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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]]

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 ...

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 H 2 CO 3 and HCO 3 - 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) [(HOCH 2 )3CNH 2 ]. 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 o...

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 + , Hg 2+ and Pb 2+ form insoluble chlorides since they have low values of K sp . Group-2 as insoluble sulphide in acidic medium H 2 S <========> H + + HS - ; K 1 - first ionization constant HS - <========> H + + S 2 - ; K 2 – second ionization constant [S 2 - ] = K 1 K 2 [H 2 S]/[H + ] 2 K sp values of second group sulphides (PbS, CuS, SnS, HgS, As 2 S 3 , Bi 2 S 3 , Sb 2 S 3 ) are very low. In acidic buffer, [S 2 - ] is decreased due to common ion effect and this results in the precipitation of Pb 2+ , Cu 2+ 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 NH 4 OH and NH 4 Cl The concentration of OH - in a...

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 ...

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 CH 3 COOH (weak acid) is decreased by the addition of CH 3 COONa (CH 3 COO- being the common ion) CH 3 COOH CH 3 COO - + H + ……………………….. (A) CH 3 COONa -----------> CH 3 COO - + Na + In the presence of CH 3 COO - equilibrium (A) shifts in backward direction. Ionisation of H 2 S (weak acid) is decreased by the addition of HCl (H + being the common ion) H 2 S 2H + + s 2- HCl H + + Cl - Ionisation of NH 4 OH (weak base) is decreased by the addition of NH 4 Cl (NH4 + being the common ion) NH 4 OH NH 4 + + OH - NH 4 Cl --------------> NH 4 + + Cl - Solubility of a sparingly soluble salt is decreased by the addition of common ion. Presence of AgNO 3 or KCl decreases the solubility of AgCl. AgCl Ag + + Cl - AgNO 3 Ag + (common ion) + NO 3 - KCl K + + Cl - (common ion) The common ion effect is thus ba...

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...