Chapter 17 Applications of Redox

Reactions

17.1 Explain how a non-spontaneous redox reaction can be driven forward during electrolysis

17.1 Relate the movement of charge through an electrolytic cell to the chemical reactions that occur

17.1 Apply the principle of electrolysis to its applications such as chemical synthesis, refining, plating, and cleaning.

Objectives

17.2 Relate the construction of a galvanic cell to how it functions to produce a voltage and an electrical current

17.2 Trace the movement of electrons in a galvanic cell

17.2 Relate chemistry in a redox reaction to separate reactions occurring at electrodes in a galvanic cell

Objectives

Oxidation reduction reactions involve a transfer of electrons.

OIL- RIG Oxidation Involves Gain Reduction Involves Loss LEO-GER Lose Electrons Oxidation Gain Electrons Reduction

Review

Galvanic Cells (Voltaic) = A battery which uses spontaneous chemical processes to produce electricity◦ The amount of electricity depends on how bad the

atoms (molecules) want the electrons or want to give them up.

The Cell

Corrosion: The oxidation of metals over time from being oxidized by surrounding oxidizing agents (such as oxygen). ◦ Generally very slow, but some are more quickly

oxidized (depending on activity of metal as a solid)

Corrosion

How to stop corrosion: Sacrificial Anodes. Since some metals corrode easier than others, we have the metal we want safe (steel, iron) in contact with a metal that is more easily oxidized (like zinc). ◦ The zinc gets oxidized first, and loses electrons

and takes the hit instead of the iron or steel.

Sacrificial Anodes

Running a reaction backwards ◦ Forcing electrons onto the atoms/molecules

Separating Atoms◦ Give everyone an octet without each other

Used to separate metals from their salts

Electrolysis

Using electrolysis to place aqueous metals onto a surface. ◦ This is how jewelry is plated in gold and silver,

how silver ware is coated in silver, but not completely out of silver ware.

Electroplating

Moving electrons is electric current.

8H++MnO4-+ 5Fe+2 +5e-

® Mn+2 + 5Fe+3 +4H2O Helps to break the reactions into half

reactions.

8H++MnO4-+5e- ® Mn+2 +4H2O

5(Fe+2 ® Fe+3 + e- ) In the same mixture it happens without

doing useful work, but if separate

Applications

H+

MnO4-

Fe+2

Connected this way the reaction starts Stops immediately because charge builds

up.

Reducing Agent

Oxidizing Agent

e-

e-

e- e-

e-

e-

Anode Cathode

Oxidizing agent pushes the electron. Reducing agent pulls the electron. Unit is the volt(V)

Cell Potential

Zn+2

SO4-2

1 M HCl

Anode

0.76

1 M ZnSO4

H+

Cl-

H2 in

Cathode

1 M HCl

H+

Cl-

H2 in

Standard Hydrogen Electrode This is the reference

all other oxidations are compared to

Eº = 0 º indicates standard

states of 25ºC, 1 atm, 1 M solutions.

Zn(s) + Cu+2 (aq) ® Zn+2(aq) + Cu(s) The total cell potential is the sum of the

potential at each electrode.

Eº cell = EºZn® Zn+2 + Eº Cu+2 ® Cu

We can look up reduction potentials in a table.

One of the reactions must be reversed, so change the sign.

Cell Potential

Determine the cell potential for a galvanic cell based on the redox reaction.

Cu(s) + Fe+3(aq) ® Cu+2(aq) + Fe+2(aq)

Fe+3(aq) + e-® Fe+2(aq) Eº = 0.77 V

Cu+2(aq)+2e- ® Cu(s) Eº = 0.34 V

Cu(s) ® Cu+2(aq)+2e- Eº = -0.34 V

2Fe+3(aq) + 2e-® 2Fe+2(aq) Eº = 0.77 V

Cell Potential

solid½Aqueous½½Aqueous½solid Anode on the left½½Cathode on the right Single line different phases. Double line porous disk or salt bridge. For the last reaction Cu(s)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)

Line Notation

Rusting - spontaneous oxidation. Most structural metals have reduction

potentials that are less positive than O2 .

◦ If you are more positive on the chart, you can oxidize anything below you (the product)

Corrosion

Water Rust

Iron Dissolves- Fe ® Fe+2

e-

Salt speeds up process by increasing conductivity

Coating to keep out air and water. Galvanizing - Putting on a zinc coat Has a lower reduction potential, so it is

more. easily oxidized. Alloying with metals that form oxide coats. Cathodic Protection - Attaching large pieces

of an active metal like magnesium that get oxidized instead.

Preventing Corrosion

1.0 M Zn+2

e- e-

Anode Cathode

1.10

Zn Cu1.0 M Cu+2

Cathode (Reduction)Half-Reaction

Standard PotentialE° (volts)

Li+(aq) + e- Li(s) -3.04

K+(aq) + e- K(s) -2.92

Fe2+(aq) + 2e- Fe(s) -0.41

Pb2+(aq) + 2e- Pb(s) -0.13

Cu+(aq) + e- Cu(s) 0.52

I2(s) + 2e- 2I-(aq) 0.54

Ag+ (aq) + 1e- Ag (s)

0.80

Pt+2 (aq) + 2e- Pt (s)

1.23

Cl2(g) + 2e- 2Cl-(aq) 1.36

1. Write the chemical shorthand for a Lead and Lithium battery.

2. What is the cell potential for a Iron and Platinum battery?

3. Which element (and charge) is the best oxidizing agent?

4. Which element (and charge) is the best reducing agent?