Crushing and Grinding

The Ore obtained from mining is in the form of big lumps. These are broken into smaller pieces with the help of crushers known as Jaw Crushers. This process is known as Crushing. The crushed ore is then converted into fine powder with the help of stamp mill or ball mill. This process is known as Pulverisation of Ore.

Metallurgical Operations

The branch of science which deals with the extraction of metal from its ore is called Metallurgy. There is no common process available for the extraction of all the metals because they differ in their chemical and physical properties and also the impurities associated with them. However, there are certain operations or procedures common in the metallurgy of all metals. These operations or procedures are called Metallurgical Operations. The various steps involved in metallurgical operations are;

1.   Crushing and Grinding of Ore (Pulverisation)

2.   Concentration of Ore.

3.   Extraction of Metal from Oxide Ore.

4.   Refining of Crude Metal.

Types of Minerals

The Various types of minerals depend upon the nature of the chemical compounds that contains the given metal. These are:

1. OXIDES:  Oxides are those minerals in which a metal occurs in combination with oxygen. Examples of such minerals are,

            (a).       Iron as Hematite (Fe₂O₃), Magnetite (Fe₂O₃. FeO), Chromate (Fe₃O₄Cr₂).

            (b).      Aluminium as Bauxite                        (Al₂O₃ . 2H₂O)

            (c).       Manganese as Pyolusite                      (MnO₂)

            (d).      Tin as Tin stone or Cassiterite             (SnO₂)

            (e).       Titanium as Rutile                               (TiO₂)

            (f).       Copper as Cuprite                               (Cu₂O)

2. SULPHIDES:  Sulphide are those minerals in which a metal occurs in combination with sulphur.

            For Examples:–        

            (a).       Iron as Iron Pyrites                 (FeS₂)

            (b).      Copper as Copper Pyrites       (CuFeS₂)

            (c).       Copper as Copper Glance       (Cu₂S)

            (d).      Mercury as Cinnabar               (HgS)

            (e).       Zinc as Zinc Sulphide             (ZnS)

            (f).       Lead as Lead Sulphides          (PbS)

3. CARBONATES:  These are those minerals in which a metal occurs as a carbonates.

            For Example:

            (a).       Calcium as Calcite                  (CaCO₃)

            (b).       Magnesium as Dolomite         (CaCO₃ . MgCO₃)

            (c).       Zinc as Calamine                     (ZnCO₃)

            (d).      Iron as Siderite                        (FeCO₃)          

4. SULPHATES:  The minerals in which a metal occurs as sulphates are sulphate minerals.

            For Example

            (a).       Barium as Barytes                   (BaSO₄)

            (b).      Lead as Anglesite                   (PbSO₄)

            (c).       Calcium as Gypsum                (CaSO₄ . 2H₂O)

            (d).      Magnesium as Epsom Salt      (MgSO₄ . 7H₂O)

            (e).       Sodium as Glaber’s salt          (Na₂SO₄ . 10H₂O)

5. HALIDES:  These are minerals in which a metal occurs in combination with a halogen. Examples of halide minerals are:

            (a).       Sodium as common salt          (NaCl)

            (b).      Aluminium as Cryolite            (Na₃AlF₆)

            (c).       Calcium as Fluorite                 (CaF₂)

            (d).      Silver as Horn Silver               (AgCl)

            (e).       Magnesium as Carnallite         (KCl. MgCl₂ . 6H₂O)

6. PHOSPHATES:  The minerals in which a metal occurs as a phosphate are called phosphate minerals. Example;

                        Calcium as hydroxyapatite [Ca₃(PO₄)₅(OH)₂]

7. SILICATES:  The minerals in which the metal occurs as a silicate are called Silicate Minerals.

            For example.

            (a).       Silicon as Quartz (SiO₂)

            (b).      Calcium and Magnesium as Asbestos (CaSiO₃.MgSiO₃)

            (c).       Aluminium as China Clay (Al₂O₃.SiO₂.2H₂O)

Minerals and Ores

Majority of metals are present in earth’s crust in the combined form as oxides, sulphides, carbonates, phosphates etc. The Combined State of the metals is called Minerals. Thus a mineral may be defined as the Combined State in which a metal occurs naturally in the crust of earth. Minerals are generally associated with many earthy and rocky impurities which are called Gangue or Matrix. All the minerals of a particular metal cannot be treated for its isolation, because the process may be very tedious in certain cases and may not be profitable in some other cases. The minerals from which a metal can be profitably and conveniently extracted is called Ore.

