Module 4: Electrochemistry

Determine the oxidation number (charge) of each element in the following compound: HBr

H = 1+, Br = 1-

Determine the oxidation number (charge) of each element in the following compound: H₂Se

Each H = 1+ Se = 2-

Determine the oxidation number (charge) of each element in the following compound: CuCl₂

Cu = 2+, Each Cl = 1-

Determine the oxidation number (charge) of each element in the following compound: B₂O₃

Each B = 3+, Each O = 2-

Determine the oxidation number (charge) of each element in the following compound: HNO₃

H = 1+, N = 5+, each O = 2-

Define a redox reaction (oxidation-reduction reaction).

A type of reaction involving the transfer of electrons between two species.

Define oxidation.

loss of electrons

Define reduction.

gain of electrons

Analyze the following redox reaction by identifying which reactant will be reduced and which reactant will be oxidized.

Ag (s)   +   Au¹⁺ (aq)   ⇌   Ag¹⁺ (aq)   +   Au (s)

Ag is oxidized (loses electrons) while Au¹⁺ was reduced (gains electrons).

Analyze the following redox reaction by identifying which product was reduced and which product was oxidized.

Al³⁺ (aq)   +   Cr (s)   ⇌   Al (s)   +   Cr³⁺ (aq)

Cr was oxidized (lost electrons) while Al³⁺ was reduced (gained electrons).

Write the half-reactions of the redox reactions below For each, label which substance is oxidized and which substance is reduced in each reaction.

Cd (s) +  Zn²⁺ (aq)  →  Cd²⁺ (aq)  +  Zn (s)

Cd (s)  ^→  Cd²⁺ (aq)  +  2e^-

Zn²⁺ (aq) + 2e^-  →  Zn (s)

Oxidized = Cd (cadmium); Reduced = Zn (zinc)

Write the half-reactions of the redox reactions below For each, label which substance is oxidized and which substance is reduced in each reaction.

Bi (s) +  Al³⁺ (aq)  →  Bi³⁺ (aq)  +  Al (s)

Bi (s) → Bi³⁺ (aq)  +  3e^-

Al³⁺ (aq) + 3e^-  →  Al (s)

Oxidized = Bi (bismuth); Reduced = Al (aluminum)

List the 7 steps used to balance a redox reaction using the half-reaction method. Then, write the balanced equation of the redox reaction below using this method. Show all steps.

Au(s)  +  CrO₄^2- (aq)   Au⁺ (aq) +  CrO₂ (aq)

1. Write skeletal equations for the oxidation and reduction half-reactions.

2. Balance each half-reaction for all elements except H and O.

3. Balance each half-reaction for O by adding H₂O(l).

4. Balance each half-reaction for H by adding H⁺(aq).

5. Balance each half-reaction for charge by adding electrons.

6. If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.

7. Add the two half-reactions and simplify.

8. If the reaction takes place in a basic medium, add OH⁻ (aq) ions the equation obtained in step 7 to neutralize the H⁺ (aq) ions (add in equal numbers to both sides of the equation) and simplify.

 

1. Ox: Au(s) -> Au⁺(aq) 

  Redux: CrO₄^2-(aq) -> CrO₂

3. Ox: Au(s) -> Au⁺(aq) 

  Redux: CrO₄^2-(aq) -> CrO₂+ 2H₂O(l)

4. Ox: Au(s) -> Au⁺(aq) 

  Redux: 2H⁺+ CrO42-(aq) -> CrO₂+ 2H₂O(l)

5. Ox: Au(s) -> Au⁺(aq) +e^- 

  Redux: 2e^- + 2H⁺+ CrO42-(aq) -> CrO₂ + 2H₂O(l)

6. Ox: (Au(s) -> Au⁺(aq) +e^-) x2

  Redux: 2e^- + 2H⁺+ CrO42-(aq) -> CrO₂ + 2H₂O(l)

7. 2Au(s) + 2H⁺+ CrO₄^2- -> 2Au⁺(aq) +  CrO₂ + 2H₂O(l)

Electrochemical cell

a device that contains the reactants and products of a redox system but prevents physical contact between the reactants. The transfer of electrons is indirect and requires an external circuit.

