Chemistry

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Electrochemical Principle Of Metallurgy

Similar to thermodynamic principles, electrochemical principles also find applications in metallurgical process. The reduction of oxides of active metals such as sodium, potassium etc., by carbon is thermodynamically not feasible.

Such metals are extracted from their ores by using electrochemical methods. In this technique, the metal salts are taken in a fused form or in solution form. The metal ion present can be reduced by treating it with some suitable reducing agent or by electrolysis.

Gibbs free energy change for the electrolysis process is given by the following expression

ΔG° = -nFE°

Where n is number of electrons involved in the reduction process, F is the Faraday and E° is the electrode potential of the redox couple.

If E° is positive then the ΔG is negative and the reduction is spontaneous and hence a redox reaction is planned in such a way that the e.m.f of the net redox reaction is positive. When a more reactive metal is added to the solution containing the relatively less reactive metal ions, the more reactive metal will go into the solution. For example,

Cu (s) + 2Ag⁺ (aq) → Cu²⁺ (aq) + 2Ag (s)
Cu²⁺ (aq) + Zn (s) → Cu(s) + Zn²⁺ (aq)

Electrochemial Extraction of Aluminium – Hall – Heroult Process:

In this method, electrolysis is carried out in an iron tank lined with carbon which acts as a cathode. The carbon blocks immersed in the electrolyte act as a anode. A 20% solution of alumina, obtained from the bauxite ore is mixed with molten cryolite and is taken in the electrolysis chamber.

About 10% calcium chloride is also added to the solution. Here calcium chloride helps to lower the melting point of the mixture. The fused mixture is maintained at a temperature of above 1270 K. The chemical reactions involved in this process are as follows.

Ionisaiton of alumina Al₂O₃ → 2Al³⁺ + 3O^2-
Reaction at cathode 2Al³⁺ (melt) + 6e^– → 2Al (l)
Reaction at anode 6O^2- (melt) → 3O₂ + 12e^–

Since carbon acts as anode the following reaction also takes place on it.

C(s) + O^2- (melt) → CO + 2e^–
C(s) + 2O^2- (melt) → CO₂ + 4e^–

Due to the above two reactions, anodes are slowly consumed during the electrolysis. The pure aluminium is formed at the cathode and settles at the bottom. The net electrolysis reaction can be written as follows.

4Al³⁺ (melt) + 6O^2- (melt) + 3C (s) → 4Al (l) + 3CO₂ (g)

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Thermodynamic Principle of Metallurgy

As we discussed, the extraction of metals from their oxides can be carried out by using different reducing agents. For example, consider the reduction of a metal oxide M_xO_y. \(\frac{2}{y}\)M_xO_y (s) → \(\frac{2x}{y}\)M (s) + O₂ (g) ………. (1)

The above reduction may be carried out with carbon. In this case, the reducing agent carbon may be oxidised to either CO or CO₂.

C + O₂ → CO₂ (g) …………. (2)
2C + O₂ → 2CO (g) ……………. (3)

If carbon monoxide is used as a reducing agent, it is oxidised to CO₂ as follows,

2CO + O₂ → 2CO₂ (g) ………………. (4)

A suitable reducing agent is selected based on the thermodynamic considerations. We know that for a spontaneous reaction, the change in free energy (ΔG) should be negative. Therefore, thermodynamically, the reduction of metal oxide [equation (1)] with a given reducing agent [Equation (2), (3) or (4)] can occur if the free energy change for the coupled reaction. [Equations (1) & (2), (1) & (3) or (1) & (4)] is negative. Hence, the reducing agent is selected in such a way that it provides a large negative ΔG value for the coupled reaction.

Ellingham Diagram

The change in Gibbs free energy (ΔG) for a reaction is given by the expression.
ΔG = ΔH – TΔS ………….. (1)

where, ΔH is the enthalpy change, T the temperature in kelvin and ΔS the entropy change. For an equilibrium process, ΔG^° can be calculated using the equilibrium constant by the following expression
ΔG^° = – RT lnK_P

Harold Ellingham used the above relationship to calculate the ΔG^° values at various temperatures for the reduction of metal oxides by treating the reduction as an equilibrium process.

He has drawn a plot by considering the temperature in the x-axis and the standard free energy change for the formation of metal oxide in y-axis. The resultant plot is a straight line with ΔS as slope and ΔH as y-intercept. The graphical representation of variation of the standard Gibbs free energy of reaction for the formation of various metal oxides with temperature is called Ellingham diagram.

