Standard-State Reduction Half-Cell Potential

Standard-State Reduction Half-Cell Potential

The standard-state cell potentials for some common half-reactions are given in the table below.

Standard-State Reduction Potentials, E^o_red

		Half-Reaction		Eored
		K+ + e- K		-2.924		Best
		Ba2+ + 2 e- Ba		-2.90		reducing
		Ca2+ + 2 e- Ca		-2.76		agents
		Na+ + e- Na		-2.7109
		Mg2+ + 2 e- Mg		-2.375
		H2 + 2 e- 2 H-		-2.23
		Al3+ + 3 e- Al		-1.706
		Mn2+ + 2 e- Mn		-1.04
		Zn2+ + 2 e- Zn		-0.7628
		Cr3+ + 3 e- Cr		-0.74
		S + 2 e- S2-		-0.508
		2 CO2 + 2 H+ + 2 e- H2C2O4		-0.49
		Cr3+ + e- Cr2+		-0.41
		Fe2+ + 2 e- Fe		-0.409
		Co2+ + 2 e- Co		-0.28
		Ni2+ + 2 e- Ni		-0.23
		Sn2+ + 2 e- Sn		-0.1364
		Pb2+ + 2 e- Pb		-0.1263
		Fe3+ + 3 e- Fe		-0.036
		2 H+ + 2 e- H2		0.0000...
		S4O62- + 2 e- 2 S2O32-		0.0895
Oxidizing		Sn4+ + 2 e- Sn2+		0.15
power		Cu2+ + e- Cu+		0.158		Reducing
increases		Cu2+ + 2 e- Cu		0.3402		power
		O2 + 2 H2O + 4 e- 4 OH-		0.401		increases
		Cu+ + e- Cu		0.522
		I3- + 2 e- 3 I-		0.5338
		MnO4- + 2 H2O + 3 e- MnO2 + 4 OH-		0.588
		O2 + 2 H+ + 2 e- H2O2		0.682
		Fe3+ + e- Fe2+		0.770
		Hg22+ + 2 e- Hg		0.7961
		Ag+ + e- Ag		0.7996
		Hg2+ + 2 e- Hg		0.851
		H2O2 + 2 e- 2 OH-		0.88
		HNO3 + 3 H+ + 3 e- NO + 2 H2O		0.96
		Br2(aq) + 2 e- 2 Br-		1.087
		2 IO3- + 12 H+ + 10 e- I2 + 6 H2O		1.19
		CrO42- + 8 H+ + 3 e- Cr3+ + 4 H2O		1.195
		Pt2+ + 2 e- Pt		1.2
		MnO2 + 4 H+ + 2 e- Mn2+ + 2 H2O		1.208
		O2 + 4 H+ + 4 e- 2 H2O		1.229
		Cr2O72- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O		1.33
		Cl2(g) + 2 e- 2 Cl-		1.3583
		PbO2 + 4 H+ + 2 e- Pb2+ + 2 H2O		1.467
		MnO4- + 8 H+ + 5 e- Mn2+ + 4 H2O		1.491
		Au+ + e- Au		1.68
		H2O2 + 2 H+ + 2 e- 2 H2O		1.776
		Co3+ + e- Co2+		1.842
Best		S2O82- + 2 e- 2 SO42-		2.05
oxidizing		O3(g) + 2 H+ + 2 e- O2(g) + H2O		2.07
agents		F2(g) + 2 H+ + 2 e- 2 HF(aq)		3.03

     There is no need to remember that reducing agents become stronger toward the upper right corner of this table, or that the strength of the oxidizing agents increases toward the bottom left corner. All you have to do is remember some of the chemistry of the elements at the top and bottom of this table.
Take a look at the half-reaction at the top of the table.

K+ + e- K

Eored = -2.924 V

     What do we know about potassium metal? Potassium is one of the most reactive metalsit bursts into flame when added to water, for example. Furthermore, we know that metals are reducing agents in all of their chemical reactions. When we find potassium in this table, we can therefore conclude that it is listed among the strongest reducing agents.
Conversely, look at the last reaction in the table.

F2 + 2 e- 2 F-

Eored = 3.03 V

     Fluorine is the most electronegative element in the periodic table. It shouldn't be surprising to find that F₂ is the strongest oxidizing agent in the above table.
Referring to either end of this table can also help you remember the sign convention for cell potentials. The previous section introduced the following rule: Oxidation-reduction reactions that have a positive overall cell potential are spontaneous. This is consistent with the data in the above table. We know that fluorine wants to gain electrons to form fluoride ions, and the half-cell potential for this reaction is positive.

F2 + 2 e- 2 F-

Eored = 3.03 V

     We also know that potassium is an excellent reducing agent. Thus, the potential for the reduction of K⁺ ions to potassium metal is negative

K+ + e- K

Eored = -2.924 V

      but the potential for the oxidation of potassium metal to K⁺ ions is positive.

K K+ + e-

Eoox = -(-2.924 V) = 2.924 V