PHARMACEUTICAL ANALYSIS
Physicochemical Concept required for Analysis
Dr. M. Shahar Yar Lecturer
Dept. of Pharmaceutical Chemistry Faculty of Pharmacy
Jamia Hamdard Hamdard Nagar New Delhi- 110062
(04.10.2007) CONTENTS
Introduction
Electrolytic Dissociation
Modern concept of Acid and Bases Chemical Equilibrium
pH and Buffer Solution Solubility Product Common Ion Effect Hydrolysis of Salt
Keywords
Chemical equilibrium, pH and buffer action, Solubility product, Common ion effect
Introduction
The term physicochemical properties refer the influence of organic functional group (present within the molecules) on properties of molecule Viz. Water solubility, partition coefficient, chemical bonding, chelation etc. All these properties ultimately influence the absorption, distribution, metabolism and excretion (ADME).
Electrolytic Dissociation
Certain substances (electrolytes) have ability to conduct electricity when they are dissolved in the water. In the term of physical chemistry electrolytes are defined as the electrovalent substances that form ions in solution, which conduct an electric current.
In 1887 Arrhenius suggested that this ability to conduct electricity was due to the fact that, in solution electrolytes undergo dissociation into positively and negatively charged fragments, which are called as ions. The phenomenon of decomposition of an electrolyte by passing the electric current through its solution is termed as electrolysis. The process of electrolysis is carried in an apparatus called as electrolytic cell. The cell contains water-solution of an electrolyte in which two metallic rods are dipped which is known as electrodes. In a cell positive ions move towards the negative electrode, termed as cathode, and the negative ions move towards the positive electrode, termed as anode.
The passage of ions and subsequent neutralization of the ionic charge at the electrode, bring about conduction of electric current through the solution.
Example: - Let us consider the electrolysis of HCl, in solution, HCl is ionized, HCl H+ + Cl-
In the electrolytic cell Cl- ions will move towards the anode and H+ ion will move toward the
cathode .The following reaction will take place at the electrodes,
At cathode:
H+ + e-
H (Reduction)
As you see, each hydrogen ion picks up an electron from the cathode to a hydrogen atom.
Pairs of hydrogen atoms then united to form molecules of hydrogen gas H2.
At anode:
Cl- Cl + e-
(Oxidation)
After the chloride ion loses its electron to the anode. Pairs of chlorine atoms unite to form chloride gas, Cl2.
The net effect of the process is the decomposition of HCl into hydrogen and chlorine gases.
The overall reaction is:
2 HCl H2 + Cl2 (decomposition)
Strong electrolytes: A strong electrolyte is a substance that gives a solution in which almost all the molecules are ionized. The solution itself called as strong electrolytic solution. Such solutions are good conductors of electricity and have high value of equivalent conductance even at low concentrations. The strong electrolytes are:
The strong bases e.g. NaOH, KOH, Ca (OH)2 etc.
The strong acids e.g. HCl, H2SO4, HNO3, HBr and HI.etc.
Practically all salts (NaCl, KCl etc.) are strong electrolytes etc.
Weak electrolytes- A weak electrolyte is a substance that gives a solution in which only a small proportion of the solute molecules are ionized. Such a solution is called a weak
electrolytic solution that has low value of equivalent conductance. The weak electrolytes are:
The weak acids e.g. acetic acid, oxalic acid, sulphurous acid etc.
The weak bases e.g. alkyl amine etc
Salts e.g. mercury (II) chloride etc
Non electrolytes: A covalent substances, which furnish neutral molecules in solution, their water solution does not conduct an electric current.
Examples: sugar, alcohol and glycerol.
Modern Concept of Acid and Bases
Different theories proposed to define the acid & bases are given below:
1. Savante Arrhenius (1857-1927): According to Savante Arrhenius, acids are those substances, which form hydrogen ions (H+) and bases are those substances, which form hydroxide ions (OH- ), when they are dissolved in water.
e.g When HCl is dissolved in water it gives H+ ions:
HCl H+ (aq) + Cl- (aq) HCl is a very strong acid
When NaOH dissolves in water it gives OH- ion:
NaOH Na+ + OH-
NaOH is a very strong base.
This explanation though useful, is very limited. It even has a few problems. We now know that H+ ion is so reactive and it does not exist as such in aqueous solutions. Instead H+ reacts with water to give hydronium, H3O+ ion. However there are many bases that do not contain hydroxide ions. These react with water to produce OH– ions. For example the base NH3 does not contain OH- ion but reacts with water to produce OH- ion.
NH3 (g)+ H2O (l) NH4+(Aq) + OH-(from water)
Problems with Arrhenius' Theory
1) The solvent has no role to play in Arrhenius' theory. An acid is expected to be an acid in any solvent. This was found not to be the case. For example, HCl is an acid in water, behaving in the manner Arrhenius expected. However, if HCl is dissolved in benzene, there is no dissociation, the HCl remaining as undissociated molecules. The nature of the solvent plays a critical role in acid-base properties of substances.
2) All salts in Arrhenius' theory should produce solutions that are either acidic or basic. This is not the case. If equal amounts of HCl and ammonia react, the solution is slightly acidic. If equal amounts of acetic acid and sodium hydroxide are reacted, the resulting solution is basic.
Arrhenius had no explanation for this.
3) The need for hydroxide as the base led Arrhenius to propose the formula NH4OH as the formula for ammonia in water. This led to the misconception that NH4OH is the actual base, not NH3.
4) H+, a bare proton, does not exist for very long in water, since proton affinity of H2O is about 799 kJ/mol.
2.Bronsted /Lowry Theory: Bronsted (1879-1947) / Lowry (1874-1936)
In 1923, Bronsted (Danish) and Lowry (English) published independent papers on the same subject. Unlike the Arrhenius theory, their approach was not limited to aqueous solutions but for all proton (H+) containing systems.
ACID: Substance that can donate proton (H+).
BASE: Substance that can accept proton (must contain lone pair of electrons).
Acids may be cations, neutral molecules, or anions, while bases may be anions or neutral molecules. Just as reduction process must always accompany an oxidation process, a proton donor (acid) must accompany a proton acceptor (base). Once an acid transfers its proton it becomes the conjugate base (CB) and once a base accepts the proton it becomes the conjugate acid (CA). Since protons are always transferred in the Arrhenius concept, all Arrhenius acid/base reactions are also Bronsted-Lowry acid/base reactions.