Observations from the Ellingham Diagram

1. For most of the metal oxide formation, the slope is positive. It can be explained as follows. Oxygen gas is consumed during the formation of metal oxides which results in the decrease in randomness. Hence, ΔS becomes negative and it makes the term, TΔS positive in the straight line equation.

2. The graph for the formation of carbon monoxide is a straight line with negative slope. In this case ΔS is positive as 2 moles of CO gas is formed by the consumption of one mole of oxygen gas. It indicates that CO is more stable at higher temperature.

3. As the temperature increases, generally ΔG value for the formation of the metal oxide become less negative and becomes zero at a particular temperature. Below this temperature, ΔG is negative and the oxide is stable and above this temperature ΔG is positive. This general trend suggests that metal oxides become less stable at higher temperature and their decomposition becomes easier.

4. There is a sudden change in the slope at a particular temperature for some metal oxides like MgO, HgO. This is due to the phase transition (melting or evaporation).

Applications of the Ellingham Diagram:

Ellingham diagram helps us to select a suitable reducing agent and appropriate temperature range for reduction. The reduction of a metal oxide to its metal can be considered as a competition between the element used for reduction and the metal to combine with oxygen.

If the metal oxide is more stable, then oxygen remains with the metal and if the oxide of element used for reduction is more stable, then the oxygen from the metal oxide combines with elements used for the reduction. From the Ellingham diagram, we can infer the relative stability of different metal oxides at a given temperature.

1. Ellingham diagram for the formation of Ag₂O and HgO is at upper part of the diagram and their decomposition temperatures are 600 and 700 K respectively. It indicates that these oxides are unstable at moderate temperatures and will decompose on heating even in the absence of a reducing agent.

2. Ellingham diagram is used to predict thermodynamic feasibility of reduction of oxides of one metal by another metal. Any metal can reduce the oxides of other metals that are located above it in the diagram. For example, in the Ellignham diagram, for the formation of chromium oxide lies above that of the aluminium, meaning that Al₂O₃ is more stable than Cr₂O₃.

Hence aluminium can be used as a reducing agent for the reduction of chromic oxide. However, it cannot be used to reduce the oxides of magnesium and calcium which occupy lower position than aluminium oxide.

3. The carbon line cuts across the lines of many metal oxides and hence it can reduce all those metal oxides at sufficiently high temperature. Let us analyse the thermodynamically favourable conditions for the reduction of iron oxide by carbon.

Ellingham diagram for the formation of FeO and CO intersects around 1000 K. Below this temperature the carbon line lies above the iron line which indicates that FeO is more stable than CO and hence at this temperature range, the reduction is not thermodynamically feasible. However, above 1000 K carbon line lies below the iron line and hence, we can use coke as reducing agent above this temperature. The following free energy calculation also confim that the reduction is thermodynamically favoured.

From the Ellingham Diagram at 1500 K,

2Fe (s) + O₂ (g) → 2FeO (g) ΔG₁ = – 350 kJ mol^-1 ……………. (1)
2C (s) + O₂ (g) → 2CO (g) ΔG₂ = – 480 kJ mol^-1 …………….. (2)

Reverse the reaction (1)

2FeO (s) → 2Fe (s) + O₂ (g) – ΔG₁ = + 350 kJ mol^-1 …………….. (3)

Now couple the reactions (2) and (3)

2FeO (s) + 2C → 2Fe (l, s) + 2CO (g) ΔG₃ = – 130 kJ mol^-1 …………….. (4)

The standard free energy change for the reduction of one mole of FeO is, ΔG₃/2 = – 65 kJ mol^-1

Limitations of Ellingham Diagram

1. Ellingham diagram is constructed based only on thermodynamic considerations. It gives information about the thermodynamic feasibility of a reaction. It does not tell anything about the rate of the reaction. More over, it does not give any idea about the possibility of other reactions that might be taking place.

2. The interpretation of ΔG is based on the assumption that the reactants are in equilibrium with the products which is not always true.

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Extraction of Crude Metal

The extraction of crude metals from the concentrated ores is carried out in two steps namely

1. Conversion of the ore into oxides of the metal of interest and
2. Reduction of the metal oxides to elemental metals. In the concentrated ore, the metal exists in positive oxidation state and hence it is to be reduced to its elemental state.
3. We can infer from the principles of thermodynamics, that the reduction of oxide is easier when compared to reduction of other compounds of metal and hence, before reduction, the ore is first converted into the oxide of metal of interest.
4. Let us discuss some of the common methods used to convert the concentrated ore into the oxides of the metal of interest.

