Hydrolysis of Salts Many drugs are formulated as salts, a process that involves taking the weak acid or base and reacting it with base o...
Hydrolysis of Salts
Many drugs are formulated as salts, a process that involves taking the weak acid or base and reacting it with base or acid, respectively, in order to generate an ionic compound, or salt. When salts are dissolved in water, the ions dissociate completely and associate with water molecules to form solvated anions and cations. One common error is to confuse low solubility with low percent dissociation, but these two processes are totally different.
Barium sulfate, for example, has very low water solubility, but whatever amount does dissolve is 100% ionized into Ba+2 and SO4-2. The same applies to many drugs formulated as salts. They may have varying degrees of water solubility, but whatever amount is dissolved in water is 100% ionized into the component ions. Indeed, the decision to formulate a drug as a salt, usually adding one more step in the manufacturing process, comes from the need to obtain greater solubility in body fluids. The salt is invariably more soluble than the parent compound, although some exceptions will be discussed in a later chapter.
Another caveat to add at this point is that whether the drug is administered as a free acid or base, or as its salt, dissolution in water will bring about the corresponding acid-base equilibrium. The equilibrium constant will be satisfied, regardless of whether we begin with the parent compound or one of its salts. The reaction of potassium acetylsalycilate (aspirin, potassium salt) and water is shown below, abbreviating the acid as HA, the salt as KA.
Eq.2.7. Once in solution, A- will pick up H+ (from water dissociation) to form HA. We can write this reaction as:
Eq.2.8. But as soon as H+ is withdrawn from the solution this way, the H2O dissociates some more to replace that H+. This reaction we have seen before:
Eq.2.9. Adding both these reactions gives us the net reaction for the hydrolysis of A-:
Eq.2.10. This is the so-called hydrolysis reaction, and its equilibrium constant, commonly designated by Kh (hydrolysis constant), can be written as follows:
Eq.2.11. We leave out the water from the denominator because as usual its activity stays constant. Multiplying the numerator and denominator by [H+] yields a value for Kh equal to Kw/Ka.
The hydrolysis reaction above will be similar for all salts of drugs that are weak acids and have been reacted with a strong base (i.e. KOH) to form a salt. It also explains why an aqueous solution of the salt of a weak acid is slightly basic, i.e., OH- is generated upon hydrolysis. This last point is not significant in physiological fluids since the latter are commonly buffered, so that they can resist pH changes.
A similar situation arises when the salt we are considering is formed from a parent drug which is a weak base (i.e. epinephrine), and a strong acid (HCl).
Let’s abbreviate epinephrine as RNH2 and its chloride salt as RNH3+Cl- or RNH2.HCl. Upon dissolution of this salt in water there is complete ionization of the salt, as follows:
Eq.2.12. The Cl- does not affect the water dissociation equilibrium (Eq.2.9), but the RNH3+ can because some of it can combine with OH- to produce RNH2 and H2O by the reaction:
Eq.2.13. Adding this equation and the water dissociation we get the following net reaction:
Eq.2.14. This last equation looks like a simple dissociation of a protonic acid, and this is the way we will consider the ionization equilibrium of basic drugs in physiological fluids. In this manner, we can view acids and bases on a similar framework of ionization equilibrium, a perspective which will be convenient for our purposes.
The equilibrium expression for Eq.2.14. has the form:
Eq.2.15. Multiplying numerator and denominator by [OH-] we get:
Eq.2.16. Although it would be most proper to call K the hydrolysis constant for the amine RNH2, for the purposes of our discussion we will call it the dissociation constant (Ka) of the cationic form of the amine, or RNH3+.
Equation 2.16 expresses the inverse relationship between the base strength (Kb) of the parent compound, and the acid strength of its protonated, cationic form:
A strong base (large Kb) has a protonated, cationic form which is a weak acid (low Ka)
A weak base (small Kb) has a protonated, cationic form which is a strong acid (high Ka)
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