BASIC PHARMACEUTICAL CHEMISTRY PHARM 155/151
CONTENT Equilibria in Electrolytes Ionization equilibrium constant Ionic strength Activity coefficient Debye-Huckel theory Solubility equilibria and activity coefficient Common ion effect Acids and bases Solvent effect on acidic strength Ionization constant Ionic product of water Conductance EMF Heat of neutralization Hydrolysis
Electrolytes The thermodynamics of electrolyte solutions is important for a large number of chemical systems, such as acid-base chemistry, biochemical processes and electrochemical reactions. Electrolytes are substances whose aqueous solutions are conductors of electricity. They are capable of producing ions in solution, whereas nonelectrolytes do not have this property. Compounds which produce a large number of ions in solution are called strong electrolytes. Compounds which produce a small number of ions in solution are weak electrolytes. Soluble compounds that produce no dissolved ions are called nonelectrolytes.
Strong electrolytes because they exist almost completely in the ionic form in moderately concentrated aqueous solutions. Example NaCl and HCl Inorganic acids such as HCl, HNO 3, H 2 SO 4, and HI; inorganic bases as NaOH and KOH of the alkali metal family and Ba(OH) 2 and Ca(OH) 2 of the alkaline earth group; and most inorganic and organic salts are highly ionized and belong to the class of strong electrolytes. Weak electrolytes, there is oppositely directed arrows in equation indicating that equilibrium between the molecules and the ions. Example CH 3 COOH Most organic acids and bases and some inorganic compounds, such as H 3 BO 3, H 2 CO 3, and NH 4 OH, belong to the class of weak electrolytes. Even some salts (lead acetate, HgCl 2, HgI, and HgBr) and the complex ions Hg(NH 3 ) 2+, Cu(NH 3 ) 4 2+, and Fe(CN) 6 3- are weak electrolytes. Electrolytes
DRUGS AND IONIZATION Some drugs, such as anionic and cationic antibacterial and antiprotozoal agents, are more active when in the ionic state. Other compounds, such as the hydroxybenzoate esters (parabens) and many general anesthetics, bring about their biologic effects as nonelectrolytes. The sulfonamides, are thought to exert their drug action both as ions and as neutral molecules.
Chemical Equilibrium Solutions of weak electrolytes, such as weak acids, are characterized by dissociated and undissociated species existing together in dynamic equilibrium. Characteristics of True Chemical Equilibria They show no macroscopic evidence of change A dynamic balance of forward and reverse processes exists within them. They are the same regardless of the direction from which they are approached.
Chemical Equilibrium Equilibrium constants may be written for dissociations, associations, reactions, or distributions.
Equilibrium Constant Typical Form of Equilibrium Constant However, this is not strictly correct non-ideal solutions: In concentrated solutions like seawater Ideal solutions: In infinitely dilute solutions where ionic interactions can be ignored Chemical Equilibrium
In all of the equilibria that you have learnt, we have assumed that the equilibria occur in ideal solutions. some solutions are quite ideal in moderate concentrations, whereas others approach ideality only under extreme dilution. When ion concentrations become high enough so that ions interact, the interactions affect solution equilibria. Equilibria are affected not only when the concentrations of the ions involved in an equilibrium become high, but also when the concentrations of spectator ions become high. Spectator ions are not involved directly in equilibria, but they affect the environment of the ions that are part of an equilibrium. Chemical Equilibrium
Chemical Activity Oxygen has a partial negative charge and hydrogen partial positive charge Example: A solution of calcium and sulfate in water The hydration of Ca 2+ and SO4 2- ions to induces shielding which affects the ability of Ca 2+ and SO 4 2- to meet and react (and precipitate as a solid in this case). Hydration Ions do not act as independent particles in solvent (water) Surrounded by a shell of solvent molecules
Activity If other ions like Na + and Cl - are added to CaSO 4 solution, they are attracted to the ions of opposite charge The added ions effectively increase the amount of electrostatic shielding, thus decreasing the ability of Ca 2+ and SO4 2- to interact. Therefore, gypsum or CaSO 4.2H 2 O, will appear more soluble in seawater than in freshwater. These interactions result in non-ideal solutions. Ions with higher charge are more effective than ions with lower charge at this shielding effect. In such solutions, equilibrium constants are expressed in terms of this effective concentration, which is formally called the activity, which is the concentration available for reaction.
