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Not to be confused with Electrochemical series.

In chemistry, a reactivity series (or reactivity series of elements) is an empirical, calculated, and structurally analytical progression[1] of a series of metals, arranged by their "reactivity" from highest to lowest.[2][3][4] It is used to summarize information about the reactions of metals with acids and water, single displacement reactions and the extraction of metals from their ores.[5]

Table[edit]

Metal Ion Reactivity Extraction
Caesium Cs Cs+ reacts with cold water Electrolysis (a.k.a. electrolytic refining)
Rubidium Rb Rb+
Potassium K K+
Sodium Na Na+
Lithium Li Li+
Barium Ba Ba2+
Strontium Sr Sr2+
Calcium Ca Ca2+
Magnesium Mg Mg2+ reacts very slowly with cold water, but rapidly
in boiling water, and very vigorously with acids
Beryllium Be Be2+ reacts with acids and steam
Aluminium Al Al3+
Titanium Ti Ti4+ reacts with concentrated mineral acids pyrometallurgical extraction using magnesium,
or less commonly other alkali metals, hydrogen or calcium in the Kroll process
Manganese Mn Mn2+ reacts with acids; very poor reaction with steam smelting with coke
Zinc Zn Zn2+
Chromium Cr Cr3+ aluminothermic reaction
Iron Fe Fe2+ smelting with coke
Cadmium Cd Cd2+
Cobalt Co Co2+
Nickel Ni Ni2+
Tin Sn Sn2+
Lead Pb Pb2+
Antimony Sb Sb3+ may react with some strong oxidizing acids heat or physical extraction
Bismuth Bi Bi3+
Copper Cu Cu2+ reacts slowly with air
Tungsten W W3+ may react with some strong oxidizing acids
Mercury Hg Hg2+
Silver Ag Ag+
Gold Au Au3+[6][7]
Platinum Pt Pt4+

Going from the bottom to the top of the table the metals:

  • increase in reactivity;
  • lose electrons (oxidize) more readily to form positive ions;
  • corrode or tarnish more readily;
  • require more energy (and different methods) to be isolated from their compounds;
  • become stronger reducing agents (electron donors).

Defining reactions[edit]

There is no unique and fully consistent way to define the reactivity series, but it is common to use the three types of reaction listed below, many of which can be performed in a high-school laboratory (at least as demonstrations).[6]

Reaction with water and acids[edit]

The most reactive metals, such as sodium, will react with cold water to produce hydrogen and the metal hydroxide:

2 Na (s) + 2 H2O (l) →2 NaOH (aq) + H2 (g)

Metals in the middle of the reactivity series, such as iron, will react with acids such as sulfuric acid (but not water at normal temperatures) to give hydrogen and a metal salt, such as iron(II) sulfate:

Fe (s) + H2SO4 (l) → FeSO4 (aq) + H2 (g)

There is some ambiguity at the borderlines between the groups. Magnesium, aluminium and zinc can react with water, but the reaction is usually very slow unless the metal samples are specially prepared to remove the surface passivation layer of oxide which protects the rest of the metal. Copper and silver will react with nitric acid; but because nitric acid is an oxidizing acid, the oxidizing agent is not the H+ ion as in normal acids, but the NO3− ion.

Comparison with standard electrode potentials[edit]

The reactivity series is sometimes quoted in the strict reverse order of standard electrode potentials, when it is also known as the "electrochemical series".[8]

The following list includes the metallic elements of the first six periods. It is mostly based on tables provided by NIST.[9][10] However, not all sources give the same values: there are some differences between the precise values given by NIST and the CRC Handbook of Chemistry and Physics. In the first six periods this does not make a difference to the relative order, but in the seventh period it does, so the seventh-period elements have been excluded. (In any case, the typical oxidation states for the most accessible seventh-period elements thorium and uranium are too high to allow a direct comparison.)[11]

Hydrogen has been included as a benchmark, although it is not a metal. Borderline germanium, antimony, and astatine have been included. Some other elements in the middle of the 4d and 5d rows have been omitted (Zr–Tc, Hf–Os) when their simple cations are too highly charged or of rather doubtful existence. Greyed-out rows indicate values based on estimation rather than experiment.