Occurrence of Elements

Elements occur in nature in two states; Native State and Combined State.

1. NATIVE STATE   

The Elements are said to be in native state if they are found in their elementary form. Generally, less active elements are found in Native State. Oxygen, nitrogen, noble gases are some non – metals which occur in native state with the common examples of metals are gold, silver, copper, platinum etc.

2. COMBINED STATE   

The elements are said to occur in the Combined State if they are found in nature in the form of their compounds. Generally the reactive elements occur in the form of their compounds. For Example, the reactive metals are found in the form of their oxides, carbonates, Sulphide and silicates etc.

Sea as a Source of Elements

Sea water contains a large number of elements. The most common among them are sodium, potassium, calcium, chlorides and sulphate. The main source of these elements is the chemical weathering of igneous rocks followed by the extraction of soluble salts with water.

It may be noted that although sea water is a treasure house of valuable elements, yet only four elements namely: Chlorine, bromine, magnesium and sodium are commercially recovered from sea water.

In recent years new sources has been discovered on the ocean floor which is called Manganese Nodules. These nodules are rock like objects consisting of oxides of Mn, Fe and small amounts of CO, Cu and N. It is believed that marine organisms play a major role in the formation of such wodules.

Earth as a Source of Elements

Earth is a source of a large number of elements. The different parts of earth contains different elements. Earth can be broadly divided into three main parts; namely Atmosphere, Hydrosphere and Lithosphere.

1.         ATMOSPHERE

It consists of a gaseous mixture surrounding the earth. The main constituents of atmosphere are; nitrogen (» 78%), oxygen (» 21%) and other gases such as CO₂, He, Ne, Ar, Kr, Xe (less than 1%).

2.         HYDROSPHERE

It covers about 80% of earth’s surface and constitutes streams, lakes, rivers and oceans. There are large number of elements present in sea water. These are present in the form of their dissolved salts.

3.         LITHOSPHERE

It is a solid phase of the earth. It consists of different type of rocks such as igneous rocks, sedimentary rocks and metamorphic rocks. These rocks from a source of large number of elements.

The distribution of main elements on the surface of the earth in terms of their relative abundance is as follows;

Preparation of Ozone

2O₂ + O₂ ↔ 2O₃ ; DH = +288KJ

Effect of concentration
On increasing the concentration of oxygen, the equilibrium will shift toward right i.e. forward reaction is favoured. If the concentration of ozone is increased, the equilibrium will shift towards left i.e. backward reaction is favoured.

Effect of Temperature
The formation of ozone is endothermic reaction i.e. forward reaction is endothermic whereas backward reaction is exothermic reaction. So on increasing temperature forward reaction will be favoured and if temperature is decreased, backward reaction is favoured.

Effect of Pressure
On increasing pressure, forward reaction is favoured because it involves decreases in number of moles. If pressure is decreased, the backward reaction is favoured, because it involves increase in number of moles.

Relation Between Kp and Kc

Application of Le–Chatelier’s Principle

HABER’S PROCESS

One of the most important industrial reactions is the synthesis of ammonia from nitrogen and hydrogen, by Haber’s Process.

N₂    +   3H₂   ↔   2NH₃;    DH   =   – 92 KJ/mol.

Effect of Concentration:

According to Le – Chatelier’s Principle increases in concentration of N₂ or H₂ (reactants) will shift the equilibrium towards right i.e. forward reaction is favourable and hence production of NH₃ will be more. But if concentration of NH₃ is increased, then equilibrium will shift towards left i.e. backward reaction.

Effect of Temperature:

The above reaction is exothermic in forward direction and hence endothermic in backward direction. Thus according to Le – Chatliers Principle, increase in temperature will favor endothermic reaction and hence reaction will proceed in backward direction i.e. equilibrium will shift towards left. If temperature is lowered, the reaction will proceed in forward direction, because it is exothermic and hence equilibrium will shift towards right.

Effect of Pressure:

The forward reaction i.e. formation of NH₃ is accompanied with decreases in number of moles and backward reaction is accompanied with increase in number of moles.

According to Le – Chatelier’s Principle increase in pressure favour that reaction which is accompanied with decrease in number of moles and hence on increasing pressure equilibrium will shift toward right i.e. forward reaction. But if pressure is decreased, the favourable reaction is one which involves increase in number of moles. So on decreasing the pressure the reaction will favour backward reaction i.e. equilibrium will shift to left.

