Lecture 11Lecture 11

Oxidation & Reduction I

Suggested reading: Chapter 4.1-4.5

http://www.youtube.com/watch?v=vJbllSaj_L4

2Ag+ + Cu Cu2+ + 2Ag

2AgNO3 + Cu(0) Cu(NO3)2 + 2Ag(0)

9,10 Bis(phenylethynyl)anthracene

dioxetanedione (a d )diphenyl oxalate peroxyacid ester )phenol

Sodium Chlorate Gummy bears! (Sucrose)

Photosynthesis

http://chemwiki.ucdavis.edu/Biological_Chemistry/Photosynthesis/Photosynthesis overview/Photosynthesis_overview

Photosynthesis

http://chemwiki.ucdavis.edu/Biological_Chemistry/Photosynthesis/Photosynthesis overview/Photosynthesis_overview

http://room114.wikispaces.com/Cellular+Respiration

Photosynthesis Electroplating

http://chemwiki.ucdavis.edu/Biological_Chemistry/Photosynthesis/Photosynthesis overview

http://www.cartoonbarry.com/2007/10/gold_plated_apple_macbook_pro.html

/Photosynthesis_overview

http://room114.wikispaces.com/Cellular+Respiration

Photosynthesis Electroplating

http://chemwiki.ucdavis.edu/Biological_Chemistry/Photosynthesis/Photosynthesis overview

http://www.cartoonbarry.com/2007/10/gold_plated_apple_macbook_pro.html

/Photosynthesis_overview

http://room114.wikispaces.com/Cellular+Respiration

Geology

Photosynthesis Electroplating

http://chemwiki.ucdavis.edu/Biological_Chemistry/Photosynthesis/Photosynthesis overview

http://www.cartoonbarry.com/2007/10/gold_plated_apple_macbook_pro.html

/Photosynthesis_overview

http://room114.wikispaces.com/Cellular+Respiration

Geology Medicine

Redox Reactions

Electron gain = reductionElectron gain reductionElectron loss = oxidation

El t li d i tElectron supplier = reducing agentElectron remover = oxidizing agent

“LEO the lion says GER”

Redox & Oxidation Numbers

• Often, it is difficult to keep track of where electrons have come from and gone to, since transfer of electrons is usually accompanied by transfer of atoms

• Analyze redox reactions according to changes in a y e ed ea t a d g t a ge oxidation number

• Oxidation increase in oxidation number• Oxidation = increase in oxidation number

• Oxidation number = the charge an atom would have if the more electronegative atom in a bond completely acquired the electrons of the bond completely

Assigning Oxidation Numbers (ξ)

Metals are more electropositive F always wants an p

than HF always wants an

extra electron

Assigning Oxidation Numbers

ξ(S) in Hydrogen sulfide (H2S)

• Overall charge is zero so Overall charge is zero, so 2ξ(H) + ξ(S) = 0

• ξ(H) = +1 with nonmetalsξ(S) 2• ξ(S) = -2

ξ(Mn) in permanganate ion (MnO4-)

• ξ(Mn) + 4ξ(O) = -1ξ(Mn) 4ξ(O) 1• ξ(O) = -2

• ξ(Mn) = +7

Redox Half Reactions

Two conceptual reactions in which the electron loss and gain are displayed Two conceptual reactions in which the electron loss and gain are displayed explicitly (not necessarily the real behavior of electron transfer)

B l i d tiBalancing redox reactions:1. write the unbalanced half reactions for the two species as

reductions2. Balance the elements other than H3. Balance O atoms by adding H2O to the other side of the arrow4. If the solution is acidic balance the H atoms by adding H+. If 4. If the solution is acidic, balance the H atoms by adding H . If

basic, balance the H atoms by adding OH- to one side and H2O to the other.

5 B l th h b ddi -5. Balance the charge by adding e-

6. Multiply each half-reaction by a factor to ensure that the numbers of e- match.

7. Subtract one half-reaction from the other and cancel redundant terms.

Redox Half Reactions

Oxidation of Fe2+ by permanganate (MnO -) Oxidation of Fe2 by permanganate (MnO4 )

Fe3+ (aq) + e- Fe2+ (aq)

Oxidation is a reverse reduction

( ) 2 ( )MnO4-(aq)Mn2+(aq)

MnO4-(aq)Mn2+(aq) + 4H2O Balance O with H2O

MnO4-(aq)+8H+Mn2+(aq) + 4H2O Balance H with H+

MnO -(aq)+8H++5e-Mn2+(aq) + 4H O Balance charge with e-MnO4 (aq)+8H +5eMn (aq) + 4H2O Balance charge with e

