# Electrochemical series

The electrochemical series is a listing of redox pairs according to their standard electrode potential ( redox potential under standard conditions ). It is also known as the redox series , especially for metals .

The redox behavior of a substance can be derived from the electrochemical series.

Every redox reaction can be described by two pairs and the direction of reactions can be predicted from the electrochemical series.

## Interpretation and meaning

### Metals

In metals, the metal itself and its associated form ion a redox couple   ( Red ... ... ox + z e - ). For example: ${\ displaystyle \ rightleftharpoons}$

${\ displaystyle \ mathrm {Cu} ^ {2 +} + 2 \ mathrm {e} ^ {-} \ rightleftharpoons \ mathrm {Cu}}$
Galvanic cell (here: Daniell element)

Cu is the reduced form ("Red ...") and Cu 2+ is the oxidized form ("... ox"). The redox potential is a measure of the readiness of the ions to accept the electrons. The ions of the noble metals accept electrons more readily than the ions of base metals, which is why the redox potential of the Cu / Cu 2+ pair with +0.35 V is significantly more positive than that of the Zn / Zn 2+ pair with −0 under standard conditions , 76 V. And that in turn means that Zn is one of the less noble metals and is a stronger reducing agent, i.e. its reactant is reduced and itself oxidized and emits electrons.

(“Under standard conditions” means that the concentration - more precisely: activity - of the ions must be 1 mol / l so that the redox potential assumes the values ​​in the table . This restriction is necessary because equilibrium reactions are involved. According to LeChatelier's principle , a a higher concentration of metal ions also a greater willingness to be reduced to the metal and therefore a higher redox potential. The Nernst equation describes this relationship mathematically.)

Redox potentials themselves cannot be measured. In contrast, the difference between two electrode potentials can be measured . An electrode under standard conditions is created simply by immersing a metal in a solution that contains its ions at a concentration of 1 mol / l. If two such electrodes are connected electrically (ion bridge), a galvanic cell is created and a voltage can be measured between the metals. This voltage is equal to the difference between the standard electrode potentials that belong to the redox pairs in the electrode spaces and are tabulated in the electrochemical series. For the example of the combination of the redox pairs Cu / Cu 2+ and Zn / Zn 2+ , a Daniell element with a voltage of 1.11 V is created.

### Ion / gas electrodes (normal hydrogen electrode)

Gaseous hydrogen and its two protons are also a redox couple:

${\ displaystyle 2 \ mathrm {H} ^ {+} + 2 \ mathrm {e} ^ {-} \ rightleftharpoons \ mathrm {H} _ {2}}$

Electrodes for redox pairs with gaseous substances are made by immersing an inert metal (Pt) in a 1 mol / l solution of ions (H + ) and flushing it with the associated gas (H 2 ) at a pressure of 1 bar. In a special case, a normal hydrogen electrode is created . This electrode is easy to set up and provides a constant, reproducible potential. Since the redox couple H 2 / H + also describes the effect of acids (it always appears when metals are dissolved in acids: e.g. Mg + 2 H +  → Mg 2+  + H 2 ), it became the standard potential of the normal hydrogen electrode is defined as zero for practical reasons.

All other standard potentials are therefore the voltages that are measured in a galvanic cell when the normal hydrogen electrode on the left and the electrode of the redox pair on the right are connected. (In each case under standard conditions!)

### Applications

The electrochemical voltage series allows the calculation of the maximum voltages that batteries and accumulators can deliver. Conversely, these are the voltages that must be applied at least to drive electrolysis or to charge the batteries.

