Standard reduction potentials are very useful in chemistry. They are also known as standard cell potentials, or standard electrode potentials. They are measured in volts, and they tell you how likely an element or ion is to be reduced by gaining electrons. We explain the concepts clearly, and give you a list of standard reduction potentials.
For example, fluorine, at the top of the table shown at the bottom of this article, has an extremely high affinity for electrons and will literally rip them off almost any other molecule it comes in contact with. The fluorine molecule will gain 2 electrons, one for each atom, with the result being two fluoride ions.
This table is often consulted when trying to determine a powerful enough reducing or oxidizing agent for a redox reaction, or determine which metals will displace others. They also represent the voltage of an electrochemical half-reaction.
All of the reactions listed in the table below are half-reactions, which means they need another half-reaction that transfers electrons in the opposite direction, to make a complete reaction.
What is a standard reduction potential?
Let’s define what a standard reduction potential is. Spontaneous redox (oxidation-reduction reactions) can produce electrical energy. A voltaic (galvanic) cell is an electrochemical cell that can capture this energy, using two half-cells, an anode, a cathode, and a salt bridge – similar to the photo at the top of this page.
The SRP is measured for a half-reaction by using hydrogen half-cell, known as a SHE (standard hydrogen electrode). The hydrogen half-cell is an arbitrary electrode that serves as a reference, and the cell potential for the hydrogen half-cell is set at zero.
What standard potentials tell you, and don’t tell you
When you calculate the standard potential for a reaction, a positive result tells you that the reaction should be spontaneous. It does not mean it will be, and it doesn’t tell you how fast the reaction will occur.
For example, based on calculations, aluminum should displace hydrogen from water. However, under normal conditions, it does not due to a passive aluminum oxide micro-coating that forms. But when you mix aluminum with gallium, you allow water to bypass the oxide coating, and you can see the reaction occur.
You can also use the standard potential of a reaction, to calculate the equilibrium constant for the reaction. You can calculate the Gibbs free energy change of the reaction. And you can use the Nernst equation to calculate the potential, free energy change, and equilibrium constant under non-standard conditions, such as the reactants and the products being a different concentration.
Reduction Potential Tips
Here are some useful tips when reading this table:
- All standard reduction potentials are relative to hydrogen, which is assigned a value of zero. The potentials for all other reactions are measured by using what is known as a standard hydrogen electrode.
- All reactions are listed as reduction reactions, where an element or ion gains electrons and takes on a more negative charge.
- Note that elemental fluorine, at the top of the table, is the most powerful oxidizer, and that nothing in the table can oxidize fluoride to a positive ion.
- Note that elemental lithium, at the bottom of the list of standard reduction potentials, is the most powerful reducer. You can even see lithium reduce potassium ions to potassium metal. When lithium acts a reducing agent, you reverse the equation and the charge listed in the table, giving the reaction a potential of +3.05 volts.
- The more positive the potential, the more likely the half-reaction is to occur
- These potentials are all measured in standard conditions (25 Celsius, 1 atm pressure, 1 molar concentration solution).
- Reactions with H+ ions in the equation take place in acidic solutions. Reactions with OH- ions take place in a basic solution.
- To get the potential of the reverse reaction, known as an “oxidation potential”, simply reverse the sign of the potential. For example, standard oxidation potential for the half reaction of fluoride ions to elemental fluorine has a potential of -2.87 volts (which means it is very difficult to make this reaction occur).
- This table can be used to predict if a metal will replace another metal in solution. For example, to see if zinc metal will replace copper ions in solution, producing elemental copper, check if the potential of Zn->Zn+2 (+0.76) added to the potential of Cu+2->Cu (+0.34) is greater than zero (it is!)
Standard Reduction Potentials for common half-reactions
Common oxidizing and reducing agents
If you scan down the list, you will notice many common strong oxidizing agents in the lab near the top of the list, such as hydrogen peroxide, the peroxydisulfate ion, the permanganate ion, and the hypochlorite ion. The higher up in the list, the more powerful of an oxidizing agent the element/ion on the left side of the half-reaction is.
You’ll also see the nitrate ion, NO3– in the list. The nitrate ion is a medium-powered oxidizing agent, which is why nitric acid can react with many metals like copper that do not react with a non-oxidizing acid like hydrochloric acid.
Further down the list, you will see common reducing agents such as zinc metal, the tin (II) ion, and both the sulfite and thiosulfate ion which have similar reducing capability. The further down the list of standard reduction potentials, the more powerful the reducing agent on the right side of the half-reaction is.
