Structures of functional groups

In alcohols, the oxygen of the -OH group is attached to sp³ hybridised carbon by a sigma bond formed by the overlap of sp³ hybrid orbital of carbon with an sp³ hybrid orbital of oxygen. The figure shown below illustrates the bonding in methanol.
The C-O-H bond angle in alcohol is slightly less than the tetrahedral angle (109⁰28). It is due to the repulsion between the unshared electron pairs of oxygen.

In phenols, the –OH group is attached to sp² hybrid carbon of an aromatic ring. The C-O-H bond angle in phenol is 109⁰. The carbon-oxygen bond length (136pm) in phenol is slightly less than that in methanol (142pm). This is due to partial double bond character on account of the conjugation of unshared electron pair of oxygen with the aromatic ring.

In ethers, the four electron pairs, ie; the two bond pairs and two lone pairs of electrons around oxygen are arranged approximately in a tetrahedral arrangement. The c-o-c bond angle (111.7⁰ in methoxy methane) is slightly greater than the tetrahedral angle (109⁰28) due to the repulsive C-O bond length (141 pm) in ethers is almost the same as in alcohols (142 pm in methanol).

pH of Buffer solution

The pH of acidic and basic buffer can be calculated by Henderson – Hasselbalch equations. Consider an acidic buffer HA + A^-
HA <=======> H⁺ + A^-
K_a = [H⁺] [A^-] / [HA]
[H⁺] = K_a [HA]/[A^-]
[H⁺] = K_a [acid]/[salt]
There fore pH = -log[H⁺]
pH = pK_a + log [salt]/[acid]
when, [salt]/[acid] = 1 ,
pH = pK_a
Since pK_a of an acid is a constant at constant temperature, the pH of the buffer is constant. Thus buffer capacity is maximum in a solution containing equivalent amount of acid and its salt.
The pH of basic buffer is also given by Henderson – Hasselbalch equation
BOH <=======> B⁺ + OH^-
K_b = [B⁺][OH^-]/[BOH]
[OH^-] = K_b [BOH]/[B⁺]
pOH = pK_b + log [salt]/[base]
pH = 14 – pOH
= 14 – [pK_b + log [salt]/[base]]

Buffer action

The property of a buffer solution to resist change in its pH value even when small amounts of the acid or the base are added to it is called buffer action.
Consider the acidic buffer solution containing acetic acid and sodium acetate. They dissociate as
CH3COONa <=======> CH3COO_- + H₊
CH3COONa <=======> CH3COO_- + Na₊
When a few drops of an acid, HCl is added to this buffer solution, the H₊ ions combine with CH3COO_- ions to form weakly ionized molecules of CH3COOH.
CH3COO_- + H₊ <=======> CH3COOH
Thus H₊ ion concentration does not change and hence the pH of the solution remains constant.
When a few drops of base, NaOH is added to the buffer solution, hydroxyl ions of the base neutralize the acid, forming salt and water.
Similarly, in a basic buffer solution of NH4OH and NH4Cl, they dissociates as
NH4OH <======> NH4₊ + OH_-
NH4Cl ----------> NH4₊ + Cl_-
When a few drops of a base added, the OH_- ions given by it combine with NH4₊ ions to form the weakly ionized NH4OH.
NH4₊ + OH_- ---------> NH4OH
Thus the OH_- ion concentration or the pH of the solution remains unaffected.
When a small amount of an acid is added, the H₊ ions given by it combines with the OH_- ions already produced by NH4OH.
H₊ + OH_- --------> H2O
Therefore the H₊ ions concentration or the pH of the solution remains unaffected.
The buffer capacity of a buffer solution is defined as the number of moles of acid or base added per liter of the solution to change the pH by one unit.

Buffer solutions

Maintenance of PH in blood and in intracellular fluids is absolutely crucial to the processes that occur in living organisms. This is primarily because the functioning of enzymes is sharply pH dependent. The normal pH value of blood plasma is 7.4 and several illness or death can result from sustained variations of a few tenths of pH unit. Also many medical and cosmetic formulations require that these must be kept and administered at a particular pH. There are solutions which resist the change in pH on addition of small amount of acid or alkali and are called Buffer solution. For example a mixture of H₂CO₃ and HCO₃- is a natural buffer system which maintains the pH of blood. A buffer that is widely used in clinical laboratory and in biochemical studies in the physiological pH range is prepared from tris amino methane (hydroxy methyl) (THAM) [(HOCH₂)3CNH₂].
In order for a solution to act as a buffer it must have two components, one of which is able to neutralize acid and the other able to neutralize the base. Common buffer solutions are mixtures containing a
Weak acid and its conjugate base (one of its salt) called acidic buffer
eg:- CH₃OOH/CH₃COONa, H₂CO₃/Na₂CO₃, Boric acid/borax
Weak base and its conjugate acid (one of its salt) called basic buffer
Eg:- NH4OH/NH4Cl, Zinc hydroxide/ zinc chloride, Glycine/ glycine hydrochloride.

