In chemistry, the reactivity series is a series of metals, in order of reactivity from highest to lowest. It is used to determine the products of single displacement reactions, whereby metal A will replace another metal B in a solution if A is higher in the series.

Different Methods of definition

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In the UK a reduced version of the below series is taught as part of the GCSE chemistry course, leading to various mnemonics being invented to aid memory. The activity series taught in the US is defined by the ease of oxididation and corresponds to the ordering of the table of standard electrode potentials. This is markedly different from the below table.

A reactivity series of common metals

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Here a series of some of the most common metals, listed in descending order of reactivity.

Metals	Metal Ion	Reactivity
K	K+	reacts with water
Na	Na+
Li	Li+
Ba	Ba2+
Sr	Sr2+
Ca	Ca2+
Mg	Mg2+	reacts with acids
Al	Al3+
Mn	Mn2+
Zn	Zn2+
Cr	Cr2+
Fe	Fe2+
Cd	Cd2+
Co	Co2+
Ni	Ni2+
Sn	Sn2+
Pb	Pb2+
H2	H+	included for comparison
Sb	Sb2+	highly unreactive
Bi	Bi2+
Cu	Cu2+
Hg	Hg2+
Ag	Ag+
Au	Au3+
Pt	Pt+

A metal can replace metals listed below it in the activity series, but not above. For example, sodium is highly active and thus able to replace hydrogen from water:

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

Metals that can replace hydrogen within acids but not water are listed in the middle of the activity series, for example zinc replaces hydrogen in sulfuric acid:

Zn (s) + H₂SO₄ (aq) → ZnSO₄ (aq) + H₂ (g)

The activity series has applications in electrochemistry, where two dissimilar metals are chosen as electrodes of a battery (though the above table is not exact for this purpose. See Table of standard electrode potentials).

Significance

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The reactivity series determines qualitatively characteristics such as the reactions with water, air and acids as demonstrated above. However it is defined by the nature of the metals in single displacement reactions.

When a metal in elemental form is placed in a solution of a metal salt it may be, overall, more energetically feasible for this "elemental metal" to exist as an ion and the "ionic metal" to exist as the element. Therefore the elemental metal will 'displace' the ionic metal over time, thus the two swap places. Only a metal higher in the reactivity series will displace another.

Explanation

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The reactivity of these metals are due to the difference in stability of their electron arrangements as atoms and as ions. As they are all metals they will form positive ions when they react.

Potassium has a single outer shell electron to lose to obtain a stable "Noble gas" electron arrangement whereas the precious metals which exist in the d-block cannot form structures which are much more stable than their current state with the loss of small numbers of electrons. Metals that require the loss of single electrons to form stable ions tend to be more reactive than similar metals which require the loss of more than one electron.

Metals with a greater total number of electrons tend to be more reactive as their outermost electrons (the ones which will be lost) exist further from the positive nucleus and therefore they are held less strongly.

External links

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• Science Line Chemistry

Chemical properties

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