Ch.18 - ElectrochemistrySee 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

In voltaic/galvanic cells electrical current is created, while in an electrolytic cell electrical current is consumed to drive the reaction. 

The Ampere

The ampere or amp represents the SI unit for electrical current and represents charge per second. 

Example #1: Gold can be plated out of a solution containing Au3+ based on the following half reaction:

Au3+ (aq) + 3 e -  ---->  Au (s)

a) What mass of gold is plated by a 41 minute flow of 6.8 A current?

Example #2: A solution of Mn+5 is used to plate out Mn in an electrochemical cell. If a total of 1.13 g of Mn is plated out in a total time of 1600 seconds, what was the electrical current used? (MW of Mn is 54.94 g/mol)

Example #3: If steady current of 15 amperes is provided by a stable voltage of 12 Volts for 600 seconds, answer each of the following questions.

a)  Calculate the total charge that passes through the circuit in this time.

b)  Calculate the total number of moles of electrons that pass through the circuit in this time.

c)  Calculate the total amount of energy that passes through the circuit in this time.

d)  Calculate the power that the battery provides during this process.

 

 

Additional Problems
A current of 5.00 A is passed through an aqueous solution of chromium (III) nitrate for 30.0 min. How many grams of chromium metal will be deposited at the cathode? a) 0.027 g b) 1.62 g c) 4.85 g d) 6.33 g
What mass of lead sulfate is formed in a lead-acid storage battery when 1.00 g of Pb undergoes oxidation?
Copper can be electroplated at the cathode of an electrolysis cell by the half-reaction: Cu2+(aq) + 2 e– → Cu(s) How much time would it take for 325 mg of copper to be plated at a current of 5.6 A?
Silver can be electroplated at the cathode of an electrolysis cell by the half-reaction: Ag+ (aq) + e– → Ag (s) What mass of silver would plate onto the cathode if a current of 6.8 A flowed through the cell for 72 min?
A major source of sodium metal is the electrolysis of molten sodium chloride. What magnitude of current produces 1.0 kg of sodium metal in 1 hour?
What mass of aluminum metal can be produced per hour in the electrolysis of a molten aluminum salt by a current of 25 A?
Determine whether HI can dissolve each metal sample. If it can, write a balanced chemical reaction showing how the metal dissolves in HI and determine the minimum volume of 3.5 M HI required to completely dissolve the sample. a. 2.15 g Al b. 4.85 g Cu c. 2.42 g Ag
Determine if HNO3 can dissolve each metal sample. If it can, write a balanced chemical reaction showing how the metal dissolves in HNO3 and determine the minimum volume of 6.0 M HNO3 required to completely dissolve the sample. a. 5.90 g Au b. 2.55 g Cu c. 4.83 g Sn
The molar mass of a metal (M) is 50.9 g/mol; it forms a chloride of unknown composition. Electrolysis of a sample of the molten chloride with a current of 6.42 A for 23.6 minutes produces 1.20 g of M at the cathode. Determine the empirical formula of the chloride.
A 0.0251 L sample of a solution of Cu + requires 0.0322 L of 0.129 M KMnO4 solution to reach the equivalence point. The products of the reaction are Cu2+ and Mn2+. What is the concentration of the Cu2+ solution?
A current of 11.3 A is applied to 1.25 L of a solution of 0.552 M HBr converting some of the H+ to H2(g), which bubbles out of solution. What is the pH of the solution after 73 minutes?
A 215 mL sample of a 0.500 M NaCl solution with an initial pH of 7.00 is subjected to electrolysis. After 15.0 minutes, a 10.0 mL portion (or aliquot) of the solution was removed from the cell and titrated with 0.100 M HCI solution. The endpoint in the titration was reached upon addition of 22.8 mL of HCl. Assuming constant current, what was the current (in A) running through the cell?
Amps are equivalent to which unit? A.  Volts B.  Moles of e− C.  Coulombs D.  Volts / second E.  Coulombs / second  
If the following half-reactions are used in an electrolytic cell:             Li+(aq) + e − → Li(s)                E° = −3.04             Mg2+(aq) + 2 e − → Mg(s)       E° = −2.38 And 10 Amps are applied for 1 minute, determine which solid would form and how much would form in grams. A.  0.043 grams of Lithium B.  0.022 grams of Lithium C.  0.15 grams of Magnesium D.  0.076 grams of Magnesium E.  0.086 grams of Lithium  
Gold can be plated out of a solution containing Au 3+ based on the following half reaction: Au 3+ (aq) + 3 e -      →       Au (s)   What mass of gold is plated by a 53 minute flow of 8.2 A current?
