Metals And Non-Metals
1.
Properties of Metals and Non-Metals:
METALS:
Electrons during chemical reactions are called
metals. Thus metals are electropositive elements with relatively low ionization
energies. They are characterized by bright luster, hardness, ability to
resonate sound and are excellent conductors of heat and electricity. Metals are
solids under normal conditions except for Mercury With the exception of
hydrogen, all elements that form positive ions by losing.
Physical
Properties of Metals-
In the above depiction of the
periodic table most of the elements are metals. There are various kinds of
metals:
·
Alkaline
earth metals
·
Alkali
Metals
·
Transition
metals
·
Actinides
·
Lanthanides
A. High melting and
boiling point:
Metals have high melting and
boiling point due to their high density and solid state.
B. Lustre shininess:
Metals are shiny. It is due
to his property of metals that they are lustrous, and they reflect the light incident on its surface. Also, metals can be
polished, and this is one of the reasons why metals are used to make jewellery
and desired by women and men alike.
C. High density:
Metals have high density.
Metals are very strong and hard, exceptions being sodium and potassium. They
can be cut with a knife.
D. Good Conductors:
Metals conduct heat and
electricity. It is by virtue of this property of metals that heat, and
electricity can pass through them. Every metals is a good conductor of heat and
electricity.
E.
Non Degradable:
Metals
are non degradable in nature. Metals occur in the solid state. All metals are
solid except with an exception for mercury which is in liquid state in its
natural form.
F. Malleable:
Metals have the ability to withstand hammering
and can be made into thin sheets known as foils. For example, a sugar cube
sized chunk of gold can be pounded into a thin sheet that will cover a football
field.
G.
Ductility:
Metals can be drawn into wires. For example, 100
g of silver can be drawn into a thin wire about 200 meters long.
Chemical Properties of Metals:
A. Reaction with Oxygen:
Almost all metals combine with
oxygen to form metal oxides.
Metal + Oxygen → Metal
oxide
For example, when copper is heated in air, it
combines with oxygen to form copper(II) oxide, a black oxide.
2Cu + O2 → 2CuO
(Copper) ŕ (Copper(II)
oxide)
Similarly, aluminium forms
aluminium oxide.
4Al
+ 3O2 → 2Al2 O3 (Aluminium) ŕ (Aluminium oxide)
Metal oxides are basic in nature.
But some metal oxides, such as aluminium oxide, zinc oxide show both acidic as
well as basic behaviour. Such metal oxides which react with both acids as well
as bases to produce salts and water are known as amphoteric oxides. Aluminium
oxide reacts in the following manner with acids and bases –
Al2O3
+ 6HCl → 2AlCl3 + 3H2O Al2O3 + 2NaOH
→ 2NaAlO2 + H2O (Sodium aluminate)
Most metal oxides are insoluble in water but
some of these dissolve in water to form alkalis. Sodium oxide and potassium
oxide dissolve in water to produce alkalis as follows –
Na2O(s)
+ H2O(l) → 2NaOH(aq) K2O(s) + H2O(l)
→ 2KOH(aq)
B. Reaction with Water:
Metals react with water and
produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water
dissolve in it to further form metal hydroxide. But all metals do not react
with water.
Metal + Water → Metal oxide + Hydrogen
Metal oxide + Water → Metal hydroxide
Metals like potassium and sodium
react violently with cold water. In case of sodium and potassium, the reaction
is so violent and exothermic that the evolved hydrogen immediately catches
fire.
2K(s) + 2H2O(l) → 2KOH(aq) +
H2(g) + heat energy 2Na(s) + 2H2O(l) → 2NaOH(aq) +
H2(g) + heat energy
The reaction of calcium with
water is less violent. The heat evolved is not sufficient for the hydrogen to
catch fire.
Ca(s) + 2H2O(l)
→ Ca(OH)2(aq) + H2(g)
Calcium starts floating because the bubbles of
hydrogen gas formed stick to the surface of the metal. Magnesium does not react
with cold water. It reacts with hot water to form magnesium hydroxide and
hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking
to its surface. Metals like aluminium, iron and zinc do not react either with
cold or hot water. But they react with steam to form the metal oxide and
hydrogen.
2Al(s) + 3H2O(g) → Al2O3(s)
+ 3H2(g) 3Fe(s) + 4H2O(g) → Fe3O4
(s) + 4H2(g)
Metals such as lead, copper,
silver and gold do not react with water at all.
C. Reaction with Acids:
Acids react with most metals to form
a salt and hydrogen gas. As discussed previously, metals that are more active
than acids can undergo a single displacement reaction. For example,
zinc metal reacts with hydrochloric acid producing zinc chloride and hydrogen
gas.