Conversion of Ores into Oxides

Roasting

Roasting is the method, usually applied for the conversion of sulphide ores into their oxides. In this method, the concentrated ore is oxidised by heating it with excess of oxygen in a suitable furnace below the melting point of the metal.

Roasting also removes impurities such as arsenic, sulphur, phosphorous by converting them into their volatile oxides.

For Example

4As + 3O₂ → 2As₂O₃ ↑
S₈ + 😯₂ → 8SO₂ ↑
P₄ + 5O₂ → P₄O₁₀ ↑

Calcination

Calcination is the process in which the concentrated ore is strongly heated in the absence of air. During this process, the water of crystallisation present in the hydrated oxide escapes as moisture. Any organic matter (if present) also get expelled leaving behind a porous ore. This method can also be carried out with a limited supply of air.

For examples,
During calcination of carbonate ore, carbon dioxide is expelled

During calcination of hydrated ore, the water of hydration is expelled as vapour

Reduction of Metal Oxides

Metal oxide can be reduced to crude metal by using a suitable reducing agent like carbon, carbon monoxide, hydrogen, aluminium and other reactive metals such as sodium etc. The choice of reducing agent depends on the nature of the metal.

For example, carbon cannot be used as a reducing agent for the reactive metals such as sodium, potassium, aluminium etc. Similarly CO cannot be used to reduce oxides such as ZnO, Al₂O₃. Later in this, we study selection of suitable reducing agents by applying ellingham diagram.

Smelting

In this method, a flux (a chemical substance that forms an easily fusible slag with gangue) and a reducing agent such as carbon, carbon monoxide (or) aluminium is added to the concentrated ore and the mixture is melted by heating at an elevated temperature (above the melting point of the metal) in a smelting furnace. For example the oxide of iron can be reduced by carbon monoxide as follows.

Fe₂O₃(s) + 3CO (g) → 2Fe(s) + 3CO₂(g)↑

In this extraction, a basic flux, limestone (CaO) is used. Since the silica gangue present in the ore is acidic in nature, the limestone combines with it to form calcium silicate (slag).

In the extraction of copper from copper pyrites, the concentrated ore is heated in a reverberatory furnace after mixing with silica, an acidic flux. The ferrous oxide formed due to melting is basic in nature and it combines with silica to form ferrous silicate (slag). The remaining metal sulphides Cu₂S and FeS are mutually soluble and form a copper matte.

2CuFeS₂(s) + O₂(g) → 2Cu₂(l, s) + 2SO₂ (g)
2Cu₂O(l) + Cu₂S (l) → 6Cu(l) + SO₂(g)

The matte is separated from the slag and fed to the converting furnace. During conversion, the FeS present in the matte is first oxidised to FeO. This is removed by slag formation with silica. The remaining copper sulphide is further oxidised to its oxide which is subsequently converted to metallic copper as shown below.

2Cu₂S (l, s) + 3O₂ (g) → 2Cu₂O (l, s) + 2SO₂(g)
2Cu₂O(l) + Cu₂S (l) → 6Cu (l) + SO₂ (g)

The metallic copper is solidified and it has blistered appearance due to evolution of SO₂ gas formed in this process. This copper is called blistered copper.

Reduction By Carbon:

In this method the oxide ore of the metal is mixed with coal (coke) and heated strongly in a furnace (usually in a blast furnace). This process can be applied to the metals which do not form carbides with carbon at the reduction temperature.

Examples:

ZnO (s) + C (s) → Zn (s) + CO (g)↑
Mn₃O₄ (s) + 4C (s) → 3Mn (s) + 4CO (g)↑
Cr₂O₃ (s) + 3C (s) → 2Cr (s) + 3CO (g)↑

Reduction By Hydrogen:

This method can be applied to the oxides of the metals (Fe, Pb, Cu) having less electropositive character than hydrogen.

Ag₂O (s) + H₂ (g) → 2Ag (s) + H₂O (l)
Fe₃O₄ (s) + 4H₂ (g) → 3Fe (s) + 4H₂O (l)

Nickel oxide can be reduced to nickel by using a mixture of hydrogen and carbon monoxide (water gas)

2NiO (s) + CO (g) + H₂ (g) → 2Ni (s) + CO₂ (g) + H₂O (l)

Reduction By Metal:

Metallic oxides such as Cr₂O₃ can be reduced by an aluminothermic process. In this process, the metal oxide is mixed with aluminium powder and placed in a fire clay crucible. To initiate the reduction process, an ignition mixture (usually magnesium and barium peroxide) is used.

