METALS AND CHEMISTRY OF METALS
‘Metal’ is
derived from the Greek word ‘metallon’, which directly translates to ‘mine’ or
‘quarry’. Generally, metals are usually defined as solid materials (element,
compound or alloy) that are typically hard, shiny and possess features such as
the ability to conduct electricity and heat, etc. the meaning of the term
‘metal’ differs for various communities. For instance, astronomers use the
blanket term ‘metal’ for convenience to describe all elements other than
hydrogen and helium (the main components of stars , which in turn comprise most
of the visible matter in the universe) collectively. Therefore, in astronomy
and physical cosmology, the metallicity of an object is the proportion of its
matter made up of chemical elements other than hydrogen and helium. In
addition, many elements and compounds that are not normally classified as
metals become metallic under high pressure, these are known as metallic
allotropes of non-metals. Many metals such as aluminium foil have a shiny
silvery appearance.
Generally, all
the elements in the periodic table can be classified either as metals or
non-metals, based on their properties (physical and chemical). Although there
is no clear-cut division between the two groups, metals show certain
characteristics which distinguish them from non-metals. In this case, the
metallic bracket covers semi-metals (metalloids).
Most elements are
metals and currently, 91 out of the 118 total elements in the periodic table
are classified as metals. On the periodic table, metals are separated from
non-metals by a zig-zag line stepping through carbon, phosphorus, selenium,
iodine and radon. These elements and those to the right of them are non-metals.
Elements just to the left of the line may either be metals or metalloids
(semi-metals). Metalloids are elements or substances that possess some metallic
properties; that is, they are intermediate between metals and non-metals. Examples
of such metals (metalloids) include boron and silicon.
STRUCTURE AND BONDING OF METALS
Some elementary
features used to distinguish between metals and non-metals are their physical
and chemical properties as well as their structure and bonding. The two former
will be discussed later.
Structure may
generally be defined as the shape, pattern or lattice arrangement of metals. In
fact, the lattice arrangement which depends on the type of bonding between
atoms or molecules is directly responsible for the structural framework of any
element, compound or more generally, substance.
The atoms of
metallic substances are closely positioned to the neighbouring atoms in one of
two common arrangements. The first arrangement is known as body centred cubic.
In this arrangement, each atom is positioned at the centre of eight others. The
other is known as face centred cubic. Here, each atom is positioned in the
centre of the six faces of the cube. The on-going arrangements of atoms in
these structures form a crystal. Some metals adopt both structures depending on
the temperature. Other types of packing structures are primitive cubic
structure, the hexagonal closest packing (hcp), body centred tetragonal, etc.
Find attached, the structures of fcc and bcc
At room
temperature, elements like Li, Na, K, Rb, Ba, V, Cr and Fe have the bcc
structure. Silver exhibits fcc structure. Atoms of metals readily lose their
outermost shell electron(s), resulting in a free flowing cloud of electrons
within their otherwise solid arrangement. This provides the ability of metallic
substances to easily transmit heat and electricity. While this flow of
electrons occur, the solid characteristic of the metal is produced by
electrostatic interactions between its
atom and electron cloud. This type of bond is called a metallic bond.
METALS IN NATURE
As loosely stated
previously, three quarters of all known chemical elements are metals. The most
abundant varieties in the earth’s crust are aluminium, iron, calcium, sodium,
potassium and magnesium. The vast majority of metals found in ores (mineral
bearing substances), but, a few such as copper, gold, platinum and silver
frequently occur in the free state because they do not readily react with other
elements.
Metals are
usually crystalline solids. In most cases, they have a relatively simple
crystal structure distinguished by a close packing of atoms and a high degree
of symmetry. Typically, the atoms of metals contain less than half the full
complement of electrons in their outermost shell. Because of these
characteristic, metals tend not to form compounds with each other. They .do
however combine more readily with non-metals (example, oxygen and sulphur),
which generally have more than half the maximum number of valence electrons.
Metals differ widely in their chemical reactivity. The most reactive include
lithium, potassium, radium; whereas those of low reactivity are gold, silver,
palladium and platinum. We shall make more reference to the nature of metals
when we discuss the extraction of metals.
PROPERTIES/CHARACTERISTICS OF METALS
Properties of
metals simply refer to the things that are common to metals. It simply entails
certain factors or conditions which metals share. Properties or characteristics
of metals can be sub-divided into physical and chemical properties.