Zn(s)+2HCl(aq)→ZnCl2(aq)+H2(g)
Potassium
reacts with dilute hydrochloric acid to give potassium chloride and hydrogen
gas.
2K+2HCl⟶2KCl+H2
D.
Reaction with Bases:
Bases also react
with certain metals like zinc or aluminum for example to produce hydrogen gas.
For exanple, sodium hydroxide reacts with zinc and water to form sodium zincate
and hydrogen gas.
Zn(s)+2NaOH(aq)+2H2O(l)→Na2Zn(OH)4(aq)+H2(g).
Sodium aluminate and hydrogen gas are
formed when sodium hydroxide reacts with aluminium metal.
2NaOH + 2Al + 2H2O ⇨ 2NaAlO2 + 2H2
NON-METALS:
Elements that tend to gain electrons to form
anions during chemical reactions are called non-metals. These are
electronegative elements with high ionization energies. They are non-lustrous,
brittle and poor conductors of heat and electricity (except graphite).
Non-metals can be gases, liquids or solids.
Physical Properties of
Non-Metals:
A. Melting and Boiling Points:
The melting points of non-metals are generally lower
than metals, but are highly variable.
B.
Luster:
These have no metallic luster and do not reflect
light.
C. Low density:
Non-metals have low density as compare to metals.
Some Non-metals are solid in nature but with low density.
D.
Conduction:
They
are poor conductors of heat and electricity.
E.
Degradable:
Non-metals
are degradable in nature.
F.
Non-Malleable
and Ductile:
Non-metals are very brittle, and cannot be
rolled into wires or pounded into sheets.
G.
Physical
State:
Most of the non-metals exist in two of the
three states of matter at room temperature: gases (oxygen) and solids (carbon).
Only bromine exists as a liquid at room temperature.
Chemical Properties of Non-Metals:
A. Reaction
of non-metals with oxygen: They
react with oxygen to form acidic oxides or neutral oxides.
For example, Carbon forms acidic carbon dioxide on reacting
with oxygen.
C
+ O2
Carbon Carbon
Dioxide
In the same way sulphur reacts with
oxygen of air to form acidic sulphur dioxide.
S
+ O2
Sulphur
Sulphur
Dioxide
B. Reaction
of non-metals with water: They
do not react with water (steam) to evolve hydrogen gas.
C. Reaction
of non-metal with acids: They
do not react with acids because they are negative charged electron.
D. Reaction
of non-metal with salt solution: They
do not react with salt solution but displaces less reactive non-metal from the
salt.
E. Reaction of non-metals
with chlorine:
Non-metals react with
chlorine to form covalent chlorides. For example,
H2
+ Cl2
Hydrogen Chlorine Hydrogen
chloride
P4 + 6Cl2
Phosphorous
Chlorine Phosphorous
trichloride
F.
reaction of non-metals with hydrogen:
Non-metals
react with hydrogen to form covalent Hydrides. For example,
H2
+ S
Hydrogen sulphur Hydrogen
sulphide
N2 + 3H2
Nitrogen
Hydrogen Ammonia
2. Reactivity Series:
Reactivity
Series of Metals:
The order of intensity of reactivity
is known as reactivity series. Reactivity of element decreases on moving from
top to bottom in the given reactivity series.
In the reactivity series, copper, gold, and silver are at the
bottom and hence least reactive. These metals are known as noble metals.
Potassium is at the top of the series and hence most reactive.
Reactivity of some metals are given in descending order
K > Na > Ca > Mg > Al
> Zn > Fe > Pb > Cu
Reaction of metals
with solution of other metal salts:
Reaction of metals with solution of other metal salt is
displacement reaction. In this reaction more reactive metal displace the less
reactive metal from its salt.
Metal A + Salt of
metal B ŕ Salt of metal A +
Metal B
Examples:
Iron displaces copper from copper sulphate solution.
Fe + CuSO4 ŕ FeSO4 + Cu
Similarly, aluminium and zinc displace copper from the solution
of copper sulphate.
2Al + 3CuSO4 ŕ Al2(SO4 )3 +
3Cu
Zn + CuSO4 ŕ ZnSO4 + Cu
In all the above examples, iron, aluminium and zinc are more
reactive than copper. That’s why they displace copper from its salt solution.
Reactivity Series
of Non-Metals:
3. Formation of Ionic
Compounds:
The crystalline
solids formed by neatly packed ions of opposite charge. Ionic compounds are
usually formed when metals react with non-metals.
In other
words, ionic compounds held together by ionic bonds as classed as ionic
compounds. Elements can gain or lose electrons in
order to attain their nearest noble
gas configuration. Formation of ions (either
by gaining or losing electrons) for the completion of octet helps them gain
stability.