BaO₂ + Mg → BaO + MgO

During the above reaction a large amount of heat is evolved (temperature up to 2400°C, is generated and the reaction enthalpy is : 852 kJ mol^-1) which facilitates the reduction of Cr₂O₃ by aluminium power.

Active metals such as sodium, potassium and calcium can also be used to reduce the metal oxide

B₂O₃ + 6Na → 2B + 3Na₂O
Rb₂O₃ + 3Mg → 2Rb + 3MgO
TiO₂ + 2Mg → Ti + 2MgO

Auto-Reduction:

Simple roasting of some of the ores give the crude metal. In such cases, the use of reducing agents is not necessary. For example, mercury is obtained by roasting of its ore cinnabar (HgS).

HgS (s) + O₂ (g) → Hg (l) + SO₂↑

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Concentration Of Ores

Generally, the ores are associated with nonmetallic impurities, rocky materials and siliceous matter which are collectively known as gangue. The preliminary step in metallurgical process is removal of these impurities. This removal process is known as concentration of ore.

It increases the concentration of the metal of interest or its compound in the ore. Several methods are available for this process and the choice of method will depend on the nature of the ore, type of impurity and environmental factors. Some of the common methods of ore concentration are discussed below.

Gravity Separation or Hydraulic Wash

In this method, the ore having high specific gravity is separated from the gangue that has low specific gravity by simply washing with running water. Ore is crushed to a finely powdered form and treated with rapidly flowing current of water.

During this process the lighter gangue particles are washed away by the running water. This method is generally applied to concentrate the native ore such as gold and oxide ores such as haematite (Fe₂O₃), tin stone (SnO₂) etc.

Froth Flotation

This method is commonly used to concentrate sulphide ores such as galena (PbS), zincblende (ZnS) etc. In this method, the metallic ore particles which are preferentially wetted by oil can be separated from gangue. In this method, the crushed ore is suspended in water and mixed with frothing agent such as pine oil, eucalyptus oil etc. A small quantity of sodium ethyl xanthate which acts as a collector is also added.

A froth is generated by blowing air through this mixture. The collector molecules attach to the ore particle and make them water repellent. As a result, ore particles, wetted by the oil, rise to the surface along with the froth. The froth is skimmed off and dried to recover the concentrated ore. The gangue particles that are preferentially wetted by water settle at the bottom.

When a sulphide ore of a metal of interest contains other metal sulphides as impurities, depressing agents such as sodium cyanide, sodium carbonate etc are used to selectively prevent other metal sulphides from coming to the froth. For example, when impurities such as ZnS is present in galena (PbS), sodium cyanide (NaCN) is added to depresses the flotation property of ZnS by forming a layer of zinc complex Na₂[Zn(CN)₄] on the surface of zinc sulphide.

Leaching

This method is based on the solubility of the ore in a suitable solvent and the reactions in aqueous solution. In this method, the crushed ore is allowed to dissolve in a suitable solvent, the metal present in the ore is converted to its soluble salt or complex while the gangue remains insoluble. The following examples illustrate the leaching processes.

Cyanide Leaching

Let us consider the concentration of gold ore as an example. The crushed ore of gold is leached with aerated dilute solution of sodium cyanide. Gold is converted into a soluble cyanide complex. The gangue, aluminosilicate remains insoluble.

4Au (s) + 8CN^– (aq) + O₂ (g) + 2H₂O (l) → 4[Au(CN)₂]^– (aq) + 4OH^–(aq)

Recovery of Metal of Interest from the Complex by Reduction:

Gold can be recovered by reacting the deoxygenated leached solution with zinc. In this process the gold is reduced to its elemental state (zero oxidation sate) and the process is called cementation.

Zn (s) + 2[Au(CN)₂]^– (aq) → [Zn(CN)₄]^-2 (aq) + 2Au (s)

Ammonia Leaching

When a crushed ore containing nickel, copper and cobalt is treated with aqueous ammonia under suitable pressure, ammonia selectively leaches these metals by forming their soluble complexes viz. [Ni(NH₃)₆]²⁺, [Cu(NH₃)₄]²⁺, and [Co(NH₃)₅H₂O]³⁺ respectively from the ore leaving behind the gangue, iron(III) oxides/hydroxides and aluminosilicate.