But if water is not involved (HCl & NH3) the reaction can be explained by Bronsted/Lowry concept and not Arrhenius.
Example:
HCl + NH3 NH4+ + Cl- acid base CA CB
Concept of conjugate acid and conjugate base: A conjugate pair is an acid-base pair that differs by one proton in their formulas (remember: proton, hydrogen ion, etc.).
A conjugate pair always consist one acid and one base.
HCl + H2O H3O+ + Cl¯
Here is the one conjugate pair from the first example reaction:
HCl and Cl¯
Usually, HCl is called an acid and Cl¯ is called its conjugate base, we can also speak Cl¯ as a base and HCl its conjugate acid.
The other conjugate pair is:
H2O and H3O+
Water is the base, since it is minus a proton compared to H3O+, which is the conjugate acid to water.
Remember conjugate pairs differ by only one proton. If proton is taken away (or added to it), we get the formula of other.
Here are some more conjugate acid-base pairs to look for:
H2O and OH¯
HCO3¯ and CO32¯ H2PO4¯ and HPO42¯ HSO4¯ and SO42¯ NH4+ and NH3
CH3NH3+ and CH3NH2 HC2H3O2 and C2H3O2¯
This last one is special. Because it is used so often, it has an abbreviation: Acetic acid's (HC2H3O2) abbreviation is HAc and the acetate ion's (C2H3O2¯) is Ac¯.
Problems with the Theory : This theory works very nicely in all protic solvents (water, ammonia, acetic acid, etc.), but fails to explain acid base behavior in aprotic solvents such as benzene and dioxane. That job will be left for a more general theory, such as the Lewis theory of Acids and Bases.
Bronsted/Lowry expands Arrhenius theory to include any proton transfer (water not a compulsory requirement).
3. Lewis Theory: Gilbert Lewis (1875-1946) just as the Arrhenius theory did not support observations of acid-base behavior in non-aqueous systems; the Bronsted-Lowry model excluded non-protonated systems. Lewis suggested his theory in a 1923 book "Thermodynamics and the Free Energy of Chemical Substances" and fully developed the theory in 1938.
ACID: Substances that can accept a pair of electrons from another atom to form a new bond.
BASE: Substances that can donate a pair of electrons to another atom to form a new bond.
It turns out that it may be more accurate to say that "Lewis acids" are substances that are electron-deficient and "Lewis bases" are substances, which are electron-rich.
The product of Lewis acid-base reaction referred to as adduct. The proton itself can act as Lewis acid. Lewis expands acid/base reactions to include many substances without H in formula.
F3B + :NH3 F3B:NH3 Explained by Lewis but not Arrhenius or BL acid base adduct.
Several categories of substances can be considered as Lewis acids:
1) Positive ions
2) Having incomplete octet in the valence shell 3) Polar double bonds (one end)
4) Expandable valence shells
Several categories of substances can be considered as Lewis bases:
1) Negative ions
2) One of more unshared pairs in the valence shell 3) Polar double bonds (the other end)
4) The presence of a double bond Example:
The hydroxide ion donates a pair of electrons for covalent bond formation, thus OH- is a Lewis base in this reaction. The hydrogen ion accepts the pair of electrons so it is acting as a Lewis acid. Shown below is an example of a Lewis acid-base reaction that cannot be viewed as a Brønsted-Lowry acid-base reaction.
The BF 3 is the Lewis acid and the N (CH3)3 is the Lewis base. Both of the electrons in the covalent bond formed by a Lewis acid-base reaction come from the same atom (in the above example, the nitrogen donates both electrons). Such bonds are called coordinate covalent bonds.
All Bronsted/Lowry acid/base reactions are also Lewis acid/base reactions.
PROBLEMS:Identify the following given example according to the above mentioned theories Which theory can explain the following?
(1) Question Answers
HI + H2O H3O+ + I- Explained by all 3 theories HI + NH3 NH4+ + I- Explained by BL & Lewis I2 + NH3 NH3I+ + I- Explained by Lewis
I2 + Cl ICl + I Cannot be explained by any of the theories
X- + Y+ Y:X Explained by Lewis but not Arrhenius or BL
H2 + Cl2 2HCl Cannot be explained by any of the theories!
(2) If (a)-(e) are based on the following equation:
XY + X2Z X3Z+ + Y- Then answer the following:
(a) What is needed in the equation to be Arrhenius acid-base reaction?
(b) What is needed in reaction to be Bronsted/Lowry acid-base reaction?
(c) What is needed in reaction to be Lewis acid-base reaction?
(d) Using the equation, why is Bronsted-Lowry more general than Arrhenius?
(e) Using the equation, why is Lewis more general than Bronsted-Lowry?
Answers: (a) X = H & Z = O (b) X = H (c) Z must possess lone pair in X2Z (d) X2Z not limited to H2O and X not limited to H (e) X not limited to H
4. The solvent-system definition: This definition is based on a generalization of the earlier Arrhenius definition to all auto dissociating solvents. In all such solvents there is a certain concentration of a positive species, solvonium cations and negative species, solvate anions, in equilibrium with the neutral solvent molecules. For example:
2H2O H3O+ (hydronium) + OH- (hydroxide) 2NH3 NH4+ (ammonium) + NH2− (amide) or even some aprotic systems
N2O4 NO+ (nitrosonium) + NO3− (nitrate)
2SbCl3 SbCl2+(dichloroantimonium)+SbCl4- (tetrachloroantimonate) A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is an acid and one causing the reverse is a base. Thus, in liquid ammonia, KNH2
(supplying NH2-) is a strong base, and NH4NO3 (supplying NH4+) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds (supplying SO2+) behave as acids, and sulfites (supplying SO32−) behave as bases.
Here are some non-aqueous acid-base reactions in liquid ammonia
2NaNH2 (base) + Zn(NH2)2 (amphiphilic amide) Na2[Zn(NH2)4] 2NH4I (acid) + Zn(NH2)2 (amphiphilic amide) [Zn(NH3)4)]I2 Nitric acid can be a base in liquid sulphuric acid:
HNO3 (base) + 2H2SO4 NO2+ + H3O+ + 2HSO4-
And things become even stranger in the aprotic world, for example in liquid N2O4: AgNO3 (base) + NOCl (acid) N2O4 + AgCl
Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, HClO4 is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid.