Physical properties
Physical
properties as the name implies refers to the properties of metals that can be
easily observed by the mere human eye. They can further be explained as
properties that are dependent on:
I.
The arrangement of their atoms or
molecules in crystal lattices when in the solid state; and
II.
The bond that binds the atoms of
molecules in the solid, liquid or gaseous state.
Most metals are
solids at room temperature and exist as crystal lattices in which their atoms
are held together by strong metallic bonds. Thus, metals have the following
physical properties:
i.
High melting and boiling points:
Metals generally usually have high melting and boiling point. This means that
they must attain a very high temperature before boiling or melting can occur in
the case of liquids or solids respectively. It is noteworthy that helium has
the lowest melting point of about -2720c and tungsten has the
highest melting point of about 35000C, though it is sometimes argued
to be carbon. It is also noteworthy that mercury, sodium and potassium which
are all metals have relatively low melting point of -390C, 970C
and 630C respectively.
ii.
Characteristic lustre: This property
of metals means that most of them are shiny solids at room temperature (except
mercury which is a shiny liquid metal/element)
iii.
Malleability: This may simply be
defined as the ability of solid metals to be hammered into sheets.
iv.
Ductility: This can be defined as the
ability of metals to be drawn into thin wires under stress without cleaving.
This can also be referred to as plastic deformation due to atomic
rearrangement. This atomic rearrangement is due to an applied force/work. This
force in turn can be tensile, compressive, shear or torsion.
v.
Sonorous: This means that they give
off sound notes when hit.
vi.
Hard but not brittle: This means that
whilst they are very strong/hard materials, they are not fragile or brittle
(breakable)
vii.
They have great tensile strength: This
simply means that by virtue of their strong/hard nature, they can withstand
relatively high tension.
viii.
Relatively high density: This means that they
are usually heavy for their size. Although most metals have higher densities
than most non-metals, there is wide variations in their densities, for
instance, lithium has the least density, while osmium has the highest density.
ix.
Good conductors of heat and
electricity: The electrical and thermal (heat) conductivities of metals
originate from the fact that their outer electrons are delocalized, that is;
they easily lose their valence electrons. This situation can be visualized by
seeing the atomic structure of a metal as a collection of atoms embedded in a
sea of highly mobile electrons. The electrical conductivity, as well as the
electron’s contribution to the heat capacity and heat conductivity of metals
can be calculated from the free electron model, which does not take into
account the detailed structure of the ion lattice.
x.
Opaque: This can simply be referred to
as the ability of metals not to allow the passage of light (not radioactive)
through them.
CHEMICAL PROPERTIES OF METALS
Chemical
properties of elements generally are dependent on the number of valence
electrons present in their atomic structure. An important property that is
determined by valence electron is the tendency of the atoms of an element to
ionize. This ionization pattern is used to define metals and non-metals. Based
on its chemical properties, metals can be defined as an element whose atoms
ionize by electron loss while a non-metal is an element whose atoms ionize by
electron gain.
Metals are
usually inclined to form cations (positive net charge) due to electron loss,
while most non-metals usually form anions (negative net charge) due to electron
gain. Hydrogen, which is usually considered as a non-metal, is the only exception
to this as it is usually an electron donor.
Some
characteristic chemical properties of metals are as follows:
i.
Ionization behaviour:
Based on the fact
that metallic atoms have few valence electrons, they have a great tendency to
ionize and form positive ions by losing electrons, that is, they are
electropositive, e.g.
Na
Na+
(univalent) + e-
Zn Zn2+(divalent) + 2e-
Al Al3+(trivalent)
+ 3e-
By this virtue of
metals, they majorly react or form compounds with non-metals (negative ions or
radicals). Some of these radicals include:
S
+ 2e- S2-(divalent)
Cl2
+ 2e- 2Cl-(univalent)
These negative
ions enter into chemical combination with positive ions or groups to form
electrovalent compounds. Those atoms with four or five valence electrons share
electrons during chemical reactions to form covalent compounds.
ii.
Reduction and oxidizing agents:
Metals are
reducing agents by definition because they tend to lose/donate their electrons
readily during chemical reactions. On the other hand, non-metals are oxidizing
agents because they accept electrons readily during chemical reactions. Thus,
in the formation of sodium, it acts as a reducing agent while oxygen (non-metal)
acts as an oxidizing agent.
2Na(s)+1/2O2(g) Na2+O2-(s)
iii.