It is also possible to calculate the direction and strength of the reaction . If two redox pairs are mixed in a reaction solution, the reduction will take place for the pair with the higher redox potential and the oxidation for the pair with the lower redox potential. If you dive z. For example, a zinc sheet in a CuSO 4 solution, zinc is oxidized due to its lower redox potential (−0.76 V) and goes into solution as zinc ions, whereas copper ions (+0.35 V) are reduced at the same time and deposit as a copper coating on the zinc sheet. (This popular example disregards the requirement for standard conditions. For example, a copper sheet that is immersed in a ZnSO 4 solution will be coated a little with zinc because initially no zinc is present and the Cu 2+ concentration is zero. The effect can be calculated with the Nernst equation , but is immeasurably small, so that the example has a certain justification.) A measure of the strength of the reaction is the Gibbs energy (free enthalpy) of the associated reaction, which after

${\ displaystyle \ Delta G ^ {\ circ} = - z \, F \, \ Delta E ^ {\ circ}}$

can be calculated. This is the number of exchanged electrons,  = 96.485 C mol −1 the Faraday constant and Δ E ° the difference between the standard potentials . ${\ displaystyle z}$${\ displaystyle F}$${\ displaystyle (\ Delta E ^ {\ circ} = E_ {red} ^ {\ circ} -E_ {ox} ^ {\ circ})}$

The reduced form of a redox couple with a very negative standard potential is a very strong reducing agent because it tends to donate electrons (e.g. sodium). In contrast, the oxidized form of a redox couple with a very positive standard potential is a strong oxidizing agent (e.g. fluorine as the strongest known oxidizing agent, i.e. with the highest standard potential) because it tends to accept electrons. Contact corrosion of metals with different standard electrode potentials can result. The electrochemical series is thus a listing of oxidizing agents according to oxidation strength or, at the same time, a reverse listing of reducing agents according to reducing strength.

In addition, the electrochemical series contains a gradation of metals (“very noble metal”, “noble metal”, “less noble metal”, “base metal”, “very base metal”) according to their endeavors to be oxidized in acids. The standard potentials of noble metals have a positive sign, while those of the base metals have a negative sign. Base metals therefore dissolve in acids because acids contain H + . (The arguments for example Zn / Cu apply analogously.)

Noble metals, on the other hand, only dissolve in oxidizing acids.

## Electrochemical series

(Standard potentials at 25 ° C; 101.3 kPa; pH = 0; ion activities 1 mol / l)