Formation of sodium chloride (NaCl):
In sodium chloride; sodium is a metal (alkali metal) and
chlorine is non-metal.
Atomic number of sodium = 11
Electronic configuration of sodium: 2, 8, 1
Number of electrons in outermost orbit = 1
Valence electrons = Electrons in outermost orbit = 1
Atomic number of chlorine = 17
Electronic configuration of chlorine: 2, 8, 7
Electrons in outermost orbit = 7
Therefore, valence electrons = 7
Sodium has one valence electron and chlorine has seven valence
electrons. Sodium requires losing one electron to obtain stable configuration
and chlorine requires gaining one electron in order to obtain stable electronic
configuration. Thus, in order to obtain stable configuration sodium transfers
one electron to chlorine.
After loss of one electron sodium gets one positive charge (+)
and chlorine gets one negative charge after gain of one electron. Sodium
chloride is formed because of transfer of electrons. Thus, ionic bond is formed
between sodium and chlorine. Since, sodium chloride is formed because of ionic
bond, thus it is called ionic compound. In similar way; potassium chloride
(KCl) is formed.
Formation of
Magnesium Chloride (MgCl2):
The atomic number of magnesium is 12
Electronic configuration of magnesium: 2, 8, 2
Number of electrons in outermost orbit = 2
Valence electron = 2
Atomic number of chlorine = 17
Electronic configuration of chlorine: 2, 8, 7
Electrons in outermost orbit = 7
Therefore, valence electrons = 7
Magnesium loses two electrons in order to obtain stable
electronic configuration. Each of the two chlorine atoms gains one electron
lost by magnesium to obtain stable electronic configuration. The bonds so
formed between magnesium and chlorine are ionic bonds and compound (magnesium
chloride) is an ionic compound.
Atomic number of calcium is 20.
Electronic configuration of calcium: 2, 8, 8, 2
Number of electrons in outermost orbit = 2
Valence electron = 2
Valence electrons of chlorine = 7
Formation of Calcium Chloride:
Calcium loses two electrons in order to achieve stable
electronic configuration. Each of the two chlorine atoms on the other hand
gains one electron losing from calcium to get stability. By losing of two
electrons calcium gets two positive charges over it. Each of the chlorine atoms
gets one positive charge over it.
The bonds formed in the calcium chloride are ionic bonds and
compound (calcium chloride) is an ionic compound. In similar way; Barium
chloride is formed.
Formation of
Calcium oxide (CaO):
Valence electron = 2
Atomic number of oxygen is 8
Electronic configuration of oxygen is: 2, 6
Number of electrons in outermost orbit = 6
Valence electron = 6
Calcium loses two electrons and gets two positive charges over
it in order to get stability. Oxygen gains two electrons; lost by calcium and
thus gets two negative charges over it.
Bond formed between calcium oxide is ionic bond. Calcium oxide
is an ionic compound. In similar way; magnesium oxide is formed.
Due to
the presence of the strong force of attraction between the positive and
negative ions, ionic compounds are solids and are hard to break. They generally
break into pieces when pressure is applied, hence they are considered brittle.
Due to
the presence of electrostatic forces of attraction between ions, a large amount
of energy is required to break the ionic bonds between the atoms. Thus, ionic
compounds have high melting and boiling points.
Ionic
compounds are generally soluble in polar solvents such as water whereas solubility tends to decrease in
non-polar solvents such as petrol, gasoline, etc.
Ionic
compounds do not conduct electricity in the solid-state but are good conductors
in a molten state. Conduction of electricity involves the flow of charge from
one point to another. In the solid-state, as the movement of ions is not possible,
ionic compounds don’t conduct electricity. Whereas in the molten state, ionic
compounds conduct electricity as electrostatic forces of attraction between the
ions are overcome by the heat released.
4.
Basic Metallurgical Processes:
Metallurgy is defined as a process that is used for the extraction of metals in
their pure form. The compounds of metals mixed with soil, limestone, sand, and
rocks are known as minerals. Metals are commercially extracted from minerals at
low cost and minimum effort. These minerals are known as ores. A substance
which is added to the charge in the furnace to remove the gangue (impurities)
is known as flux. Metallurgy deals with the process of purification of metals
and the formation of alloys.
The
various steps used in metallurgy are given below:
1.
Enrichment or dressing ore
2.
Conversion of the enriched ore into the oxide of
metal
3.
Extraction of metal from the metal oxide
4.
