Chapter 20

A chemical reaction can perform two types of work:

1.Produce a gas to perform PV work2.Use movement of electrons from

redox reactions to perform electrical work

A voltaic (galvanic) cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants.

By physically separating the reduction half of a redox reaction from the oxidation half, we create a flow of electrons through an external circuit.

Used to accomplish electrical work

Two solid metals that are connected by the external circuit are called electrodes. Anode: Cathode:

Electrodes may or may not participate in the reaction Zn/Cu Pt or other conducting material

Each of the two components of a voltaic cell is called a half-cell Oxidation half-cell Reduction half-cell

For a voltaic cell to work, the solutions in the two half-cells must remain electrically neutral Need migration of ions

Salt bridge or porous glass barrior

Voltaic Cells

Anode Acceptor of electrons: Oxidation

Cathode Source of electrons: Reduction

Anions always migrate toward the anode and cations toward the cathode.

Voltaic (Galvanic) Cell

The following redox reaction is spontaneous: Cr2O7

2-(aq)+ 6I-(aq) 2Cr3+(aq) + 3I2(s)

A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution is suspended in each solution, and the two conductors are connected with wires through a voltmeter to detect an electric current. The resultant voltaic cell generates an electric current. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron migration, the direction of ion migration, and the signs of the electrodes.

Describing a Voltaic Cell

The two half-reactions in a voltaic cell are Zn(s) Zn2+(aq)

+ 2e-ClO3

-(aq) + 6H+(aq) + 6e- Cl-(aq) + 3H2O(l)

(a) Indicate which reaction occurs at the anode and which at the cathode. (b) which electrode is consumed in the cell reaction? (c) Which electrode is positive?

Cell EMF Under Standard ConditionsWhy do electrons transfer

spontaneously during redox reactions?

Electrons flow from the anode of a voltaic cell to the cathode because of a difference in potential energy. Potential energy higher at the anode

Electrons flow spontaneously toward the electrode with the more positive electrical potential.

Cell EMF Under Standard Conditions

The difference in potential energy per electrical charge between two electrodes is measured in units of volts.

1V = 1 (J/C) Where V (volts), J (joule), and C

(coulomb)

Cell EMF Under Standard Conditions

The potential difference between two electrodes provides a driving force that pushes electrons through the external circuit. Electromotive Force (emf)

Emf of a cell is denoted as Ecell (the cell potential)

For spontaneous reactions, the cell potential will be positive

Standard EMF

Emf depends on The particular cathode and anode half reactions Concentrations of the reactants and products Temperature

Tabulated values of standard reduction potentials denoted Eo

red to calculate Eocell

Eocell = Eo

red (cathode) – Eored (anode)

Standard Emf

Indirectly measure the standard reduction potential of a half-reaction

Reference point: 2H+ (aq, 1 M) + 2e- H2 (g, 1

atm)

Assigned a standard reduction potential of exactly zero volts

Called a standard hydrogen electrode (SHE)

Standard Emf

When determining standard reduction potentials from other half-reactions, write the reaction as a reduction even though it is “running in reverse” as an oxidation reaction. Whenever an electrical potential is

assigned to a half-reaction, write the reaction as a reduction.

Eored are intensive properties

Calculating Eored from

Eocell

For the Zn-Cu2+ voltaic cell, we have

Zn + Cu2+ Zn2+ + Cu Eocell =

1.10V

Given that Eored of Zn2+ to Zn is -0.76

V, calculate the Eored for the

reduction of Cu2+ to Cu

A voltaic cell is based on the half-reactions:

In+ In3+ + 2e-Br2 + 2e- 2Br-

The standard emf for the cell is 1.46V and Eo

red for the reduction of bromine is +1.06V. Using this information, calculate Eo

red for the reduction of In3+ to In+.

Calculating Eored from

Eocell

Calculating Eocell from

Eored

Using the standard reduction potentials listed in Table 20.1, calculate the standard emf for the following voltaic cells:

1.Cr2O72- + 14H+ + 6I- 2Cr3+ + 3I2 +

7H2O2.2Al + 3I2 2Al3+ + 6I-

Standard EMF

For each of the half-cells in a voltaic cell, the standard reduction potential provides a measure of the driving force for reduction to occur.

The more positive the value of Eored,

the greater the driving force for reduction under standard conditions.

The more positive Eored value

identifies the cathode

Determining Half-Reactions at ElectrodesA voltaic cell is based on the

following two standard half-reactions:

Cd2+ + 2e- Cd Sn2+ + 2e- SnBy using your chart, determine (a)

the half-reaction that occurs at the cathode and the anode, and (b) the standard cell potential

Use Eored values to understand

aqueous reaction chemistryThe more positive the Eo

red value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced and, therefore, to oxidize the other species. Better oxidizing agent

Strengths of Oxidizing and Reducing Agents

The half-reaction with the smallest reduction potential is most easily reversed as an oxidation. The more negative the Eo

red, the stronger the ability to act as the reducing agent

Strengths of Oxidizing and Reducing Agents

Strengths of Oxidizing and Reducing Agents

Determining the Relative Strengths of Oxidizing AgentsUsing standard reduction potentials:Rank the following ions in order of

increasing strength as oxidizing agents: NO3

-, Ag+, Cr2O72-

Rank the following species from the strongest to the weakest reducing agent: I-, Fe, Al

Free Energy and Redox Reactions

Determining the spontaneity of redox reactions

Eo = Eored (reduction process) –

Eored (oxidation

process)

Spontaneous or Not?

