Pharmaceutical Solutions Lecture Notes PDF

Pharmaceutical solutions Acids and bases Ions in solution Assuming solutions in water  Acids and Bases An acid is a proton (H+) donor HA → H+ + A− A base is a proton (H+) acceptor B + H+ → BH+ The above more strictly known as Brønsted acids and bases Le...

Pharmaceutical solutions Acids and bases Ions in solution Assuming solutions in water  Acids and Bases An acid is a proton (H+) donor HA → H+ + A− A base is a proton (H+) acceptor B + H+ → BH+ The above more strictly known as Brønsted acids and bases Lewis acids and bases (a more general definition) – An acid is an electron-pair acceptor – A base is an electron-pair donor  Protons (H+) in water Protons are highly solvated in water Represented by the hydronium ion H3O+ HCl + H2O → H3O+ + Cl− Acid: HCl; base: H2O  Conjugate acids and bases AH + H2O → H3O+ + A− B + H2O → BH+ + HO− AH: acid; B: base A− : conjugate base; BH+: conjugate acid Acids can be neutral, cationic or anionic Bases can be anionic, neutral or cationic Examples  Conjugate acids of bases as acids For any base B, can consider the conjugate acid, BH+, as an acid in its own right E.g: Considering NH4+ as an acid:  General equilibrium for proton transfer Can generally write an equation for the process of an acid transferring a proton to a base (in water) e.g: Equilibrium constant: 𝑎 𝐻3𝑂 + 𝑎 𝐵𝑎𝑠𝑒 𝐾= 𝑎 𝐴𝑐𝑖𝑑 𝑎 𝐻2𝑂  Activity and concentration Equilibrium constants, K, for acid-base equilibria determined by the activities, a, of the components But, provided dealing with dilute solutions… …can replace activities by concentrations as an approximation… …and can assume activity/concentration of water ([H2O]) is constant  Acidity of aqueous solutions Determined by [H3O+] Wide range of concentrations, e.g. 1 mol L−1 to 10−14 mol L−1 Sørensen (1909) defined: pH = −log10 [H3O+] ‘p’ for ‘power’ 10−2 mol L−1 : pH 2; 10−10 mol L−1 : pH 10 Now many functions of the form: pX = −log10 X  Acid dissociation in water Generally considering dilute aqueous solutions [H2O] relatively constant Ka is the acidity constant for acid AH  pKa pKa = −log10Ka Measure of acid strength Strong acid: high Ka, low pKa Weak acid: low Ka, high pKa Compound pKa at 25 °C H2O 15.74 H3O+ −1.74 CH3CO2H 4.76 NH4+ 9.24  Polyprotic acids E.g: H3PO4 [H3O+] in aq. H3PO4 dominated by first dissociation Note: species such as H2PO4- and HPO42- capable of acting as acid or base  Bases in water pKb = −log10Kb  pKa of conjugate acid as a measure of base strength If B a strong base, conjugate acid BH+ a weak acid and vice versa Hence can use pKa of BH+ as a measure of basicity of B  Dissociation of water (at 25 °C) Also known as the autoprotolysis of water Kw is known as the autoprotolysis constant [H3O+][HO−] = 10−14 (−log10[H3O+]) + (−log10[HO−]) = −log10(10−14) pH + pOH = 14 Neutrality at pH 7 At 37 °C, Kw = 13.68, neutrality at pH 6.84