A solution of Mn+5 is used to plate out Mn in an electrochemical cell. If a total of 3.70 g of Mn is plated out in a total time of 1 seconds, what was the electrical current used? (MW of Mn is 54.94 g/mol)  
The following half-reactions combine to make an electrolytic cell: K+ (aq) + e- → K (s)                                       E° = −2.93 Li+ (aq) + e- → Li (s)                                      E° = −3.05 If a current of 1.5 Amps is placed on this cell for 5 minutes, which of the following describes the mass of solid formed (in grams) and the voltage required of the current. [1 Amp = 1 coul/s;    96500 coul = 1 mol e-] A.  0.0324 g;  - 0.12 V             B. 0.000539 g ; - 0.12 V                      C. 0.00466 g;  0.12 V                    D. 0.0324 g;  5.98 V                           E.  0.00466 g;  5.98 V  
How many seconds will be required to deposit 6.23 g of Ag from a solution that contains Ag+ ions if a current of 0.35 A is used? a. 1.7 x 106 b. 1.6 x 104 c. 2.0 x 103 d. 2.1 x 102
A vanadium electrode is oxidized electrically. If the mass of the electrode decreases by 114 mg during the passage of 650 coulombs, what is the oxidation state of the vanadium product?a) +1b) +2c) +3d) +4
Determine the optimum mass ratio of Zn to MnO 2 in an alkaline battery.
A given amount of electric charge deposits 2.159 g of silver from an Ag + solution. What mass of copper from a Cu2+ solution will be deposited by the same amount of electric charge?A. 0.635 gB. 1.97 gC. 2.54 gD. 127 g
If each of these ions were reduced to metal with one coulomb, which would yield the greatest mass?A. Cu2+B. Ag+C. Hg2+D. Cu+
A battery relies on the oxidation of magnesium and the reduction of Cu 2+. The initial concentrations of Mg2+ and Cu2+ are 1.0 x 10–4 M and 1.5 M, respectively, in 1.0-liter half-cells.b. What is the voltage of the battery after delivering 5.0 A for 8.0 h?
A battery relies on the oxidation of magnesium and the reduction of Cu 2+. The initial concentrations of Mg2+ and Cu2+ are 1.0 x 10–4 M and 1.5 M, respectively, in 1.0-liter half-cells.c. How long can the battery deliver 5.0 A before going dead?
A rechargeable battery is constructed based on a concentration cell constructed of two Ag/Ag+ half-cells. The volume of each half-cell is 2.0 L and the concentrations of Ag + in the half-cells are 1.25 M and 1.0 x 10–3 M.a. How long can this battery deliver 2.5 A of current before it goes dead?
A rechargeable battery is constructed based on a concentration cell constructed of two Ag/Ag+ half-cells. The volume of each half-cell is 2.0 L and the concentrations of Ag + in the half-cells are 1.25 M and 1.0 x 10–3 M.c. Upon recharging, how long would it take to redissolve 1.00 x 10 2 g of silver at a charging current of 10.0 amps?
How many grams of Cu metal (63.55 g/mol) will be plated out of a Cu(NO3 )2 solution by passing a current of 3.00 amperes for 11 minutes? (1 Faraday = 96,485 Coulombs)1. 2.608 g2. 0.652 g 3. 1.304 g4. 0.0721 g
For the following battery: Cd(s) | CdCl 2(aq) || Cl -(aq) | Cl2(l) | C(s) a) Write the reduction half reaction occuring at the C(s) electrode. (Include physical states of reactants and products.) b) Calculate the mass of Cl 2 consumed if the battery delivers a constant current of 713 A for 30.0 min.
A current of 0.5 amp flows through an electrochemical cell that evolves Cl 2 gas by the oxidation of Cl1-. What number of moles of Cl 2 gas are generated in a period of 30 min?A) 900 moleB) 9.3 x10-3 moleC) 4.6 x10-3 moleD) 1.6 x 10-4 moleE) 7.8 x 10-5 mole 
How many minutes will be required to deposit 1.00 g of chromium metal from an aqueous CrO42- solution using a current of 6.00 amperes?(A) 186 min(B) 30.9 min(C) 15.4 min(D) 5.15 min
If the following half-reactions are used in an electrolytic cell:Li +(aq) +  e−  →  Li(s)          E  0 = − 3.04Mg 2+(aq) + 2e−  →  Mg(s)   E 0 = − 2.38And 10 Amps are applied for 1 minute, determine which solid would form and how much would form in grams. A. 0.043 grams of LithiumB. 0.022 grams of LithiumC. 0.15 grams of MagnesiumD. 0.076 grams of MagnesiumE. 0.086 grams of Lithium  
Oxygen Supply in Submarines Nuclear submarines can stay under water nearly indefinitely because they can produce their own by the electrolysis of water. How many liters of O2 at 298 K and 1.00 bar are produced in 3.00 hr in an electrolytic cell operating at a current of 0.0200 A?
A current of 4.02 A is passed through a Ni(NO3)2 solution. How long (in hours) would this current have to be applied to plate out 6.20 g of nickel?