Redox reactions are some of the most exciting reactions in chemistry. We hope you consult this table often!
What is the strongest oxidizing agent?
Fluorine, ozone, hydrogen peroxide, and the permanganate ion are considered some of the strongest oxidizing agents. That said, elemental fluorine is almost never used as an oxidizing agent due to how dangerous it is.
List of Standard Reduction Potentials
Standard Reduction Potential Table
|F2(g) + 2e-1 ———> 2F-1(aq)||+2.87|
|O3(g) + 2H+1(aq) + 2e-1 ———> O2(g) + H2O(l)||+2.08|
|S2O82-(aq) + 2e-1 ———> 2 SO42-(aq)||+2.05|
|Au1+(aq) + e-1 ———> Au(s)||+1.83|
|Co3+(aq) + e-1 ———> Co2+(aq)||+1.82|
|H2O2(aq) + 2 H+1(aq) + 2e-1 ———> 2 H2O(l)||+1.77|
|MnO4-1(aq) + 4 H+1(aq) + 3e-1 ———> MnO2(s) + 2 H2O(l)||+1.695|
|PbO2(s) + SO42-(aq) + 4 H+1(aq) + 2e-1 ———> PbSO4(s) + 2 H2O(l)||+1.69|
|2 HOCl(aq) + 2 H+1(aq) + 2e-1 ———> Cl2(g) + 2 H2O(l)||+1.63|
|Mn3+(aq) + e-1 ———> Mn2+(aq)||+1.51|
|MnO4-1(aq) + 8 H+1(aq) + 5e-1 ———> Mn2+(aq) + 4 H2O(l)||+1.49|
|ClO3-1(aq) + 12 H+1(aq) + 10e-1 ———>Cl2(g) + 6 H2O(l)||+1.49|
|PbO2(s) + 4 H+1(aq) + 2e-1 ———> Pb2+(aq) + 2 H2O(l)||+1.46|
|BrO3-1(aq) + 6 H+1(aq) + 6e-1 ———> Br-1(aq) + 3 H2O(l)||+1.44|
|Ce4+(aq) + e-1 ———> Ce3+(aq)||+1.44|
|Au3+(aq) + 3e-1 ———> Au(s)||+1.42|
|Cl2(g) + 2e-1 ———> 2 Cl-1(aq)||+1.36|
|Cr2O72-(aq) + 14 H+1(aq) + 6e-1 ———> 2 Cr3+(aq) + 7 H2O(l)||+1.33|
|O3(g) + H2O(l) + 2e-1 ———> O2(g) + 2 OH-1(aq)||+1.24|
|MnO2(s) + 4 H+1(aq) + 2e-1 ———> Mn2+(aq) + 2 H2O(l)||+1.23|
|O2(g) + 4 H+1(aq) + 4e-1 ———> 2 H2O(l)||+1.23|
|Pt2+(aq) + 2e-1 ———> Pt(s)||+1.20|
|IO3-1(aq) + 5H+1(aq) + 4e-1 ———HIO(aq) + 2 H2O(l)||+1.13|
|Br2(aq) + 2e-1 ———> 2 Br-1(aq)||+1.07|
|NO3-1(aq) + 4 H+1(aq) + 3e-1 ———> NO(g) + 2 H2O(l)||+0.96|
|NO3-1(aq) + 3 H+1(aq) + 2e-1 ———> HNO2(g) + H2O(l)||+0.94|
|2 Hg2+(aq) + 2e-1 ———> Hg22+(aq)||+0.91|
|HO2-1(aq) + H2O(l) + 2e-1 ———> 3 OH-1(aq)||+0.87|
|2 NO3-1(aq) + 4 H+1(aq) + 2e-1 ———> 2 NO2(g) + 2H2O(l)||+0.80|
|Ag+1(aq) + e-1 ———> Ag(s)||+0.80|
|Fe3+(aq) + e-1 ———> Fe2+ (aq)||+0.77|
|O2(g) + 2H+1(aq) + 2e-1 ———> H2O2(aq)||+0.69|
|I2(s) + 2e-1 ———> 2 I-1(aq)||+0.54|
