Document 6566814

pH, acid and bases, buffer solu2ons Acids and Bases •  A significant majority of biological transforma2ons are catalyzed by acids or bases •  Acids are proton donors and bases are proton acceptors (or hydroxide donors) •  Acids differ in their ability to donate protons •  Strong acids (e.g. HCl or H2SO4) and strong bases (e.g. KOH, NaOH) completely dissociate into anions and ca2ons •  Weak acids and weak bases (eg, Ca[OH]2) dissociate only par2ally in solu2ons Acid-­‐Base Defini2ons Bronsted-­‐Lowry Acid and Bases: •  An acid is a substance that donates a proton (H+) and a base is a substance that accepts a proton Lewis acids and bases: •  Lewis acid is a compound that accepts an electron pair from a base; lewis base is a compound that donates an electron pair to an acid – a more general defini2on and also includes metal ca2ons such as Mg2+, Fe3+, and Zn2+ as lewis acids Electrophiles are lewis acids and nucleophiles are lewis bases; However the term acid and base is generally used when the electrons are donated to H+ and the term electrophile-­‐nucleophile is generally used when electrons are donated to a carbon atom hUp://science.uvu.edu/ochem/wp-­‐content/images/B/bronsted4.png hUp://employees.csbsju.edu/cschaller/principles%20chem/acidity/Lewisabc.gif Water as an acid and a base • 
Water can act both as an acid and a base •  Water has ability to ionize slightly, into hydronium and hydroxide ions H2O + H2O H3O+ + OH− •  For dissocia2on of water, K is the dissociaBon constant ü  The probability that a hydrogen in pure water will exist as a hydrogen ion (H+) is 1.8 × 10−9, ü  The molar concentra2on of H+ ions (or of OH− ions) in pure water is the product of the probability, 1.8 × 10−9, 2mes the molar concentra2on of water, 55.56 mol/L. The result is 1.0 × 10−7 mol/L, Thus, dissocia2on constant for water is: We can neglect changes in water concentra2on and Instead of K, we can use the simplified term “ion product”, Kw at 25 °C What is pH? •  pH is the nega2ve log of the hydrogen ion concentra2on: pH = −log [H+ ] •  To calculate the pH of a solu2on: •  1) Calculate hydrogen ion concentra2on [H+] 2) Calculate the base 10 logarithm of [H+] 3) pH is the negaBve of the value found in step 2 •  For example, for pure water at 25°C, •  pH=−log[H+]=−log(10−7) =−(−7)=7.0 •  Low pH, high H+ concentra2on, high acidity •  High pH, low H+ (or high OH), high basicity Calcula2on of pH of Solu2ons •  What is the pH of a solu2on in which HCl (hydrochloric acid) concentra2on is 3.2 × 10− 4 mol/L? •  Answer: Since strong acid and bases dissociate completely in solu2on, the concentra2on of hydronium ion is equal to the concentra2on of hydrochloric acid, HCl H+ + Cl-­‐ so; pH = −log [H+ ] = −log (3.2×10−4) = −log (3.2)−log (10−4) = (−0.5) + 4.0 = 3.5 pH and pOH •  pOH is equal to −log [OH−] and it is derived from the defini2on of Kw: Kw =[H+][OH−]=10−14 log [H+ ] + log [OH− ] = log 10−14 pH + pOH = 14 (always true in a aqueous solu2on) (pOH is not commonly used, pH is generally preferred) Example •  What is the pH of a solu2on whose hydroxide ion concentra2on is 4.0 × 10− 4 mol/L? Answer: [OH−]=4.0×10−4 pOH = −log [OH− ] = −log (4.0×10−4) = −log (4.0)−log (10−4) = (−0.60) + 4.0 = 3.4 •  If pOH is 3.4, than pH = 14 – 3.4 = 10.6 pH = 10.6 pH of Weak Acid and Base Solu2ons •  Weak acid and bases do not dissociate completely into their ions, so the concentra2on of hydronium ion or hydroxide ion is related to the dissocia2on constant of the weak acid and base •  Many biomolecules contain func2onal groups such as carboxylic acids, amino groups, phosphate groups, which have weak acidic and basic characteris2cs •  Because of this, the influence of intracellular pH is important for the structure and func2on of these biomolecules hUp://scienceaid.co.uk/chemistry/physical/images/ka.png Conjugate Acid and Conjugate Base •  R—CH2 —COOH R—CH2 —COO− pKa = 4−5 •  (Acid) (conjugate base) •  R—CH —NH+ R—CH —NH pKa =9−10 (conjugate acid) (base) •  pKa is the nega2ve log of the dissocia2on constant of the acid (Ka); pKa= −logK •  pKa is used to express the relaBve strengths of acids and bases •  The stronger the acid is, the lower is its pKa value •  For any weak acid, its conjugate is a strong base pKa •  When the concentraBons of weak acid and its conjugate are the same then, pKa = pH •  Ka =[H+] −log