17.1 explain how a non-spontaneous redox reaction can be driven forward during electrolysis 17.1...
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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?
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