8–1

Ibrahim Barry Chapter 20-1

Chapter 20

Electrochemistry

8–2

Chapter 20-2

Overview

• Half-Reactions – Balancing oxidation – reduction in acidic and basic solutions

• Voltaic cells– Construction of voltaic cells– Notation for voltaic cells– Electromotive force (EMF)– Standard cell potentials– Equilibrium constants from EMFs– Concentration dependence of EMF

• Electrolytic cells– Aqueous electrolysis– Stoichiometry of electrolysis

8–3

Chapter 20-3

Electrochemistry

• Electrochemistry - Field of Chemistry dealing with transfer of electrons from one species to another.

E.g. Zn in CuSO4(aq).

• Electrochemical cell - combination of two half reactions to produce electricity from reaction.

E.g. Danielle cell: Zn and Cu electrodes in salts of these ions.

8–4

Chapter 20-4

Balancing Redox Reactions: Oxidation Number Method

• Determine oxidation # for each atom- both sides of equation.• Determine change in oxidation state for each atom.• Left side: make loss of electrons = gain. • Balance on other side.• Insert coefficients for atoms that don't change oxidation state.

E.g. Balance

FeS(s)+CaC2(s)+CaO(s)Fe(s)+CO(g)+CaS(s)

• In acidic or basic solution balance as above, then balance charge with H+ or OH on one side and water on other side.

• Cancel waters that appear on both sides at end.• E.g. Balance which occurs in acidic solution:

242 CrOClCrOClO

8–5

Chapter 20-5

Balancing: Half-Reaction Method

• Write unbalanced half reactions for the oxidation and the reduction

• Balance the number of elements except O and H for each. • Balance O's with H2O to the deficient side. • Balance H's with H+ to the hydrogen deficient side

– Acidic: add H+

– Basic: add H2O to the deficient side and OH to the other side.

• Balance charge by adding e to the side that needs it. • Multiply each half-reaction by integers to make electrons

cancel.• Add the two half-reactions and simplify.

E.g. Balance: – Acidic: Zn(s) + VO2+(aq) Zn2+(aq) + V3+(aq).– Basic: Ag(s) + HS(aq) + (aq) Ag2S(s) + Cr(OH)3(s).

24CrO

8–6

Chapter 20-6

Galvanic (Voltaic) and Electrolytic Cells

• Cell reaction – Redox reaction involved in electrochemical cell.

• Voltaic (galvanic) cell – reaction is spontaneous and generates electrical current.

• Electrolytic – non-spontaneous reaction occurs due to passage of current from external power source. E.g. charging of batteries.

8–7

Chapter 20-7

Galvanic Cell 2

• anode – electrode where oxidation occurs. • cathode – electrode where reduction occurs.• salt bridge – ionic solution connecting two half-

cells (half-reactions) to prevent solutions from mixing.E.g. Which is the anode and cathode in the cell to the right? Write the half–reactions.

Cd(s) + 2Ag+(aq) 2Ag(s) + Cd2+(aq)• Sign of electrodes (current flows from anode to

cathode):

E.g. determine direction of electron flow for the reaction for a galvanic cell made from Ni(s) and Fe(s). The reaction is:

2Fe3+(aq) + 3Ni 2Fe(s) + 3Ni2+(aq)

Electrode Sign Description 1. Cathode: + cations migrate to it.

cations reduced 2. Anode: anions migrate to it.

cations made.

8–8

Chapter 20-8

Shorthand Notation for Galvanic Cells

• Shorthand way of portraying electrodes in a voltaic electrochemical cell.

• Redox couple-oxidized and reduced forms of same element when it is involved in electrochemical reaction. Shorthand: Ox/Red E.g. Cu2+/Cu, Zn2+/Zn.– 2 couples required for electrochemical reaction.

• Shorthand rules: – Anode reaction-left; reduced form first.– Cathode-right; oxidized form first.– Vertical line drawn between different phases including reaction of

gases at metal electrode.– Double vertical drawn where salt bridge separates two half-

reactions.E.g. draw cell diagram for

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s).Fe3+(aq) + H2(g) Fe2+(aq) + 2H+(aq).

8–9

Chapter 20-9

Electrical Work

• Earlier w = PV • In electrochemistry electrical pressure = potential

difference; w = Eq or charge times the electrical pressure.

• Units: CoulombVolts = Joules (SI Units 1J = 1CV) ;

Also want to relate to # moles.– 1 mole e = 1 Faraday = 1 F

• F = qeN = 1.602x1019C6.022x1023/mol = 9.65x104C/mol e.

8–10

Chapter 20-10

Cell Potentials for Cell Reactions: Spontaneity of Redox Reactions

G vs E:

G E G = nFE. Use this to calculate G for electrochemical reaction when cell voltage known.

E.g. Determine G for Zn/Cu cell if E = 1.100V

G E At Equilibrium 0 0 Spontaneous + Not spontaneous +

8–11

Chapter 20-11

Standard Reduction Potentials

• As with thermodynamic quantities, we list cell potentials at standard state = 1M at 1 atm and usually 25C.

• Cell potential is the sum of half-cell potentials using Hess’ law.

• Half-cell reactions for Daniell cell were

• Potential at each electrode initially determined relative to SHE = Standard hydrogen electrode.