Refining or Purification of the metal
1. Enrichment or dressing of Ores:
Ores mined from the earth are usually contaminated
with large amounts of impurities such as soil, sand, etc., called gangue. The
impurities must be removed from the ore prior to the extraction of the metal.
The processes used for removing the gangue from the ore are based on the
differences between the physical or chemical properties of the gangue and the
ore. Different separation techniques are accordingly employed. Now, we shall discuss the different
processes which are used for enrichment of different types of ores.
A. Hydraulic
Washing:
This method is used for the enrichment of those
ores which are heavier than gangue particles present in them. In this method, a
stream of water is passed through crushed and finely powered ore.
The Lighter gangue particles are washed away with
water while the heavier ore particles are left behind. Oxide ores of tin and
lead are concentrated by this method.
B.
Froth Floatation:
This method is used
for concentration of sulphide ores of copper, lead and zinc. In this method,
powdered ore is put in a tank full of water. And then, some Pine oil is added
to it. In the tank the particles of sulphide ore are wetted by pine oil whereas
the gangue particles are wetted by water. Then air is passed through this
mixture.
This results in the
agitation of water in tank, which cause the sulphide ore particles to stick
with the oil and rise to the surface in the form of froth. The gangue particles
being heavier remain behind at the bottom of water tank. The froth is separated
and concentrated sulphide ore is obtained from it.
C.
Magnetic Separation:
This method is used for concentration
of magnetic ores of iron (magnetite and chromite) and manganese (pyrolusite) by
removing non-magnetic impurities present in them. This process involves the use
of a magnetic separation.
A magnetic separator consists of a
leather belt which moves over two rollers. Out of two rollers one roller has a
magnet in it. In this method, the finely powdered magnetic ore is dropped over
the moving belt at one end. When the powdered ore falls down from the moving
belt at the other end having a magnetic roller, the particles of ore are
attracted by the magnet and form a separate heap from the non- magnetic
impurities.
D. Chemical
Separation:
This method is based
on the differences in some chemical properties of the gangue and the ore. For
example, the impure ore of metal aluminium (bauxite or aluminium oxide) is
concentrated by Baeyer’s process.
BAEYER’S PROCESS:
In Baeyer’s process, the
finely powdered bauxite ore is treated with hot sodium hydroxide solution to
form a water soluble compound called sodium aluminate.
Al2O3
+ 2NaOH
Bauxite Sodium hydroxide
Sodium aluminate Water
The gangue present in bauxite does
not react in sodium hydroxide sol, so the gangue can be separated by the
process of filtration. After filtration, the filtrate containing solution of
sodium aluminate is acidified with HCl to form precipitates of aluminium
hydroxide.
NaAlO2
+ HCl + H2O
Sodium aluminate Hydrochloric acid
Aluminium hydroxide Salt
The precipitates of aluminium
hydroxide are then filtered, washed, dried and heated to get pure aluminium
oxide.
2Al(OH)3
Aluminium Hydroxide
Aluminium Oxide
It should be noted that pure aluminium oxide is also
known as alumina.
2. Conversion of the enriched ore into
the oxide of metal:
After concentration of ores, the sulphide
or carbonate ores of some metals are converted into metal oxides because it is
easier to obtain metals from their metal oxides as compared to metal sulphides
or metal carbonates. The sulphide ores of metals can be converted into their
oxides by roasting while the carbonate ores of metals can be converted into
their oxides by calcinations.
Roasting:
It may be defined as the process of
strongly heating a sulphide ore in the presence of air to convert it into metal
oxide. For example,
2ZnS + 3O2
Zinc
sulphide Zinc
Oxide
2PbS + 3O2
Lead
sulphide Lead
Oxide
Calcinations:
It may be defined as the process of strongly heating a
carbonate ore in the absence of air to convert it into metal oxide. For example,
CaCO3
Calcium
carbonate Calcium
oxide
3. Extraction of metal from the metal
oxide:
A.
Reduction by Heat (pyrometallurgy):
The oxides of metals which are present at
the bottom of reactivity series can be reduced to metals by action of heat
alone e.g. mercury oxide can be reduced to mercury metal by heating it to a
temperature of about 300°C.
2HgO
Mercury
oxide Mercury
metal
B. Reduction by Coke (smelting):
The
Oxides of Metals like Zn, Fe, Cu, Ni, Sn and Pb are usually reduced by using
carbon as reducing agent. In this process, coke is mixed with roasted ore and
heated to a high temperature in a furnace. Coke reduces the metal oxides into
free metal. For example,
ZnO + C
Zinc
oxide Carbon Zinc
Carbon monoxide
PbO + C
Lead
oxide Carbon Lead
Carbon monoxide
C. Reduction by Aluminium (aluminotherapy):
Oxides
of manganese and chromium metals are reduced to metals with the help of
Aluminium.