Application of solubility product and common ion effect

Qualitative analysis of cations is largely based on the principle of solubility product and common ion effect. Cations are separated in to six groups depending on the solubility of their salts.
Group-1 as insoluble chlorides
Only Ag⁺, Hg²⁺ and Pb²⁺ form insoluble chlorides since they have low values of K_sp.
Group-2 as insoluble sulphide in acidic medium
H₂S <========> H⁺ + HS^- ;
K₁ - first ionization constant
HS^- <========> H⁺ + S^2-;
K₂ – second ionization constant
[S^2-] = K₁K₂ [H₂S]/[H⁺]²
K_sp values of second group sulphides (PbS, CuS, SnS, HgS, As₂S₃, Bi₂S₃, Sb₂S₃) are very low. In acidic buffer, [S^2-] is decreased due to common ion effect and this results in the precipitation of Pb²⁺, Cu²⁺ etc of second group as their sulphides. Third and fourth group sulphides have high value of K_sp, hence they remain soluble.
Group- 3 as insoluble hydroxide in basic buffer of NH₄OH and NH₄Cl
The concentration of OH^- in ammoniacal solution decreases when NH₄Cl is added to it, because of the common ion effect. Thus only for least soluble hydroxide of Fe³⁺, Al³⁺, Cr³⁺ etc. ionic product exceeds the corresponding solubility products, hence only these ions are precipitated. Hydroxides of successive remain soluble due to high K_sp values.

The solubility product constant (Ksp)

The solubilityof ionic solids in water varies depending on a number of factors like lattice enthalpy of the salt and tha solvation enthalpy of the ions in a solution. As a general rule, for a salt to be able to dissolve in a particular solvent, its solvation enthalpy must be greater than its lattice enthalpy. Each salt has its characteristic solubility, which depends on temperature. We can classify salts on the basis of their solubility in three categories.
Soluble - Solubility > 0.1 M
Slightly soluble - 0.01 M < solubility < 0.1M
Sparingly soluble – solubility < 0.01M
We have now consider the equilibrium between the sparingly soluble ionic salt and its saturated aqueous solution. A solution which remains in contact with excess of the solute is said to be saturated. The amount of a solvent () in 100 ml or 1L) to form a saturated solution at a given temperature is termed as the solubility of the solute in the solvent at that temperature.
For a sparingly soluble salts like AgCl,PbI₂, BaSO₄ etc. Ionisation is very small and concentration of the salt may be considered as constant. Thus, for AgCl,
AgCl <========> Ag⁺ + Cl^-
K = [Ag⁺][Cl^-]/[AgCl]
For a pure solid substance the concentration remains constant and we can write
K_sp = K[AgCl]
=[Ag⁺][Cl^-]
Where K_sp is the solubility product constant or solubility product.
If S represents solubility of AgCl (in mol L^-1)
Then [Ag⁺]= [Cl^-] = S molL^-1
K_sp = [Ag⁺][cl^-] = S²
S = √K_sp
A solubility product constant expression is the product of the chemical equation for a solubility equilibrium, with each term raised to the power given by the coefficient in the chemical equation.
Example:-
For BaSO₄ (binary solute giving two ions)
BaSO₄ <=======> Ba²⁺+ SO₄2^-
Ksp = [Ba²⁺][SO₄2^-] = S2
For PbI₂ (Ternary solute giving three ions)
PbI₂ <=======> Pb²⁺ + 2I^-
Ksp = [Pb²⁺][I^-]2 = (s)(2s)2 = 4S³
For Al(OH)₃ (Quaternary solute giving four ions)
Al(OH)₃ <=======> Al³⁺ + 3OH^-
K_sp = [Al³⁺][OH^-]³ = (S)(3S)³ = 27Ss4
For a solute AxBy, (giving (x+y) ions)
K_sp = [A^y+]^x [B^x-]^y = (xS)^x (yS)^y = X^xY^y(S)^(x+y)
Precipitation reactions
For the sparingly soluble solute AB,
AB <=======> A⁺ +B^-
Q is the equilibrium constant with the given [A⁺] and [B^-]
Q = [A⁺][B^-]
When applied to solubility product, Q is generally called the ionic product. Precipitation occurs if Q>K_sp, then solution is said to be super saturated and precipitation cannot occur if Q_sp (a case of unsaturated solution) and a solution is just saturated if Q=K_sp.

Common ion effect

The decrease in the ionization of a weak electrolyte by the presence of a common-ion from a strong electrolyte, is called the common ion effect. Ionisation of CH₃COOH (weak acid) is decreased by the addition of CH₃COONa (CH₃COO- being the common ion)
CH₃COOH <======> CH₃COO^- + H⁺ ………………………..(A)
CH₃COONa -----------> CH₃COO^- + Na⁺
In the presence of CH₃COO^- equilibrium (A) shifts in backward direction.
Ionisation of H₂S (weak acid) is decreased by the addition of HCl (H⁺ being the common ion)
H₂S <=======> 2H⁺ + s^2-
HCl <=======> H⁺ + Cl^-
Ionisation of NH₄OH (weak base) is decreased by the addition of NH₄Cl (NH4 ⁺ being the common ion)
NH₄OH <=======> NH₄⁺ + OH^-
NH₄Cl --------------> NH₄⁺ + Cl^-
Solubility of a sparingly soluble salt is decreased by the addition of common ion. Presence of AgNO₃ or KCl decreases the solubility of AgCl.
AgCl <=======> Ag⁺ + Cl^-
AgNO₃ <========> Ag⁺ (common ion) + NO₃^-
KCl <=======> K⁺ + Cl^-(common ion)
The common ion effect is thus based on Le-Chatelier’s principle in which the stress on the equilibrium that results from raising one of the product concentrations is relieved by shifting the equilibrium to left.

