Electrochemistry

Electrochemistry

Electrochemistry is the study of the interconversion of electrical and chemical energy.

Electrochemical cells may be:

(1) voltaic cells - in which a spontaneous reaction generates electrical energy

(2) electrolytic cells in which electrical energy is used to bring about a nonspontaneous reaction.

All electrochemical reactions are redox reactions.

In an electrochemical cell, these two reactions occur at two different electrodes (metal plates or wires).

Reduction occurs at the cathode and oxidation occurs at the anode.

In the cell, anions move to the anode and cations move to the cathode.

Voltaic (Galvanic) Cells

Any spontaneous reaction can serve as a source of energy in a voltaic cell.

The cell is designed so that oxidation occurs at one electrode (anode) and reduction at the other (cathode).

The electrons produced at the anode move through an external circuit (where they do electrical work) as they are transferred to the cathode.

Example: Zn Cu+2 Cell

spontaneous redox rxn:

Zn(s) + Cu+2(aq) ------( Zn+2(aq) + Cu(s)

To turn this rxn into a voltaic cell the electrons given off by Zn must be redirected through an electrical circuit before they reduce Cu+2 ions to Cu atoms.

The voltaic cell consists of two half-cells (in beakers):

(1) a zinc anode dipping into a solution containing Zn+2 ions

(2) a Cu cathode dipping into a solution of Cu+2 ions

The external circuit consists of a voltmeter with leads to the anode and cathode.

A salt bridge (an inverted glass U-tube filled with a salt solution and plugged with glass wool at each end) connects the two beakers and facilitates the movement of ions.

The salt bridge (which completes the circuit) is filled with a salt solution (KNO3 or Na2SO4) not involved in the electrode reaction. It prevents direct contact between the Zn atoms and the Cu+2 ions.

Zn Zn+2 Cu+2 Cu (shorthand notation)

anode rnx (oxidation) is shown to the left

= salt bridge separating the half-cells

cathode rxn (reduction) is shown to the right

= phase boundary between electrode and (aq)

The zinc anodes oxidation half-rxn produces electrons (Zn(s) --( Zn+2(aq) + 2e-) and pumps them into the external circuit. This is the negative pole of the cell.

At the positive pole or the Cu cathode, the electrons are consumed by the reduction half-rxn (Cu+2(aq) + 2e- --( Cu(s)). This electrode pulls electrons from the external circuit.

Zn+2 ions (surplus) build at the Zn electrode. The Cu electrode become deficient as Cu+2 ions are consumed. Cations move toward the Cu cathode and anions move toward the zinc anode to maintain neutrality (salt bridge).

Other Examples of Salt Bridge Cells

(1) Ni Ni+2 Cu+2 Cu

(2) Zn Zn+2 H+ H2 Pt

Anode: Zn(s) --( Zn+2(aq) + 2e-

Cathode: 2 H+(aq) + 2e- ---( H2(g)

Pt is used as an inert electrode, as no metal is present in the cathode reaction and a conductor is needed.

Example

When chlorine gas is bubbled through a water solution of NaBr, a spontaneous redox reaction occurs:

Cl2(g) + 2 Br-(aq) ---( 2 Cl-(aq)+ Br2(l)

(a) What is the cathode reaction?

(b) Which way do the electrons move in the external circuit?

(c) Which way do anions move within the cell? Cations?

Standard Voltages

Cell voltage is a measure of the driving force behind the spontaneous reaction in a voltaic cell.

Cell voltage is an intensive property. It is independent of the number of electrons passing through the cell. It depends on the nature of the redox rxn and the concentration of the species involved.

The standard voltage (E0) for a given cell is that measured when the current flow is essentially zero, all ions and molecules in solution are at a concentration of 1 M and all gases are at a pressure of a 1 atm.

Standard voltage (overall rxn) = sum of standard voltage for each half rxn

E0 = E0red + E0ox

Standard half-cell voltages are obtained from a list of standard potentials (see text). The potentials are E0red. For E0ox just change the sign of E0red.

Half-reaction potentials are based on a value of zero assigned for the reduction of H+ to H2 gas (S.H.E. standard hydrogen electrode).

This means that the standard voltages for forward and reverse rxns are equal in magnitude but opposite in sign.

E0red = -0.762 V, E0ox = +0.762 V

Strength of Oxidizing and Reducing Agents

The strength of an oxidizing agent is directly related to the standard voltage for its reduction.