Alkali Leaching

In this method, the ore is treated with aqueous alkali to form a soluble complex. For example, bauxite, an important ore of aluminum is heated with a solution of sodium hydroxde or sodium carbonate in the temperature range 470 – 520 K at 35 atm to form soluble sodium meta-aluminate leaving behind the impurities, iron oxide and titanium oxide.

Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2Na[Al(OH)₄] (aq)

The hot solution is decanted, cooled, and diluted. This solution is neutralised by passing CO₂ gas, to the form hydrated Al₂O₃ precipitate.

2Na[Al(OH)₄] (aq) + 2CO₂(g) → Al₂O₃.3H₂O(s) + 2NaHCO₃(aq)
The precipitate is fitered of and heated around 1670 K to get pure alumina Al₂O₃.

Acid Leaching

Leaching of sulphide ores such as ZnS, PbS etc., can be done by treating them with hot aqueous sulphuric acid.

2ZnS (s) + 2H₂SO₄ (aq) + O₂(g) → 2ZnSO₄(aq) + 2S(s) + 2H₂O

In this process the insoluble sulphide is converted into soluble sulphate and elemental sulphur.

Magnetic Separation

This method is applicable to ferromagnetic ores and it is based on the difference in the magnetic properties of the ore and the impurities. For example tin stone can be separated from the wolframite impurities which is magnetic. Similarly, ores such as chromite, pyrolusite having magnetic property can be removed from the non magnetic siliceous impurities.

The crushed ore is poured on to an electromagnetic separator consisting of a belt moving over two rollers of which one is magnetic. The magnetic part of the ore is attracted towards the magnet and falls as a heap close to the magnetic region while the nonmagnetic part falls away from it as shown in the figure 1.2.

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General Trends In Properties Of P -Block Elements

We already learnt that the properties of elements largely depends on their electronic configuration, size,  ionisation enthalpy, electronegativity etc. Let us discuss the general trend in such properties of various p-block elements.

Electronic Configuration and Oxidation State:

The p-block elements have a general electronic configuration of ns², np^1-6. The elements of each group have similar outer shell electronic configuration and differ only in the value of n (principal quantum number). The elements of group 18 (inert gases) have completely filled p orbitals, hence they are more stable and have least reactivity.

The elements of this block show variable oxidation state and their highest oxidation state (group oxidation state) is equal to the total number of valance electrons present in them. Unlike s-block elements which show only positive oxidation state, some of the p-block elements show negative oxidation states also.

The halogens have a strong tendency to gain an electron to give a stable halide ion with completely filled electronic configuration and hence – 1 oxidation state is more common in halogens. Similarly, the other elements belonging to pnictogen and chalcogen groups also show negative oxidation states.

General Electronic Configurations and Oxidation States of P-Block Elements

The tendency of an element to form a cation by loosing electrons is known as electropositive or metallic character. This character depends on the ionisation energy. Generally on descending a group the ionisation energy decreases and hence the metallic character increases.

Figure 2.1 P-Block Elements With Their Ionisation Enthalpies, Electronegativity and Metallic Nature.

In p-block, the elements present in lower left part are metals while the elements in the upper right part are non metals. Elements of group 13 have metallic character except the first element boron which is a metalloid, having properties intermediate between the metal and nonmetals. The atomic radius of boron is very small and it has relatively high nuclear charge and these properties are responsible for its nonmetallic character.

In the subsequent groups the non-metallic character increases. In group 14 elements, carbon is a nonmetal while silicon and germanium are metalloids. In group 15, nitrogen and phosphorus are non metals and arsenic & antimony are metalloids. In group 16, oxygen, sulphur and selenium are non metals and tellurium is a metalloid. All the elements of group 17 and 18 are non metals.

Ionisation Enthalpy:

We have already learnt that as we move down a group, generally there is a steady decrease in ionisation enthalpy of elements due to increase in their atomic radius. In p-block elements, there are some minor deviations to this general trend. In group 13, from boron to aluminium the ionisation enthalpy decreases as expected. But from aluminium to thallium there is only a marginal difference.

This is due to the presence of inner d and f-electrons which has poor shielding effect compared to s and p-electrons. As a result, the effective nuclear charge on the valance electrons increases. A similar trend is also observed in group 14.

The remaining groups (15 to 18) follow the general trend. In these groups, the ionisation enthalpy decreases, as we move down the group. Here, poor shielding effect of d- and f-electrons are overcome by the increased shielding effect of the additional p-electrons. The ionisation enthalpy of elements in successive groups is higher than the corresponding elements of the previous group as expected.