5. The Usanovich definition: The most general definition is that of the Russian chemist Mikhail Usanovich, and can basically be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This tends to overlap the concept of redox (oxidation-reduction), and so is not highly favoured by chemists. This is because redox reactions focus more on physical electron transfer processes, rather than bond making/bond breaking processes, although the distinction between these two processes is somewhat ambiguous.
6 The Lux-Flood definition: This definition, proposed by german chemist Hermann Lux in 1939, further improved by Håkon Flood circa 1947 and now commonly used in modern geochemistry and electrochemistry of molten salts, describes an acid as an oxide ion acceptor and a base as an oxide ion donor. For example:
MgO (base) + CO2 (acid) MgCO3 CaO (base) + SiO2 (acid) CaSiO3
NO3- (base) + S2O72- (acid) NO2+ + 2SO42-
7. The Pearson definition: In 1963 Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species which are small, have high charge states, and are weakly
polarizable. 'Soft' applies to species which are large, have low charge states and are strongly polarizable. Acids and bases interact and the most stable interactions are hard-hard and soft-soft.
This theory has found use in both organic and inogranic chemistry.
Chemical Equilibrium
Reversible reaction is a chemical reaction that may proceed in both the forward and reverse directions. In other words, the reactant and product of one reaction may play reverse roles, without adding chemicals.
A + B C + D (Forward Reaction)
A + B C + D (Backward Reaction)
Such a reaction is represented, by writing a pair of arrows between the reactants and Product:
A + B C + D
The products C and D are produced from reactants A and B, but C and D can react to form A and B.
In an irreversible reaction the equilibrium states are shifted so close to either the products or the reactants that the reaction effectively does not have an equilibrium between the products and the reactants. Hence, irreversible reactions can be viewed as an extreme, "special case" of reversible reactions. Irreversible reactions are often called "spontaneous" or "favorable". These reactions are usually entropically driven, as opposed to thermodynamically driven. In an irreversible reaction, there is generally a great increase in entropy
Reversible reaction is shown by the sign a half-arrow to the right (forward reaction), and, a half-arrow to the left (backward reaction).
Most reactions are not reversible (irreversible) and have the usual complete arrow only pointing to the right.
Example: The thermal decomposition of ammonium chloride
¾ On heating strongly above 340oC, the white solid ammonium chloride thermally decomposes into a mixture of two colorless gases ammonia and hydrogen chloride.
¾ On cooling the reaction is reversed and solid ammonium chloride reforms.
i) This is an example of sublimation but here it involves both physical and chemical changes.
ii) When a substance sublimes it changes directly from a solid into a gas without melting and on cooling reforms the solid without condensing to form a liquid.
¾ Ammonium chloride + heat ammonia + hydrogen chloride NH4Cl(s NH3(g) + HCl(g)
Note:
¾ Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction.
¾ Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones.
¾ The decomposition is endothermic (heat absorbed or heat taken in) and the
formation of ammonium chloride is exothermic (heat released or heat given out).
¾ This means if the direction of chemical change is reversed, the energy change is also reversed.
¾ Ammonium fluoride, ammonium bromide (>450oC) and ammonium iodide (>550oC), with a similar formula, all sublime in a similar physical-chemical way when heated, so the equations will be similar i.e. just swap F, Br or I for the Cl.
i) Similarly, ammonium sulphate also sublimes when heated above 235oC and thermally decomposes into ammonia gas and sulphuric acid vapour.
(NH4)2SO4(s) NH3(g) + H2SO4(g)
¾ Ammonium nitrate does not undergo a reversible sublimation reaction; it melts and then decomposes into nitrogen (I) oxide gas (dinitrogen oxide) and water vapour.
NH4NO3(s) N2O(g) + 2H2O(g)
¾ This is very different reaction, in fact it is an irreversible redox reaction. The nitrate ion, NO3-, or any nitric acid formed, HNO3, act as an oxidizing agent If the products are cooled, ammonium nitrate is NOT reformed.
Chemical equilibrium: The state of a reversible reaction when the two opposing reactions occur at the same rate and the concentrations of reactants and products do not change with time.
Furthermore, the true equilibrium of a reaction can be attaind from the both sides.Thus the equilibrium concentrations of the reactants and products are the same whether we start with reactants or products.
Charecteristics of chemical equilibrium: Before we take up the mathematical study of the chemical equilibrium, let us understand the chemical charecteristics of the chemical equilibrium.
1. Consistency of concentration: the concentration of the various species in the reaction mixture becomes constant after the establishment of equilibrium in a closed vessel at the
constant temperature. The reaction mixture at this stage is called as equilibrium mixture and the concentration at this stage is known as equilibrium concentration, which is represented by square bracket with subscript eq, []eq.
2. Equilibrium can be initiated from either side: The state of equilibrium of a reversible reaction can be approached in both of the direction , either reactants side or products side.
Example : H2(g) + I2(g) 2HI(g)
3. Equilibrium cannot be attained in open vessel: Equilibrium is only attained in closed vessels where no part of the reactant or product is allowed to escape out. However, the equilibrium can be attained when all the reactants products are in the same phase i.e, ethanol, ethanoic acid.
4. A catalyst cannot change of the equilibrium point: Addition of the catalyst in a reaction mixture cannot change the equilibrium point. This enhances the rate of the reaction and equilibrium is achieved earlier.
5. Value of equilibrium constant does not depend upon the initial concentration of reactants.
6. At equilibrium i.e. ∆G = 0
Factors Affecting an Equilibrium Position:
1. Changes in Concentration
a) On increasing the concentration of a reactant equilibrium is shifted towards the products (or right hand) side because the rate of the forwards reaction is increased.
b) On increasing the concentration of a product equilibrium is shifted towards the reactant (or left hand) side because the rate of the reverse reaction is speeded up.
c) On decreasing the concentration of a reactant (by removal or by compounding it with something else or by precipitation) equilibrium is shifted towards the reactants (or left hand) side because the forward reaction is slowed down. The reverse reaction will 'overtake' the forward reaction.
d) On decreasing the concentration of a product equilibrium is shifted towards the products (or right hand) side because the reverse reaction is slowed and the forwards reaction 'overtakes'.
2. Changes in Pressure: Consider:
N2O4(g) 2NO2(g)
dinitrogen tetra oxide nitrogen dioxide (very light yellow, almost colorless) (brown)
Increased pressure will cause more collisions to take place between gaseous molecules (pressure is a measure of the number of particles per unit volume).