The
process of reduction of a metal oxide to the metal with the help of aluminium
is called aluminotherapy.
3MnO2 + 4Al
Manganese
dioxide Manganese
metal Aluminium oxide
D. Electrolytic Reduction:
The Oxides of metals which are quite high
in reactivity series can be reduced to metals by electrolytic reduction. For
example, sodium and magnesium metals are obtained by electrolytic reduction of
their chloride solutions in molten state.
2NaCl
Sodium
chloride
Sodium metal Chlorine gas
MgCl2
Magnesium
chloride
Magnesium metal Chlorine gas
During
this process, chlorine gas is liberated at anode while metals (sodium or
magnesium) deposit at cathode.
4. Refining or Purification of the
metal:
The metal obtained by above methods is
usually impure. So, it is to be purified. The method used for refining of metal
depends on the nature of metal and impurities present in it. Some common
methods which are used for purification of impure metals are:
A. Distillation
Method:
This method is useful for purification of those
volatile metals which have low boiling points such as zinc and mercury. In this
method, the impure metal is heated to its boiling point in a vessel. The
vapours of metal thus formed are
collected
and cooled in a separate vessel to get pure metal. The impurities being
non-volatile remain behind.
B. Liquation Method:
By this method those metals can
be purified which have low melting point. In this method the block of Impure
metal is placed on the top side of sloping floor of a furnace and heated
gently. Due to high temperature the fusible metal melts and flows down to the
bottom of sloping floor while the non-fusible impurities remain behind on the
floor. Finally the pure metal is collected from the bottom of sloping floor.
C.
Electrolytic Refining:
Many metals like Cu, Zn, Pb, Cr, Ni,
Ag, and Au are refined electrolytically for refining of an impure metal by
electrolysis.
We shall understand
electrolytic refining of metals by taking the example of refining of copper. In
case of copper, a thick block of impure copper is made anode and a thin block
of pure metal is made cathode and copper sulphate solution is used as an
electrolyte. On passing electric current, pure copper metal from the
electrolyte solution deposits on the cathode. At the same time an equal amount
of impure copper dissolves from anode into the electrolyte solution. The soluble
impurities settle down in the solution below the anode and are called as anode
mud.
5. Corrosion and It’s Prevention :
In the corrosion
process, iron metal acts as the anode in a galvanic cell and is oxidized to Fe2+;
oxygen is reduced to water at the cathode. The relevant reactions are as
follows:
The Fe2+ ions
produced in the initial reaction are then oxidized by atmospheric oxygen to
produce the insoluble hydrated oxide containing Fe3+, as represented
in the following equation:
The sign and
magnitude of E° for the corrosion process indicate that there is a strong driving force
for the oxidation of iron by O2 under standard conditions (1 M
H+). Under neutral conditions, the driving force is somewhat less
but still appreciable (E = 1.25 V at pH 7.0). Normally, the reaction of
atmospheric CO2 with water to form H+ and HCO3− provides
a low enough pH to enhance the reaction rate, as does acid rain. Automobile
manufacturers spend a great deal of time and money developing paints that
adhere tightly to the car’s metal surface to prevent oxygenated water, acid,
and salt from coming into contact with the underlying
metal. Unfortunately, even the best paint is subject to scratching or
denting, and the electrochemical nature of the corrosion process means that two
scratches relatively remote from each other can operate together as anode and
cathode, leading to sudden mechanical failure.
Small Scratches in a Protective Paint Coating
Can Lead to the Rapid Corrosion of Iron. Holes in a protective coating allow
oxygen to be reduced at the surface with the greater exposure to air (the
cathode), while metallic iron is oxidized to Fe2+(aq) at the less
exposed site (the anode). Rust is formed when Fe2+(aq) diffuses to a
location where it can react with atmospheric oxygen, which is often remote from
the anode. The electrochemical interaction between cathodic and anodic sites
can cause a large pit to form under a painted surface, eventually resulting in
sudden failure with little visible warning that corrosion has occurred.
Methods to Prevent Rusting of Iron:
Rusting of iron can be prevented by
cutting off the contact between the metal and air. Some methods which are used
to prevent rusting of iron are
1.
Rusting of iron can be prevented by
applying paints, oils and grease over the surface of iron.
2.
Rusting of iron can be prevented by
galvanization. Galvanization is the process of depositing a thin layer of zinc
metal on iron articles.
3.
Rusting of iron can be prevented by
coating its surface with other metals like tin, nickel and chromium.
4.
Rusting of iron can be prevented by
alloying it with chromium and nickel to make stainless steel.