Buffers

Plan for Today

• Review Homework 7-6

• Buffers

• Lewis Acid/Base Reactions for Synthesis

Buffers

• Buffer Definition: A buffer is an acid-base system that resists changes in pH.

• Buffer Region: Observed in a graph where the pH does not change substantially despite additions of acid or base.

Henderson-Hasselbalch Equation

• The pH of a buffer solution can be determined using the Henderson-Hasselbalch relationship:

  • \text{pH} = \text{pKa} + \log\left(\frac{[\text{base}]}{[\text{acid}]}
    ight)

• This equation is most useful when:

  • 0.10 < \frac{[\text{conjugate base}]0}{[\text{conjugate acid}]0} < 10

  • Where [conjugate base]0 and [conjugate acid]0 refer to the initial concentrations in the solution.

  • Assumes minimal change in initial concentrations as equilibrium is approached.

Buffer Composition

• A buffer contains:

  • A weak acid (HA) and its conjugate base (A−).

  • Alternatively, a weak base and its conjugate acid.

Reaction with Strong Base

• If a strong base (B−) is added to the buffer:

  • Reaction occurs with the weak acid:

  • $\text{HA} + \text{B}^- \to \text{HB} + \text{A}^-$

  • Where HA is the weak acid and the product is a weaker acid.

Reaction with Strong Acid

• If a strong acid (HX) is added to the buffer:

  • Reaction occurs with the conjugate base:

  • $\text{A}^- + \text{HX} \to \text{HA} + \text{X}^-$

Acid Addition to Water

• Inquiry about the pH of different solutions:

  • Pure water: pH = 7

  • 0.01 M HCl: Select options:

  • A. 1

  • B. 2

  • C. 7

  • D. 12

Acid Addition to a Buffer

• Adding 0.01 mol of HCl to 1 L buffer with 1.00 M acetic acid and 1.00 M acetate:

  • Which part of the buffer reacts?

  • Options:

    • A. AcOH

    • B. –OAc

    • C. Water

    • D. No reaction

  • Effect on pH after addition:

  • Options:

    • A. Increase (a bit)

    • B. Increase (a lot)

    • C. Decrease (a bit)

    • D. Decrease (a lot)

    • E. No reaction

Determining pH Changes

• Adding HCl leads to:

• Reaction with the basic component (-OAc) of the buffer:

  • Increase in AcOH, decrease in –OAc.

  • The equilibrium constant (Ka) of acetic acid: $1.7 × 10^{-5}$.

Buffer Calculations

• Using the Henderson-Hasselbalch equation to calculate:

  • New pH after reaction:

  • Given:

  • $\text{AcOH + H}<em>2\text{O} \rightleftharpoons \text{H}</em>3\text{O}^+ + \text{–OAc}$

Adding Base to Water

• Inquiry about pH of solutions:

  • Pure water: pH = 7

  • 0.01 M NaOH: Select options:

  • A. 2

  • B. 12

  • C. 7

  • D. 14

Base Addition to a Buffer

• Adding 0.01 mol NaOH to 1 L buffer with 1.00 M acetic acid and 1.00 M acetate:

  • Identify reacting part:

  • Options:

    • A. Acid (acetic acid)

    • B. Base (acetate)

  • What happens to the pH?

  • Options:

    • A. Increase

    • B. Decrease

  • New pKa of acetic acid: 4.77.

• New pH calculations are needed after the addition.

• Reaction specifics:

  • NaOH reacts with acidic part (AcOH), leading to:

  • Decrease in AcOH, increase in AcO−.

Preparing a Buffer

• Two methods for buffer preparation:

  1. Mix a weak acid with its conjugate base and use the Henderson-Hasselbalch equation to find the ratio of [A−]/[HA].

  2. Start with a weak acid and titrate with a strong base to achieve a desired pH.

Buffer Capacity

• Definition: The buffer's capacity indicates the amount (mol) of acid/base that can be added before pH changes by ±1 from the pKa.

Buffer Range

• Definition: Range of pH over which the buffer remains effective.

• Typically within ±1 of the pKa of the acid.

• Examples of buffering options for various pHs:

  • pH 4:

  • pH 7:

  • pH 10:

  • Acid and corresponding Ka and pKa specifications:

  • A. HF, Ka = 7.2 × 10–4, pKa = 3.14

  • B. CH3COOH, Ka = 1.76 × 10–5, pKa = 4.75

  • C. H2CO3 (first), Ka = 4.3 × 10–7, pKa = 6.37

  • D. H2CO3 (second), Ka = 4.8 × 10–11, pKa = 10.32

  • E. NH4Cl, Ka = 5.6 × 10–10, pKa = 9.24

Buffer Capacity Examples

• Example Case 1: Addition of 0.1 mol HCl to 1 L of buffer (1.00 M acetic acid and 1.00 M acetate) with given $Ka = 1.7 × 10^{-5}$:

• Calculate new pH: Options:

  • A. 7

  • B. 4.77

  • C. 4.68

  • D. 4.86

• After addition, check if the system remains a buffer.

• Example Case 2: Addition of 1.0 mol HCl to the same buffer:

• Calculate new pH, similar considerations apply as with Example Case 1.

• For both examples after adding HCl, determine if the system remains a buffer.

Lewis Acid/Base Theory

• Extends beyond H+ transfer.

• Definition of Lewis Acids and Bases:

  • Lewis Bases: Electron-rich species that can donate electron density.

  • Lewis Acids: Electron-poor species that can accept electron density.

  • Examples of Lewis Acids: Fe3+, Ca2+, Pb2+, Ni2+, Al3+, BF3, H+.

Arrow Pushing in Lewis Acid/Base Reactions

• Procedure:

  1. Draw Lewis structures to identify acids and bases.

  2. Use curved arrows to indicate electron movement from base to acid.

  3. Draw products with appropriate formal charges.

Reaction Mechanism Example

• Forming the product from HCl and NH3:

• Involves electron donation from NH3 (nucleophile) to H+ (electrophile).

Electrophiles and Nucleophiles

• Electrophiles: Molecules with partial positive charge that accept electrons (e.g., CH3Br when the positive end is carbon).

• Nucleophiles: Molecules with partial negative charge that donate electrons (e.g., OH−).

• Multi-step process can be illustrated with curved arrows.

Review of Key Concepts

• Summary of Lewis Acid and Base roles:

  • Nucleophile: Lewis Base, donates electrons.

  • Electrophile: Lewis Acid, accepts electrons.

Upcoming Tasks

• By Saturday, 11/8 at 8:30 AM:

  • Complete and upload Week 10 Discussion packet. All are required to submit, regardless of meeting status.

  • By Monday, 11/10 at 8:30 AM:

  • Complete and upload Homework 7-7. Will review in class on Tuesday.