• An increase in pressure will favour the reaction involving the most particles
• An increase in pressure favours 2NO2 molecules reacting more than one N2O4 reacting
• The reverse reaction is speeded up
• The equilibrium shifts to the left hand side i.e. a lighter color is produced (after an expected initial darkening caused by the original color being 'compressed').
If the pressure is decreased (by expanding a syringe of gas, for example), the initial colour thins but rapidly darkens. The equilibrium has shifted to the right hand side because the reverse reaction has slowed down - its collisions have been reduced in frequency.
N.B. Pressure changes only matter if there is a different number of a gas molecule on each side.
Pressure changes are irrelevant if there are no gas molecules in the reaction, e.g.
H2(g) + I2 2 HI(g)
2 moles of gas 2 moles of gas Here is no change in equilibrium position.
3. Change in Temperature: Generally, an increase in temperature speeds up a reaction, but ∆H values must be considered.An endothermic reaction is helped more by a temperature rise than an exothermic one, e.g.
endothermic→
N2O4(g) 2NO2(g)
dinitrogen tetra oxide ←exothermic nitrogen dioxide (very light yellow, almost
colorless) (brown)
¾ Increasing temperature will favour the forward reaction more than the reverse, and so the colour darkens and the equilibrium shifts to the right hand side.
¾ Decreasing temperature slows the endothermic reaction the most, and so the forwards
reaction slows drastically and the reverse reaction wins, which means that the colour lightens as the equilibrium shifts to the left hand side.
4. The Effect of Catalysis on Equilibrium: A catalyst provides an easier path for the reaction;
the path for the reverse reaction is made equally easier. A catalyst will not shift an equilibrium position because both rates are equally increased and only equilibrium is achieved earlier.
Law of Mass Action : The Law of Mass Action first expressed by two Norwegiean chemists Waage and Guldberg in 1864. It states that the rate of a chemical reaction is proportional to probability that the reacting molecules will be found together in a small volume. Historically, Berthollet was one of the first to develop a theory of mass action.
The law of mass action can be summarized as: The rate of a chemical reaction is directly proportional to the product of the effective concentrations of each participating molecule. Or The rate of chemical reaction is proportional to the active masses of the reactants.
The term active mass is molar concentration or no. of moles per liter. It is expressed by enclosed the substances concentration in square brackets.
Molecular collision theory: (To explain the law of mass action) generally it is assumed a chemical reaction is the result of the collision between the reacting molecules. In which some of the collision are inactive. The chemical changes produced in the chemical reaction is proportional to the collision, actually takes place. Thus, at the fixed temperature the rate of reaction is determined by the number of collision between reactant molecules present in unit volume.
Let us consider four boxes of one cubic centimeter volume; consisting different number of reactant molecules A and B. they undergo the collision and form the product C and D, the rate of reaction is governed by the number of possible collision between them.
Suppose the first box contain one molecule of each A and B. the possibility of collision at any instant is 1x1= 1. in the second box, the number of molecules each of A and B is 2. The possibility of the collision at any instant is 2x2=4. in the third box, the no. of molecule of A is 2 and no. of molecules of B is 3. The possibilities of the collision at any instant are 2 x 3=6. In the 4th box, the number of molecules of A is 3 and B is 3. The possibilities of collision at any instant are 3x3=9. So in general we can say that possibilities of collision between the reacting species A and B are equal to the product of the number of molecules of each reacting species per unit volume. Since the rate of reaction is determined by the molecular impacts, it is proportional to moles per unit volume i.e, molar concentration Thus we can write
Rate of reaction ά [A] [B]
= k [A] [B]
From the above consideration, it stands that the rate of reaction is proportional to the molar concentration of each reactant.
Determination of equilibrium constant by using equilibrium law: -
In chemistry, the equilibrium constant is a quantity characterizing a chemical equilibrium in a chemical reaction. It is a useful tool in determining the concentration of various reactants and products in a system where chemical equilibrium occurs.
A typical equilibrium situation is as below:
where A and B are reactant chemical species, C and D are product species, and a, b, c, and d are the stoichiometric coefficients of the respective reactants and products.
Equilibrium occurs when the forward reaction rate equals the backward reaction rate.
The equilibrium constant is defined as:
The law of chemical equilibrium states that this constant (K) is independent of initial concentrations or most reaction conditions, except temperature. [J] represents the chemical activity of the species at equilibrium, under given reaction conditions. Sometimes the activity can be replaced by molar concentration or partial pressure without significant error, yielding Kc
or Kp respectively. K is dimensionless, while Kc and Kp are not necessarily so.
Derivations: For single-step reactions, where the law of mass action is valid, we can write Rate of forward reaction = Rate of backward reaction
where kf and kb are the forward and backward reaction rate constants, respectively. Cross- dividing yields
Since the rate constants are constant by definition, the law of chemical equilibrium is valid. But it should be kept in mind that this derivation is valid only for single-step elementary reactions (which are very rare) and not for general reactions where the law of mass action does not apply.
Using this treatment for general reactions is an example of getting the right answer using the wrong method. The law of chemical equilibrium however holds for all reactions. The general derivation requires advanced treatment using thermodynamics.
The general definition of the equilibrium constant/concentration may be defined as the product of equilibrium concentration of the product devided by the product of equilibrium concentration of the reactant, with each concentration term raised to the power equal to the cofficient of the substance in the balanced equation.
Applications of equilibrium constant: Equilibrium constants can be defined for many physical/chemical processes. Examples include the acidity constant (the equilibrium constant for the dissociation of protons from acids) and the solubility constant (the equilibrium constant for precipitating out of solution).
The key consequence of the equilibrium constant value is the position of the associated equilibrium. If its value is larger than unity, the equilibrium is said to lie to the right (of the arrow) indicating a greater concentration of products relative to reactants; values less than 1 correspond to an equilibrium that favours the reactants. Hence knowledge of the equilibrium constant under a given set of conditions is essential if a reaction is to be exploited in any practical context.
For example, in the Haber process for the formation of ammonia, the value of K is around 30 at standard pressures and temperatures for the process.
Application of law of mass action: Electrolyte may be classified as either strong or weak electrolytes. The nature of electrolytes depends on the extent of their dissociation or ionizing into their ions in solution. Strong electrolyte almost completely dissociated even in moderately
concentrated solution and hence do not constitute equilibrium system. Oppositely weak electrolyte is only incompletely dissociated even in favarable ionization condition of dilute solution; therfore, an equilibrium, which can consider in terms of law of mass action, is reached between undissociated molecules and ions.
