Reactivity series Totally Explained

Everything about Reactivity Series totally explained

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.

Differing definitions

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

A reactivity series of common metals

Here is 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⁺
Ba	Ba²⁺
Sr	Sr⁺
Ca	Ca²⁺
Li	Li⁺
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	Sb³⁺	highly unreactive
Bi	Bi³⁺
Cu	Cu²⁺
Hg	Hg²⁺
Ag	Ag⁺
Pt	Pt⁺
Au	Au³⁺

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

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

Simplified Version

The simplified version that's taught in the GCSE and GCE 'O' Level chemistry course, as the basic, are listed below. Higher education and standard level are required to study more metals as shown above.

Metals	Metal Ion	Reactivity
Li	Li⁺	Going from bottom to top, the metals: reactivity increases lose electrons more readily form positive ions more readily become stronger reducing agents A metal 'high up' in the reactivity series: reacts vigorously and quickly with chemicals readily gives up electrons in reactions to form positive ions is corroded easily A metal 'low down' in the reactivity series: doesn't react vigorously and quickly with chemicals doesn't readily give up electrons in reactions to form positive ions isn't corroded easily
K	K⁺
Ca	Ca²⁺
Na	Na⁺
Mg	Mg²⁺
Al	Al³⁺
Zn	Zn²⁺
Fe	Fe²⁺
Sn	Sn²⁺
Pb	Pb²⁺
H₂	H⁺
Cu	Cu²⁺
Ag	Ag⁺
Au	Au³⁺

Significance

The reactivity series determines qualitatively characteristics such as the reactions with water, air and acids as demonstrated above. However it's 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. What is really occurring when the two metals swap is that the metal higher on the chain is acting as a reducing agent, and transferring one or more of its electrons to the other metal, which since it's lower on the chain than the first metal is more apt to be oxidized, receive the electron. It follows the basic reaction form of 2Li+Cu²⁺--->2Li⁺+Cu
Lithium metal is the most reactive of all metals....it is the one with the highest ΔH of reaction gives away one of its electrons.... the reason that K metal or Cs metal seem more reactive when placed in water is because their atoms are larger meaning their electrons are held farther out from the nucleus, making it easier for them to give them up at a faster rate than lithium...which means that they release more energy in a smaller amount of time then lithium can because they can let go of their electrons faster than lithium ..but over all, a mole of lithium releases a lot more energy than a mole of cesium when giving away its electrons.

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