Ionic Atmosphere Cation surrounded by anions and anions are surrounded by cations - Effective charge is decreased - Shields the ions and decreases attraction Net charge of ionic atmosphere is less than ion - ions constantly moving in/out of ionic atmosphere Activity Each ion see less of the other ions charge and decreases the attraction Each ion-plus-atmosphere contains less net charge and there is less attraction between any particular cation and anion
Chemical Equilibrium This effect of the ionic environment within the solution is known as the ionic strength and may be represented as . The ionic strength of an electrolyte solution takes into account all the ionic species and their concentrations. = ½ (c 1 z c 2 z 2 2 +…) = ½ c i z i 2 where: C i is the concentration of the ith species and z i is its charge Ionic Strength (m) Measure of the total concentration of ions in solution - More highly charged an ion is the more it is counted - Sum extends over all ions in solution
= ½ c i z i 2 = ½ {0.01(+1) (-1) (+1) (-2) 2 } = ½ {0.08} = 0.04M 1.Calculate the ionic strength of M Na 2 SO What is the ionic strength of a solution that is M in KNO 3 and M Na 2 SO 4 ? Exercise μ= 0.5 Σcizi 2 = 0.5 ([Na + ](+1) 2 + [SO4 -2 ](-2) 2 ) = 0.5 ((0.02)(1) + (0.01)(4)) = M
What is the ionic strength of a.0.01M KCl b.0.01M BaSO 4 c.0.01M Na 2 SO 4 Exercise The ionic strength of a 1:1 electrolyte such as KCl is the same as the molar concentration; µ of a 1:2 electrolyte such as Na 2 SO 4 is three times the concentration; µ for a 2:2 electrolyte is four times the concentration.
Compare the ionic strength of Lake water and seawater μ (sw) μ ( LW ) The ionic strength of seawater is about 500 times larger than that of lake water
Equilibria in Electrolytes Effect of Ionic Strength on Solubility Equilibria Involving Ionic Compounds are Affected by the Presence of All Ionic Compounds in the Solution Knowing the ionic strength is important in determining solubility Example: K sp = 1.3x Explain. If Hg 2 (IO 3 ) 2 is placed in pure water, up to 6.9x10 -7 M will dissolve, but if M KNO 3 is added, up to 1.0x10 -6 M Hg 2 (IO 3 ) 2 will dissolve.
Equilibria in Electrolytes Activity: Effective concentration of ions in solution is related to the ideal concentration by the activity coefficient ( ). Activity coefficients serve as correction factors so that equilibrium constants measured in an ideal solution can be used to predict equilibria in non-ideal solutions where: A C is activity of C [C] is concentration of C C is activity coefficient of C “Real” Equilibrium Constant Using Activity Coefficients
Equilibria in Electrolytes Activity and activity coefficient In any real solution interactions occur between the components which reduce the effective concentration of the solution. The activity is a way of describing this effective concentration. The ratio of the activity to the concentration is called the activity coefficient, γ, that is, γ= activity/concentration. Therefore, for an ideal solution γ= 1 When drugs are salts they ionise in solution and the activity of each ion is the product of its activity coefficient and its molar concentration.
Equilibrium Constant and Activity “ Real” Equilibrium Constant Using Activity Coefficients is always ≤ 1 Activity coefficient measures the deviation from ideal behavior - If =1, the behavior is ideal and typical form of equilibrium constant is used Activity coefficient depends on ionic strength - Activity coefficient decrease with increasing ionic strength - Approaches one at low ionic strength Equilibria in Electrolytes
The extended Debye-Huckel Equation Activity Coefficients of Ions Only valid for concentrations ≤ 0.1M In theory, α is the diameter of hydrated ion where: is the activity coefficient is ion size (pm) z is the ion charge is the ionic strength Equilibria in Electrolytes
Activity Activity Coefficients from Debye-Hϋckel Equation
The extended Debye-Huckel Equation P. Debye and E. Huckel derive a theoretical expression that permits the calculation of activity coefficients of ion from their charge and their average size. This equation takes the form where, = activity coefficient of the species X Z = charge on the species X = ionic strength of the solution = effective diameter of the hydrated ion X in nanometer (10 -9 m) Equilibria in Electrolytes Only valid for concentrations ≤ 0.1M The constants 0.51 and 3.3 are applicable to aqueous solution at 25 o C;
1. What is the activity coefficient of Mg +2 in a 3.3 mM solution of Mg(NO 3 ) 2, using both equations? Exercise ionic strength : = ½ {(3.3mM)(+2) 2 + (6.6mM)(-1) 2 } = ½ { } = ½{19.8} = 9.9 mM = M log = (-0.51)(+2) 2 (0.0099) 1/2 / {1 +(800)(0.0099) 1/2 /305} log = /1.261 = = = log = (-0.51)(+2) 2 (0.0099) 1/2 / { (0.8)(0.0099) 1/2 } log = =
Exercise 1. Calculate the activity coefficient of a M aqueous solution of sodium phenobarbital at 25 C, which has been brought to ionic strength of 0.09 by addition of sodium chloride. (Mean effective ionic diameter for sodium phenobarbital is 0.2 nm) log = (-0.51)(0.09) 1/2 / { (0.2)(0.09) 1/2 } log = = 0.76
The ideal solution is one that is infinitely dilute Most real solutions deviate widely from ideality To simplify calculations, activity coefficients is often neglected and molar concentra tions are applied in equilibrium. For some, the error introduced by the assumption of unity for the activity coefficient is not large enough to lead to false conclusions. However the disregard of activity coefficients may introduce significant numerical error in calculations. Significant discrepancies occur when the ionic strength of the solution is large (>0.01) when the ions involved have multiple charges. For analytical chemistry, normally dilute solutions are used when concentrations become important to our calculations. Summary
Physical Pharmacy, David Attwood and Alexander T Florence, Pharmaceutical Press Physical Pharmacy, Alfred Martin, 4 th Edition. Martin’s physical pharmacy and pharmaceutical sciences, 6 th Edition References