Introduction:

Elements are generally broadly categorized into metals and non-metals however there is no clear line of division between the two groups. It is acknowledged that elements that easily lose electrons show metallic characteristics. The characteristics represented by metals employed to differentiate them from non-metals. Metals, in common, encompass a quite different appearance, the metallic lustre or the manner in which the light is reflected on its surface.

The other way in which metals are dissimilar from non-metals is in the electrical and thermal conductivities. Metals are usually good conductors whereas non-metals are poor or non conductors.

The periodic categorization of elements illustrates that in the Periodic Table, metallic character reduces and non-metallic character increases from left to right across any period. This is due to the reason of decrease in the atomic size and increase in nuclear charge across a period prevents the elements from losing electrons readily. In Period 3 for illustration, metallic character decreases from sodium to magnesium to aluminium whereas the remaining elements silicon, phosphorus, sulphur, chlorine and argon are non-metals. The metallic character of elements increases down the group of the Periodic Table as the tendency to lose electrons increases as the atomic size increases down the group. For illustration, in Group IA the order of metallic character is He < Mg < Ca < Sr < Ba.

Metals are of great practical significance to man. The growth of automobiles, jet motors, airplanes, skyscrapers and various other objects characteristic of modem civilization and general progress in technology is an outcome of progress in the science of metals.

Physical Properties of Metals:

The physical properties of metals can be illustrated by employing the nature of bond (that is, metallic bonding) between the atoms. A few of these characteristics are as follows:

1) Metals encompass high electrical and thermal conductivity that is, good conductors, due to the mobile nature of their valence electrons.

2) Metals comprise a lustrous shiny appearance. All the metals apart from copper and gold are silvery in color. The high lustre of metals is due to the reason of the valence electrons which can absorb energy from light and then re-emit it when the electron drops back from its excited state to its original energy level.

3) Metals are malleable that is, can be hammered to sheets example: gold leaves; and are ductile that is, can be drawn into wire. Metals exhibit such features because of the quite uniform attraction between the mobile valence electrons and the ions formed. Metals can thus be stretched or drawn out into different shapes as the ions and electrons can move to other positions without breaking up the metallic structure.

4) Metals are hard and strong because of the binding action of the electrons which hold the atoms in a close-packed lattice structure. Some metals, however, do not show this property, e.g. sodium and potassium are light, soft metals and mercury is a liquid at room temperature.

Non-metals usually exist as covalent molecules held altogether by weak Vander Waal's forces and they don't exhibit such metallic characteristics. Carbon, a non-metal, is an exemption as it exists as diamond which is a hard solid and as graphite which is a good conductor. Several elements such as silicon and boron represent some metallic characteristics and are termed to as metalloids - intermediate among metals and non-metals.

Chemical Properties of Metals:

The chemical properties of any element are regulated by the number of valence electrons present in their atomic structure. The tendency to lose electrons based on the ionization energy. As it is easier to eliminate an electron from a large atom than from a small one, metallic character increases as we descend the groups and decreases from left to right of the period, of the periodic table. This ionization pattern is employed to state metals and non-metals.

1) Ionization ability:

Metals are electropositive due to their tendency to lose electrons whenever supplied by energy

M → M⁺ + e^-

The stronger the tendency, the more electropositive and more metallic an element is. Metals are usually electron donors and non-metals are electron acceptors. Metallic atoms encompass some valence electrons, between one to three.

Na → Na⁺ + e^- (Univalent)

Ca → Ca²⁺ + 2e^- (divalent)

Al → Al³⁺ + 3e^- (trivalent)

Whenever metals ionize this manner, the positive ions formed chemically combine with negative ions, made by non-metals whenever they gain electrons, to give ionic or electrovalent compounds.

2) Reducing ability:

Electron donors are the reducing agents. Metals are reducing agents as they tend to donate their electrons readily throughout chemical reaction. The donated electrons are accepted through oxidizing agents generally non-metals.

2Ca           +          O₂            →           2Ca²⁺O^2-

Reducing          Oxidizing

agent                   agent

3) Reaction with acids:

Metals that are more electropositive than hydrogen react by mineral acids example: HCl, H₂SO₄, HNO₃ and transfer the hydrogen ion, H⁺ of the acid.

Zn + 2HCl → ZnCl₂ + H₂

The reaction can be simplified to the oxidation and reduction (that is, Redox) processes.

Zn → Zn²⁺ + 2e^- oxidation

2H⁺ + 2e^- → 2H → 2H₂ reduction

4) Formation of oxides:

Some of the metals react with oxygen to form basic oxides that are ionic compounds. Basic oxides that are water soluble form alkalis example:

4Na + O₂ → 2Na₂O(s)

Na₂O(s) + H₂O → 2Na⁺OH^- (aq)

The rate of reaction based on the reactivity of the metals. Sodium quickly tarnishes in air, magnesium and calcium tarnish in moist air whereas iron, aluminium, zinc and tin are oxidized in moist air to form the oxides.