Application of Kohlrausch's law

Determination of λ0M (limiting molar conductivity) for weak electrolytes

It is not possible to determine the λ⁰M for a weak electrolyte by the extrapolation of λ _M versus √c plot. But it can be calculated from Kohlrausch’s law. Consider the weak electrolyte CH3COOH. According to Kohlraush’s law,

λ⁰_CH3COOH = λ⁰_H+ + λ⁰_CH3COO- ………………………………… (Equation 1)

We can experimentally determine the λ⁰M of strong electrolytes such as HCl,CH3COONa and NaCl.
From Kohlraush’s law,

λ⁰_HCl = λ⁰_H+ + λ⁰_Cl- ……………………………………. (Equation 2)

λ⁰_CH3COONa = λ⁰_Na+ + λ⁰_CH3COO- ………………………………… (Equation 3)

λ⁰_NaCl = λ⁰_Na+ + λ⁰_Cl- ……………………………………. (Equation 4)

Substracting equation 4 from the sum of equations 2 and 3, we get,

λ⁰_HCl + λ⁰_CH3COONa - λ⁰_NaCl = λ⁰_H+ + λ⁰_Cl- + λ⁰_CH3COO- + λ⁰_Na+ - λ⁰_Na+ - λ⁰_Cl-

= λ⁰_H+ + λ⁰_CH3COO-

= λ⁰_CH3COOH

Thus by measuring the molar conductance values of NaCl, HCl and CH3COONa, one can easily determine the λ⁰_M (limiting molar conductivity)

Determination of degree of dissociation of weak electrolytes

The molar conductivity of a weak electrolyte depends up on the extent of dissociation. As the dilution increases, the degree of dissociation increases. The degree of dissociation of a weak electrolyte can be calculated by the relation :
Degree of dissociation α = λ^C_M / λ⁰_M
Where λ^C_M is the molar conductivity at concentration ‘C’ and λ⁰_M is the limiting molar conductance of the weak electrolyte. The limiting molar conductance can be calculated by Kohlrausch’s law.


Related article Kohlrausch's law definition

Kohlrausch's law definition

Kohlrausch's law of indendent migration of ions states that the molar conductance of infinite dilution (limiting molar conductivity ) is the sum of the individual contributions of the anions and cations of the electrolyte.

It can be given in the mathematical form as:
λ⁰_M = v₊λ⁰₊ + v_-λ⁰_-
λ⁰₊ and λ⁰_- are the limiting molar conductivity of cation and anion respectively.
For NaCl,
λ⁰_M = λ⁰_Na+ + λ⁰_Cl-
And for Al2(SO4)3,
λ⁰ _Al2(SO4)3 =2 λ⁰_Al3+ + λ⁰_{SO4 2-}

Phosphine gas (PH3)

In the laboratory phosphine is prepared by heating white phosphorus with concentrated caustic alkali solution in an inert atmosphere of oil gas or CO2.
P4 + 3NaOH + 3H2O ------------> 3NaH2PO2 (Sodium hypophosphite) + PH3 (phosphine)
Metal phosphides on hydrolysis form phosphine
Ca3P2 + 6H2O ------------> 3Ca(OH)2 + 2PH3
A pure sample of phosphine can also be prepared by heating phosphorous acid
4H3PO3 ---------------> 3H3PO4 + PH3
Phosphine has pyramidal structure and is a weaker base than NH3

Oxides Of Phosphorus

The main oxides of phosphorus are phosphorus trioxide (P4O6) and phosphorus pentoxide (P4O10). Phosphorous trioxide is regarded as the anhydride of phosphorous acid (H3PO4). phosphorus trioxide (P4O6) is prepared by heating phosphorous in limited supply of oxygen. P4O10 is prepared by burning white phosphorous in excess of air or oxygen.
P4 + 3O2 -------------> P4O6
P4 + 5 O2 ------------> P4O10
P4O10 has great affinity for water and hence it is used as a dehydrating agent. It can dehydrate HNO3 and H2SO4 to yield N2O5 and SO3 respectively.
2H2SO4 + P4O10 ------------> 2SO3 + 4HPO3
P4O6 and P4O10 dissolves in water to give phosphorus acid and orthophosphoric acid respectively.
P4O6 + 6H2O -------------> 4H3PO3
P4O10 + 6H2O --------------> 4H3PO4