Le-Chatelier’s Principle

In 1884, the French Chemist Henery Le-Chatelier and Braun proposed a generalization known as Le – Chatelier’s Principle which states that, “If a system under equilibrium be subjected to a change in temperature, pressure or concentration, then the equilibrium shifts itself in such a way so as to undo or neutralize the effect of the change.”

1.         Effect of Concentration

According to Le-Chatlier’s Principle, when the concentration of one of the substance in a system at equilibrium is increased, then the equilibrium will shift so as to use up the substance added. Suppose at equilibrium if one of the reactant is added, the equilibrium will shift in forward direction i.e. reactant is consumed. On the other hand, if concentration of product is increased, the equilibrium will shift in the backward direction, because it consumes the product.

Thus increase in concentration of any of the reactants shifts the equilibrium towards forward direction and increase in concentration of any of products shifts the equilibrium towards backward direction.

2.         Effect of Temperature

According to Le – Chatlier’s Principle, when temperature of the system is changed (increased or decreased), the equilibrium shifts in such a direction so as to undo the effect.

If the temperature is increased, the equilibrium will get shifted in direction which is accompanied by decrease in temperature i.e. endothermic reaction. If the temperature is decreased, the equilibrium shifts in that direction which is accompanied by increase in temperature i.e. exothermic reaction.

Thus, Increase in temperature, favours endothermic reaction.

Decrease in temperature, favours exothermic reaction.

3.         Effect of Pressure

According to Le – Chatlier’s Principle, increase of external pressure should effect the equilibrium in such a way as to reduce pressure. This means increase in pressure favours the reaction which is accompanied with decrease in number of moles. Whereas, decrease in pressure favours the reaction which is accompanied with increase in number of moles.

The change of pressure has no effect on those equilibria where number of moles of reactants and number of moles of products is same.

Law of Chemical Equilibrium

Law of chemical equilibrium is direct result of law of mass action applied to reversible chemical reaction.

Consider a general reversible reaction,

aA    +   bB    ↔   cC  +  dD.

Applying law of mass action to both reactions i.e. forward reaction and backward reaction.

ü    FOR FORWARD REACTION

aA  +  bB  →  cC  +  dD

\      Rate of forward reaction, r_f   µ  [A]^a[B]^b 

or            r_f   =  K_f [A]^a[B]^b   ––––––– (i).                  

Where [A], [B] are molar conc. of A and B, and K_f is rate constant for forward reaction.

ü    FOR BACKWARD REACTION  

cC  +  dD    →  aA    +    bB

\     Rate of backward reaction, r_b  µ   [C]^c[D]^d

or             r_b   =   K_b[C]^c[D]^d  ––––––– (ii).

Where [C], [D] and K_b are molar concentrations of A, B and rate constant of backward reaction respectively.

ü    AT EQUILIBRIUM

Rate of forward reaction, r_f   =   Rate of backward reaction, r_b

Using (i) and (ii) we get,

\ r_f   =   r_b

K_f [A]^a[B]^b   =    K_b[C]^c[D]^d

K_f / K_b = [C]^c[D]^d / [A]^a[B]^b 

or         K_C  = [C]^c[D]^d / [A]^a[B]^b, where K_c  =  K_f / K_b

 ‘K_C’ is known as Equilibrium Constant in terms of concentrations.

From the above equation, Law of chemical equilibrium may be stated as, “At constant temperature, for a reversible chemical reaction in equilibrium, the ratio of product of molar concentrations of products to that of molar concentration of reactants each term is raised to a power of stoichiometric coefficient as represented by balanced chemical reaction.”

In a chemical reaction if reactants and products are in gaseous phase, then partial pressures of reactants and products are considered instead of molar concentrations. Under such conditions equilibrium constant is represented by ‘K_P’ (i.e. equilibrium constant in terms of partial pressure), For example

aA (g)    +    bB (g)   ↔   cC (g)   +    dD (g)

K_P    = [P_C] ^c [P_D] ^d / [P_A] ^a [P_B] ^b

Where P_A, P_B, P_C and P_D are partial pressures of A, B, C and D.

The concentration ratio

[C]^c[D]^d / [A]^a[B]^b  is called Concentration Quotient or Reaction Quotient.

Law of Mass Action

The law of mass action was given by two Norwegian Chemists Gulberg and Waage in 1867 which states,

 “At constant temperature the rate of a chemical reaction is directly proportional to the product of the molar concentrations of reactants each term raised to the power equal to the Stoichiometric co-efficient as represented by balanced chemical equation.”

Let us consider a chemical reaction

aA    +   bB    →  Product.

Where a, b are stoichiometric co-efficients.