Reaction with acids to liberate or
displace hydrogen:
Metals which are
usually more electropositive than hydrogen readily displace hydrogen ion h+
from an acid. Non-metals cannot displace the hydrogen from acids. Most
generally, metals, on reaction with acids liberate hydrogen whilst forming a
salt; e.g.:
Zn(s)+2H+(aq)+2Cl(aq)- Zn2+(Cl-)2(aq)+H2(g)
Zn(g)+H2+SO42- Zn22++SO42-(aq)+H2(g)
Also, it is
noteworthy that it is noteworthy that in some special cases, for instance, when
copper reacts with dilute tetraoxosulphate(vi) acid, water is being produced (H2O);
while Cu2+ and SO42- will remain in
solution as the salt CuSO4
iv.
Reaction with oxygen:
Metals are
usually inclined to form cations through electron loss, reacting with oxygen in
the air to form oxides over various time scales (for instance, iron rusts over
years, while potassium burns in seconds). Most metals react with oxygen to form
basic oxides which are usually ionic compounds. Soluble basic oxides (with the
ability to dissolve in H2O) form alkalis, e.g.;
Ca(s)+1/2O2(g) Ca2+O2-(s)(Calcium
oxide)
Ca2+O2-(s)+H2O(l) Ca2+(OH-)2(aq)
Some metals like
aluminium and zinc however form amphoteric oxides. Amphoteric oxides are oxides
that can act either as an acid or base. They are usually of medium
electronegativity.
Example:
4Al
+ 3O2 2Al2O3 (Aluminium
oxide)
Other examples of
oxidizing of metals include:
4Na
+ O2 2Na2O (Sodium oxide)
The transition
metals (such as iron, copper, zinc and nickel) are slower to oxidize because
they form passivating layer of oxide that protects the interior. Others, like
palladium, platinum and gold, do not react with the atmosphere at all. Some
metals form a barrier layer of oxide on their surface which cannot be
penetrated by further oxygen molecules and thus retain their shiny appearance
and good conductivity for many decades (like aluminium, magnesium, some steels
and titanium). The oxides of metals are generally basic, as opposed to those of
non-metals, which are acidic. Blatant exceptions are largely oxides with very
high oxidation states, such as CrO3, Mn2O7 and
OsO4, which have strictly acidic reactions.
Painting,
anodizing or plating metals are good ways of preventing corrosion. However, a
more reactive metal in the electrochemical series must be chosen for coating,
especially when chipping of the coating is expected. Water and the two metals
form an electrochemical cell, and if the coating is less reactive than the
coatee, the coating actually promotes corrosion.
Non-metallic
oxides are covalent compounds formed by electron sharing. The non-metallic
oxides in which the non-metal exercises its maximum valency dissolve in water
to give an acidic solution, due to the formation of hydrogen ions, H+.
Examples include carbon(iv)oxide[CO3] and sulphur(vi)oxide[SO3]
CO2(g)
+ H2O(l) H2CO3(aq) 2H+(aq)
+ CO32+(aq)
Other
non-metallic oxides are either acidic, e.g. sulphur(iv)oxide[SO2],
or neutral, e.g. carbon(ii)oxide[CO]
v.
Few metals form compounds (hydrides)
with hydrogen:
When hydrogen
reacts with certain metals covalently (i.e. sharing of electrons), they form
compounds referred to as hydrides. But few very reactive metals force hydrogen
to accept electron to form their hydrides, which are salt-like solids.
Na(s)
+ 1/2H2 NaH
The hydride
reacts with cold water to liberate hydrogen
H-
+ H2O OH-(aq) + H2(g)
When in molten
form, these hydride of metals act as electrolytes, liberating hydrogen at the
anode.
vi.
Reactions of metals with water:
The reactions of
water with metals vary for a range of metals. Some have violent reactions,
while others are somewhat mild. Some examples of these reactions include:
-
Place a piece of sodium metal on the
surface of tap water in a trough; it darts about on the surface of the water.
But, when a piece of paper interfaces the water and sodium piece, its movement
is restricted. Thus, the sodium metal melts into a silvery white ball and then
bursts into a yellow flame and finally explodes with a cracking sound. The sodium
metal combines vigorously with water, liberating hydrogen and forming the
hydroxide and heat is given off.