Related article oxyacids of phosphorus

Extraction of sulphur

Sulphur occurs in nature in the elemental forms, as metal sulphides and as sulphates. Sulphur also occurs as H2S present in natural gas.
Sulphur is extracted by the following methods
1. Frasch process
In this process, sulphur is extracted from underground deposits by pumping super heated steam down the beds to melt the sulphur and then blown out the molten sulphur with compressed air.
2. Extraction of sulphur from natural gas
Natural gas contains a large amount of hydrogen sulphides (H2S). Hydrogen sulphide is first absorbed in monomethanolamine and then converts H2S into sulphur by the following reactions.
2H2S + 3O2 ------------> 2SO2 + 2H2O
2H2S + SO2 ------------> 3/8 S8 + 2H2O (at 673K and Fe2O3 as catalyst)

Related article extraction of aluminium
extraction of copper

Oxides of Xenon ( XeO3 and XeO4)

Xenon trioxide (XeO3)
XeO3 prepared by the slow hydrolysis of XeF6
XeF6 + 3H2O ------------> XeO3 + 6HF
Xenon trioxide is soluble in water and its aqueous solution is weakly acidic.
XeO3 + H2O <--------> H+ + HXeO4 – Xenate ion
XeO3 has pyramidal structrure in which Xe is in sp3 hybridisation.


Xenon tetroxide (XeO4)
It is prepared by treating barium perxenate (Ba2XeO6) with anhydrous sulphuric acid.
Ba2XeO6 + 2H2SO4 -----------> XeO4 + 2BaSO4 + 2H2O
Xenon tetroxide is highly unstable and has tetrahedral structure.

Related article fluorides of xenon
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Hydrogen sulphide (H2S)

Hydrogen sulphide is prepared in laboratory by the action of dilute HCl or dilute H2SO4 on ferrous sulphide in Kipp’s apparatus.

FeS + H2SO4 ---------->4 + H2S

Physical properties

It is colourless gas with the smell of rotten eggs. It is denser than air and is soluble in water.

Chemical properties

1. Combustibility

Hydrogen sulphide burns in limited supply of oxygen to form sulphur. In presence of excess of oxygen, it gives sulphur dioxide.

2. Acidic property

H2S is a weak dibasic acid and forms two types of salts namely bisulphides and sulphides

NaOH + H2S ----------> NaHS + H2O

2NaOH + H2S ----------> Na2S + 2H2O

3. Action with metals

When H2S is passed over hot metals the sulphides and hydrogen are formed

Cu + H2S -----------> CuS + H2

4. Hydrogen sulphide in qualitative analysis

H2S precipitates metal sulphides having characteristic colours from metal salt solutions in acidic or alkaline medium. Therefore H2S is used in qualitative analysis to identify metal ions belonging to different classes.

In acidic solution, it precipitates group 2 cations (Hg2+, Cu2+, Pb2+, Bi3+, Sb2+, Sn2+, As3+, Cd2+ and Sn4+) as their sulphides with characteristic (Co2+, Ni2+,Zn2+ and Mn2+) as their sulphides with characteristic colours.

Related article Sulphur

Sulphur dioxide and Sulphur trioxide images

Related article sulphur dioxide

Pictrures of oxyacids of chlorine

For more details about oxyacids of chlorine

Sodiumthiosulphates (Na2S2O3.5H2O)

It is prepared by boiling sulphur with an aqueous solution of sodium sulphate
Na2SO3 + 1/8 S8 ------> Na2S2O3
Properties
Sodium thiosulphate is a water soluble crystalline substance.
a) Reaction with iodine
Sodium thiosulphate is oxidised by iodine to form sodium tetra thionate. This reaction is the basis of iodometric titrations.
2Na2S2O3 + I2 ---------> Na2S4O6 + 2NaI
b) Reaction with chlorine
Sodium thiosulphate can remove excess chlorine by forming HCl. Hence it is used as an antichlor.
Na2S2O3 + Cl2 + H2O -----------> Na2SO4 + 2HCl + S
c) Reaction with silver halides
Silver halides dissolve in sodium thiosulphate solution due to formation of a complex sodium argento thiosulphate. This reaction is the basis of its use in photography as fixer.
AgCl + Na2S2O3 -------------> NaAgS2O3 + NaCl

Oxides of sulphur

The stable oxides of sulphur are sulphur dioxide (SO2) and sulphur trioxide (SO3)
Sulphur when burnt in air form sulphur dioxide.
S8 + 8O2 -------------> 8SO2
SO2 is a gas at room temperature. It exists as descrete SO2 molecule with angular structure. SO3 is an acidic oxide and exist as planar triangular molecule in gas phase. In solid state, SO3 exist either as cyclic trimer or linear chain polymer.

Bleaching powder (CaOCl2)

Bleaching powder is a mixture of calcium hypochlorite, Ca(OCl)₂ and calcium chloride CaCl₂. Thus, its chemical composition is Ca(OCl)₂ + CaCl₂ or can be written as CaOCl₂ or Ca(OCl)Cl (Calcium chloro hypochlorite).

Bleaching powder is manufactured by passing chlorine gas over dry slaked lime at 40⁰ C.