Z Sym Element Reaction E° (V)
3 Li lithium Li+ + e− → Li −3.04
55 Cs caesium Cs+ + e− → Cs −3.03
37 Rb rubidium Rb+ + e− → Rb −2.94
19 K potassium K+ + e− → K −2.94
56 Ba barium Ba2+ + 2 e− → Ba −2.91
38 Sr strontium Sr2+ + 2 e− → Sr −2.90
20 Ca calcium Ca2+ + 2 e− → Ca −2.87
11 Na sodium Na+ + e− → Na −2.71
57 La lanthanum La3+ + 3 e− → La −2.38
39 Y yttrium Y3+ + 3 e− → Y −2.38
12 Mg magnesium Mg2+ + 2 e− → Mg −2.36
59 Pr praseodymium Pr3+ + 3 e− → Pr −2.35
58 Ce cerium Ce3+ + 3 e− → Ce −2.34
68 Er erbium Er3+ + 3 e− → Er −2.33
67 Ho holmium Ho3+ + 3 e− → Ho −2.33
60 Nd neodymium Nd3+ + 3 e− → Nd −2.32
69 Tm thulium Tm3+ + 3 e− → Tm −2.32
62 Sm samarium Sm3+ + 3 e− → Sm −2.30
61 Pm promethium Pm3+ + 3 e− → Pm −2.30
66 Dy dysprosium Dy3+ + 3 e− → Dy −2.29
71 Lu lutetium Lu3+ + 3 e− → Lu −2.28
65 Tb terbium Tb3+ + 3 e− → Tb −2.28
64 Gd gadolinium Gd3+ + 3 e− → Gd −2.28
70 Yb ytterbium Yb3+ + 3 e− → Yb −2.19
21 Sc scandium Sc3+ + 3 e− → Sc −2.09
63 Eu europium Eu3+ + 3 e− → Eu −1.99
4 Be beryllium Be2+ + 2 e− → Be −1.97
13 Al aluminium Al3+ + 3 e− → Al −1.68
22 Ti titanium Ti3+ + 3 e− → Ti −1.37
25 Mn manganese Mn2+ + 2 e− → Mn −1.18
23 V vanadium V2+ + 2 e− → V −1.12
24 Cr chromium Cr2+ + 2 e− → Cr −0.89
30 Zn zinc Zn2+ + 2 e− → Zn −0.76
31 Ga gallium Ga3+ + 3 e− → Ga −0.55
26 Fe iron Fe2+ + 2 e− → Fe −0.44
48 Cd cadmium Cd2+ + 2 e− → Cd −0.40
49 In indium In3+ + 3 e− → In −0.34
81 Tl thallium Tl+ + e− → Tl −0.34
27 Co cobalt Co2+ + 2 e− → Co −0.28
28 Ni nickel Ni2+ + 2 e− → Ni −0.24
50 Sn tin Sn2+ + 2 e− → Sn −0.14
82 Pb lead Pb2+ + 2 e− → Pb −0.13
1 H hydrogen 2 H+ + 2 e− → H2 0.00
32 Ge germanium Ge2+ + 2 e− → Ge +0.1
51 Sb antimony Sb3+ + 3 e− → Sb +0.15
83 Bi bismuth Bi3+ + 3 e− → Bi +0.31
29 Cu copper Cu2+ + 2 e− → Cu +0.34
84 Po polonium Po2+ + 2 e− → Po +0.6
44 Ru ruthenium Ru3+ + 3 e− → Ru +0.60
45 Rh rhodium Rh3+ + 3 e− → Rh +0.76
47 Ag silver Ag+ + e− → Ag +0.80
80 Hg mercury Hg2+ + 2 e− → Hg +0.85
46 Pd palladium Pd2+ + 2 e− → Pd +0.92
77 Ir iridium Ir3+ + 3 e− → Ir +1.0
85 At astatine At+ + e− → At +1.0
78 Pt platinum Pt2+ + 2 e− → Pt +1.18
79 Au gold Au3+ + 3 e− → Au +1.50

The positions of lithium, sodium and gold are changed on such a series.