Element in the redox
couple, the oxidation
state of which changes
oxidized form z  e - reduced form Standard
potential
E °
Fluorine (F) F 2 + 2 e - 2 F - +2.87 V
Sulfur (S) S 2 O 8 2− + 2 e - 2 SO 4 2− +2.00 V.
Oxygen (O) H 2 O 2 + 2 H 3 O + + 2 e - 4 H 2 O +1.78 V
Gold (Au) Au + 2 e - Au +1.69 V
Gold (Au) Au 3+ + 3 e - Au +1.50 V
Gold (Au) Au 3+ + 2 e - Au + +1.40V
Chlorine (Cl) Cl 2 + 2 e - 2 Cl - +1.36V
Chromium (Cr) Cr 6+ + 3 e - Cr 3+ +1.33V
Oxygen (O) O 2 + 4 H + + 4 e - 2 H 2 O +1.23V
Platinum (Pt) Pt 2+ + 2 e - Pt +1.20 V
Iridium (Ir) Ir 3+ + 3 e - Ir +1.156 V
Bromine (Br) Br 2 + 2 e - 2 Br - +1.07V
Nickel (Ni) NiO 2 + 2 H 2 O + 2 e - Ni (OH) 2 + 2 OH - +0.98 V
Silicon (Si) SiO 2 + 4 H + + 4 e - Si + 2 H 2 O +0.857 V
Palladium (Pd) Pd 2+ + 2 e - Pd +0.85 V
Mercury (Hg) Hg 2+ + 2 e - Ed +0.85 V
Silver (Ag) Ag + 2 e - Ag +0.80 V
Iron (Fe) Fe 3+ 2 e - Fe 2+ +0.77 V
Tellurium (Te) Te 4+ + 4 e - Te +0.568 V
Iodine (I) I 2 + 2 e - 2 I - +0.53 V
Copper (Cu) Cu + 2 e - Cu +0.52 V
Oxygen (O) O 2 + 2 H 2 O + 4 e - 4 OH - +0.40 V
Iron (Fe) [Fe (CN) 6 ] 3− 2 e - [Fe (CN) 6 ] 4− +0.36 V
Copper (Cu) Cu 2+ + 2 e - Cu +0.35V
Bismuth (bi) Bi 3+ + 3 e - Bi +0.308 V
Copper (Cu) Cu 2+ 2 e - Cu + +0.16 V
Tin (Sn) Sn 4+ + 2 e - Sn 2+ +0.15 V.
Hydrogen (H) 2 H + 2 e - H 2 +0 V
Iron (Fe) Fe 3+ + 3 e - Fe −0.04 V
Tungsten (W) WO 2 + 4 H + + 4 e - W + 2 H 2 O −0.119 V
Lead (Pb) Pb 2+ + 2 e - Pb −0.13 V
Tin (Sn) Sn 2+ + 2 e - Sn −0.14 V
Molybdenum (Mo) Mon 3+ + 3 e - Mon −0.20 V
Nickel (Ni) Ni 2+ + 2 e - Ni −0.23 V
Cobalt (Co) Co 2+ + 2 e - Co −0.28 V
Thallium (Tl) Tl + 2 e - Tl −0.34 V
Indium (In) In 3+ + 3 e - In −0.34 V
Cadmium (Cd) Cd 2+ + 2 e - CD −0.40 V
Iron (Fe) Fe 2+ + 2 e - Fe −0.41 V
Sulfur (S) S. + 2 e - S 2− −0.48 V
Gallium (Ga) Ga 3+ + 3 e - Ga −0.549 V
Chromium (Cr) Cr 3+ + 3 e - Cr −0.76 V
Zinc (Zn) Zn 2+ + 2 e - Zn −0.76 V
water 2 H 2 O + 2 e - H 2 + 2 OH - −0.83 V
Boron (B) B (OH) 3 + 3 H 3 O + + 3 e - B + 6 H 2 O −0.89 V
Chromium (Cr) Cr 2+ + 2 e - Cr −0.91 V
Niobium (Nb) Nb 3+ + 3 e - Nb −1.099 V
Vanadium (V) V 2+ + 2 e - V −1.17 V
Manganese (Mn) Mn 2+ + 2 e - Mn −1.18 V
Titanium (Ti) Ti 3+ + 3 e - Ti −1.21 V
Uranium (U) UO 2 2+ + 4 H + + 6 e - U + 2 H 2 O −1.444 V
Zirconium (Zr) Zr 4+ + 4 e - Zr −1.45 V
Hafnium (Hf) Hf 4+ + 4 e - Hf −1.55 V
Aluminum (Al) Al 3+ + 3 e - Al −1.66 V
Uranium (U) U 3+ + 3 e - U −1.66 V
Titanium (Ti) Ti 2+ + 2 e - Ti −1.77 V
Beryllium (Be) Be 2+ + 2 e - Be −1.85 V
Scandium (Sc) Sc 3+ + 3 e - Sc −2.077 V
Neodymium (Nd) Nd 3+ + 3 e - Nd −2.323 V
Magnesium (Mg) Mg 2+ + 2 e - Mg −2.362 V
Yttrium (Y) Y 3+ + 3 e - Y −2.372 V
Cerium (Ce) Ce 3+ + 3 e - Ce −2.483 V
Lanthanum (La) La 3+ + 3 e - La −2.522 V
Sodium (Na) Na + 2 e - N / A −2.71 V
Calcium (Ca) Ca 2+ + 2 e - Approx −2.87 V
Strontium (Sr) Sr 2+ + 2 e - Sr −2.89 V
Barium (Ba) Ba 2+ + 2 e - Ba −2.92 V
Potassium (K) K + 2 e - K −2.92 V
Cesium (Cs) Cs + 2 e - Cs −2.92 V
Rubidium (Rb) Rb + 2 e - Rb −2.98 V
Lithium (Li) Li + 2 e - Li −3.04 V