Redox Half Reactions

Oxidation of Fe2+ by permanganate (MnO -) Oxidation of Fe2 by permanganate (MnO4 )

( ) 2 ( )

Fe3+ (aq) + e- Fe2+ (aq)

MnO4-(aq)+8H++5e-Mn2+(aq) + 4H2O

Ensure the number of electrons match:

5Fe3+ (aq) + 5e- 5Fe2+ (aq)

Ensure the number of electrons match:

MnO4-(aq)+8H++5e-Mn2+(aq) + 4H2O

Subtract one half reaction from the other:Subtract one half reaction from the other:

MnO4-(aq)+8H+- 5Fe3+Mn2+(aq) + 4H2O - 5Fe2+ (aq)

MnO4-(aq)+8H+ + 5Fe2+Mn2+(aq) + 4H2O + 5Fe3+ (aq)

Standard Potentials

000 STHG

Standard state: equilibrium state at atmospheric pressure Standard state: equilibrium state at atmospheric pressure and the temperature in question

)(ln0 TKRTGr

R=8 3145 J/mol*K (gas constant)R=8.3145 J/mol K (gas constant)K=equilibrium constant

Thermodynamic criterion of spontaneity at constant temperature and pressure is ΔrGo<0

K>1

Standard Potentials

• The overall chemical equation in redox reactions is the difference of two reduction half reactionsdifference of two reduction half reactions

• The standard Gibbs energy of the overall reaction is the d ff f h f h h fdifference of the Gibbs energies of the half reactions

• Since only the difference matters, we can choose one half-y ,reaction to have ΔrGo=0 and report all other values relative to it.

H+(aq) + e-½ H2 (g) ΔrGo=0

Measuring Gibbs Energies

Galvanic cell: electrochemical cell where a chemical reaction Galvanic cell: electrochemical cell where a chemical reaction generates an electric current.

The potential difference between electrodes in each half-cell is pmeasured and converted to ΔrGo via:

Δ Go= νFEΔrGo=-νFE

• ν = stoichiometric coefficient of electrons transferred when h h lf b dthe half-reactions are combined

• F = faraday’s constant = 96.48 kC/mol

Spontaneous Reactions can be put to Work

http://www.youtube.com/watch?v=0oSqPDD2rMA

Reduction of Zn2+

Zn2+(aq) + H2(g) Zn(s)+2H+ (aq) Δ Go=147 kJ/mol Zn (aq) H2(g) Zn(s) 2H (aq) ΔrG =147 kJ/mol

electrons

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

Z l d i Zn electrode is dissolved, H2 gas is evolved

Anode (oxidation)

Cathode (reduction)

http://www.youtube.com/watch?v=ckyqMuR7dZE&feature=related

Standard cell potential: Reduction of Zn2+

H+(aq) + e-½ H2 (g) Eo(H+,H2)=0V

Zn2+(aq) + 2e- Zn(s)

( q) 2 (g) ( , 2)

Eo(Zn2+,Zn)=-0.76V

The standard cell potential E is the difference between the two standard potentials of the half reactions: p

2H+(aq)+Zn(s) Zn2+(aq) + H2(g) E ll=+0.76V2H (aq) Zn(s) Zn (aq) H2(g) Ecell 0.76V

Zinc tends to dissolve in acids

Interpreting standard potentials

M ( ) ½ H ( ) M( ) H ( ) M+(aq) + ½ H2(g) M(s) + H+(aq)

Interpreting standard potentials

M ( ) ½ H ( ) M( ) H ( ) M+(aq) + ½ H2(g) M(s) + H+(aq)

Standard potentials at 298K

Electrochemical seriesOx/Red couple with strongly positive Eo (Ox is strongly oxidizing)Ox/Red couple with strongly negative Eo (Red is strongly reducing)

The Nernst Equation

T j d h d f i i i l To judge the tendency of a reaction to run in a particular direction (the cell need not be in equilibrium)

QRTGG rr ln0

)/( FGE rcell )/(00 FGE r

QF

RTEEcell ln0

A bB C dDdc[D][C])]products([ m

iQaA+bBcC+dD ba[B][A][ ][ ]

)]reactants([)]p([

n

j

iQ

The Nernst Equation

Example: What is the dependence of the H+/H2 couple on pH?

A 1 b d T 25oCAssume: pressure = 1 bar and T=25oC.