2Na + 2H2O 2NaOH + H2
-
Repeating the above experiment with
potassium (K), the reaction is more rapid and violent than that of sodium. The
heat produced in this reaction is enough to ignite the hydrogen liberated which
then burns with a lilac flame, and the molten potassium ball explodes. This
reaction is dangerous.
-
If we put calcium metal on top of
water in a beaker, it sinks to the bottom of the beaker, producing
effervescence and liberating bubbles of gas, thus, forming calcium hydroxide
which is slightly soluble in water.
Ca + H2O Ca(OH)2 + H2
Reaction of water (steam) on Mg, Al,
Zn, Fe, Pb and Cu:
Magnesium, aluminium,
zinc, iron, lead and copper on passing over steam in their red hot state, their
oxides are formed and hydrogen is liberated in each case.
COMPARISON OF THE PROPERTIES OF DIFFERENT METALS
Metals differ in
their ability to lose their valence electron(s) in their atoms to form positive
ions. As a result, they can be arranged in a series according to their
comparative tendencies to give up their valence electron(s) (i.e. according to
their electropositivity). It is noteworthy that metals are reduction agents and
as such are oxidized (i.e. lose electrons). The series of their arrangement by
this virtue is called the replacement or electromotive or electrochemical
series. Most textbooks refer to it generally as the electrochemical series. The
relative position of a metal in this series indicates the chemical activity of
the metal as well as the chemical properties of its compounds. A similar series
is the activity series of metals.
Metals towards
the beginning of the series, like caesium and lithium, are more readily
oxidized than those towards the end, like silver and gold. In general, a metal
will replace any other metal or hydrogen in a compound that it precedes in the
series and under ordinary circumstances; it will be replaced by any metal or
hydrogen that it follows.
Find below a
comprehensive tabulation, critically examining elements in the electrochemical
series as well as their distinguishing properties
Find attached, the activity series of metals, to help us understand the electrochemical series in details
EXTRACTION OF METALS
Metals are often extracted from the earth’s
crust by means of mining, resulting in ores that are relatively rich sources of
the requisite elements. Ore is located by prospecting techniques, followed by
the exploration and examination of deposits. Mineral sources are generally
divided into surface mines, which are mined by excavation using heavy
equipment, and subsurface mines.
More than 80 of
the known elements are metals. They are widely distributed in the earth’s
crust. The ore form in which metals exist in nature is related to its
reactivity. The most reactive metals e.g. sodium and potassium, are found as
chlorides or trioxocarbonate(iv)s, which are very stable compounds; the
moderately reactive metals e.g. zinc and lead, are found as oxides or
sulphides; while the least reactive metals like gold and silver are found in
the uncombined forms of more or less definite chemical composition known as
minerals. Often, these minerals are found mixed with earthly materials called
ores. Generally, ores are usually concentrated and changed to oxides before
extraction.
Once the ore is mined, the metals must be extracted,
usually by chemical or electrolytic reduction. Pyrometallurgy uses high
temperatures to convert ore into raw metals, while hydrometallurgy employs
aqueous chemistry for the same purpose. The methods used depend on the metal
and their contaminants. When a metal ore is an ionic compound of that metal and
a non-metal, the ore must usually be smelted (heated with a reducing agent) to
extract the pure metal. Many common metals, such as iron, are smelted using
carbon as a reducing agent. Some metals, such as aluminium and sodium, have no
commercially practical reducing agent, and are extracted using electrolysis
instead. Sulfide and trioxocarbonate(iv) ores are not reduced directly to the
metal but are roasted in air to convert them to oxides.
ALLOYS
An alloy is a mixture of two or more elements in which
the main component is a metal. Most pure metals are either too soft, brittle or
chemically reactive for practical use. Combining different ratios of metals as
alloys modifies the properties of pure metals to produce desirable
characteristics. The aim of making alloys is generally to make them less
brittle, harder, resistant to corrosion, or have a more desirable color and
luster. Of all the metallic alloys in use today, the alloys of iron (steel,
stainless steel, cast iron, tool steel, alloy steel) make up the largest
proportion both by quantity and commercial value. Iron alloyed with various
proportions of carbon gives low, mid and high carbon steels, with increasing
carbon levels reducing ductility and toughness. The addition of silicon will
produce cast irons, while the addition of chromium, nickel and molybdenum to
carbon steels (more than 10%) results in stainless steels.