Law of mass action has vast application like determination of ionic product of water, the hydrogen ion exponents and the degree of ionization of weak acids and bases.
Le Chatelier's Principle: Le Chatelier's principle states that when a system in chemical equilibrium, is disturbed by a change of temperature, pressure, or a concentration, the equilibrium composition of the system shifts in a way that tends to counteract this change of variable. There are three ways that can affect the outcome of the equilibrium using Le Chatelier's principle are as follows:
• Changing concentrations by adding or removing products or reactants to the reaction vessel.
• Changing partial pressure of gaseous reactants and products.
• Changing the temperature.
These actions change equilibrium differently; therefore it is must to determine what needs to happen for the reaction to get back in equilibrium.
Example involving change of concentration: In the equation 2NO (g) + O2 (g) 2NO2 (g)
If you add more NO (g) the equilibrium shifts to the right producing more NO2 (g) If you add more O2 (g) the equilibrium shifts to the right producing more NO2(g)
If you add more NO2 (g) the equilibrium shifts to the left producing more NO(g) and O2(g) Example involving pressure change: In the equation
2SO2 (g) + O2 (g) 2SO3 (g)
an increase in pressure will cause the reaction to shift in the direction that reduces pressure, that is the side with the fewer number of gas molecules. Therefore an increase in pressure will cause a shift to the right, producing more products. (A decrease in volume is one way of increasing pressure)
Example involving temperature change: In the equation
N2 (g) + 3H2 (g) 2NH3 + 91.8 kJ
An increase in temperature will cause a shift to the left because the reverse reaction uses the excess heat. An increase in forward reaction would produce even more heat since the forward reaction is exothermic. Therefore the shift caused by a change in temperature depends upon whether the reaction is exothermic or endothermic.
pH & Buffer Solution
The pH of a solution is a measure of the molar concentration of hydrogen ions in the solution and as such is a measure of the acidity or basicity of the solution. The letters pH stands for "power of hydrogen" and the numerical value is defined as the negative base 10 logarithm of the molar concentration of hydrogen ions.
pH = -log10[H+]
The letters of its name are derived from the absolute value of the power (p) of the hydrogen ion concentration (H). The product of the concentrations in water of H+ and OH− (the hydroxide ion) is always about 10−14. The strongest acid solution has about 1 mole/litre of H+ (and about 10−14 of OH−), for a pH of 1. The strongest basic solution has about 10−14 moles/litre of H+ (and about 1 of OH−), for a pH of 14. A neutral solution has about 10−7 moles/litre of both H+ and OH−, for a pH of 7.
The ionic product of water at 250C is approximately equal to 1 x 10-14. Therefore,
Kw = [H+] [OH-] = 1 X 10-14 OR
[H+] 2 = 1x 10-14 since, [H+] = [OH-] [H+] = 1 x 10-7 = [OH-]
It means that in pure water the concentration of H+ and OH- ions is 10-7 g ion per liter each.
Therefore, both the degree of acidity and alkalinity of a solution can be expressed quantitatively by hydrogen ion concentration.
In neutral solution, [H+] = [OH-] = 10-7 In acidic solution, [H+] > 10-7
In alkaline solution, [H+] < 10-7
The above method of expressing the acidity or alkalinity of a solution has given place to another still more simple and convenient method, introduced by SORENSOEN in 1909. in this method hydrogen ion concentration is expressed in term of hydrogen ion exponent, i.e., a number obtained by giving a positive value to the negative power of 10 in the expression 1 x 10-n. this was originally represented as pH but now most suitably written as pH.
Thus, pH of solution is numerically equal to the negative power to which 10 must be raised in order to express the hydrogen ion concentration in the solution or, pH is the negative log of hydrogen ion concentration in the solution.
The measurement of the pH of a sample can be done by measuring the cell potential of that sample in reference to a standard hydrogen electrode, as in the accepted procedure for measuring standard electrode potentials. This procedure would give a value of zero for a 1 Molar solution of H+ ions, so that defines the zero of the pH scale. The cell potential for any other value of H+ concentration can be obtained with the use of the Nernst equation. For a solution at 25°C this gives
Ecell = -0.0592 log10 [H+] or
pH = Ecell/0.0592
In practice, the pH is not usually measured in this way because it requires hydrogen gas at standard pressure and the platinum electrode used in the standard hydrogen electrode is easily fouled by the presence of other substances in the solution (Ebbing). Fortunately, other electrode configurations can be calibrated to read the H+ ion concentration. Laboratory pH meters are often made with a glass electrode consisting of a silver wire coated with silver chloride immersed in dilute hydrochloric acid. The electrode solution is separated from the solution to be measured by a thin glass membrane. The potential, which develops across that glass membrane, can be shown to be proportional to the hydrogen ion concentrations on the two surfaces. In the measurement instrument, a cell is made with the other electrode commonly being a mercury-mercury chloride electrode. The cell potential is then linearly proportional to the pH and the meter can then be calibrated to read directly in pH.
Simple Glass Electrode
Examples of pH Values: The pH of a solution is a measure of the molar concentration of hydrogen ions in the solution and as such is a measure of the acidity or basicity of the solution.
The letters pH stands for "power of hydrogen" and numerical value for pH is just the negative of the power of 10 of the molar concentration of H+ ions.
The usual range of pH values encountered is between 0 and 14, with 0 being the value for concentrated hydrochloric acid (1 M HCl), 7 the value for pure water (neutral pH), and 14 being the value for concentrated sodium hydroxide (1 M NaOH). It is possible to get a pH of -1 with 10
M HCl, but that is about a practical limit of acidity. At the other extreme, a 10 M solution of NaOH would have a pH of 15.
Numerical examples from Shipman, Wilson and Todd
In pure water, the molar concentration of H+ ions is 10-7 M and the concentration of OH- ions is also 10-7 M. Actually, when looked at in detail, it is more accurate to classify the concentrations as those of [H3O]+ and [OH]-. The product of the positive and negative ion concentrations is 10-14 in any aqueous solution at 25°C.
An important example of pH is that of the blood. Its nominal value of pH = 7.4 is regulated very accurately by the body. If the pH of the blood gets outside the range 7.35 to 7.45 the results can be serious and even fatal.
If you measure the pH of tap water with a pH meter, you may be surprised at how far from a pH of 7 it is because of dissolved substances in the water. Distilled water is necessary to get a pH near 7.