Ca(OH)₂ + Cl₂  ← CaOCl₂ + H₂O

Properties

1. It is an yellowish white powder with strong smell of chlorine and is soluble in water. Aqueous solution of bleaching powder contains chloride and hypochlorite ions.

CaOCl₂  ↔ Ca ²⁺ + Cl^- + ClO^-

2. In presence of small quantities of dilute acids, it liberates nascent oxygen. Thus it act as oxidizing agent.

2CaOCl₂ + H2SO₄  → CaCl₂ + CaSO₄ + 2HClO

HClO ← HCl + [O]

3. When bleaching powder is treated with excess of dilute acid or CO2, whole of the chlorine present in the molecules is evolved. The amount chlorine so liberated is called available chlorine. A good sample of bleaching powder contains 35-38% available chlorine.

CaCl₂ + 2HCl → CaCl₂ + H₂O + Cl₂

CaOCl₂ + CO₂  → CaCO₃ + Cl₂

Nitrogen

Nitrogen is the first element of group 15 of the periodic table and has the electronic configuration 1s2 2s2 2p3. Like oxygen and hydrogen, nitrogen exists in its elemental forms as a diatomic molecule. Nitrogen forms a variety of compounds in all oxidation states ranging from -3 to +5.
The common oxidation states are -3, +3 and +5. Molecular nitrogen comprises 78% by volume of the atmosphere. Dinitrogen is a colourless, odourless, tasteless gas. It has two stable isotopes, 14N and 15N. It is non toxic. Dinitrogen is chemically unreactive a at ordinary temperature. At higher temperatures dinitrogen directly combines with some metals and nonmetals to form ionic and covalent nitrides. Some examples are

6Li + N2 --------> 2Li3N
3Mg + N2 --------> Mg3N2

In laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.

NH4Cl + NaNO2 -------> N2 + 2H2O + NaCl

Uses of dinitrogen
Dinitrogen is mainly used in the manufacture of ammonia and industrial nitrogen chemicals (Eg:- calcium cyanide). It is also finds use where the presence of an inert atmosphere is required. Liquid nitrogen is used as a refrigerant to preserve biological materials.

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Transuranium elements

Uranium is the last element occurring in nature. The elements coming after uranium are called transuranium elements. They are all synthetic in nature and are man made. They are radioactive and most of them have short half lives.

For example :-

²³⁸U₉₂ + ¹n₀ → ²³⁹ U₉₂ → ²³⁹ Np₉₃ + ⁰e_-1

Uranium → Neptunium

²³⁹ Np₉₃ → ²³⁹ Pu₉₄ + ⁰e_-1

Neptunium → Plutonium

( Here the Uranium(238,92) is abundant and non fissionable, It can be made into fissionable by reaction given above )

The element beyond actinides in the periodic table form atomic number 104 to 112 are called transactinide elements. These elements are rutherfordium (Rf), dubnium (Db), Seaborgium (Sg) etc. The nuclear reactions for the preparation of some of the transactinide elements are given below.

²⁴⁹Cf₉₈ + ¹²C₆ → ²⁵⁷Rf ₁₀₄+ 4 ¹n₀

Californium → Rutherfordium

²⁴⁹Cf₉₈ + ¹⁵N₇ → ²⁵⁷Db ₁₀₄+ 4 ¹n₀

Californium → Dubinium

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Chemicals in food

Many chemical substances such as flavourings, sweeteners, dyes, antioxidants, fortifiers, emulsifiers and antifoaming agents are added to food for their preservation and enhancing their appeal. With the exception of preservatives, fortifying agents, antioxidants and artificial sweeteners. The other classes of chemicals are added either for ease in processing or for cosmetic purposes. In real sense, these substance have no nutritive value.

Preservatives
Usually food is most appetizing when it comes from the production line in the food processing plant. But, during storage and distribution undesirable changes occur in flavour, colour, texture and appetitic appeal. To delay these changes many chemical substances are added to food. These chemicals are called food preservatives. They prevent spoilage of food due to microbial growth. commonly used preservative is sodium benzoate (C6H5COONa).In order to preserve colourless food materials potassium metabisulphite or sodium metabisulphite is used.

Artificial sweetening agents
These are another types of food additives which are non-nutritive substituents for sugar. The first popular artificial sweetener was saccharin. It was marketed as its water-soluble sodium or calcium salt. Saccharin is approximately 300 times sweeter than cane sugar.
Besides saccharin, the other commonly marketed artificial sweeteners are aspartame, alitame, sucralose etc. Aspartame is unstable at cooking temperatures, limiting its use as a sugar substitute to cold foods and soft drinks. Alitame is more stable than aspartame during cooking. One potential problem with alitame and similar type of high-potency sweeteners is the difficulty in controlling sweetness of food.

Antioxidants
These are the important food additives. Antioxidants are compounds which retard the action of oxygen on the food. It helps on the preservation of the food materials. These are more reactive towards oxygen than the food products. They also reduce the rate of involvement of free radicals in the ageing process of food. The two commonly used antioxidants are butylated hydroxy toluene (BHT) and butylated hydroxy anisole (BHA). Vitamin E is a natural antioxydent.