Standard electrode potentials offer a quantitative measure of the power of a reducing agent, rather than the qualitative considerations of other reactive series. However, they are only valid for standard conditions: in particular, they only apply to reactions in aqueous solution. Even with this proviso, the electrode potentials of lithium, sodium and gold – and hence their positions in the electrochemical series – appear anomalous. The order of reactivity, as shown by the vigour of the reaction with water or the speed at which the metal surface tarnishes in air, appears to be

caesium>potassium > sodium > lithium > alkaline earth metals,

i.e., alkali metals>alkaline earth metals

the same as the reverse order of the (gas-phase) ionization energies. This is borne out by the extraction of metallic lithium by the electrolysis of a eutectic mixture of lithium chloride and potassium chloride: lithium metal is formed at the cathode, not potassium.[1]

Comparison with electronegativity values[edit]

The image shows a periodic table extract with the electronegativity values of metals.[12]

Wulfsberg[13] distinguishes:
   very electropositive metals with electronegativity values below 1.4
   electropositive metals with values between 1.4 and 1.9; and
   electronegative metals with values between 1.9 and 2.54.

From the image, the group 1–2 metals and the lanthanides and actinides are very electropositive to electropositive; the transition metals in groups 3 to 12 are very electropositive to electronegative; and the post-transition metals are electropositive to electronegative. The noble metals, inside the dashed border (as a subset of the transition metals) are very electronegative.

See also[edit]

References[edit]

  1. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 82–87. ISBN 978-0-08-022057-4.
  2. ^ France, Colin (2008), The Reactivity Series of Metals
  3. ^ Briggs, J. G. R. (2005), Science in Focus, Chemistry for GCE 'O' Level, Pearson Education, p. 172
  4. ^ Lim Eng Wah (2005), Longman Pocket Study Guide 'O' Level Science-Chemistry, Pearson Education, p. 190
  5. ^ "Metal extraction and the reactivity series - The reactivity series of metals - GCSE Chemistry (Single Science) Revision - WJEC". BBC Bitesize. Retrieved 2023-03-24.
  6. ^ a b Activity series at the Wayback Machine (archived 2019-05-07)
  7. ^ Wulsberg, Gary (2000). Inorganic Chemistry. p. 294. ISBN 9781891389016.
  8. ^ Periodic table poster at the Wayback Machine (archived 2022-02-24) by A. V. Kulsha and T. A. Kolevich gives:

    Li > Cs > Rb > K > Ba > Sr > Ca > Na > La > Y > Mg > Ce > Sc > Be > Al > Ti > Mn > V > Cr > Zn > Ga > Fe > Cd > In > Tl > Co > Ni > Sn > Pb > (H) > Sb > Bi > Cu > Po > Ru > Rh > Ag > Hg > Pd > Ir > Pt > Au
  9. ^ Standard Electrode Potentials and Temperature Coefficients in Water at 298.15 K, Steven G. Bratsch (NIST)
  10. ^ For antimony: Antimony - Physico-chemical properties - DACTARI
  11. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
  12. ^ Aylward, G; Findlay, T (2008). SI Chemical Data (6 ed.). Milton, Queensland: John Wiley & Sons. p. 126. ISBN 978-0-470-81638-7.
  13. ^ Wulfsberg, G (2018). Foundations of Inorganic Chemistry. Mill Valley: University Science Books. p. 319. ISBN 978-1-891389-95-5.

External links[edit]

source: https://en.wikipedia.org/wiki/Reactivity_series

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