Electronegativity

As we move down the 13^th group, the electronegativity first decreases from boron to aluminium and then marginally increases for Gallium, there after there is no appreciable change. Similar trend is also observed in 14^th group as well. In other groups, as we move down the group, the electro negativity decreases. This observed trend can be correlated with their atomic radius.

Anomalous Properties of the First Elements:

In p-block elements, the first member of each group differs from the other elements of the corresponding group. The following factors are responsible for this anomalous behaviour.

1. Small size of the first member
2. High ionisation enthalpy and high electronegativity
3. Absence of d orbitals in their valance shell

The first member of the group 13, boron is a metalloid while others are reactive metals. Moreover, boron shows diagonal relationship with silicon of group 14. The oxides of boron and silicon are similar in their acidic nature. Both boron and silicon form covalent hydrides that can be easily hydrolysed. Similarly, except boron trifluoride, halides of both elements are readily hydrolysed.

In group 14, the first element carbon is strictly a nonmetal while other elements are metalloids (silicon & germanium) or metals (tin & lead). Unlike other elements of the group carbon can form multiple bonds such as C=C, C=O etc. Carbon has a greater tendency to form a chain of bonds with itself or with other atoms which is known as catenation. There is considerable decrease in catenation property down the group (C>>Si>Ge≈Sn>Pb).

In group 15 also the first element nitrogen differs from the rest of the elements of the group. Like carbon, the nitrogen can from multiple bonds (N=N, C=N, N=O etc…). Nitrogen is a diatomic gas unlike the other members of the group. Similarly in group 16, the first element, oxygen also exists as a diatomic gas in that group. Due to its high electronegativity it forms hydrogen bonds.

The first element of group 17, flourine the most electronegative element, also behaves quiet differently compared to the rest of the members of group. Like oxygen it also forms hydrogen bonds. It shows only -1 oxidation state while the other halogens have +1, +3, +5 and +7 oxidation states in addition to -1 state. The flourine is the strongest oxidising agent and the most reactive element among the halogens.

Inert Pair Effect:

We have already learnt that the alkali and alkaline earth metals have an oxidation state of +1 and +2 respectively, corresponding to the total number of electrons present in them. Similarly, the elements of p block also show the oxidation states corresponding to the maximum number of valence electrons (group oxidation state).

In addition they also show variable oxidation state. In case of the heavier post-transition elements belonging to the groups (13 to 16), the most stable oxidation state is two less than the group oxidation state and there is a reluctance to exhibit the group oxidation state.

Let us consider group 13 elements. As we move from boron to heavier elements, there is an increasing tendency to have +1 oxidation state, rather than the group oxidation state, +3. For example Al⁺³ is more stable than Al⁺¹ while Tl⁺¹ is more stable than Tl⁺³. Aluminium (III) chloride is stable whereas thallium (III) chloride is highly unstable and disproportionates to thallium (I) chloride and chlorine gas.

This shows that in thallium the stable lower oxidation state corresponds to the loss of np electrons only and not ns electrons. Thus in heavier posttransition metals, the outer s electrons (ns) have a tendency to remain inert and show reluctance to take part in the bonding, which is known as inert pair effect. This effect is also observed in groups 14, 15 and 16.

Allotropism in P-Block Elements:

Some elements exist in more than one crystalline or molecular forms in the same physical state. For example, carbon exists as diamond and graphite. This phenomenon is called allotropism (in greek ‘allos’ means another and ‘trope’ means change) and the different forms of an element are called allotropes. Many p-block elements show allotropism and some of the common allotropes are listed in the table.

Table 2.2 : Some of common allotropes of p-block elements

Element	Most Common Allotropes
Boron	Amorphous boron, α-rhombohedral boron, β-rhombohedral boron, γ-orthorhombic boron, α-tetragonal boron, β-tetragonal boron
Carbon	Diamond, Graphite, Graphene, Fullerenes, Carbon nanotubes
Silicon	Amorphous Silicon, Crystalline Silicon
Germanium	α-germanium, β-germanium
Tin	Grey tin, white tin, rhombic tin, sigma tin
Phosphorus	White phosphorus, Red phosphorus, Scarlet phosphorus, Violet phosphorus, Black phosphorus.
Arsenic	Yellow arsenic, gray arsenic & Black arsenic
Anitimony	Blue-white antimony, Yellow, Black
Oxygen	Dioxygen, ozone
Sulphur	Rhombus sulphur, monoclinic sulphur
Selenium	Red selenium, Gray selenium, Black selenium, Monoclinic selenium
Tellurium	Amorphous & Crystalline