Edible colour
Edible colours used for food are essentially dyes. These are used to colour food products. Some of the azodyes using as food dyes are suspected to be dangerous for asthma patients and young children. However, natural dyes like carotene are safe food dyes.

Refining of metals

The metals prepared by different methods contain impurities. The methods used for the purification of metals are called refining. The refining method depends on the nature of metal and the nature of impurities. Some common methods are as follows
1. Distillation
Volatile metals like zinc and mercury are purified by boiling the impure metal to get vapors of the pure metal which is condensed and collected.
2. Liquation
Low melting metals like tin and lead are purified by this method. The impure metal is melted on the sloping floor of a furnace. The metal melts and flows down leaving behind the high melting impurities.
3. Poling
Impure metal is melted and stirred with green logs of wood. The impurities rise to the surfaces, get oxidised and removed as gases (CO2) or slag. The metal may get oxidised (eg:- Cu to Cu2O). The hydrocarbons in green wood reduces the metal oxide to the metal. Example:- Refining of impure Cu and Sn.
4. Cupellation
Impure silver and gold contain base metals like lead and bismuth as impurities. These are removed by heating the metal placed in a cupel (boat shaped crucible made of cement or bone ash) in a reverberatory furnace in a current of hot air. The impurities are oxidised and carried away by the current of air. The process is stopped when a shining bright surface appears.
5. Electrolytic refining
The method is based on the process of electrolysis. The crude metal is made the anode a thin sheet of pure metal the cathode. The electrolyte is the solution of a salt of the metal. On passing electricity the metal dissolves from the anode and an equal number of metal ions of the solution gets deposited at the cathode. The impurities settle down below the anode mud (eg:- refining of Cu using CuSO4 solution as electrolyte).
At anode Cu ------> Cu 2+ + 2electron
At cathode Cu 2+ + 2electron -------> Cu
6. Zone refining
Metals of very high purity can be obtained by this method. The impure metal rod is mounted horizontally and heated by a circular electric heater at one end in an atmosphere of inert gas to form a thin molten zone. By slowly moving the heater, the molten zone is moved from one end of the rod to the other end. Pure metal crystallises while impurities pass into the molten zone. By repeated the passes of the molten zone very high purity can be attained at one end. The other end is discarded. Ge, Si and Ga used in semi conductors are refined in this manner.
7. Vapour phase refining
By this chemical reaction the metal is converted to a compound, which forms a vapour, which is decomposed to get pure metal.
8. Monds process
Crude nickel is heated with carbon monoxide to form volatile nickel carbonyl. The impurities remain as solids. The vapours on further heating decomposes to give pure nickel.
Ni + 4Co -------> Ni(CO)4
Ni(CO)4 --------> Ni + 4CO
9. Van Arkel de Boer method
Crude titanium or zirconium is heated with iodine to get vapours of the tetraiodide. The vapours are then decomposed on a tungsten filament kept at high temperature on which the pure metal gets deposited. The pure metal is then peeled off from the filament after cooling.
Ti + 2I2 ---525k----> TiI4 ----1675k----> Ti + 2I2
Zr + 2I2 ----870k---> ZrI4 ----2075k----> Zr + 2I2

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Mercury halides

Mercury forms halides in the two oxidation states, +1 and +2.
1. Mercury(1)Chloride(Hg2Cl2)
Mercury(1) chloride or mercurous chloride is known as calomel.
Preparation
Mercury(1)chloride is prepared by heating a mixture of mecury(2)chloride and mercury in iron vessel.
HgCl2 + Hg -------> Hg2Cl2
It can also be obtained by reduction of mercury(2)chloride by reducing agents like tin(2)chloride in limited quantity.
2HgCl2 + SnCl2 --------> HgCl2 + SnCl4

Properties
When heated, mercury(1)chloride decomposes into mercury(2)chloride and mercury.
Hg2Cl2 -------> HgCl2 + Hg
The action of aqueous ammonia on the solid mercury(1)chloride gives a mixture of black finely divided mercury and white mercury amino chloride. This reaction is an example of disproportion reaction.
Hg2Cl2 + 2Nh3 -------> Hg(NH2)Cl + Hg + NH4Cl

Uses of mercury(1)chloride
Calomel is used in making standard calomel electrodes used as secondary reference electrode. It is also used as a purgative in medicines.

2. Mercury(2)chloride (HgCl2)
Mercury(2)chloride or mercury chloride is known as corrosive sublimate.
Preparation
Mercury(2)chloride may be prepared by heating the metal in chlorine gas.
Hg + Cl -------> HgCl2
It is also prepared by heating a mixture of mercury(2)sulphate and sodium chloride in the presence of traces of MnO2
HgSO4 + 2NaCl -------> HgCl2 + Na2SO4
Manganese dioxide prevents the formation of mercury(1)chloride.