Galvanic cell

(also called a voltaic cell) is an electrochemical in which a spontaneous redox reaction takes place.

Half-cell

contains the redox conjugate pair of a single reactant (two per galvanic cell).

Anode

the cell that contains the oxidation reaction

Cathode

the cell that contains the reduction reaction

Salt bridge

A tube filled with inert electrolyte solution that connects the anode and the cathode

Inert electrode

provides or accepts electrons in the redox reaction

Active electrode

constructed from a member of the redox reaction

Write the cell schematic (cell notation) of each galvanic cell whose net reaction is provided below, omitting concentrations and charges. Note which half-reaction is the oxidation reaction and which is the reduction reaction.

Ca (s) +  Co²⁺ (aq)  →  Ca²⁺ (aq)  +  Co (s)

A) oxidation half-reaction: Ca (s) →  Ca²⁺ (aq)  +  2e^-
B) reduction half-reaction: Co²⁺ (s) + 2e^-  →  Co (s)
C) schematic: Ca (s) I Ca²⁺ (aq) II  Co²⁺ (aq)  I  Co (s)

Write the cell schematic (cell notation) of each galvanic cell whose net reaction is provided below, omitting concentrations and charges. Note which half-reaction is the oxidation reaction and which is the reduction reaction.

1. Al (s) +  Fe³⁺ (aq)  →  Al³⁺ (aq)  +  Fe (s)

A) oxidation half-reaction: Al (s) →  Al³⁺ (aq)  +  3e^-
B) reduction half-reaction: Fe³⁺ (aq) + 3e^-  →  Fe (s)
C) schematic: Al (s)  I  Al³⁺ (aq) II  Fe³⁺ (aq)  I  Fe (s) 

Cell potential

The difference in potential between two half-cells.

Standard cell potential E^∘_cell=E^∘_cathode​−E^∘_anode​

The difference in potential between two half-cells when measured under standard-state conditions.

Standard Hydrogen Electrode

A simplified reference for sharing cell potential measurements with an assigned potential of 0 volts. Contains a stream of hydrogen gas and occurs at standard-state conditions.

Electrode potential

The potential measured for a half-cell comprised of X acting as a cathode and the SHE acting as the anode.

Standard electrode potential- When the half-cell X is under standard state conditions

The potential measured for a half-cell comprised of X acting as a cathode and the SHE acting as the anode when the half-cell X is under standard-state conditions.

Calculate the expected standard cell potential using the following information:

Cell Reaction: K (s)  +  Li⁺ (aq)    K⁺ (aq)  +  Li (s)

Anode Half-reaction: K (s)    K⁺ (aq)  +  e^-                          E˚ = -2.931 V

Cathode Half-reaction: Li⁺ (aq)  +  e^-    Li (s)                        E˚ = -3.04 V

E˚_cell = E˚_cathode - E˚_anode

E˚_cell = (-3.04 V) – (-2.931 V)

E˚_cell = - 0.109 V

Interpret the values of each standard cell potential or standard electrode potentials. Use this information to determine if the reaction spontaneous or nonspontaneous?

1. E˚_cell = +0.34 V

2. E˚_cell = -0.29 V

3. E˚_cathode = 0.99 V, E˚_anode = +1.9 V

1. Spontaneous (positive E˚_cell)

2. Nonspontaneous (negative E˚_cell)

3. Nonspontaneous (E˚_cathode < E˚_anode)

Is the Reaction spontaneous or nonspontaneous? Are products or reactants more abundant at equilibrium?

K	ΔG˚	E˚cell	Summary
> 1	< 0	> 0

Reaction is spontaneous under standard conditions.

Products are more abundant at equilibrium.

Is the Reaction spontaneous or nonspontaneous? Are products or reactants more abundant at equilibrium?