|NiO2(s) + 2 H2O(l) + 2e-1 ———> Ni(OH)2 + 2 OH-1(aq)||+0.49|
|SO2(aq) + 4 H+1(aq) + 4e-1 ———> S(s) + 2 H2O(l)||+0.45|
|O2(g) + 2 H2O(l) + 4e-1 ———> 4 OH-1(aq)||+0.401|
|Cu2+(aq) + 2e-1 ———> Cu(s)||+0.34|
|Hg2Cl2(s) + 2e-1 ———> 2 Hg(l) + 2 Cl-1(aq)||+0.27|
|PbO2(s) + H2O(l) + 2e-1 ———> PbO(s) + 2 OH-1(aq)||+0.25|
|AgCl(s) + e-1 ———> Ag(s) + Cl-1(aq)||+0.2223|
|SO42-(aq) + 4H+1(aq) + 2e-1 ———> H2SO3(aq) + H2O(l)||+0.172|
|S4O62-(aq) + 2e-1 ———> 2 S2O32-(aq)||+0.169|
|Cu2+(aq) + e-1 ———> Cu+1(aq)||+0.16|
|Sn4+(aq) + 2e-1 ———> Sn2+(aq)||+0.15|
|S(s) + 2H+1(aq) + 2e-1 ———> H2S(g)||+0.14|
|AgBr(s) + e-1 ———> Ag(s) + Br-1(aq)||+0.07|
|2 H+1(aq) + 2e-1 ———> H2(g)||0.00|
|Pb2+(aq) + 2e-1 ———> Pb(s)||-0.13|
|Sn2+(aq) + 2e-1 ———> Sn(s)||-0.14|
|AgI(s) + e-1 ———> Ag(s) + I-1(aq)||-0.15|
|Ni2+(aq) + 2e-1 ———> Ni(s)||-0.25|
|Co2+(aq) +2e-1 ———> Co(s)||-0.28|
|In3+(aq) + 3e-1 ———> In(s)||-0.34|
|Tl+1(aq) + e-1 ———> Tl(s)||-0.34|
|PbSO4(s) + 2e-1 ———> Pb(s) + SO42-(aq)||-0.36|
|Cd2+(aq) + 2e-1 ———> Cd(s)||-0.40|
|Fe2+(aq) + 2e-1 ———> Fe(s)||-0.44|
|Ga3+(aq) + 3e-1 ———> Ga(s)||-0.56|
|PbO(s) + H2O(l) + 2e-1 ———> Pb(s) + 2 OH-1(aq)||-0.58|
|Cr3-(aq) + 3e-1 ———> Cr(s)||-0.74|
|Zn2+(aq) + 2e-1 ———> Zn(s)||-0.76|
|2 H2O(l) + 2e-1 ———> H2(g) + 2 OH-1(aq)||-0.83|
|Fe(OH)2(s) + 2e-1 ———> Fe(s) + 2 OH-1(aq)||-0.88|
|Cr2+(aq) + 2e-1 ———> Cr(s)||-0.91|
|N2(g) + 4 H2O(l) + 4e-1 ———> N2O4(aq) +4 OH-1(aq)||-1.16|
|V2+(aq) + 2e-1 ———> V(s)||-1.18|
|ZnO2-1(aq) + 2 H2O(l) + 2e-1 ———> Zn(s) + 4OH-1(aq)||-1.216|
|Ti2+(aq) + 2e-1 ———> Ti(s)||-1.63|
|Al3+(aq) + 3e-1 ———> Al(s)||-1.66|
|U3+(aq) + 3e-1 ———> U(s)||-1.79|
|Sc3+(aq) + 3e-1 ———> Sc(s)||-2.02|
|Er3+(aq) + 3e-1 ———> Er(s)||–2.33|
|Ce3+(aq) + 3e-1 ———> Ce(s)||–2.34|
|Pr3+(aq) + 3e-1 ———> Pr(s)||-2.35|
|La3+(aq) + 3e-1 ———> La(s)||-2.36|
|Y3+(aq) + 3e-1 ———> Y(s)||-2.37|
|Mg2+(aq) + 2e-1 ———> Mg(s)||-2.37|
|Na+1(aq) + e-1 ———> Na(s)||-2.71|
|Ca2+(aq) + 2e-1 ———> Ca(s)||-2.76|
|Sr2+(aq) + 2e-1 ———> Sr(s)||-2.89|
|Ba2+(aq) + 2e-1 ———> Ba(s)||-2.90|
|Cs+1(aq) +e-1 ———> Cs(s)||-2.92|
|K+1(aq) + e-1 ———> K(s)||-2.92|
|Rb+1(aq) + e-1 ———> Rb(s)||-2.93|
|Li+1(aq) + e-1 ———> Li(s)||-3.05|