The Chemical Activity Series

An Activity Series is a list of substances ranked in order of relative reactivity. For example, magnesium metal can knock hydrogen ions out of solution, so it is considered more reactive than elemental hydrogen:

Mg(s) + 2 H⁺(aq) = H₂(g) + Mg²⁺(aq)

Zinc can also displace hydrogen ions from solution:

Zn(s) + 2 H⁺(aq) = H₂(g) + Zn²⁺(aq)

so zinc is also more active than hydrogen. But magnesium metal can remove zinc ions from solution:

Mg(s) + Zn²⁺(aq) = Zn(s) + Mg²⁺(aq)

The reaction goes nearly to completion. Magnesium is more active than zinc, and the activity series including these elements would be Mg > Zn > H. The following activity series built up in a similar way. The most active metals are at the top of the table; the least active are at the bottom. Each metal is able to displace the elements below it from solution (or, using the language of electrochemistry, each metal can reduce the cations of metals below it to their elemental forms).

The metal activity series. Most active (most strongly reducing) metals appear on top, and least active metals appear on the bottom.
displace H2 from water, steam, or acids	Li	2 Li(s) + 2 H2O = 2 LiOH(aq) + H2(g)
	K	2 K(s) + 2 H2O = 2 KOH(aq) + H2(g)
	Ca	Ca(s) + 2 H2O = Ca(OH)2(s) + H2(g)
	Na	2 Na(s) + 2 H2O = 2 NaOH(aq) + H2(g)
displace H2 from steam or acids	Mg	Mg(s) + 2 H2O(g) = Mg(OH)2(s) + H2(g)
	Al	2 Al(s) + 6 H2O(g) = 2 Al(OH)3(s) + 3 H2(g)
	Mn	Mn(s) + 2 H2O(g) = Mn(OH)2(s) + H2(g)
	Zn	Zn(s) + 2 H2O(g) = Zn(OH)2(s) + H2(g)
	Fe	Fe(s) + 2 H2O(g) = Fe(OH)2(s) + H2(g)
displace H2 from acids only	Ni	Ni(s) + 2 H+(aq) = Ni2+(aq) + H2(g)
	Sn	Sn(s) + 2 H+(aq) = Sn2+(aq) + H2(g)
	Pb	Pb(s) + 2 H+(aq) = Pb2+(aq) + H2(g)
	H2
can't displace H2	Cu
	Ag
	Pt
	Au

The activity series is a useful guide for predicting the products of metal displacement reactions. For example, placing a strip of zinc metal in a copper(II) sulfate solution will produce metallic copper and zinc sulfate, since zinc is above copper on the series. A strip of copper placed into a zinc sulfate solution will not produce an appreciable reaction, because copper is below zinc on the series and can't displace zinc ions from solution.

The series works well as long as the reactions being predicted occur at room temperature and in aqueous solution. It isn't difficult to find reactions that are at odds with the metal and nonmetal activity series under other conditions. There are other complications too. For example, aluminum would be expected to displace hydrogen from steam, but in fact it won't unless the aluminum oxide film on its surface is scrubbed off. Copper can't displace hydrogen from acids, but it does react with acids like nitric and sulfuric because they can act as oxidizing agents.

It might be expected that metals with lower ionization energies and lower electronegativities would be more active, since they would be expected to more easily lose electrons in a displacement reaction. But while ionization energy and electronegativity do affect a metal's ranking in the series, other factors have a strong and complex influence on relative activity , obscuring the relationship.

Activity series can be devised for nonmetals as well. Since nonmetallic elements tend to accept electrons in redox reactions, the nonmetal activity series is arranged so that the most powerful oxidizing agents are considered most active (whereas in the metal series, the most powerful reducing agents are the most active):

The nonmetal activity series. Most active (most strongly oxidizing) nonmetals appear on top, and least active nonmetals appear on the bottom.
F2	strongest oxidizing agent
Cl2
O2
Br2
I2
S
red P	weakest oxidizing agent

For example, the series predicts that Cl₂ will displace Br^- and I^- from solution, because Cl₂appears above Br₂ and I₂:

Cl₂(g) + 2 Br^-(aq) = 2 Cl^-(aq) + Br₂
Cl₂(g) + 2 I^-(aq) = 2 Cl^-(aq) + I₂(s)
Br₂ + 2 Cl^-(aq) = no reaction
I₂(s) + 2 Cl^-(aq) = no reaction

Metal Purification

Metal Production-Purification

The term metal production refers to all of the processes involved in the conversion of a raw material, such as a metallic ore, to a final form in which the metal can be used for some commercial or industrial purpose. In some instances, metal production involves relatively few steps since the metal already occurs in an elemental form in nature. Such is the case with gold, silver, platinum, and other so-called noble metals. These metals normally occur in nature uncombined with other elements and can therefore be put to some commercial use with comparatively little additional treatment.

In the majority of cases, however, metals occur in nature as compounds, such as the oxide or the sulfide, and must first be converted to their elemental state. They may then be treated in a wide variety of ways in order to make them usable for specific practical applications.

Metal Production - Purification

In most cases, metals and their ores occur in the ground as part of complex mixtures that also contain rocks, sand, clay, silt and other impurities. The first step in producing the metal for commercial use, therefore, is to separate the ore from waste materials with which it occurs. The term ore is used to describe a compound of a metal that contains enough of that metal to make it economically.

Metal Reduction

Metal Reduction

Reduction reactions are used in the process of producing metals from ores. The reduction of metals was originally understood to be the reactions used to obtain metals from their oxides by using substances having greater attraction for oxygen than the metal. The simplest example is the production of iron from its protoxide:

FeO + C = Fe + Co

This reaction takes place in blast furnaces.

The possibility of reducing metals is determined by the free energy of the reaction

MeO + R = Me + RO

Where MeO is the metal oxide and R is the reducing agent. If in this reaction (at constant temperature and pressure) the total free energy for Me and RO is less than for MeO and R, the process proceeds from left to right, with formation of metal. The process is facilitated, or made easier, if the final product, which is metal, is present in the dissolved state (solid or liquid), since dissolution is accompanied by a decrease in free energy. This explains why, in the reduction of metals, some particularly stable oxides yield the corresponding alloys as end products. Thus, the reduction of metals requires the presence of a definite thermodynamic stimulus. In addition, great importance also attaches to the kinetic conditions of reduction, which are determined by crystallochemical changes (in the case of solid oxides), the mechanism of the chemical reactions at the phase boundaries, and the mass-transfer conditions for the reagents—for example, diffusion.

In a more general chemical sense, the reduction of metals consists of the addition of electrons to an atom or group of atoms. Therefore, reduction of metals also includes processes in which metals are obtained at a cathode by electrolysis of salt melts or solutions—for example, in the case of copper:

Cu+++ 2e = Cu

where e is an electron.

In technology the most important examples of such processes are the production of aluminum by electrolysis of alumina from a melt and the production of copper from aqueous solutions of CuSO4. In nonferrous metallurgy, reduction of metals is carried out in the production of metals from sulfides, chlorides, and other compounds. Since the electrons given off by the reducing agent are necessary for reduction, reduction processes are inseparably connected with oxidizing processes.