The more positive E0red, the stronger the oxidizing agent.

The more positive E0ox, the stronger the reducing agent.

The strongest oxidizing agents are the weakest reducing agents.

Example

Consider the following species in acidic solution: MNO4-, I-, NO3-, H2S, and Fe+2. Using published standard voltages (a) classify each as an oxidizing agent and/or reducing agent (b) arrange the oxidizing agents in order of increasing strength (c) do the same with the reducing agents

Example

Using published standard voltages, calculate E0 for a voltaic cell in which the reaction is

2 Ag+(aq) + Cd(s) ---( 2 Ag(s) + Cd+2(aq)

Spotaneity of Redox Rxns

If the calculated voltage for a redox rxn is a positive quantity, the reaction will be spontaneous. If the voltage is negative, it will not be.

Example

Using standard potentials, decide whether at standard concentrations (a) the reaction

2 Fe+3(aq) + 2 I-(aq) ---( 2 Fe+2(aq) + I2(s) will occur (b) Fe(s) will be oxidized to Fe+2 by treatment with hydrochloric acid (c) a redox rxn will occur when the following species are mixed in acidic solution: Cl-, Fe+2, Cr+2, I2.

Relations Between E0, G0, and K

G0 = -nFE0

G0 = standard free energy change (gases at 1 atm and solutions at 1 M)

E0 = standard voltage

n = # of moles of e- transferred

F = Faraday constant = 9.648 x 104 J/mole V

In a spontaneous rxn E0 > 0, G0 < 0, K > 1.

Redox reactions eventually reach equilibrium (like all rnxs).

Most redox rxns go to completion or not at all.

G0 = -RT ln K

G0 = -nFE0

So, RT ln K = nFE0

E0 = RT ln K

nF

RT/F @ 25 0C = 0.0257 V

E = (0.0257 V) ln K @ 250C

n

Example

For the reaction

3 Ag(s) + NO3-(aq) + 4 H+(aq) --( 3 Ag+(aq) + NO(g) + 2 H20(l)

Use published values to calculate (at 250C) (a) G0 and (b) K.

Effect of Concentration on Voltage

When the concentration of a reactant or product changes, the voltage changes:

(1) Voltage increases if the concentration of a reactant is increased or that of a product is decreased. Makes rxn more spontaneous.

(2) Voltage decreases if the concentration of a reactant is decreased or that of a product is increased. Makes rxn less spontaneous.

A voltaic cell is dead when the concentration of reactants decreases until the redox rxn is at equilibrium, and there is no voltage produced.

Nernst Equation

The Nernst equation relates cell voltage and concentration.

G = G0 + RT ln Q

G = -nFE and G0 = -nFE0

-nFE = -nFE0 + RT ln Q

E = E0 RT ln Q

nF

E = E0 (0.0257 V) ln Q @ 25 0C

n

E = cell voltage, E0 = standard voltage, n = moles of e- exchanged, Q = reaction quotient

Example

Consider a voltaic cell in which the following occurs: O2(g) + 4 H+(aq) + 4 Br-(aq) --( 2 H2O + 2 Br2(l)

Calculate the cell voltage, E, when O2 is at 1.0 atm pressure, [H+] = [Br-] = 0.10 M.

Example

Consider a voltaic cell in which the reaction is

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

It is found that the voltage is +0.560 V when [Zn+2] = 1.0 M, PH2 = 1.0 atm. What must be the concentration of H+ in the H2-H+ half-cell?

Electrolytic Cells

In an electrolytic cell, a nonspontaneous redox rxn is made to occur by pumping electrical energy into the system.

The battery acts as an electron pump pushing the electrons into the cathode and removing them from the anode. Anions and cations must be present in the solution.

A redox reaction consumes electrons at the cathode and liberates them at the anode to maintain electrical neutrality. This process is electrolysis.

Quantitative Relationships

From the balanced half-equations you can determine the relationship between the amount of electricity passed through a cell and the amounts of substances produced at the electrodes.

Example: Cu+2(aq) + 2 e- ---( Cu(s)

2 mole e- ---( 1 mole of Cu (63.55 g of Cu)

Electrical Units

The coulomb, C, (quantity of electrical charge) is related to the charge carried by a mole of electrons through the Faraday constant.

1 mole e- = 9.648 x 104 C

The rate of current flow is measure in amperes, A.

1 A = 1 C/s

Work is required to transfer electrical charge within an electric field. The potential difference is the change in potential energy of a charge when work is done on it.