Use standard reduction potentials to determine whether the following reactions are spontaneous under standard conditions.

1.Cu + 2H+ Cu2+ + H2

2.Cl2 + 2I- 2Cl- + I2

Spontaneous or Not?

Use standard reduction potentials to understand the activity series of metals

Activity series of metals: strongest reducing agent at the top

Calculate standard emf forNi + 2Ag+ Ni2+ + 2Ag

EMF and ΔG Relationship between G and EMF:

ΔG = -nFEWhere n = number of electrons transferred n

the reaction, G = Gibbs free energy, E = EMF, and F = Faraday’s constant

Faraday’s constant is the quantity of electrical charge on one mole of electrons (a faraday)

1 F = 96,485 C/mol = 96,485 J/V-mol

Determining ΔGº and K

(a) Use standard reduction potentials to calculate ΔGº and K at 298K for the reaction:

4Ag + O2 + 4H+ 4Ag+ + 2H2O

(b) Suppose the reaction in part (a) was written: 2Ag + ½ O2 + 2H+ 2Ag+ + H2O What are values of Eº, ΔGº, and K when the reaction is written this way?

For the reaction 3 Ni2+ + 2Cr(OH)3 + 10OH-

3Ni + 2CrO42-

+ 8H2O

(a) What is the value of n? (b) Given that ΔGº equals +87 kJ/mol, calculate K at a temperature of 298K

Determining n and K

Cell EMF Under Nonstandard Conditions As a voltaic cell is discharged , the

reactants of the reaction are consumed and the products are generated The concentrations of these substances

changes EMF drops until E = 0, and the

concentration of reactants and products are at equilibrium

How does cell emf depend on the concentration of reactants and products?

The Nernst Equation

Dependence of cell emf on concentration

Nernst Equation:

At 298K with units of volts, the equation simplifies to:

The Nernst Equation

The Nernst equation helps us understand why the emf of a voltaic cell drops as the cell discharges

Increasing the concentration of reactants or decreasing the concentration of products increases the driving force (higher emf)

Decreasing the concentration of reactants or increasing the concentration of the products decreases the driving force (lower emf)

Voltaic Cell EMF Under Nonstandard ConditionsCalculate the emf at 298K generated

by:

Cr2O72-(aq)+ 14H+(aq) + 6I-(aq)

2Cr3+(aq) + 3I2(s) + 7H2O(l)

When [Cr2O72-] = 2.0M, [H+] = 1.oM,

[I-] = 1.0M, [Cr3+] = 1.0x10-5M

Calculate the emf at 298K generated by:

2Al (s)+ 3I2(s) 2Al3+ (aq) + 6I- (aq)

When [Al3+ ]= 4.0x10-3M and [I- ]0.010M

Voltaic Cell EMF Under Nonstandard Conditions

Calculating Concentrations in a Voltaic cell If the voltage of the following cell is

+0.45V at 298K when [Zn2+] = 1.0M and PH2= 1atm, what is the concentration of H+?

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)

Batteries and Fuel Cells

A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells.

When cells are connected in series, the battery produces a voltage that is the sum of the emfs of the individual cells. Multiple cells in series Multiple batteries in series

Batteries and Fuel Cells

The substances that are oxidized at the anode and reduced at the cathode determine the emf of the battery. The usable life of the battery depends on

the quantities of these substances. Need a porous barrier between anode

and cathode compartments Primary and Secondary batteries

Batteries and Fuel Cells

Lead-Acid Battery (12-v car battery, 6 voltaic cells in series that each produce 2V)

Alkaline Battery (most common primary battery)

Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries (secondary batteries)

Lead-Acid Batteries

Corrosion

Corrosion reactions are spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound. Oxidation is a thermodynamically

favored process in air at room temperature

CorrosionPrevent corrosion by forming a

protective oxide layer that is impermeable to O2 and H2O

Examples: Al3+ forms protective Al2O3 layer Mg

Corrosion of IronRusting of iron requires both oxygen

and water pH of solution, presence of salts, contact

with metals more difficult to oxidize than iron, and stress on the iron can accelerate rusting

Corrosion of Iron

Cathodic protection: protecting a metal from corrosion by making it the cathode in an electrochemical cell.

The metal that is oxidized while protecting the cathode is called the sacrificial anode.

Corrosion of Iron

Cathodic protection