Properties
Mercury(2)chloride is a white crystalline solid, but from aqueous solution it crystallizes into colourless needles. It is a covalent compound sparingly soluble in water. Mercury(2)chloride gives a white precipitate on reduction with stannous chloride, SO2, formaldehyde etc, which changes to grey on standing owing to the formation of metallic mercury.
2HgCl2 + SnCl2 -------> Hg2Cl2 + SnCl4
Hg2Cl2 + SnCl2 -------> 2Hg + SnCl4
Mercury(2)chloride reacts with aqueous ammonia to form infusible white precipitate of mercury amino chloride.
HgCl2 + 2NH3 --------> Hg(NH2)Cl + NH2Cl
Gaseous ammonia or ammonium chloride on reaction with Mercury(2)chloride forms fusible white precipitate of diammine Mercury(2)chloride.
HgCl + 2NH3 -------> Hg(NH3)2Cl2

Uses of mercury(2)chloride
Mercury(2)chloride is used in the preparation of mercuric iodide.

3. Mercury(2)Iodide (HgI2)
Preparation
Mercury(2)iodide is obtained as a scarlet precipitate on addition of potassium iodide to a solution of mercury(2)chloride.
2KI + HgCl2 --------> HgI2 + 2KCl

Properties
Mercury(2)iodide readily dissolves in excess of potassium iodide solution due to the formation of potassium tetra iodo mercurate(2)Complex K2[HgI4].
HgI2 + 2KI -------> K2[HgI4]
This potassium tetraiodo mercurate(2)complex forms light yellow crystals of K2[HgI4].2H2O. The complex dissolves in potassium hydroxide solution to give Nessler's reagent which forms a brown precipitate or colouration with ammonia due to the formation of the iodide of Million's base, Hg2NI.H2O

Uses of mercury iodide
Mercury(2)iodide is used for preparing Nessler's reagent and in the treatment of skin infection.

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Related articles halides

Silver nitrate (Lunar caustic) AgNO3

Preparation
Silver nitrate is prepared by dissolving silver in dilute nitric acid.

3Ag + 4HNO3 ------> 3AgNO3 + 2H2O + NO

Properties
Silver nitrate on heating decomposes to form silver, nitrogen dioxide and oxygen.

2AgNO3 ---------> 2AgNO2 + O2

AgNO2 ---------> Ag + NO2

Silver nitrate is also decomposed by organic matter, such as glucose, paper, skin and cork. It has also a caustic and destructive effect on organic tissues.

Uses of silver nitrate
Large quantities of silver nitrate are used in the production of light sensitive plates, film and papers. In the laboratory it is used as a reagent for the detection of halide ions. It is used in making inks and hair dyes. In small doses, silver nitrate is used as a medicine for nervous diseases.

Silver halides
Silver fluorides may be prepared by the action of hydrofluoric acid on silver(1) oxide.

Ag2O + 2HF -------> 2AgF + H2O

Silver chloride, silver bromide and silver iodide are prepared by the action of silver nitrate on sodium halide.

Ag+ + x+ --------> AgX
(Where X = Cl, Br or I)

Ie Ag+ + Cl- --------> AgCl

Properties

Silver fluoride is soluble in water whereas the silver chloride, silver bromide and iodide are insoluble in water.
Silver chloride is highly soluble in ammonia solution due to the formation of Diammine silver(1) chloride complex [Ag(NH3)2]Cl.
Silver bromide is slightly soluble in ammonia solution and silver iodide insluble in ammonia solution.
All the silver halides dissolve in thiosulphate to form thiosulphate complex of silver and in cyanide solution to form dicyano complexes of silver(1).

Uses of silver halides
Silver chloride is used in photography for making printing paper. Silver bromide is used for the production of films and plates in photography and silver iodide used for the production of colloidal emulsion plates in photography.

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Copper sulphate penta hydrate (CuSO4 5H2O)

Copper sulphate penta hydrate is known as blue vitirol and is the most common oxosalt of copper(2)
Preparation
Copper sulphate is prepared industrially by blowing a current of air through copper scrap and dilute sulphuric acid.

2Cu + 2H2O + O2 --------> 2CuSO4 + 2H2O

The crude copper(2) sulphate solution obtained contain iron(2) sulphates as impurity Dilute nitric acid is added to oxidize iron(2) to iron(3) sulphate which remains in solution after crystallization and CuSO4 5H2O crystallizes out.

The crystalline copper(2) sulphate, CuSO4 5H2O has the structure in which four water molecules are coordinated to the central copper cation in square planar structure. The fifth water molecule is held by hydrogen bonds between a sulphate anion and a coordinated water molecule. The fifth hydrogen bonded water molecule is deeply embedded in the crystal lattice and hence not easily removed.

Properties

1. Copper sulphate penta hydrate is a blue coloured crystalline solid, soluble in water.

2. Action of heat
On heating, copper sulphate penta hydrate gives trihydrate at 305 kelvin, monohydrate at 373 kelvin. On further heating it forms white anhydrous copper sulphate at 573 kelvin. The white anhydrous copper sulphate decomposes to give copper oxide and sulphur trioxide on heating to 673 kelvin.