The SI unit is the volt, V, which is defined as the potential difference between two points if one joule of work is required to move one coulomb of charge from one point to another.

The amount of electrical energy is measured in joules, J. An energy of one joule results when one C moves through a potential difference of one V.

1 J = 1 C V

Power is measured in watts.

1 W = 1 J/s

1 kWh = 3.6 x 106 J = 3.6 x 103 kJ

Example

Chromium metal can be electroplated from a water solution of potassium dichromate; the reduction half-reaction is

Cr2O7-2(aq)+ 14 H+(aq)+ 12 e- --( 2 Cr(s) + 7 H2O(l)

(a) How many grams of chromium will be plated by 1.00 x 104 C?

(b) How long will it take to plate one gram of Cr using a current of 6.00 A?

(c) If the applied voltage is 4.5 V, how many kilowatt hours of electrical energy are required to plate 1.00 g of Cr?

Cell Reactions (Water Solution)

The reduction half-rxn occurs at the cathode.

The reaction may be:

(1) the reduction of a cation to the corresponding metal where a metal object serves as the cathode (electroplating)

Ex: Ag+(aq) + e- --( Ag(s) E0red = +0.799 V

(2) the reduction of a water molecule to hydrogen gas when the cation in solution is difficult to reduce

2H2O + 2 e- ---( H2(g) + 2OH-(aq) E0red = - 0.828 V

At the anode the half-rxn may be:

(1) the oxidation of an anion to the corresponding nonmetal

2I-(aq) ----( I2(s) + 2e- E0ox= - 0.534V

(2) the oxidation of a water molecule to oxygen gas when the anion cannot be oxidized

2H2O ---( O2(g) + 4 H+(aq) + 4 e- E0ox = - 1.299 V

Commercial Cells

Commercial voltaic cells provide electrical energy for many common devices.

Electrolysis of NaCl(aq)

Anode: 2 Cl-(aq) --( Cl2(g) + 2e-

Cathode : H2O(l) + 2e- -( H2(g) + 2OH-(aq)

Cl2 used for plastics, purifying water. H2 used for preparing ammonia.

Primary (Nonrechargeable) Voltaic Cells

Zn-MnO2 dry cell Leclanch cell (1.5 V)

Anode: Zn(s) ---( Zn+2(aq) + 2 e-

Cathode:

2 MnO2(s) + 2 NH4+(aq) + 2e- -( Mn2O3(s) + 2NH3(aq) + H2O(l)

An alkaline dry cell uses a KOH paste (in place of NH4Cl), which prevents the formation of ammonia gas if too large a current is drawn from the cell. It provides more current and lasts longer.

Zn(s) + 2 MnO2(s) -( ZnO(s) + Mn2O3(s)

Mercury Cell used in hearing aids, watches, cameras. This type of cell does not involve ions in solution, so there are no concentration changes as the current is drawn (constant voltage 1.3 V). The anode is Zn-Hg amalgam but the reacting species is Zn.

Anode: Zn(s) + 2OH-(aq) --( Zn(OH)2(s) + 2 e-

Cathode: HgO(s) + H2O + 2e- -( Hg(l) + 2OH-(aq)

Storage (Rechargeable) Voltaic Cells

A storage cell can be recharged repeatedly.

The products of the reaction are deposited on the electrodes. Passing a current through the cell you can reverse the reactions and restore the cell to its original condition.

12 V lead storage battery used in cars, delivers large amounts of energy for short periods, nearly constant voltage for life of cell.

Anode: Pb(s)+ HSO4-(aq) -( PbSO4(s) + H+(aq) + 2e-

Cathode: PbO2(s) + 3H+(aq) + HSO4-(aq) + 2e- -( PbSO4(s) + 2 H2O

As the cell is used, PbSO4 is deposited on the plates of the cell covering Pb and PbO2. Energy from the car engine causes the reversal of the reaction.

Fuel Cells

A fuel cell is a voltaic cell in which a fuel, usually H, is oxidized at the anode. At the cathode, oxygen is reduced.

Anode: 2 H2(g) + 4OH-(aq) --( 4 H2O + 4e-

Cathode: O2(g) + 2H2O + 4e--( 4OH-(aq)

Hydrogen fuel cells are promising for cars (3 kg of H can propel a car 300 mi + only H2O produced). Storage on the vehicle is difficult (volume and pressure) and cost per kJ is high.