4Fe + 3O₂ → 2Fe₂O₃ (rusting method)

Non-metallic oxides are the covalent compounds that are acidic example: CO₂, SO₃ or neutral example: CO and oxides.

Occurrence of Metals in Nature:

More than eighty (80) of the identified elements are metals. Metals are broadly distributed in the earth's crust either as compounds or in the free metallic state. The state in which metals exist in nature based on the reactivity of the metals.

Chemically active metals like sodium, potassium, calcium and magnesium are generally found combined with other elements. They exist as chlorides, trioxocarbonate (IV), that are stable compounds. Base metals like zinc, lead and copper, are merely moderately reactive and therefore exist in nature as oxides or sulphides. The so-called noble metals like silver and gold are unreactive and thus exist in the free or uncombined state.

The majority of metals are found as minerals in combined form having almost fixed composition and mixed by earthy materials example: sand. These natural forms of existence of metals are termed as ores. Example: Bauxite and cryolite for aluminium; limestone and marble for calcium.

Chemically active metals are mostly employed as alloys. The base metals are found in on a daily basis use, due to their cheapness example: iron, tin, zinc; whereas the noble metals are mostly utilized due to their pleasing appearance and as monetary standard due to their high cost.

Electrochemical or Activity Series:

Metals usually have few valence electrons that they readily donate or give up to form positive ions. The tendency to donate electrons differs from metal to metal and as an outcome metals can be arranged in sequence according to their comparative tendencies to give up their valence electrons (that is, their electropositivity). This is the electrochemical or activity series.

The relative position of a metal in the series points out that the chemical reactivity of the metal and the chemical properties of its compounds. The most active metals form very stable compounds and their ores are chlorides or trioxocarbonates (IV) example: potassium and calcium. The base metals form compounds of intermediate stability and their ores are oxides or sulphides or at times trioxocarbonates (IV) example: iron and zinc. The least active metals, the noble metals generally take place in Free State.

Table shows the activity series of metals and their properties.

Extraction of Metals:

The procedure of extraction of a metal from its ores based on the ease by which the compounds of the metal can be reduced to the metal. Reducing agents or electron donors are employed to supply the electrons needed to transform the metallic ion to the metal. Compounds of chemically active metals can't be reduced through common reducing agents like carbon and hydrogen; instead they are reduced via electrolysis. The ores of less active base metals are usually first decomposed to their oxides before reduction with carbon.

Mineral ores are generally found related by earthy materials and thus a preliminary preparation is often needed before extracting, the metal. The ores are preliminarily made by:

1) Concentration: The ore might be separated from the earthy material through washing in a stream of water to leave behind the heavier ores example: tin ore or shaken by oil and water, and the one floated example: zinc ore or magnetic separation example: copper ore.

2) Roasting: Oxide ores are simpler to handle throughout metal extraction than sulphides or trioxocarbonates (IV) ores. All the non oxide ores are generally transformed to the oxide by roasting it in air.

Electrolytic reduction:

The most electropositive metals at the top of the activity series example: K, Na Ca and Mg are obtained from their ores through electrolytic reduction. The ores are generally stable chlorides and trioxocarbonate (IV) and their electrolysis is a very powerful Redox method. The ore is utilized as the electrolyte and the cathode acts as the reducing half-cell that supplies electrons to transform the metallic ion to the free metal.

Electrolytic methods are costly to install and maintain. They are utilized only whenever the thermal and chemical reduction of the ores are hard to accomplish. 

Chemical reduction:

The less reactive base metals, example: lead, zinc, iron and tin are extracted from their ores via chemical reduction. This reduction procedure can be achieved via:

i) Reducing oxides of metals with coke (carbon) or carbon (II) oxide

ZnO + C → Zn + CO

Fe₂O₃ + 3C → 2Fe + 3CO

ii) Oxidation of sulphide ores to the oxide which is then reduced via coke

2PbS + 3O₂ → 2PbO + 2SO₂  

2PbO + C → 2Pb + CO

In both the methods, the carbon joins with the oxide ions to discharge the electrons needed for the reduction of the metal ions. 

C + O^2- → CO + 2e^-

Zn²⁺ + 2e^- → Zn

Fe²⁺ + 2e^- → Fe

Pb²⁺ + 2e^- → Pb

Thermal reduction:

Non-oxide ores are reduced through heating the ore in air to give the free metals. Mercury is obtained through heating mercury (II) sulphide in air

HgS + O₂ → Hg + SO₂

Table: Activity series of metals and methods of metal extraction

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