High School Science - Redox Reactions

Redox Reactions

INPUT • Corrosion occurs readily on boats due to the flow of electrons out of the metals turning the metal into ions. A corrosive free propeller

A propeller that has corroded in the water

INPUT When two metals are submerged into an electrolytic solution and touched together, electricity is produced by flowing electrons out of one of the metals. You can see the propeller has lost some of its metal. No zinc is used here

Electrons leave the zinc anode but do not leave the other metal.

To keep the metal in the propeller from corroding, the reactive metal of zinc is oxidized and turned into ions, instead of the propeller. Zinc was used here

INPUT

Oxidized Zinc This particular piece of zinc is made to go around the propeller shaft to protect the shaft and the propeller.

This is what the corroded zinc looks like. This corroded zinc was located on the shaft of a propeller.

The zinc on this shaft and propeller was allowed to completely oxide. This allowed oxidation to occur to the propeller and shaft it was supposed to protect.

Redox Reactions • Redox reactions occur when there is a gain and loss of electrons from different reactants. • Reduction and oxidation go hand in hand. • If one compound loses 10e-, than another has to gain 10e-. • The following reactions show metals reacting and exchanging electrons with each other and with non-metals:

Balancing Redox Reactions For instance, if you put copper metal into a solution of silver nitrate (AgNO3), you will see the copper begin to dissolve as dark silver solid begins to appear. For redox reactions, start with half of the reaction at a time.

22[ 22 Ag+(aq)

Cu(s)  Cu2+(aq) + 2 e- copper loses electrons Balance the electrons + 22 e  2 Ag(s) ] silver ions pick up the electrons

2Ag+(aq) + Cu(s)  2 Ag(s) + Cu2+(aq)

Add together

Electrons were gained by the Ag and lost by the Cu The Ag was reduced (+1 → 0) and the Cu was oxidized (0 → +2)

Student Practice Even aluminum can “rust” or oxidize. This is a redox reaction, because any time something gets oxidized, something else must be reduced. (The electrons have to go somewhere!) How many electrons Balance the transfer in each half electrons. reaction?

4 [ Al(s)  Al3+(aq) + 3e- ]

3 [ O2(g) + 4e-  2O2-(aq) ] 3 O2(g) +Final 4 Al(s) answer?  4 Al3+ (aq) + 6 O2- (aq) How did you do?

Student Practice Now try these on your own. Cadmium ripping off iron’s electrons:

1)

Feo → Fe2+ + 2e 2e - + Cd2+ → Cdo Cd2+ + Feo → Fe2+ + Cdo

-

Our trouble-making reaction. Iron going to rust:

2)

4 [ Feo → Fe3+ + 3e 3 [ 4e - + O2 → 2O2- ]

-

]

4Feo + 3O2 → 4Fe3+ + 6O2Just These one don’t more stay thing apart. …

2 Fe2O3 (s) (RUST)

Bring it Together

Redox reactions can be used to generate electricity by having the electrons that are pulled from one metal to another pass through an electric device.

We will take a closer look at how the redox reaction is used to generate electricity.

The Battery • Let’s take a practical look at how the oxidation and reduction of metals as they gain and lose electrons can be used to our own advantage as a human race. • If a light or electronic device is placed between the metals as the electrons pass from one metal to the other, then the energy can be used. • Use the chart on the following slide to determine whether an element is more or less reactive than another. Elements near the bottom of the chart with negative numbers are more reactive and readily give up their electrons to the elements above

STANDARD REDUCTION POTENTIALS

Fe3+ + 3e-



Fe(s)

-0.06

IN AQUEOUS SOLUTION AT 25° C

Pb2+ + 2 e-



Pb(s)

-0.13

Sn2+ + 2 e-



Sn(s)

-0.14

E° (Volt)

Ni2+ + 2 e-



Ni(s)

-0.25

Half-Reactions

.

F2 (g) + 2 e-



2 F-

2.87

Co2+ + 2 e-



Co(s)

-0.28

Co3+ + e-



Co2+

1.82

TI + + e-



TI(s)

-0.34

Au3+ + 3 e-



Au (s)

1.50

Cd 2+ + 2 e-



Cd(s)

-0.40

Cl2 (g) + 2 e-



2 Cl-

1.36

Cr 3+ + e-



Cr2+

-0.41

Fe 2+ + 2e-



Fe(s)

-0.44

Cr 3+ + 3 e-



Cr(s)

-0.74

O2 (g) + 4 H + 4 e



2 H2O (l)

1.23

Br2 (l) + 2 e-



2 Br-

1.07

Zn 2+ + 2 e-



Zn(s)

-0.76

2 Hg2+ + 2 e-



Hg22+

0.92

Mn 2+ + 2 e-



Mn(s)

-1.18

Hg2+ + 2 e-



Hg (l)

0.85

Al 3+ + 3 e-



Al(s)

-1.66

Ag+ + e-



Ag(s)

0.80

Be 2+ + 2 e-



Be(s)

-1.70

Hg2 + 2 e



2 Hg(l)

0.79

Mg 2+ + 2 e-



Mg(s)

-2.37

Fe3+ + e-



Fe2+

0.77

Na + + e-



Na(s)

-2.71

I2 (s) + 2 e-



2 I-

0.53

Ca 2+ + 2 e-



Ca(s)

-2.87

Cu+ + e-



Cu(s)

0.52

Sr 2+ + 2 e-



Sr(s)

-2.89

Cu2+ + 2 e-



Cu(s)

0.34

Ba 2+ + 2e-



Ba(s)

-2.90

Cu2+ + e-



Cu+

0.15

Rb + + e-



Rb(s)

-2.92

Sn4+ + 2 e-



Sn2+

0.15

K + + e-



K(s)

-2.92

S(s) + 2 H + 2 e



H2S (g)

0.14

Cs + + e-



Cs(s)

-2.92

2 H+ + 2 e-



H2 (g)

0.00

Li + + e-



Li(s)

-3.05

+

2+

-

+

-

-

ee-

K+ K+ NO3-

Zn

Zn2+ NO3NO3-

eeZn2+

NO3Zn2+ NO3-

NO3Zn2+

K+ Salt Bridge

Zinc metal

Salt Bridge

NO3-

NO3-

SO42- Cu2+ Cu2+ SO42SO42Cu2+ SO42-

Cu2+

2+

Zn2+ NO3NO3-

Zn(NO3)2

Cu Cu

NO3Zn NO3 NO3-

Copper metal

Cu2+ SO42-

Cu2+ SO42-

CuSO4