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Ionization of Water, Weak Acids, and Weak Electolytes Adopted from Nilsen and Cox – Lehninger principles of biochemistry (sixth edition)  Ionization of water Breaking down of water as a compound into its constituent ions Pure water- slightly ionized With a slight tendency t...

Ionization of Water, Weak Acids, and Weak Electolytes Adopted from Nilsen and Cox – Lehninger principles of biochemistry (sixth edition)  Ionization of water Breaking down of water as a compound into its constituent ions Pure water- slightly ionized With a slight tendency to undergo reversible ionization to yield a hydrogen ion (a proton) and a hydroxide ion  Ionization of water Measures by its electrical conductivity Pure water- carries electrical current as H3O+ migrates towards the cathode and OH- towards the anode Small degree of ionization at equilibrium (2 of every 109 molecules in pure water at 250C) Expressed at equilibrium constant  Ionization of water Pure water at 250C : conc. water 55.5 M (grams of H2O in 1L /MW: (1,000 g/L)(18.015 g/mol)) Expressed at equilibrium constant (Show calculations for pH of water)  Ionization of Water Keq = [H+][OH-] / [H2O] = 1.8 x 10-16 M Concentration of water - one liter = 1,000g Mole Wt Water = 18.015 [H2O] = 55.5 M Kw = [H+][OH-] = Keq x [H2O] = 1 x 10-14 M2 for pure water [H+] = [OH-] so, [H+] = 10-7 M pH is negative log [H+] , for pure water = 7.0 pH and Buffers  + The pH Scale Designates the H and - OH Concentrations The term pH is defined by the expression pH = log _1_ = -log [H+] [H+] The symbol p denotes “ negative logarithm of” Neutral solution at 250C : conc. of [H+]: 1.0 x 10-7 M  Calculation for neutral solution at 0 25 C pH = log ____1____ = 7.0 1.0 x 10-7 * The concentration of H+ must be expressed in terms of molar (M)  A difference of 1 unit of pH means that the solution has 10x the H+ concentration of the other This does not tell us the absolute magnitude of difference. The pH of most aqueous liquids.  Approximation of pH A. Indicator dyes Litmus Phenolphthalein Phenol red -undergo color changes as a proton dissociates from the dye molecule  Accurate determination of pH pH meter - A glass electrode - Selectively sensitive to H+ concentration - Insensitive to Na+, K+ and other cations - A signal form a glass electrode is placed in a test solution is amplified and compared with a signal generated by a solution with an accurate pH Weak Acids and Bases Have Characteristic Acid Dissociation Constants Strong acids and bases completely ionized in dilute aqueous solutions Acids- proton donors Bases- proton acceptors Conjugate acid-base pair- a proton donor and its corresponding proton acceptor Example: acetic acid (CH3COOH) and acetate anion ( CH3COO-)  The stronger the acid, the greater its tendency to lose its protons. Keq- the tendency of any acid [HA] to lose a proton and form its conjugate base[A-] for the reversible reaction  Ionization constants/ Acid dissociation constant Ka equilibrium constants for ionization reactions Stronger acids have larger dissociation constants ( phosphoric acids, carbonic acids) Weaker acids have smaller dissociation constants ( monohydrogen phosphate) Stronger acid = lower pKa A  Weak Acids HA ↔H + + A- K e = [H+][A-] / [HA] = Ka Henderson-Hasselbalch Equation Rearranges Ka pH = pKa + log ( [A-] / [HA] ) when pKa = pH … [A-] = [HA] Titration Curved Reveal the pKa of Weak Acids Titration- used to determine the amount of an acid in a given solution A measure volume of the acid is titrated with a solution of a strong base (NaOH), of known concentration. Strong base is added in increments until the acid is neutralized (pH paper/pH meter) Titration curve- a plot of pH against the amount of NaOH added reveals the pKa of the weak acid The titration curve of acetic acid. Weak acids have different pKas Buffering against pH Changes in Biological Systems  Cells and organisms maintain a specific and constant cytosolic pH (near 7), keeping the biomolecules in their optimal ionic state. The pH of extracellular fluids is tightly regulated in multicellular organisms. Constant pH is achieved by biological buffers: mixtures of weak acids and their conjugate bases. Buffers are Mixtures of Weak Acids and Their Conjugate Bases Buffers- aqueous systems that tend to resist changes in pH when small amounts of acid (H+) or base (OH-) are added A buffer system consists of a weak acid and a conjugate base (ex. Acetic acid and acetate ion) Maximal buffering power is where the concentration of the proton donor is equals to that of the proton acceptor. A  The Henderson-Hasselbalch Equation Relates pH, pKa, and Buffer concentration restates the expression for the ionization constant of an acid. calculates for pKa, given pH and molar ratio of proton donor and acceptor Calculates the molar ratio of proton donor and acceptor, given pH and pKa  Weak Acids or Bases Buffer Cells and Tissues against pH changes Weak acids and bases in the cytoplasm: high protein conc.(AA with functional groups) Histidine side chain: pKa: 6.0 Histidine in proteins buffers effectively near neutral pH Histidine side chain is either protonated or unprotonated from near neutral pH  Ionization of histidine. The amino acid histidine, a component of proteins, is a weak acid. The pKa of the protonated nitrogen of the side chain is 6.0. Show sample calculation for the ionization of histidine  Untreated Diabetes Produce Life-Threatening Acidosis Normal blood plasma pH: 7.35-7.45 pH optimum- the maximal catalytic activity of enzymes at a characteristic pH Lack of insulin, insensitivity to insulin, untreated diabetes- forces the tissues to store FAs as their primary fuel. Dependence to FAs leads to the accumulation of high concentrations of β-hydroxybutyric acid and acetoacetic acid  Dissociation of this acid lowers the pH of the blood plasma to > 7.35 leading to acidosis. Severe acidosis: headache, drowsiness, nausea, vomiting, and diarrhea, followed by stupor, coma, and convulsions Presumption: lower pH: not functional enzymes Diagnostic test: blood and urine  Other conditions that causes acidosis: 1. fasting and starvation 2. heavy exertion (lactic acid in the blood) 3.Kidney failure- diminished capacity to regulate bicarbonate levels 4. Lung diseases (pneumonia, emphysema, asthma) reduce the capacity to dispose CO2 produced by fuel oxidation in tissues accumulating carbonic acid Enzymes have pH optima Related to their Function End of presentation