## Historical

The first series of stresses were qualitative in nature and were created in 1793 by two researchers, Alessandro Volta and, independently of that, the physicist Christoph Heinrich Pfaff . To set up his row, Volta used the sensations that two different, touching metals triggered on his damp tongue, while Pfaff set up his rows according to the varying degrees of twitching of frogs' legs. When Pfaff assumed different metals, he received different series of voltages. The voltage series of Volta and Pfaff were later summarized in simplified form for the case of dilute sulfuric acid as the electrolyte:

Alessandro Volta  Christoph Heinrich Pfaff
Zinc

tin

iron

copper
silver
gold
carbon
graphite
manganese dioxide
Zinc
Tin
Tungsten
Iron
Bismuth
Antimony
Copper
Silver
Gold
Tellurium
Platinum

A year later, in 1794, Volta had included 28 elements and connections in his series of voltages.

The physical chemist Gilbert Newton Lewis , who had already dealt with electrochemical potentials in his doctoral thesis, which he completed in 1899, determined many standard potentials with previously unattainable accuracy or for the first time, namely those of sodium, potassium, rubidium, chlorine, bromine, iodine, oxygen, mercury , Silver, thallium, lead and iron. In 1938 Wendell M. Latimer published an extensive work on the redox potentials in aqueous solutions; the Latimer diagram is named after him, which describes the potentials of various oxidation states for a given element. These potential data for an element become even clearer if they are plotted according to the method proposed by Arthur A. Frost in 1951: see Frost diagram .

## literature

• Inorganic chemistry. Buchners Verlag, 1972.
• Elements chemistry II. Klett Verlag, 2000.

## Individual evidence

1. David R. Lide (Ed.): CRC Handbook of Chemistry and Physics . 97th edition. (Internet version: 2016), CRC Press / Taylor and Francis, Boca Raton, FL, Thermochemistry, Electrochemistry, and Solution Chemistry, pp. 5-78 - 5-84.
2. Course internship in general and inorganic chemistry (in Google Books) . Retrieved January 15, 2017.
3. Trevor Pearson, Stacey Handy, Rainer Lakner Forst: Electrochemical investigations into the corrosion mechanism of decorative chrome surfaces by calcium chloride. In: Galvanotechnik , edition 08/2009, p. 1736
4. a b Detlef Wienecke-Janz (Ed.): Reorganization of Europe and Restoration . 1793–1849 (=  The Great Chronicle World History . Volume 12 ). Wissenmedia Verlag / Chronik-Verlag, Gütersloh / Munich 2008, ISBN 978-3-577-09072-8 , Technical progress and engineering services - Volta examines electricity, p. 30 ( google.de [accessed June 27, 2015]).
5. ^ Voltage series - Germany, 1793. In: History of electricity theory and electrochemistry. Ulm University, accessed on November 18, 2015 .
6. ^ Wilhelm Philipp Hauck: The galvanic batteries, accumulators and thermopiles. A description of the hydro- and thermoelectric power sources ... A. Hartleben, Vienna / Pest / Leipzig 1883, introduction, p. 13 ( Textarchiv - Internet Archive ).
7. ^ RE Kohler: Lewis, Gilbert Newton. In: Complete Dictionary of Scientific Biography. Encyclopedia.com Cengage Learning, 2015, accessed November 21, 2015 .
8. ^ Joel H. Hildebrand: Gilbert Newton Lewis. (PDF) In: Biographical Memoirs of the National Academy of Sciences. National Academy of Sciences, accessed November 21, 2015 .
9. Arthur A. Frost: Oxidation Potential-Free Energy Diagrams . In: Journal of the American Chemical Society (JACS) . tape 73 , no. 6 . ACS, June 1951, ISSN  1520-5126 , pp. 2680–2682 , doi : 10.1021 / ja01150a074 .