Section 4: The Metal Activity Series: What Can Destroy a Metal?
Crucial Chemistry
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Metals have a variety of properties that make them useful in real world applications.
According to the Chem Talk in this section, metals:
o are shiny, or can be polished to shine, so they are used in art and jewelry;
o conduct electricity, so they are used in electrical circuits;
o conduct heat, so they are used in cookware;
o can withstand high temperatures, so they are used to build strong structures; and
o can formed into different shapes, so they are used as nails, flat surfaces, or boxes.
The majority of the elements on the periodic table are metallic. Most of them do not
occur in nature as pure metals, because they are so reactive. Metals are most often found
in their ionic form as part of solid crystals or dissolved in water.
The activity series of metals provides chemists with information about how reactive
metals are in relationship to each other and to hydrogen ions.
o Metal atoms and metal ions react with each other according to their position on
the activity series.
o These reactions are called oxidation-reduction reactions, or redox.
 During oxidation, electrons are lost by neutral metal atoms to form
positively charged metal ions.
 During reduction, electrons are gained by positively charged metal ions to
form neutral metal atoms.
 There are several ways to help students remember which process is which:
 OIL RIG = Oxidation Is Loss and Reduction Is Gain
 LEO (the lion says) GER = Loss of Electrons is Oxidation and
Gain of Electrons is Reduction
Anytime oxidation occurs, reduction also occurs, so these processes are called halfreactions, which chemist represent symbolically in the form similar to a chemical
reaction.
o When magnesium oxidizes, the half reaction is Mg  Mg2+ + 2e The metal atom serves as the “reactant” and the metal ion and electrons
serve as the “products”.
o When silver reduces, the half reaction is Ag+ + 1e-  Ag
 The metal ion and electron serve as the “reactants” and the metal atom
serves as the “product”
To make a complete redox reaction, combine the oxidation half-reaction with the
reduction half-reaction. In order for this to occur, the total number of electrons on either
side of the arrow have to equal when the half-reactions are combined
o For the magnesium and silver examples, we need to multiply the entire reduction
half reaction by 2, so that there are 2 electrons on both sides when we combine
the half-reactions:
2 (Ag+ + 1e-  Ag) = 2Ag+ + 2e-  2Ag
Now add the half-reactions together and “cancel out” the electrons:
Oxidation half-reaction
Mg
 Mg2+ + 2e+
Reduction half-reaction
2Ag + 2e  2Ag______
Complete redox reaction
Mg + 2Ag+  2Ag + Mg2+
While complete redox reactions can be written in two directions, only one direction will
actually occur in nature according to the activity series.
o The most reactive metal will be the one that oxidizes
o The least reactive metal will be the one that reduces
o Mg + 2Ag+  2Ag + Mg2+ will occur in nature because Mg is more reactive
o 2Ag + Mg2+  Mg + 2Ag+ will not occur in nature because Mg is more reactive
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