Then,

According to law of mass action,

Rate of reaction, (r) µ  [A]^a[B]^b

Where [A] and [B] are molar concentration of A and B.

\        r   =   K[A]^a[B]^b

Characteristics of Chemical Equilibrium

 The various characteristics of chemical equilibrium are:

(i).      The observable properties of the system become constant at equilibrium and remain unchanged thereafter. 

(ii).      The equilibrium is obtained only in the reactions carried in a closed vessels.

(iii).     The equilibrium is dynamic in nature.

(iv).     A catalyst cannot alter the equilibrium point.

(v).      The free energy change at equilibrium state is zero.

Equilibrium in Chemical Processes

Decomposition of Calcium Carbonate

Calcium carbonate when heated to 1073K in a closed evacuated vessel, starts decomposing to yield calcium oxide and carbon dioxide gas. Carbon dioxide builds up pressure within the vessel which can be recorded on a manometer. The pressure goes on increasing as the reaction proceeds and finally becomes constant. It appears as if the reaction has come to a stop although CaCO₃ is still present. This indicates that the system has attained the equilibrium state.

At equilibrium, rate of forward reaction i.e.

CaCO₃   →  CaO  +   CO₂

becomes equal to rate of  backward reaction i.e.

CaO   +   CO₂  →  CaCO₃

CaCO₃     ↔   CaO  +  CO₂  

Equilibrium in Physical Processes

·                    Solid – Liquid Equilibrium

When a solid – liquid system at melting point is taken in a well – insulated container, then this system constitute a system in which solid is in dynamic equilibrium with liquid.

For Example:

Let us consider ice and water at 273K (melting point of ice) taken in a perfectly insulated thermos flask. It may be noted that temperature as well as masses of ice and water remain constant. This represents a dynamic equilibrium between ice and water.

Ice   ↔    Water

Since there is no change in mass of ice and water, the number of molecules going from ice into water is equal to number of molecules going from water into ice. Thus, at equilibrium,

Rate of melting   =    Rate of freezing.

·                    Liquid – Gas Equilibrium

Let us consider evaporation of water in a closed vessel. When a small amount of water is taken in an evacuated vessel at room temperature it starts evaporating. This process continues for some time as indicated by the gradual decrease in level of water.

After some time it is observed that level of water becomes constant indicating that no more water is evaporating. This indicates that a state of equilibrium has attained between water and water vapours.

H₂O (l)      ↔      H₂O (g).

The equilibrium seems to be static, but in actual practice it is dynamic. In beginning rate of evaporation is more and hence water vapours concentration increases, which in turn condenses back into the liquid. As the concentration of water vapours increases, rate of condensation also increases. At equilibrium, rate of evaporation is equal to rate of condensation.

\    Rate of evaporation   =   Rate of condensation.

Dynamic Nature of Equilibrium

The observable properties of the system become constant at equilibrium, which leads us to think that reaction stops altogether at equilibrium. But it is not true, actually at equilibrium the rate of forward reaction becomes equal to rate of backward reaction and hence there is no net change in concentration of various species. This can be illustrated by considering the reaction between H₂ and I₂ to form hydrogen iodide (HI).

When H₂ and I₂ are taken in a closed vessel maintained at 717K, hydrogen molecules combine with iodine molecules to form hydrogen iodide.

H₂ (g)   +   I₂ (g)    →  2HI (g).

The molecules of HI formed, begins to dissociate back to H₂ and I₂.

2HI (g)  →   H₂(g)    +   I₂ (g).

As the reaction progresses, the concentration of H₂ and I₂ decreases and hence rate of forward reaction slows down. On the other hand, concentration of hydrogen iodide increases and therefore rate of backward reaction increases. A stage is reached when rate of forward reaction and rate of backward reaction becomes equal and the system attains equilibrium. At equilibrium no change in concentration occurs provided the temperature of reaction mixture remains constant. Thus, at equilibrium the reaction does not stop but the system attains constant observable properties because of equal rates of forward reaction and backward reaction. Thus, equilibrium is Dynamic in Nature.

The dynamic nature of chemical equilibrium can be demonstrated by adding small amount of radioactive iodine (I₂*) to above reaction in equilibrium. It is observed that after some time hydrogen iodide also contains some radioactive iodine (I₂*). It indicates that reaction is going on even at equilibrium. If the equilibrium being static then hydrogen iodide should not contain radioactive iodine.

H₂   +    I₂*  ↔ 2HI*

Thus, it demonstrates the dynamic nature of chemical equilibrium.