krit.club logo

Redox Processes - Electrochemical cells

Grade 12IBChemistry

Review the key concepts, formulae, and examples before starting your quiz.

🔑Concepts

•

Electrochemical cells are categorized into two types: Voltaic (Galvanic) cells, which convert chemical energy into electrical energy via spontaneous reactions, and Electrolytic cells, which use electrical energy to drive non-spontaneous reactions.

•

In all electrochemical cells, oxidation occurs at the anode and reduction occurs at the cathode (AN OXAN\ OX, RED CATRED\ CAT).

•

In a Voltaic cell, the anode is the negative electrode and the cathode is the positive electrode. Electrons flow through the external circuit from the anode to the cathode.

•

A salt bridge is essential in Voltaic cells to maintain electrical neutrality by allowing the migration of ions (K+K^+ and NO3−NO_3^- are commonly used) between half-cells.

•

The Standard Hydrogen Electrode (SHESHE) serves as the universal reference with a potential of 0.00 V0.00\ V under standard conditions: 298 K298\ K, 100 kPa100\ kPa for H2(g)H_2(g), and 1.0 mol dm−31.0\ mol\ dm^{-3} for H+(aq)H^+(aq).

•

Standard Electrode Potential (E⊖E^\ominus) measures the tendency of a species to be reduced. A more positive E⊖E^\ominus value indicates a stronger oxidizing agent.

•

Cell diagrams are written in shorthand: Zn(s)∣Zn2+(aq)∣∣Cu2+(aq)∣Cu(s)Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s), where the single line (∣|) represents a phase boundary and the double line (∣∣||) represents the salt bridge.

•

The spontaneity of a reaction is determined by the Gibbs free energy change (ΔG⊖\Delta G^\ominus). A reaction is spontaneous if Ecell⊖>0E^\ominus_{cell} > 0, which results in ΔG⊖<0\Delta G^\ominus < 0.

📐Formulae

Ecell⊖=Ecathode⊖−Eanode⊖E^\ominus_{cell} = E^\ominus_{cathode} - E^\ominus_{anode}

ΔG⊖=−nFEcell⊖\Delta G^\ominus = -nFE^\ominus_{cell}

Q=I×tQ = I \times t

n(e−)=QFn(e^-) = \frac{Q}{F}

F≈9.65×104 C mol−1F \approx 9.65 \times 10^4\ C\ mol^{-1}

💡Examples

Problem 1:

Calculate the standard cell potential (Ecell⊖E^\ominus_{cell}) for a cell consisting of a Mg/Mg2+Mg/Mg^{2+} half-cell and a Ag/Ag+Ag/Ag^+ half-cell. Given: E⊖(Mg2+/Mg)=−2.37 VE^\ominus(Mg^{2+}/Mg) = -2.37\ V and E⊖(Ag+/Ag)=+0.80 VE^\ominus(Ag^+/Ag) = +0.80\ V.

Solution:

Ecell⊖=+0.80 V−(−2.37 V)=+3.17 VE^\ominus_{cell} = +0.80\ V - (-2.37\ V) = +3.17\ V

Explanation:

The half-cell with the more positive reduction potential (Ag+/AgAg^+/Ag) acts as the cathode, while the one with the more negative potential (Mg2+/MgMg^{2+}/Mg) acts as the anode. The cell potential is the difference between them.

Problem 2:

Determine the standard Gibbs free energy change (ΔG⊖\Delta G^\ominus) for the reaction: Zn(s)+Cu2+(aq)→Zn2+(aq)+Cu(s)Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s), given Ecell⊖=+1.10 VE^\ominus_{cell} = +1.10\ V and F=96500 C mol−1F = 96500\ C\ mol^{-1}.

Solution:

ΔG⊖=−(2)(96500)(1.10)=−212300 J mol−1=−212.3 kJ mol−1\Delta G^\ominus = -(2)(96500)(1.10) = -212300\ J\ mol^{-1} = -212.3\ kJ\ mol^{-1}

Explanation:

The number of electrons transferred (nn) in this redox reaction is 22 (from Zn→Zn2+Zn \rightarrow Zn^{2+}). Using the formula ΔG⊖=−nFEcell⊖\Delta G^\ominus = -nFE^\ominus_{cell}, the negative result confirms the reaction is spontaneous.

Problem 3:

An electrolytic cell contains molten NaClNaCl. Calculate the mass of sodium produced if a current of 5.0 A5.0\ A is passed through the cell for 3030 minutes. (Ar(Na)=22.99A_r(Na) = 22.99)

Solution:

Q=5.0×(30×60)=9000 CQ = 5.0 \times (30 \times 60) = 9000\ C. n(e−)=900096500≈0.0933 moln(e^-) = \frac{9000}{96500} \approx 0.0933\ mol. Since Na++e−→NaNa^+ + e^- \rightarrow Na, n(Na)=0.0933 moln(Na) = 0.0933\ mol. mass=0.0933×22.99≈2.14 gmass = 0.0933 \times 22.99 \approx 2.14\ g.

Explanation:

First, calculate the total charge (QQ) in Coulombs. Then, find the moles of electrons using Faraday's constant. Since 11 mole of electrons produces 11 mole of NaNa, the mass is found by multiplying moles by molar mass.

Electrochemical cells - Revision Notes & Key Formulas | IB Grade 12 Chemistry