ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 9 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 9 BUFFERS

HYDROLYSIS OF SALTS - Reaction of salt with water to produce hydronium ion or hydroxide ion or both (do not go to 100% completion) - Not all salts hydrolyze - The salt of a strong acid and a strong base does not hydrolyze - Neutral solution is the result - The salt of a strong acid and a weak base hydrolyzes - Acidic solution is the result

HYDROLYSIS OF SALTS - The salt of a weak acid and a strong base hydrolyzes - Basic solution is the result - The salt of a weak acid and a weak base hydrolyzes Slightly acidic, neutral, or basic, depending on relative weaknesses of acid and base

Acidic Hydrolysis positive ion of salt + H 2 O Conjugate base + H 3 O + - The hydronium ion makes the solution acidic NH H 2 O → NH 3 + H 3 O + HYDROLYSIS OF SALTS

Basic Hydrolysis negative ion of salt + H 2 O Conjugate acid + OH - - The hydroxide ion makes the solution basic F - + H 2 O → HF + OH - HYDROLYSIS OF SALTS

BUFFER SOLUTION - A mixture of a conjugate acid-base pair - Tends to resist changes in pH upon addition of an acid or a base - The resistive action is the result of equilibrium between the weak acid (HA) and its conjugate base (A - ) HA(aq) + H 2 O(l) → H 3 O + (aq) + A - (aq) - Commonly used in biological systems - Enzyme-catalyzed reactions depend on pH

BUFFER SOLUTION Examples HC 2 H 3 O 2 /C 2 H 3 O 2 - HF/F - NH 3 /NH 4 + H 2 CO 3 /HCO 3 -

- When an acid is added, the conjugate base converts the excess H 3 O + ion into its acid (conjugate base removes excess H 3 O + ) H 3 O + (aq) + A - (aq) → HA(aq) + H 2 O(l) - When a base is added, the acid converts the excess OH - ion into its conjugate base and water (acid removes excess OH - ion) HA(aq) + OH - (aq) → A - (aq) + H 2 O(l) - These reactions go to completion (large equilibrium constants) BUFFER SOLUTION

- In actual fact, the pH changes but very slightly - Large amounts of added H 3 O + or OH - may overcome the buffer action and change pH of solutions - Buffers are most effective when the ratio of acid to conjugate base is 1:1 - Buffers are less efficient in handling acids if the acid is more than the conjugate base - Buffers are less efficient in handling bases if the acid is less than the conjugate base BUFFER SOLUTION

HA(aq) + H 2 O(l) ↔ H 3 O + (aq) + A - (aq) HENDERSON-HASSELBALCH EQUATION -If [HA] = [A - ], then K a = [H 3 O + ] Taking the negative logarithm on both sides gives - logK a = - log [H 3 O + ] pK a = pH

- If [HA] ≠ [A - ] HENDERSON-HASSELBALCH EQUATION - pH changes by 1 if the ratio changes by a factor of 10 or

HENDERSON-HASSELBALCH EQUATION - C a and C b are the analytical concentrations of the acid and the conjugate base, respectively - n a and n b are the number of moles of the acid and the conjugate base, respectively C a = n a /V and C b = n b /V (concentration ratio equals mole ratio) In general or

HENDERSON-HASSELBALCH EQUATION Addition of a Strong Acid - Some buffer base is converted to the conjugate acid Addition of a Strong Base - Some buffer acid is converted to the conjugate base

PREPARING BUFFERS - Measure the amount of weak acid (HA) to be used [or weak base (B)] - Calculate the amount of strong base (OH - ) to be added [or strong acid (H + )] - This makes a mixture of HA and A - which is a buffer [or B and BH + ] Or - Add the correct proportions of the acid and its conjugate base, then check the pH

Prepare 1.00 L 0f M tris buffer solution at pH Convert mole to gram tris hydrochloride (MM = g/mol) - Weigh out sample and dissolve in a beaker with ~ 800 mL H 2 O - Add NaOH solution until pH is exactly 8.40 (continuous stirring) - Quantitatively transfer solution to a 1000 mL volumetric flask - Dilute to the mark and mix PREPARING BUFFERS

- A measure of how well a solution resists changes in pH - Increases with increasing concentration of buffer - Maximum when pH = pK a - The greater the buffer capacity the less the pH changes upon addition of H + or OH - - Choose a buffer whose pK a is closest to the desired pH - pH should be within pK a ± 1 BUFFER CAPACITY

- Acid-base indicators are highly colored weak acids or bases - The various protonated forms have different colors - Have very low concentrations in order not to interfere with analytes INDICATORS

HIn + H 2 O ↔ H 3 O + + In - - Indicators exist predominantly in the protonated form (HIn) in acidic solutions - Indicators exist predominantly in the deprotonated form (In - ) in basic solutions - The result of color changes