Meters for pH measurement can give precise numerical values, but approximate values can be obtained with various indicators. Red and blue litmus paper has been one of the common indicators. Red litmus paper turns blue at a basic pH of about 5, and blue litmus paper turns red at an acid pH of about 8. Neither changes color if the pH is nearly neutral. Litmus is an organic compound derived from lichens.
Phenolphthalein is also a common indicator, being colorless in solution at pH below 8 and turning pink for pH above 8.
Buffer solution: Buffer solutions are solutions which resist change in hydronium ion and the hydroxide ion concentration (and consequently pH) upon addition of small amounts of acid or base, or upon dilution. Buffer solutions consist of a weak acid and its conjugate base (more common) or a weak base and its conjugate acid (less common). The resistive action is the result of the equilibrium between the weak acid (HA) and its conjugate base (A−):
HA(aq) + H2O(l) H3O+(aq) + A−(aq)
Any alkali added to the solution is consumed by hydronium ions. These ions are mostly regenerated as the equilibrium moves to the right and some of the acid dissociates into hydronium ions and the conjugate base. If a strong acid is added, the conjugate base is protonated, and the pH is almost entirely restored. This is an example of Le Chatelier's principle and the common ion effect. This contrasts with solutions of strong acids or strong bases, where any additional strong acid or base can greatly change the pH.
When writing about buffer systems they can be represented as salt of conjugate base/acid, or base/salt of conjugate acid. It should be noted that here buffer solutions are presented in terms of the Brønsted-Lowry notion of acids and bases, as opposed to the Lewis acid-base theory (see acid-base reaction theories). Omitted here are buffer solutions prepared with solvents other than water.
Types of buffer solution: Buffer solutions are two types
(1) Acidic buffer solutions
:
An acidic buffer solution is simply one, which has a pH less than 7.Acidic buffer solutions are commonly made from a weak acid and one of its salts - often a sodium salt.
A common example would be a mixture of ethanoic acid and sodium ethanoate in solution. In this case, if the solution contained equal molar concentrations of both the acid and the salt, it would have a pH of 4.76. It wouldn't matter what the concentrations were, as long as they were the same.
You can change the pH of the buffer solution by changing the ratio of acid to salt, or by choosing a different acid and one of its salts.
Function: - We'll take a mixture of ethanoic acid and sodium ethanoate as typical example.
Ethanoic acid is a weak acid, and the position of this equilibrium will be well to the left:
CH3COONa CH3COO- + Na+ CH3COOH CH3COO- + H+
Adding sodium ethanoate to this adds lots of extra ethanoate ions. According to Le Chatelier's Principle, that will tip the position of the equilibrium even further to the left.
The solution will therefore contain these important things:
• Lots of un-ionized ethanoic acid;
• Lots of ethanoate ions from the sodium ethanoate;
• Enough hydrogen ions to make the solution acidic.
Other things (like water and sodium ions), which are present, aren't important to the argument.
Adding an acid to this buffer solution: The buffer solution must remove most of the new hydrogen ions otherwise the pH would drop markedly.
Hydrogen ions combine with the ethanoate ions to make ethanoic acid. Although the reaction is reversible, since the ethanoic acid is a weak acid, most of the new hydrogen ions are removed in this way.
CH3COO-(aq) + H+(aq) CH3COOH
Since most of the new hydrogen ions are removed, the pH won't change very much - but because of the equilibria involved, it will fall a little bit.
Adding an alkali to this buffer solution: Alkaline solutions contain hydroxide ions and the buffer solution removes most of these.
This time the situation is a bit more complicated because there are two processes, which can remove hydroxide ions.
Removal by reacting with ethanoic acid: The most likely acidic substance, which a hydroxide ion is going to collide with, is an ethanoic acid molecule. They will react to form ethanoate ions and water.
CH3COOH (aq) + OH- CH3COO-(aq) + H2O
Note: You might be surprised to find this written as a slightly reversible reaction. Because ethanoic acid is a weak acid, its conjugate base (the ethanoate ion) is fairly good at picking up hydrogen ions again to re-form the acid. It can get these from the water molecules. You may well find this reaction written as one-way, but to be fussy about it, it is actually reversible!.
Because most of the new hydroxide ions are removed, the pH doesn't increase very much.
Removal of the hydroxide ions by reacting with hydrogen ions: Remember that there are some hydrogen ions present from the ionization of the ethanoic acid.
CH3COOH(aq) CH3COO-(aq)
+
H+(aq)Hydroxide ions can combine with these to make water. As soon as this happens, the equilibrium tips to replace them. This keeps on happening until most of the hydroxide ions are removed.
Again, because you have equilibria involved, not all of the hydroxide ions are removed - just most of them. The water formed re-ionizes to a very small extent to give a few hydrogen ions and hydroxide ions.
(2) Alkaline buffer solutions: An alkaline buffer solution has a pH greater than 7. Alkaline buffer solutions are commonly made from a weak base and one of its salts.
A frequently used example is a mixture of ammonia solution and ammonium chloride solution. If these were mixed in equal molar proportions, the solution would have a pH of 9.25. Again, it doesn't matter what concentrations you choose as long as they are the same.
Function: - We'll take a mixture of ammonia and ammonium chloride solutions as typical.
Ammonia is a weak base, and the position of this equilibrium will be well to the left:
NH3(aq) + H2O(aq) NH4+
(aq) + OH-(aq)
Adding ammonium chloride to this adds lots of extra ammonium ions. According to Le Chatelier's Principle, that will tip the position of the equilibrium even further to the left.
The solution will therefore contain these important things:
• Lots of unreacted ammonia;
• Lots of ammonium ions from the ammonium chloride;
• Enough hydroxide ions to make the solution alkaline.
Other things (like water and chloride ions), which are present, aren't important to the argument.
Adding an acid to this buffer solution: There are two processes, which can remove the hydrogen ions that you are adding.
Removal by reacting with ammonia: The most likely basic substance, which a hydrogen ion is going to collide with, is an ammonia molecule. They will react to form ammonium ions.
NH3(aq) + H+(aq) NH4+(aq)
Most, but not all, of the hydrogen ions will be removed. The ammonium ion is weakly acidic, and so some of the hydrogen ions will be released again.
Removal of the hydrogen ions by reacting with hydroxide ions: Remember that there are some hydroxide ions present from the reaction between the ammonia and the water.