Other significant metallic alloys are those of
aluminium, titanium, copper and magnesium. Copper alloys have been known since
prehistory (bronze gave the Bronze Age its name) and have many applications
today, most importantly in electrical wiring. The alloys of the other three
metals have been developed relatively recently; due to their chemical
reactivity they require electrolytic extraction processes. The alloys of
aluminium, titanium and magnesium are valued for their high strength-to-weight
ratios; magnesium can also provide electromagnetic shielding. These materials
are ideal for situations where high strength-to-weight ratio is more important
than material cost, such as in aerospace and some automotive applications.
Alloys specially designed for highly demanding
applications, such as jet engines, may contain more than ten elements.
Examples of
alloys include:
Bronze = Copper +
Tin
Brass = Copper +
Zinc
METALLURGY
Metallurgy is the branch or domain of material science
that studies the physical and chemical behaviour of metallic elements, their intermetallic
compounds and their mixtures (i.e. alloys). Metallurgy can further be explained
by examining the application of metals generally.
Applications of metals
Some metals and metal alloys possess high structural
strength per unit mass, making them useful materials for carrying large loads
or resisting impact damage. Metal alloys can be engineered to have high
resistance to shear, torque and deformation. However the same metal can also be
vulnerable to fatigue damage through repeated use or from sudden stress failure
when a load capacity is exceeded. The strength and resilience of metals has led
to their frequent use in high-rise building and bridge construction, as well as
most vehicles, many appliances, tools, pipes, non-illuminated signs and railroad
tracks.
The two most commonly used structural metals, iron and
aluminium, are also the most abundant metals in the Earth's crust.
Metals are good conductors, making them valuable in
electrical appliances and for carrying an electric current over a distance with
little energy lost. Electrical power grids rely on metal cables to distribute
electricity. Home electrical systems, for the most part, are wired with copper
wire for its good conducting properties.
The thermal conductivity of metal is useful for containers
to heat materials over a flame. Metal is also used for heat sinks to protect
sensitive equipment from overheating.
The high reflectivity of some metals is important in
the construction of mirrors, including precision astronomical instruments. This
last property can also make metallic jewelry aesthetically appealing.
Some metals have specialized uses; radioactive metals
such as uranium and plutonium are used in nuclear power plants to produce
energy via nuclear fission. Mercury is a liquid at room temperature and is used
in switches to complete a circuit when it flows over the switch contacts. Shape
memory alloy is used for applications such as pipes, fasteners and vascular
stents (used in hospitals).
CLASSIFICATION OF METALS BASED ON THEIR GROUP ON THE PERIODIC
TABLE
Basically, the
metals on the periodic table can be classified into five categories, viz;
i.
Alkali Metals:
These are the
group 1 metals. All the elements in this group are metals except hydrogen. They
include lithium, sodium, potassium, rubidium, etc. They have one electron in
their valence or outermost shell, thus, they are very electropositive and are
very good reducing agents as they are easily oxidized.
ii.
Alkaline Earth Metals:
These are the
group 2 metals. It is noteworthy that all the elements here are metals. They
include magnesium, calcium, strontium, barium and radium. They are less
reactive than the alkali metals as they have 2 electrons in their valence
shell.
iii.
Group 3 Metals:
All the elements
in this group are metals, except boron which is a metalloid. Aluminium is a
familiar group 3 metal. It has 3 valence electrons in each atom and tends to
form covalent compounds than ionic/electrovalent ones.
iv.
Group 4 Elements:
Group 4 elements
change from non-metals (carbon) to metals(tin and lead) as we go down the
group. Since the atoms of these elements have 4 valence electrons, they tend to
form covalent compounds. Tin and lead form compounds in which they exist in the
+2 and +4 oxidation states. The stable compounds of lead are the
lead(ii)compounds which are mainly ionic.
v.
Transition Elements:
The transition
elements are all metals of economic importance. The first transition series
(scandium to zinc) are particularly important industrially. The transition
metals are found in the d-block of the periodic table between groups 2 and 3.
They occupy three rows, with ten elements in each row. The term “transition
elements” refer only to an element which has partially filled d-orbitals. For
our level, we will make reference to only the first transition series. The
first transition series contain the following metals: scandium, titanium,
vanadium, chromium, manganese, iron, cobalt, nickel, copper and zinc. Each has
a partially filled 3-d shell and a 4s shell containing 1 or 2 electrons. With
the exception of zinc and copper which have completely filled 3-d orbitals and
scandium which has no electron in its 3-d orbital.
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