Reactivity series

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Metal Ion Reactivity Extraction
Cs Cs+ reacts with water electrolysis
Rb Rb+
K K+
Na Na+
Li Li+
Ra Ra2+
Ba Ba2+
Sr Sr2+
Ca Ca2+
Mg Mg2+ reacts with acids
Al Al3+
Ti Ti4+ reacts with concentrated mineral acids Pyrometallurgical extraction using magnesium, or less commonly other alkali metals, hydrogen or calcium in the Kroll process
Mn Mn2+ reacts with acids smelting with coke
Zn Zn2+
Cr Cr3+
Fe Fe2+
Cd Cd2+
Co Co2+
Ni Ni2+
Sn Sn2+
Pb Pb2+
Sb Sb3+ may react with some strong oxidizing acids heat or
physical extraction
Bi Bi3+
Cu Cu2+
W W3+
Hg Hg2+
Ag Ag+
Pt Pt2+
Au Au3+

Going from bottom to top the metals:

  • increase in reactivity;
  • lose electrons more readily to form positive ions;
  • corrode or tarnish more readily;
  • become stronger reducing agents (electron donors).

Defining reactions[edit]


Reaction with water and acids[edit]






Single displacement reactions[edit]

An iron nail placed in a solution of copper sulfate will quickly change colour as metallic copper is deposited and the iron is converted into iron(II) sulfate:








Comparison with standard electrode potentials[edit]


Li > K > Sr > Na > Ca > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H > Cu > Ag > Hg > Pt > Au



potassium > sodium > lithium > alkaline earth metals,


Anomalous electrode potential of lithium[edit]

The standard electrode potential of a reaction Eo is related to the Gibbs free energy change ΔGo and the enthalpy change ΔHo by

\Delta G^\ominus = -nFE^\ominus = \Delta H^\ominus - T\Delta S^\ominus\,
M(s) → M+(aq) + e
M = Li Na K
ΔatHo / kJ mol−1 +162 +110 +90
Ei / kJ mol−1 +520 +496 +419
ΔhydrHo / kJ mol−1 –960 –846 –761
ΔfHo / kJ mol−1 –278 –240 –252
ΔfSo / J K−1 mol−1 +51 +73 +104
ΔfGo / kJ mol−1 –293 –261 –283
Eo / V +3.04 +2.71 +2.93
Sources: Jolly (1991),[5] Greenwood &
Earnshaw (1984).[4][6]

The same relation can be applied to the half-reactions of tabulated standard electrode potentials, as the energy associated with the electron(s) will cancel whenever a full reaction is considered and so can be ignored. The enthalpy change of the half-reaction can be further analyzed as a Born–Haber cycle, and shown to be equal to the enthalpy change of atomisation ΔatHo plus the ionization energy Ei plus the enthalpy change of hydration ΔhydrHo of the gaseous ions. There is a further contribution from the entropy change of the half-reaction, ΔSo. Values of these quantities for lithium, sodium and potassium are given in the table (note that the reaction is shown for the formation of M+, which is the reverse of the reduction reaction for which standard electrode potentials are conventionally quoted).

The high electrode potential of the Li/Li+ couple can be seen to be entirely due to the unusually large enthalpy change of hydration of the gaseous Li+ ion, which overrides the other three factors. The lithium ion is small and highly polarising, and so binds the coordinated water molecules very tightly. This can also be seen in the relatively low positive entropy change for the formation of Li+(aq): the presence of lithium ions in solution creates more order in the solvent (water) than for either sodium or potassium ions.

The reaction of metallic lithium with water is more exothermic than any other alkali metal (many times more exothermic for a given mass or volume), but it is less vigorous for kinetic reasons. The enthalpy change of atomisation is twice as high for lithium than for potassium, indicating stronger bonding in the metal (both metals have the same body-centered cubic structure). Another way to measure an element's reactivity relative to other elements, is to look at its electronegativity. This is most useful when looking at how anions might replace other anions in a reaction. Anions with greater electronegativity values are more reactive and will replace anions with lower electronegativity values.

See also[edit]



  3. ^ Lim Eng Wah (2005), Pearson Education, p. 190  Unknown parameter |unused_data= ignored (help); Missing or empty |title= (help)
  5. ^ Jolly, William L. (1991), Modern Inorganic Chemistry (2nd ed.), New York: McGraw-Hill, pp. 586–87, ISBN 0-07-032768-8 
  6. ^ The values for the enthalpy changes of hydration of Li+ (−520 kJ/mol) and Na+ (−405 kJ/mol) given in Greenwood & Earnshaw (1984) are inconsistent with the other data quoted in the same source and certainly far too low. The enthalpy changes for the gas-phase reactions Li+(g) + 6H2O → [Li(H2O)6]+ and Na+(g) + 5H2O → [Na(H2O)5]+ have been measured as −529(16) kJ/mol and −347(14) kJ/mol respectively: Rodgers, M. T.; Armentrout, P. B. (2000), "Noncovalent Metal-Ligand Bond Energies as Studied by Threshold Collision-Induced Dissociation", Mass Spectrom. Rev. 19 (4): 215, doi:10.1002/1098-2787(200007)19:4<215::AID-MAS2>3.0.CO;2-X, PMID 10986693 . The enthalpy changes of hydration quoted in the table have been calculated from the enthalpy changes of formation of the ions in solution.

External links[edit]