NH3(aq) + OH-(aq) NH4+(aq)+ OH-(aq)
Hydrogen ions can combine with these hydroxide ions to make water. As soon as this happens, the equilibrium tips to replace the hydroxide ions. This keeps on happening until most of the hydrogen ions are removed.
Again, because you have equilibria involved, not all of the hydrogen ions are removed - just most of them.
Adding an alkali to this buffer solution: The hydroxide ions from the alkali are removed by a simple reaction with ammonium ions.
Because the ammonia formed is a weak base, it can react with the water - and so the reaction is slightly reversible. That means that, again, most (but not all) of the hydroxide ions are removed from the solution.
Applications of buffer solution: Their resistance to changes in pH makes buffer solutions very useful for chemical manufacturing and essential for many biochemical processes. The ideal buffer for a particular pH has a pKa equal to the pH desired, since a solution of this buffer would contain equal amounts of acid and base and be in the middle of the range of buffering capacity.
Buffer solutions are necessary to keep the right pH for enzymes in many organisms to work.
Many enzymes work only under very precise conditions; if the pH strays too far out of the margin, the enzymes slow or stop working and can denature, thus permanently disabling its catalytic activity. A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3−) is present in blood plasma, to maintain a pH between 7.35 and 7.45.
Industrially, buffer solutions are used in fermentation processes and in setting the correct conditions for dyes used in colouring fabrics. They are also used in chemical analysis and calibration of pH meters.
Calculating pH of a buffer
HA [H+] + [A-]
The equilibrium above has the following acid dissociation constant:
Simple manipulation with logarithms gives the Henderson-Hasselbalch equation, which describe pH in terms of pKa:
In this equation
[A−] is the concentration of the conjugate base. This may be considered as coming completely from the salt, since the acid supplies relatively few anions compared to the salt.
[HA] is the concentration of the acid. This may be considered as coming completely from the acid, since the salt supplies relatively few complete acid molecules (A − may extract H + from water to become HA) compared to the added acid.
So, the equation can be written as pH= pKa + log [salt]/[acid]
This is the henderson’s equation for calculating the pH of acidic solution. For basic solution the equation can also written as:
pOH= pKb + log [salt]/[base]
pKb is dissociation constant for base. Knowing the pOH , we can calculate first pH, from the equation:
pH + pOH = 14
Maximum buffering capacity is found when pH = pKa, and buffer range is considered to be at a pH = pKa ± 1.
Buffers are important in biochemical processes. Whether they occur naturally in plasma or in the cytosol of cells, buffers assure biological reactions occur under conditions of optimal pH. They do this by controlling the hydrogen ion concentration of solutions. The word “buffer” is so common in biochemistry; it replaces the word “water” in experimental protocols. For example, a typical statement seen in publications is “the pellet was dissolved in pH 7.5 buffer”. Words such as pH, pKa, conjugate acid, conjugate base, Henderson-Hasselbalch equation are used frequently in biochemical language and every publication that describes an experiment performed “in vitro”
(Lat., in glass), must include a clear description of the buffer that was used.
Solubility Product
Solubility equilibrium represents chemical equilibrium between solid and dissolved states of a compound at saturation.
The substance that is dissolved may be an organic solid such as sugar or an ionic solid such as table salt. The main difference is that ionic solids dissociate into constituent ions when they are dissolved in water. Most commonly, water is the solvent of interest, although the same basic principles apply with any solvent.
In the case of environmental science studies of water quality, the total concentration of dissolved solids (not necessarily at saturation) is referred to as total dissolved solids.
Non-ionic compounds: Dissolution of an organic solid can be described as an equilibrium between the substance in its solid and dissolved forms:
An equilibrium expression for this reaction can be written, as for any chemical reaction (products over reactants):
where K is called the equilibrium constant (or solubility constant) and the square brackets mean molar concentration in mol/L (sometimes called molarity with symbol M). Because a notion of concentration for a solid doesn't make sense, curly brackets are used, which mean activity, around the solid. Luckily, the activity of a solid is almost always equal to one. So, a very simple expression suffices:
This statement says that water at equilibrium with solid sugar contains a concentration equal to K. For table sugar (sucrose) at 25 °C, K = 1.971 mol/L. (This solution is very concentrated;
sucrose is extremely soluble in water.) This is the maximum amount of sugar that can dissolve at 25 °C; the solution is saturated. If the concentration is below saturation, more sugar dissolves until the solution reaches saturation, or all the solid is consumed. If more sugar is present than is allowed by the solubility expression then the solution is supersaturated and solid will precipitate until the saturation concentration is reached. This process can be slow; the equilibrium expression describes concentrations when the system reaches equilibrium, not how fast it gets there.
Ionic compounds: Ionic compounds normally dissociate into their constituent ions when they are dissolved in water. For example, for calcium sulfate:
As for the previous example, the equilibrium expression is:
where K is called the equilibrium (or solubility) constant, the square brackets mean molar concentration (M, or mol/L), and curly brackets mean activity. Since the activity of a pure solid is equal to one, this expression reduces to the solubility product expression:
This expression says that an aqueous solution in equilibrium with (saturated with) solid calcium sulfate has concentrations of these two ions such that their product equals Ksp; for calcium sulfate Ksp = 4.93×10−5. If the solution contains only calcium sulfate the concentration of each ion (and the overall solubility of calcium sulfate) is
Solubility constants: Solubility constants have been experimentally determined for a large number of compounds and some are undermentioned. For ionic compounds the constants are called solubility products. Concentration units are assumed to be molar (moles per liter) unless otherwise stated. Solubility is sometimes listed in mass units such as grams dissolved per liter of water.
Some values at 25°C:
Barium carbonate: 2.60×10-9
Copper(I) chloride: 1.72×10-7
Lead(II) sulfate: 1.81×10-8
Magnesium carbonate: 1.15×10-5
Silver chloride: 1.70×10-10
Calcium hydroxide: 8.0×10-6
Solubility =√ solubility product or solubility constant S = √ Ksp
Solubility (and equilibrium) constants themselves are dimensionless (however, they may have units). The lack of units in the constant may look inconsistent, but it comes about because the use of molar concentration in the solubility expression is only an approximation to activity, a unitless quantity that is approximately equal to molarity at low concentrations
The common-ion effect refers to the fact that solubility equilibria shift in response to Le Chatelier's Principle. In the above example, addition of sulfate ions to a saturated solution of calcium sulfate causes precipitation of CaSO4 until the ions in solution again satisfy the solubility expression. (Addition of sulfate ions could be accomplished by adding a very soluble salt, such as Na2SO4.)