2H+(aq)+2e H2(g); [H+] = 1M and PH2=1atm.• Other reaction run to determine if the SHE reaction proceeds spontaneously in

direction written when connected to other half-cells.E.g. the cell potential of copper at standard state conditions relative to SHE(acting as the anode) was 0.340 V; determine the half–cell potential for Zn Zn2+ if the potential for the Daniell cell (standard state conditions) was 1.100 V.

• Half-cell reactions reported as reductions.

Zn(s) Zn2+(aq) + 2e Anode Cu2+(aq) + 2e Cu(s) Cathode

Zn(s) +Cu2+(aq) Zn2+(aq) + Cu(s)

oanode

ocathode

ocell EEE

8–12

Chapter 20-12

Using Standard Reduction Potentials

• Large negative value means oxidation strongly favored; strong reducing agent.

• Large positive value means reduction strongly favored; strong oxidizing agent.

• Relative values in table give an indication that one half-reaction favored over other. Summing half-cell reactions allow determination of standard cell potential.

• Half-cell potential intensive property independent of amount of material we don’t use stoichiometric coefficients for determining standard cell potentials.E.g. determine the cell potential of

Br2(l) + 2I(aq) I2(l) + 2Br(aq)E.g.2 determine the cell potential of

2Ag+(aq) + Cu(s) 2Ag(s) + Cu2+(aq)E.g.3 Determine cell potential:

(aq) + Fe(s) Fe2+(aq) + Mn2+(aq) (balanced?). when it is operated galvanically. Which is the oxidizing agent? reducing agent?

• E.g. 4 Determine if the reaction below is spontaneous in the direction written.Fe3+(aq) + Ag(s) ?

4MnO

8–13

Chapter 20-13

Spontaneity of Redox Reactions

G vs E:

G E G = nFE. Use this to calculate G for electrochemical reaction when cell voltage known.

• E.g. Determine G for Zn/Cu cell if E = 1.100V

G E At Equilibrium 0 0 Spontaneous + Not spontaneous +

8–14

Chapter 20-14

• Effect of Concentration on Cell EMF: The Nernst Equation

• Recall that G = Go + RTlnQ where

or at 25°C

which is called Nernst Equation.• E.g. Determine potential of Daniell cell at 25C if

[Zn2+] = 0.100 M and [Cu2+] = 0.00100 M.

ba

nm

]B[]A[

]N[]M[Q

QlnnFRT

EE o

Qlogn

0592.0EE o

8–15

Chapter 20-15

Electrochemical Determination of pH

• Electrodes can be used to determine acidity of solution by using hydrogen electrode with another one e.g. Hg2Cl2 half-cell.

E.g. determine the pH of a solution that develops a cell potential of 0.280 V (at 25°C) given the cell below

Pt(s) | H2(g) (1atm) | H+(? M) || Pb2+(1 M) | Pb(s)E.g. 2 determine the pH of a solution that develops a cell potential of 0.200 V (at 25°C) given the cell below (called a concentration cell).

Pt(s)|H2(g)(1atm)|H+(1.00M)||H+(?M)|H2(g)(1atm)|Pt(s)

pH.

05920E'

]log[H0.0592' E E

8–16

Chapter 20-16

Standard Cell Potentials and Equilibrium Constants

• Free energy and equilibrium constant for reaction studied can be determined from cell voltage.

• Cell potential can be determined from G or from equilibrium constant. – Recall: G° = nFE° = RTlnK.

E.g. Determine free energy and equilibrium constant for reaction below (unbalanced).

(aq) + Fe(s) Fe2+(aq) + Mn2+(aq)

E.g.2 Determine cell potential and equilibrium constant of Cl2/Br2 cell.

E.g.3 The following cell has a potential of 0.578 V at 25°C; determine Ksp.

Ag(s)|AgCl(s)|Cl(1.0 M)||Ag+(1.0 M)|Ag(s).

4MnO

8–17

Chapter 20-17

Quantitative Aspects of Electrolysis

• Current, i, measured units: 1 ampere = 1 coulomb per s (1 A = 1 Cs1).• Time of electrolysis, t (s), also measured. • Total charge, Q, calculated from the product:• Q = it

Charge on 1 mol of e:

• mol of e in an electrolysis obtained from balanced cell reaction:E.g. determine # mol of electron involved in the electrolysis of the following:

– Ag+(aq) + e Ag(s)– 2Cl(aq) + 2e Cl2(g)

• Amount deposited given by:

E.g. Determine amount of Cu2+ electrolyzed from solution at constant current of 6.00 A for period of 1.00 hour.

e of mol

C96500

e

C10x602.1

e of mol

e10x02.6

eN F

1923

nF

ti =

nF

Q = moles

8–18

Chapter 20-18

Electrical Work

• Maximum electrical work: G = wmax = nFE

– n = mol of electrons

– F = Faraday’s constant

– E = cell potential

– Units – Joules

• Electrical Power: 1 watt = 1 J/s. • Energy often expressed as kilowatt –hr• 1 kW*hr = 1000 W*3600 s = 3.6x106 J• E.g. determine the maximum work in kW*hr required to produce

1.00 kg of Zn from Zn2+ in a Daniell cell where the cell potential is 1.100 V for the production of Zn metal.

Strategy:– Determine n: mol of Zn2+ times 2.

– Calculate work.