Buffer Solution Presentation.

Why do we need a buffer? If we added 0.060 mol of NaOH to 1.00 L of pure water the change in pH would be from 7.00 to 12.78 pH of 0.060 M NaOH [OH–] = 0.060 M pOH = –log(0.060) = 1.22 pH = 14.00 – 1.22 = 12.78 pH = 12.78 – 7.00 = 5.78 An enormous difference! Buffers prevent wide swings in pH unless a huge excess of strong acid or base are added

Buffer Solution A solution that resists a change in its pH when small quantities of a strong acid or a strong base are added to it. Made up of two solutes: a weak acid and its conjugate base CH3COOH and CH3COO– (CH3COONa) NH4+ (NH4Cl ) and NH3

How a Buffer Works Consider acetic acid HA and its conjugate base acetate A– present in same solution in significant amounts If we add a strong acid, it reacts with the conjugate base H+(aq) + A –(aq)  HA (aq) Thus preventing a buildup of H+ If we add a strong base, it reacts with the conjugate acid OH–(aq) + HA (aq)  A –(aq) + H2O Thus preventing a buildup of OH–

pH of a Buffer Solution pH of a buffer solution containing 0.650 M acetic acid (Ka = 1.8 × 10–5) and 0.450 M sodium acetate. HA(aq) <=> A–(aq) + H+(aq) Ka = [A-][H+] [HA] Henderson – Hasselbach equation pH = pKa + log [A-]

pH of a Buffer Solution We can use initial concentrations of both weak acid and conjugate base as equilibrium concentrations pH = 4.74 + log 0.450M 0.650M pH = 4.58 The pH of a buffer is determined by only the ratio of [A-]/[HA]. A buffer solution containing 0.065M acetic acid and 0.045M sodium acetate would have the same pH, but buffer efficiency is smaller.

Buffer Capacity: β=B /pH Let us consider the change of the pH after addition of 0.060 mol of NaOH to 1L of buffer solution 0.650M HA and 0.450M NaA Acetic acid neutralizes added strong base NaOH HA(aq) + OH–(aq)  A–(aq) + H2O pH = 4.74 + log (0.450 + 0.060)M = 4.68 (0.650 – 0.060)M pH has changed by only 4.68 – 4.58 = 0.10 unit Buffer capacity: mol of added base/ change in pH β=B /pH = (0.060 mol )/ (0.10 pH unit) = 0.60 mol/pH

Effect of Dilution on Buffer Capacity What is the pH after addition of 0.006 mol of NaOH to 1L of buffer solution 0.065M HA and 0.045M NaA HA(aq) + OH–(aq)  A–(aq) + H2O pH = 4.74 + log (0.045 + 0.006)M = 4.68 (0.065 – 0.006)M pH has changed by 4.68 – 4.58 = 0.1 unit Buffer capacity has decreased 10 times after dilution β=B /pH = (0.006 mol )/ (0.10 pH unit) = 0.06 mol/pH

Buffer Capacity Buffer capacity is expressed as the amount of strong base (acid) that must be added to 1 liter of the solution to change its pH by one unit. The buffer capacity depends essentially on 2 factors: Ratio of the acid and its conjugated base. The buffer capacity is optimal when the ratio is 1:1; that is, when pH = pKa Total buffer concentration. For example, it will take more acid or base to deplete a 0.5 M buffer than a 0.05 M buffer.

Buffer Capacity The relationship between buffer capacity and buffer concentrations is given by the Van Slyke equation: β=2.303 x C x Ka x [H3O+] (Ka + [H3O+])2 where C = the total buffer concentration (i.e. the sum of the molar concentrations of acid and its conjugated base). The maximum buffer capacity occurs when pH = pKa

Effectiveness of a Buffer The buffer effectiveness refers to the maximum amount of either strong acid or strong base that can be added before a significant change in the pH will occur. The maximum amount of base that can be added is equal to the amount of weak acid present in the buffer. The maximum amount of strong acid that can be added is equal to the amount of conjugate base present in the buffer

Importance of Buffers The concepts of buffer and buffer capacity appear in varied disciplines, including analytical chemistry, and environmental chemistry, biology, physiology, medicine, dentistry, and agriculture. For example buffer in blood: HCO3– and H2CO3 Absorb acids and bases produced by metabolism Maintains a remarkably constant pH (7.4)

Buffers Important as cells live only in a very narrow pH range. A decrease of blood pH is called acidosis, an increase is called alkalosis. Respiratory and metabolic disorders occur when the body’s acid-base is out of balance. One of the treatments for an acid-base imbalance involves intravenously injecting sodium lactate to restore acid balance and electrolytes to the body.

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