At 3o5 kelvin CuSO4 5H2O ---------> CuO + SO2

At 373 kelvin CuSO4 3H2O ---------> CuO + SO2

At 573 kelvin CuSO4 H2O ---------> CuO + SO2

At 673 kelvin CuSO4 ---------> CuO + SO2

3. Copper(2) sulphate forms well defined crystalline double salts with sulphates of strongly electropositive metals of the type M2SO4.CuSO4 6H2O like (NH4)2SO4. CUSO4 6H2O. These double salts are isomorphous with the double salts of bivalent metals, Fe, Co, Ni.

4. If an aqueous solution of copper(2)sulphate is saturated with ammonia, the blue compound [Cu (NH3)4]SO4 H2O crystallizes on evaporation.

Uses of CuSO4 5H20
Copper(2)sulphate is used in copper plating and electro refining of metals. It is used as a mordant in dyeing and in electroplating. It is also used as germicide and fungicide under the name Bordeaux mixture which is a mixture of CuSO4 and Ca(OH)2.

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Related articles Potassium dichromate

Indicators in acid - base titration

The reaction between an acid and a base is called neutralization. It is very fast and the equilibrium constant for a neutralization reaction is so large that it nearly proceeds to completion. An acid-base titration is a simple and convenient volumetric method for quantitatively estimating the concentration of one, if that of the other is known. A known volume of the solution of an acid or base is transferred to a titration flask with the help of a pipette. we add indicator and start adding known volumes of the other solution in steps with the help of a burette. The point at which the reaction is observed to be complete is called the end point of the titration and is noted by the change in the colour of the indicator. For accurate estimation it is necessary for it to coincide with the equivalence point corresponding to the stoichiometric amounts of the acid and base in the neutralization reaction. A number of weak organic acids and bases which can change its colour with in its limits with variations in the pH value of the solution to which it is added act as indicators. The choice of indicator depends on the abrupt change of the pH during neutralization process near the equivalence point.

Different theories have been put forward to explain the role of indicators in the acid-base titrations's like Ostwald's ionic theory, Quinonoid theory etc. Ostwald's theory considers indicator to be a weak acid or base whose unionised forms differently coloured. In presence of acid or base, ie pH change, there is ionization of indicator and hence the colour change appears.
For example
phenolphthalein
phenolphthalein is a weak acid (PhH)

PhH <_-_-_-_-_-_-> Ph- + H+ ...........(1)
(colourless (Pink in base)
in acid)

H+ + OH- <-_-_-_-_-_-_> H2O

In presence of an acid (H+) equilibrium (1) is displaced towards the left hand side (a case of LeChatelier's principle); when strong base like NaOH is added, this equilibrium is displaced towards right hand side and there is colour change from colourless to pink when pH changes. This indicator is not suitable for titrating weak base since weak base can't furnish enough OH- that can react with H+ of the phenolphthalein and can impart pink colour only after excess of weak base is added.

Methyl orange
Methyl orange behaves like a weak base (MeOH)

MeOH <-_-_-_-_-_> Me+ + OH- .........(2)
(yellow in base) (red in acid)

OH- + H+ <-_-_-_-_-_> H2O

In presence of a base, equilibrium (2) is displaced towards left hand side and appears yellow in base solution. On the addition of strong acid, OH- of MeOH is removed and hence equilibrium (2) is displaced towards right hand side when solution appears red. Thus there is colour changes from golden yellow to red when medium changes from basic to acidic. This indicator is not used in the titration of weak acid since it will not remove OH- of the indicator and can make colour change only after excess of weak acid has been addded.

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Fluorides of xenon

The important fluorides of xenon are xenon difluoride(XeF2), Xenon tetrafluoride(Xef4) and xenon hexafluoride.

1. Xenon difluoride(XeF2)
It is prepared by heating a mixture of Xenon and fluorine in the ration 2:1 at 400 degree Celsius and 1 bar pressure in a sealed nickel tube.

Xe + F2 ---Ni----> XeF2

XeF2 undergoes hydrolysis when treated with water an d evolves oxygen.

2XeF2 + 2H2O -------> 2Xe + 4HF + O2

In XeF2, Xenon is sp3d hybridised and the molecule has linear structure as shown.

2. Xenon tetrafluoride (Xef4)
It is prepared by heating a mixture of Xe and F2 in the molecular ratio 1:5 at 400 degree Celsius and 6 atm in a sealed nickel tube.

Xe + 2 F2 ---------> XeF4

XeF4 react with water and produces explosive XeO3

6 XeF4 + 12 H2O ---------> 2 XeO3 + 4 Xe + 3O2 + 24 HF

In XeF4, Xenon is in sp3d2 hybridised state and has square planar geometry.



3. Xenon hexafluoride (XeF6)
It is prepared by heating a mixture of xenon and fluorine in the ration 1:20 at 300 degree Celsius and 60 atm in a nicked vessel.

Xe + 3 F2 --------> Xe F6

XeF6 Undergoes slow hydrolysis with atmospheric moisture producing highly explosive XeO3.

XeF6 + 3H2O --------> XeO3 + 6 HF

XeF6 molecule possesses distorted octahedral structure. The Xe atom in XeF6 is in sp3d3 hybridisation.



Related article Oxides of xenon
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