Step D: Use the potential table: choose E°_red for the cathode reaction as written; for the anode use the listed reduction potential but subtract (or flip sign) appropriately. Compute E°cell = E°cathode(red) − E°anode(red).
• Step E: Interpret physically: identify anode/cathode, electron direction, ion migration, mass change of electrodes, gas evolution, and any visible changes.

(1) Mixed Practice Sets Grouped by Skill Combination

(a) Oxidation numbers + redox identification (quick triage)

Instructions: For each, (i) mark the atoms that change oxidation number, (ii) name what is oxidized and what is reduced, (iii) state the oxidizing agent and reducing agent.

• A1. 2 Fe2+(aq) + Cl2(g) → 2 Fe3+(aq) + 2 Cl−(aq)
• A2. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
• A3. 2 I−(aq) + H2O2(aq) + 2 H+(aq) → I2(aq) + 2 H2O(l)
• A4. NO3−(aq) + 4 H+(aq) + 3 e− → NO(g) + 2 H2O(l) (already a half-reaction: identify what is reduced and the oxidizing agent form)

Checkpoint: If no oxidation number changes, it is not a redox process (do not force half-reactions).

(b) Half-reaction balancing (then net ionic) under constraints

Instructions: Balance each in the stated medium, then combine into a net ionic equation. Show an electron audit line (electrons lost = electrons gained).

• B1 (acidic). MnO4−(aq) + Fe2+(aq) → Mn2+(aq) + Fe3+(aq)
• B2 (basic). Cr2O7^2−(aq) + SO3^2−(aq) → Cr(OH)3(s) + SO4^2−(aq)
• B3 (acidic). ClO3−(aq) + I−(aq) → Cl−(aq) + I2(aq)

Checkpoint: After combining, verify (i) atoms balance, (ii) total charge balances, (iii) electrons cancel completely.

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(c) E°_cell predictions (spontaneous direction + strength)

Instructions: Use a standard reduction potential table. For each pair, (i) choose the cathode (more positive E°_red), (ii) compute E°_cell, (iii) write the spontaneous net reaction (no need to balance with H₂O/H⁺ unless required by the half-reactions you chose).

• C1. Ag+(aq)/Ag(s) paired with Cu2+(aq)/Cu(s)
• C2. Fe3+(aq)/Fe2+(aq) paired with I2(aq)/I−(aq)
• C3. Zn2+(aq)/Zn(s) paired with H+(aq)/H2(g) (standard hydrogen electrode)

Checkpoint: E°_cell > 0 predicts spontaneity under standard conditions; it does not guarantee fast kinetics or a single product in complex aqueous mixtures.

(d) Full cell notation and diagram interpretation (mapping symbols to reality)

Instructions: For each, (i) label anode/cathode, (ii) state electron flow direction, (iii) state which ions migrate through the salt bridge/membrane, (iv) predict which electrode gains mass, (v) write the net ionic reaction.

• D1. Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
• D2. Pt(s) | Fe2+(aq), Fe3+(aq) || Ag+(aq) | Ag(s)
• D3. A diagram shows a left beaker labeled Sn(s) in Sn2+(aq) connected by a salt bridge to a right beaker labeled Fe3+(aq), Fe2+(aq) with an inert electrode. The voltmeter reads +0.62 V with the positive lead on the right electrode. (Interpret what that sign implies about cathode location.)

Checkpoint: In cell notation, the anode is written on the left by convention; in a diagram, the cathode is the electrode connected to the positive terminal of a voltmeter in a spontaneous cell.

(e) Application cases (battery, corrosion, plating) with observable outcomes

Instructions: For each case, answer: (i) half-reactions, (ii) net reaction, (iii) anode/cathode identification, (iv) one observable outcome, (v) one design or safety implication.

• E1 (battery diagnosis). A galvanic cell uses Mg(s) in Mg2+(aq) and Ag(s) in Ag+(aq). After running, the magnesium strip looks thinner and the silver electrode looks thicker. Explain using half-reactions and E° reasoning.
• E2 (corrosion couple). A steel bolt (Fe) is electrically connected to a copper pipe in moist, salty air. Predict which metal corrodes faster and where reduction occurs. Include the role of dissolved O₂.
• E3 (electroplating). You plate nickel onto a spoon in aqueous NiSO4 using a Ni anode. Predict (i) what forms at the cathode, (ii) what happens to the anode mass, (iii) one competing reaction that may occur if conditions are poor, and (iv) one observation that indicates poor plating quality.
• E4 (electrolysis product prediction). Aqueous NaCl is electrolyzed with inert electrodes. Predict gases at each electrode and what happens to pH near each electrode (qualitative). Then state one reason industrial chlor-alkali cells use membranes/diaphragms.

(2) Guided Solution Templates (Show-Your-Work Formats)

Template 1: From reaction to balanced net ionic + E°_cell + observation

Use when you are given reactants and asked for everything.

