Ch.18 - ElectrochemistryWorksheetSee all chapters
All Chapters
Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
BONUS: Lab Techniques and Procedures
BONUS: Mathematical Operations and Functions
Ch.4 - Chemical Quantities & Aqueous Reactions
Ch.5 - Gases
Ch.6 - Thermochemistry
Ch.7 - Quantum Mechanics
Ch.8 - Periodic Properties of the Elements
Ch.9 - Bonding & Molecular Structure
Ch.10 - Molecular Shapes & Valence Bond Theory
Ch.11 - Liquids, Solids & Intermolecular Forces
Ch.12 - Solutions
Ch.13 - Chemical Kinetics
Ch.14 - Chemical Equilibrium
Ch.15 - Acid and Base Equilibrium
Ch.16 - Aqueous Equilibrium
Ch. 17 - Chemical Thermodynamics
Ch.18 - Electrochemistry
Ch.19 - Nuclear Chemistry
Ch.20 - Organic Chemistry
Ch.22 - Chemistry of the Nonmetals
Ch.23 - Transition Metals and Coordination Compounds

Solution: Calculate the standard potential, E°, for this reaction from its equilibrium constant at 298 K.X(s) + Y2+(aq) ⇌ X2+(aq) + Y(s)   K= 9.49 x 105

Solution: Calculate the standard potential, E°, for this reaction from its equilibrium constant at 298 K.X(s) + Y2+(aq) ⇌ X2+(aq) + Y(s)   K= 9.49 x 105

Problem

Calculate the standard potential, E°, for this reaction from its equilibrium constant at 298 K.

X(s) + Y2+(aq) ⇌ X2+(aq) + Y(s)   K= 9.49 x 105

Solution

We are asked to calculate for the standard potential (E°cell) of the reaction. We will use the Nernst Equation to calculate for the ratio. The Nernst Equation relates the concentrations of compounds and cell potential.

E°cell = cell potential, V
R = gas constant = 8.314 J/(mol
·K)
T = temperature, K
n = mole e- transferred
F = Faraday’s constant, 96485 C/mol e- 
K = equilibrium constant


Let’s first determine how many electrons were transferred in the reaction:

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