K	ΔG˚	E˚cell	Summary
= 1	= 0	= 0

Reaction is at equilibrium under standard conditions.

Reactants and products are equally abundant.

Is the Reaction spontaneous or nonspontaneous? Are products or reactants more abundant at equilibrium?

K	ΔG˚	E˚cell	Summary
< 1	> 0	< 0

Reaction is nonspontaneous under standard conditions.

Reactants are more abundant at equilibrium.

For the cell schematic below, identify the values for n and Q, and calculate the E˚_cell. Assume that the reaction is occurring under standard temperature (298 K).

Ni (s)  I  Ni²⁺ (aq, 0.30 M)  II  Co²⁺ (aq, 0.20 M)  I  Co (s)

 

Ni (s)  +  Co²⁺ (aq)  →  Ni²⁺ (aq)  +  Co (s)

 

Ni (s)  →  Ni²⁺ (aq)  +  2e^-                   E˚_cell = -0.257 V

Co²⁺ (aq)  +  2e^-  →  Co (s)                E˚_cell = -0.28V

n = total number of electrons transferred

n = 2

Q = [products]/[reactants]

Q = [0.30/0.20] = 1.5

E˚_cell = E˚_cathode - E˚_anode

E˚_cell = (-0.28 V) – (-0.257 V) = -0.023 V

Define the term “battery” in your own words and name one of the first successful batteries.

device consisting of one or more galvanic electrochemical cells that stores and converts chemical energy to electrical energy with the purpose of powering electrical devices. One of the first successful batteries was the Daniell cell.

Dry cell

non rechargeable. Chemical components: Zn can + electrolyte paste + water + graphite rod.

Alkaline battery

non rechargeable. Chemical components: alkaline electrolytes like KOH + steel can.

Nickel-cadmium battery

rechargeable. Chemical components: nickel, cadmium, KOH. Cd is toxic, requires proper disposal.

Lithium-ion battery

rechargeable. Chemical components: lithium, oxygen, metal. Commonly used for lightweight, portable/personal electronic devices such as cellphones and laptops.

Lead-acid battery

rechargeable. Chemical components: Pb and H₂SO₄ acid solution. Hazardous; requires proper disposal. Commonly used in cars.

Fuel cell

does not need to be recharged/consistently supplied with fuel. Chemical components: oxygen, hydrogen, and water. More efficient and environmentally clean energy supply. Commonly used for spacecraft. Possibly future alternative to combustion engines.

Describe each of the following methods of corrosion remediation; Paint

applied to the surface of a metal. The paint layer prevents surface oxidation of the metal.

Describe each of the following methods of corrosion remediation; alloy

Multiple metals are mixed. Metal mixture strengthens the metal and produces a passivation layer.

Describe each of the following methods of corrosion remediation; galvanization

A metal is coated with another metal. Metal layer prevents surface oxidation of the metal below it.

Describe each of the following methods of corrosion remediation; cathode protection

A galvanic cell is created and connected to a more active metal (sacrificial anode). The cathode method conducts inert electrons, preserving the desired metal.

Describe each of the following methods of corrosion remediation; sacrificial anode

A more active metal is connected and oxidized. The more active metal (anode) is oxidized/corroded over time, preserving the desired metal.

Define electrolysis

the use of a electric current to stimulate a non-spontaneous reaction in a redox system.

Define electroplating

a process using electrolysis to produce a thin coating of metal on a conducting surface.

Give an example of an application of electroplating and an example of a metal commonly used.

Corrosion resistance, strengthening, aesthetic purposes. Cd, Cr, Ni, Au, Cu, Sn, and Ag

How are galvanic cells and electrolytic cells similar and different?

Similarities: Contain cathode/anode, transfer of electrons occurs through a redox reaction

Differences: Electrolytic cells require an external power supply in the form of an electric current because the redox reaction in electrolytic cells is nonspontaneous while the redox reaction in galvanic cells in spontaneous.