General Chemistry II Acid and Base Equilibria Lecture 3

Acids and Bases - 3 Lectures Autoionization of Water and pH (completed) Defining Acids and Bases (completed) Interaction of Acids and Bases with Water (completed) Buffer Solutions (part-completed) Acid-Base Titrations Outline - 5 Subtopics

By the end of this lecture AND completion of the set problems, you should : Know the Henderson-Hasselbalch equation and its use in calculating pH values for buffer solutions. Understand pH and the effective range of buffer activity. Be familiar with acid-base titrations, titration curves for strong and weak acids and bases. Understand the selection of indicators and effective pH ranges for end-point detection. Objectives - Lecture 3 Acids and Bases

BUFFERSOLUTIONS(Quantitative)

Henderson-Hasselbalch equation At equilibrium H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq) WEAK ACIDS

H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq) EFFECTIVE BUFFER RANGE As a ‘rule of thumb’, buffering is effective from [A - ] = 0.1 [HA] to [A - ] = 10 [HA] Thus pH is buffered in the range pK a ± 1 For a buffer at a desired pH, choose an acid-conjugate base pair with a pK a as close as possible to this pH

ACID-BASE INDICATORS Acid - base indicators are weak acids where the conjugate acid is a different color than its conjugate base. The actual color of the solution depends on the fraction of the indicator in each of its two forms i.e. on the value of the pH relative to the pK a of the indicator. Bromothymol Blue pK a = 6.8 pH 90% in acid form pH > 7.6 Blue > 90% in base form pH = 6.8 Green 50% acid form 50% base form

ACID-BASE INDICATORS methyl red red to yellow bromthymol blue yellow to blue phenolphthalein colorless to pink

TITRATIONS

TITRATIONS Acid-Base Titrations are convenient procedures for accurately determining the concentration of an acid or base in aqueous solution. It relies on the neutralization reaction between the acid and base: OH - (aq) + HB (aq) B - (aq) + H 2 O(l) All of the following slides can be considered in terms of titrating strong & weak bases with a strong acid too

TITRATIONS Equivalence Point when the amount of OH - added = the amount of acid present initially. For strong acid = initial amount of H + For weak acid = initial amount of HA (prior to ionization) H 2 O(l) + HA(aq) H 3 O + (aq) + A - (aq) OH - (aq) + H 3 O + (aq) H 2 O(l) + H 2 O(l)

STRONG ACID WITH STRONG BASE H 3 O + = OH - “Point of inflexion” At 25 ° C = pH 7.00

STRONG ACID WITH STRONG BASE Addition of excess OH - gives “End point” is when the pH reaches a predefined value. Often the pK a of an indicator or “Indicator error” is the difference between the amount of OH - required to reach the equivalence point and the end-point.

STRONG ACID WITH STRONG BASE At the End point, we know [OH - ] = [H + ] For a monoprotic acid: mol OH - = mol H + [OH - ] (mol/L) x Titre (L) = [H + ] (mol/L) x Aliquot (L) Titre = volume of base added by burette to reach End Point Aliquot = volume of acid added to flask [OH - ] & Aliquot are known initially, Titre is determined Hence can calculate [H + ]

WEAK ACID WITH STRONG BASE “Point of inflexion”  7.00 WHY? HA + H 2 O H 3 O + + A -

WEAK ACID WITH STRONG BASE 1.The starting pH is high because of the low degree of ionisation of the weak acid 2.The acid reacts completely with the added base. The presence of the conjugate base suppresses the dissociation of the acid 3.On addition of half of the base, the amounts of acid and base present in the solution are equal and pH = pK a of the acid HA + H 2 O H 3 O + + A -

POLYPROTIC ACIDS 1. Many acids have more than one proton that they can donate to a base. 2. The K a values for the dissociation of the 2nd and higher protons are smaller than the first. CO 2 (aq) +H 2 O H 2 CO 3 (aq) H 2 CO 3 (aq) + H 2 O HCO 3 - (aq) + H 3 O + (aq) K a1 = 4.3 x HCO 3 - (aq) + H 2 O CO 3 2- (aq) + H 3 O + (aq) K a2 = 4.4 x The relative concentrations of the different species present depend on pH - Low pH: species is in highly protonated form High pH: species is in deprotonated form

POLYPROTIC ACIDS H 2 CO 3 pK a1 = 6.37 HCO 3 - pK a2 = CO 3 2-

Acids & Bases Gold mining frequently uses the base, cyanide (CN - ), in the extraction process. Tailings dams sometimes have high concentrations of CN -. The conjugate acid of CN -, HCN, is extremely toxic. “Acid mine drainage” (AMD) is naturally produced by the exposure of sulfide ores to water. This process is greatly exacerbated by mining the sulfide ores (for Cu, Pb, Zn etc). AMD results in streams with elevated H + concentrations and pH values in the range 1-3. A gold mining tailings dam leaks into an AMD-affected stream - would you evacuate? Would you evacuate?

Acids & Bases For HCN, pK a = 9.14 so K a = 7.2 x pH values in the range 1-3, hence [H + ] = M Would you evacuate? Hence > % of cyanide is present as HCN!! The health issue is also related to the solubility of HCN gas, BUT immediate attention is warranted!!

Acids and Bases - End of Lecture 3 After studying this lecture and the set problems, you should : Know the Henderson-Hasselbalch equation and its use in calculating pH values for buffer solutions. Understand pH and the effective range of buffer activity. Be familiar with acid-base titrations, titration curves for strong and weak acids and bases. Understand the selection of indicators and effective pH ranges for end-point detection. Objectives Covered in Lecture 3