Given: (unbalanced) ____________________________________________  Medium: acidic/basic/neutral

1. Identify oxidation/reduction:

   • Oxidized species: __________ (oxidation state ___ → ___)
   • Reduced species: __________ (oxidation state ___ → ___)
   • Reducing agent: __________ Oxidizing agent: __________

2. Write skeleton half-reactions:

   • Oxidation: __________________________________________
   • Reduction: __________________________________________

3. Balance half-reactions (atoms, then charge):

   • Oxidation (balanced): __________________________________
   • Reduction (balanced): __________________________________

4. Electron audit:

   • e⁻ lost = ________ ; e⁻ gained = ________
   • Multiply oxidation by ___ ; multiply reduction by ___

5. Combine and simplify:

   • Net ionic equation: ____________________________________
   • Check: atoms OK? ___ charge OK? ___

6. Compute E°_cell from table:

   • Cathode half-reaction (as reduction): ____________________ E°red = ____ V
   • Anode half-reaction (table reduction form): ______________ E°red = ____ V
   • E°cell = E°cathode(red) − E°anode(red) = ____ V

7. Interpretation:

   • Anode: ________ (oxidation occurs)
   • Cathode: ________ (reduction occurs)
   • Electrode mass change: anode ___ ; cathode ___
   • Expected visible changes (gas/color/solid): _______________

Template 2: From cell notation/diagram to net reaction + E°_cell

Use when you are given a cell diagram or notation.

1. Parse the notation:

   • Left side (anode compartment): __________________________
   • Right side (cathode compartment): _______________________
   • Inert electrode present? ______ (Pt/C/graphite)

2. Write half-reactions:

   • Anode oxidation: _______________________________________
   • Cathode reduction: _____________________________________

3. Balance electrons and add: Net reaction: __________________

4. Compute E°_cell: E°cell = E°cathode(red) − E°anode(red) = ____ V

5. Map to observations:

   • Electron flow (wire): from ________ to ________
   • Salt bridge ion migration: cations → ________ ; anions → ________
   • Which electrode gains mass? ________

(3) Error-Analysis Items (Diagnose and Correct)

Instructions: Each item shows a flawed student solution. Your job: (i) identify at least two specific errors, (ii) rewrite the corrected step(s), (iii) give the corrected final answer.

Error Analysis 1: Wrong sign in E°_cell

Problem: Predict E°_cell for Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s).

Flawed solution: “Cathode is Zn because it is on the left. E°cell = E°(Zn2+/Zn) + E°(Cu2+/Cu) = (−0.76) + (+0.34) = −0.42 V, so the cell is not spontaneous.”

• Find and fix: Identify the incorrect assumption about cathode location and the incorrect formula usage. Provide the corrected anode/cathode assignment and recompute E°_cell.

Error Analysis 2: Electrons not canceled (electron audit failure)

Problem: Balance in acidic solution: MnO4− + Fe2+ → Mn2+ + Fe3+.

Flawed solution: “MnO4− + 8 H+ + 5 e− → Mn2+ + 4 H2O and Fe2+ → Fe3+ + 2 e−. Add them to get MnO4− + 8 H+ + 7 e− + Fe2+ → Mn2+ + 4 H2O + Fe3+.”

• Find and fix: Point out why adding without matching electron counts is invalid. Show the correct multipliers and the correct net ionic equation.

Error Analysis 3: Misreading inert electrodes in cell notation

Problem: Interpret Pt(s) | Fe2+(aq), Fe3+(aq) || Sn4+(aq), Sn2+(aq) | Pt(s).

Flawed solution: “Because Pt is written, platinum is oxidized on the left and reduced on the right. Net reaction: Pt → Pt2+ + 2 e− …”

• Find and fix: Explain what an inert electrode does (and does not do) in notation. Write the correct half-reactions and net reaction based on the solution species.

Error Analysis 4: Electroplating product confusion in aqueous solution

Problem: Electrolyze aqueous CuSO4 using inert electrodes.

Flawed solution: “At the cathode, sodium plates out because metal ions always plate first. At the anode, sulfate turns into sulfur.”

• Find and fix: Identify the incorrect assumptions about which ions are present and which reactions compete in water. State the correct cathode and anode products and one expected observation (color change, gas, mass change).

(4) End-of-Course Checklist of Competencies (Action Statements)

• I can identify oxidized and reduced species by tracking oxidation number changes in a chemical equation.
• I can name the oxidizing agent and reducing agent from a reaction (and justify it using electron gain/loss).
• I can write and balance oxidation and reduction half-reactions in acidic or basic aqueous solution.
• I can perform an electron audit and scale half-reactions so electrons cancel exactly.
• I can combine half-reactions into a correct net ionic equation and verify both atom balance and charge balance.
• I can select appropriate half-reactions from a standard reduction potential table for a proposed cell.
• I can compute E°cell = E°cathode(red) − E°anode(red) and use the sign to predict spontaneity under standard conditions.
• I can determine anode/cathode locations from either cell notation or a diagram (including voltmeter polarity cues).
• I can write full cell notation from a diagram, including inert electrodes when needed.
• I can predict electron flow direction in the external circuit and ion migration directions in the salt bridge/membrane.
• I can connect oxidation/reduction to observable outcomes: electrode mass changes, gas evolution, solution color changes, and solid deposition.
• I can analyze galvanic corrosion scenarios (galvanic coupling, oxygen reduction sites) and predict which metal acts as the anode.
• I can predict electroplating results (what plates at the cathode, what happens at the anode) and identify common competing reactions in water.
• I can diagnose common mistakes: wrong sign in E°_cell, forgetting to cancel electrons, misusing inert electrodes, and assuming “any cation plates” in aqueous electrolysis.