Ka = −log [H+] pKa = pH •  pKa of an acid group is the pH at which the protonated and unprotonated species are present at equal concentra2ons. •  This property is important for buffer systems pKa values might change depending on the medium •  The pKa of a func2onal group is influenced by the surrounding medium (depending on whether the undissociated acid or its conjugate base is the charged species) •  The effect of dielectric constant on pKa may be observed by adding ethanol to water: The pKa of a carboxylic acid increases, whereas that of an amine decreases (because ethanol decreases the ability of water to solvate a charged species) •  The pKa values of dissocia2ng groups in the interiors of proteins thus are profoundly affected by their local environment, including the presence or absence of water. Buffer Solu2ons and Biological Buffers Buffer Solu2ons •  Solu2ons of weak acids or bases and their conjugates exhibit ability to resist a change in pH following addi2on of strong acid or base •  This is called buffering and such solu2ons are named as buffer soluBons •  Many metabolic reac2ons occur in buffered environments to prevent pH changes due to release or uptake of protons during these biological reac2ons •  Also, oxida2ve metabolism produces CO2, which if not buffered would produce severe acidosis – elevate pH (through forma2on of carbonic acid •  Therefore, maintenance of a constant pH is important for human body and this process involves buffering by phosphate, bicarbonate, and proteins, which accept or release protons to resist a change in pH Henderson-­‐Hasselbach Equa2on •  A weak acid, HA, ionizes as follows: HA H+ +A− The equilibrium constant for this dissocia=on is: Ka = [H+][A−] / [HA] If we rearrange this equa2on, it becomes: [H+]=Ka [HA] / [A−] Then, if we take the log of both sides, it becomes: •  If we replace pKa for (– log Ka) and pH for (–H+), then the equa2on becomes: With a slight modifica2on we obtain the Henderson-­‐Hasselbach equaBon The Use of Henderson-­‐Hasselbach Equa2on The Henderson-­‐Hasselbalch equa2on has an important predic2ve use in weak acid equilibrium: •  (1) When an acid is exactly half-­‐neutralized, [A−] = [HA]. At this condi2on (half-­‐
neutraliza2on), pH = pKa •  (2) For example, when the ra2o [A−]/[HA] = 100:1, pH=pKa+ log ([A−]/[HA]) pH=pKa +log (100/1) = pKa + 2 •  (3) Or, when the ra2o [A−]/[HA] = 1:10, pH=pKa +log 1/10=pKa + (−1) •  This way we can predict how much the pH of a weak acid solu2on change if a certain amount of strong acid/base is added to the solu2on – buffer capacity •  If the equa2on is evaluated at ra2os of [A−]/[HA] ranging from 103 to 10−3 and the calculated pH values are ploUed, the resul2ng graph describes the 2tra2on curve for a weak acid • 
Harpers-­‐Illustrated Biochemistry-­‐28th edi2on, Mc-­‐Graw Hill Company Proper2es of Buffers 1) A buffer solu2on resists changes in pH most effec2vely at pH values close to the pKa of the buffer system 2) A buffer system buffers most effec2vely in the pH range pKa ± 1.0 pH unit 3) The higher the concentra2on of the buffer is, the greater is its buffer capacity How to Choose the Appropriate Buffer? •  First determine at which pH we want to work at, for example we want the pH to be constant around 7.5 •  Then look at the pKa’s of common buffers (pKa tables for common buffers are easily available from internet) – choose the buffer with a pKa that is closest to pH 7.5 •  The pKa of the buffer should be within (+/-­‐) 1 of our working pH •  Higher concentra2on gives higher buffer capacity, but too high concentra2on might have other disadvantages Buffer Solu2ons •  For experiments using 2ssue extracts or en-­‐ zymes, constant pH is maintained by the addi2on of buffers such as; •  MES ([2-­‐N-­‐morpholino]ethanesulfonic acid, pKa 6.1), •  Inorganic orthophosphate (pKa2 7.2), •  HEPES (N-­‐hydroxyethylpiperazine-­‐N9-­‐2-­‐ethanesulfonic acid, pKa 6.8) •  Tris (tris[hydroxymethyl] amino-­‐ methane, pKa 8.3). •  The value of pKa rela2ve to the desired pH is the major determinant of which buffer is selected. Which factors determine the strength of acids and bases? •  The presence of charge near the dissocia2ng proton affects the pKa value •  Presence of a nega2ve charge, for example, disfavors the release of a proton from a carboxylic acid, thus increasing the pKa value