Solubility is sensitive to temperature. For example, sugar is more soluble in hot water than cool water. It occurs because solubility constants, like other types of equilibrium constant, are functions of temperature. A thermodynamic approach is required to predict how much and in what direction a particular constant changes.
Solubility product principle: Solubility product principle states that precipitation occures if ionic product of solution is more than its solubility product, i.e., Ksp.
Thus, three condition may be there,
if ionic product is lesser than solubility product, the solution is unsaturated and more electrolyte may be dissolved and hence no precipitation will occur.
if ionic product is equal to the solubility product, the solution is just saturated and it is the condition at which precipitation may start.
if ionic product is greater than solubility product, the electrolye will be precipitated in the solution.
Common Ion Effect
The common ion effect is an application of LeChatelier's Principle. If we mix a soluble salt containing an ion common to slightly soluble salt, it will affect the position of the equilibrium of the slightly soluble salt system. Adding the common ion to the salt solution increases the concentration of the common ion. According to LeChatelier's Principle, that (Common ion) will place a stress upon the slightly soluble salt equilibria (added concentration). The equilibrium will respond so as to undo the stress of added common ion. This means that the equilibria will shift so that the common ion will be reduced which means a shift to the left thus REDUCING the solubility of the slightly soluble salt system. Let's illustrate that by determining the molar solubility of a slightly soluble salt like AgCl without common ion and with common ion and then compare the molar solubilities in the two situations.
AgCl Ag+ + Cl-
HCl H+ + Cl-
The presence of a common ion suppresses the ionization of a weak acid or a weak base
Examples: - If both sodium acetate and acetic acid are dissolved in the same solution they both dissociate and ionize to produce acetate ions. Sodium acetate is a strong electrolyte so it dissociates completely in solution. Acetic acid is a weak acid so it only ionizes slightly. According to Le Chatelier's principle, the addition of acetate ions from sodium acetate will suppress the ionization of acetic acid and shift its equilibrium to the left. Thus the percent dissociation of the acetic acid will decrease and the pH of the solution will increase.
NaC2H3O2(s) Na+(aq) + C2H3O2-(aq) HC2H3O2(l) H+(aq) + C2H3O2-
(aq)
This will decrease the hydrogen ion concentration and thus the common-ion solution will be less acidic than a solution containing only acetic acid.
Example: -We can see that this must necessarily occur if we apply Le Chatelier's Principle to equilibrium such as
Adding an excess Cl- (or Na+) to a saturated solution of NaCl imposes a stress on the equilibrium, which will adjust in order to oppose the stress. A shift to the left will use up Na+ or Cl- to form solid NaCl.
Hydrolysis of Salt
Hydrolysis may also be considered as reverse of neutralization. Hydrolysis involves the interaction of the ions of a salt with water to produce hydrogen ion (H+) and hydroxyl ion (OH-) in the solution.
Definition: Salts hydrolysis may be defined in many ways.
(1)“The phenomenon in which a salt react with water to produce either acidic or alkaline solution is called salts hydrolysis.”
(2)“The reaction in which anion and cation of a salt react with H+ and OH- furnished by water to give acidic, alkaline or neutral solution is called salts hydrolysis.”
On the basis of nature salts, hydrolysis may be three types-
i. Hydrolysis of acidic salts.
ii. Hydrolysis of basic salts iii. Hydrolysis of neutral salts
Acidic, Basic, and Neutral Salts: A salt is formed between the reaction of an acid and a base.
Usually, a neutral salt is formed when a strong acid and a strong base is neutralized in the reaction:
H+ + OH- H2O
The bystander ions in an acid-base reaction form a salt solution. Most neutral salts consist of cations and anions listed below. These ions have little tendency to react with water. Thus, salts consisting of these ions are neutral salts. For example: NaCl, KNO3, CaBr2, CsClO4 are neutral salts.
Ions of Neutral Salts:
Cations Na+, K+, Rb+, Cs+ , Mg++, Ca++ , Sr++ , Ba++
Anions Cl-, Br-, I-, ClO4-, BrO4-,ClO3-, NO3-
When weak acids and bases react, the relative strength of the conjugated acid-base pair in the salt determines the pH of its solutions. The salt, or its solution, so formed can be acidic, neutral or basic. A salt formed between a strong acid and a weak base is an acid salt, for example NH4Cl.
A salt formed between a weak acid and a strong base is a basic salt, for example CH3COONa.
These salts are acidic or basic due to their acidic or basic ions as shown below.
Acidic Ions NH4+, Al3+, Pb2+, Sn2+
Transition metal ions HSO4-
, H2PO4-
Basic Ions F-, C2H3O2-, NO2-, HCO3-, CN-, CO32-, S2-, SO42-+, HPO42-, PO43-
Hydrolysis of Acidic Salts: In a reaction between a strong acid and a weak base an acid salt will form at the equivalence point. Ammonia is a weak base, and it’s salt with any strong acid gives a solution with a pH lower than 7. For example, let us consider the reaction:
HCl + NH4OH NH4+ + Cl- + H2O
In the solution, the NH4+ ion reacts with water (called hydrolysis) according to the equation:
NH4+ + H2O NH4OH + H+. The equilibrium constant can be derived from Kw and Kb.
[NH4OH][H+]
Kh = --- [NH4+] [H2O]
NH4OH NH4++OH-
[NH4+][OH-]
Kb = --- ---(1) [NH4OH]
Kw = [H+] [OH-] ---(2)
Where, Kh = hydrolysis constant of salt Kb = ionization constant of base Kw = ionic product of water On dividing Equation (2) from equation (1)
Kw [H+][OH-][NH4OH]
--- = --- Kb [NH4+] [OH-]
Kw [H+][NH4OH]
Kh = --- = --- Kb [NH4+]
Hydrolysis of Basic Salts : In a reaction between a weak acid and a strong base a Basic salt will form at he equivalence point. The basicity is due to the hydrolysis of the conjugate base of the (weak) acid used in the neutralization reaction. For example, sodium acetate formed between the weak acetic acid and the strong base NaOH is a basic salt. When the salt is dissolved, ionization takes place:
CH3COONa Na+ + CH3COO- In the presence of water, CH3COO- undergoes hydrolysis:
