Acids and Bases, pH & pKa, Buffers PDF

Acids and Bases pH & pKa Buffers All figures not specifically referenced either come from: J. M. Berg et al., Biochemistry 7 th edition. W. H. Freeman and Company, 2010. J. M. Berg et al., Biochemistry 8 t...

Acids and Bases pH & pKa Buffers All figures not specifically referenced either come from: J. M. Berg et al., Biochemistry 7 th edition. W. H. Freeman and Company, 2010. J. M. Berg et al., Biochemistry 8 th edition. W. H. Freeman and Company, 2018. Figures from: J. M. Berg et al., Biochemistry 10 th edition, W.H. Freeman MacMillan Learning, 2023 are specified Many figures come from A.L. Lehninger, Principles of Biochemistry, 6 th edition. W.H. Freeman, 2013, Permission granted to use all these figures Due to copyright laws, you CANNOT POST NOR DISTRIBUTE these lecture notes or any course material. They can only be used for personal educational purposes pertaining to this course. 1  Acid-base reaction: Proton transfer Brönsted-Lowry Acid: Proton donor Brönsted-Lowry Base: Proton acceptor Note: H2O can act both as an acid and as a base (autoionization) HA (aq) + H2O (l) H3O+ (aq) + A− (aq) acid base acid base Conjugate acid-base pairs 2  Water acting both as acid and base: Ionization of water H2O (l) + H2O (l) OH- (aq) + H3O+ (aq) 𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔 𝑂𝐻 − [𝐻 + ] E𝑞𝑢𝑖𝑙𝑖𝑏𝑟𝑖𝑢𝑚 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝐊𝐞𝐪 = 𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔 Keq = [𝐻2 𝑂] Keq determined experimentally = 1.8 x 10 –16 at 25C. [H2O] = 55.5 M (as determined from its density d = 1g/mL, Molar Mass = 18g/mol) Rearrange equation: Keq x [H2O] = [OH-][H+] Let Kw = ionic product of water (dissociation constant) at 25° C Kw = Keq x [H2O] = [OH-][H+] = (1.8 x 10-16 M)(55.5 M) = [OH-][H+] = 1.0 x 10-14 3 Relationship between [H3O+], [OH-] and Kw Kw = 𝐻 + 𝑂𝐻 − = 1 x10-14 Take the –log10 on both sides of the equation: since -log (A x B) = -log A + -log B  -log10Kw = -log10 [H+] + -log10 [OH-] Let p = a shorthand notation for log10 -log10(1x10-14) = pH + pOH 14.0 = pH + pOH = pKw convention H3O+ = H+ Note: H+ does not exist in water 4  The pH scale pH = -log10[H+] [H+] [OH-] pOH = -log10[-OH] 14.00 1.0 x 10-14 1.0 x 100 0.00 Basic 11.00 1.0 x10-11 1.0 x 10-3 3.00 7.00 1.0 x 10-7 1.0 x10-7 7.00 Neutral 3.00 1.0 x 10-3 1.0 x 10-11 11.00 0.00 1.0 x 100 1.0 x 10-14 14.00 Acidic If H+ = 1.0 mM or 1.0 x 10-3 M 14.00 = pH + pOH at 25⁰C pH = -log10[H+] If pH = 3.00 = -log(1.0 x 10-3) pOH = 14.00 – pH = 3.00 = 14.0 – 3.00 = 11.00 H3O+ remains present in a strong basic solution (and vice versa) because Kw must always be respected! 5  Equilibrium constants for acids Ka is the equilibrium constant for the dissociation of an acid: HA (aq) + H2O (l) ⇄ H3O+(aq) + A-(aq) Ka = [H3O+] [A-] pKa = -logKa [HA] An acid (HA) is strong if it readily gives up a proton The magnitude of Ka is an expression of the relative strength of the acid The larger the Ka value, the stronger the acid! Ka values are somewhat cumbersome to remember  express and determine relative acidity as pKa The smaller the pKa value, the stronger the acid! 6  Ionization constants for acids HA (aq) + H2O (l) ⇄ H3O+(aq) + A-(aq) Acid name Acid Ka pKa 1x107 Increasing acid strength Hydrochloric acid HCl -7 Hydronium ion H3O+ 1.0 0 Acetic acid CH3COOH 1.8 x 10-5 4.74 Ammonium ion NH4+ 5.6 x 10-10 9.25 Water H2O 1.8 x 10-16 15.7 Ammonia NH3 1 x 10-35 35 𝐻3 𝑂+ 𝐴− 𝐾𝑎 = 𝐻𝐴 Strong acids are ~100% ionized in water, Ka >> 1, product favored Weak acids are