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Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry , 2007 (John Wiley)      ISBN: 9 78047081 0866 . CHEM1002 [Part 2]. A/Prof Adam Bridgeman (Series 1) Dr Feike Dijkstra (Series 2) Weeks 8 – 13

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Unless otherwise stated, all images in this file have been reproduced from:

Blackman, Bottle, Schmid, Mocerino and Wille,Chemistry, 2007 (John Wiley)     ISBN: 9 78047081 0866


CHEM1002 [Part 2]

A/Prof Adam Bridgeman (Series 1)

Dr FeikeDijkstra (Series 2)

Weeks 8 – 13

Office Hours: Monday 2-3, Friday 1-2

Room: 543a

e-mail:[email protected]

e-mail:[email protected]


Acids & Bases 3

  • Lecture 2:

  • Strong/Weak Acids and Bases

  • Calculations

  • Polyprotic Acids

  • Lecture 3:

  • Salts of Acids and Bases

  • Buffer systems

  • Blackman Chapter 11, Sections 11.3-11.6

Reproduced from ‘The Extraordinary Chemistry of Ordinary Things, C.H. Snyder, Wiley, 2002(Page 245)


Is a solution of NaCN acidic or basic?

NaCN is the salt of NaOH (strong base) and HCN (weak acid). The base “wins”: pH > 7

Overall reaction is

H2O(aq) + CN–(aq) OH–(aq) + HCN(aq)

Does a solution of NH4Cl have pH > 7 or < 7?

Salt of NH4OH (weak base) and HCl (strong acid)

acid “wins”: pH < 7 and reaction is

H2O(aq) + NH4+(aq) NH3(aq) + H3O+(aq)

Salts of Weak Acids and Bases


The Common Ion Effect

  • If you add the salt of an acid to a solution of the same acid then the equilibrium will shift towards neutral.

CH3COOH(aq) + H2O(l) CH3COO-(aq)+ H3O+(aq)

  • Addition of CH3COO-:

    • By Le Chatelier’s principle the equilibrium will shift to the left to remove CH3COO- and therefore decrease [H3O+].

  • Addition of CH3COOH:

    • By Le Chatelier’s principle the equilibrium will shift to the right to remove CH3COOH and therefore increase [H3O+].


A solution containing both a weak acid and its salt withstands pH changes when acid or base (limited amounts) are added.

Buffer with equalconcentrations of conjugate base and acid

Buffer after addition of H3O+

Buffer after addition of OH-

H3O+

OH-

CH3COO-

CH3COOH

CH3COO-

CH3COOH

CH3COO-

CH3COOH

Buffer System

H2O + CH3COOH  H3O+ + CH3COO-

CH3COOH + OH- H2O+ CH3COO-


Consider change in pH of pure water (pH = 7) if we add an equal amount of 10–3 M HCl:

[H+ ] = 1 x 10–3 M

pH goes from 7 to 3!

Buffer systems and pH Change

Huge change!What about buffers?


Buffer systems and pH Change

  • Consider a buffer solution with 0.1 M each of sodium acetate (NaAc) & acetic acid (HAc):

    • What is the pH when 10–3 M HCl is added?

HAc(aq) H+(aq) + Ac–(aq)Ka = 10-4.7

0

0.1 +10-3

0.1 - 10-3

+x

-x

+x

x

0.1+10-3-x

0.1- 10-3+x

cf slide 6


x = 1.02  Ka = 0.000020 << 0.001

pH = – log x = 4.69

Buffer systems and pH Change

the pH hardly changes from 4.7!

Solution is buffered against pH change


Henderson - Hasselbalch Equation

  • For a buffer solution, which contains similar concentrations of a conjugate acid/base pair of a weak acid:

  • The dissociation of HA or protonation of A- does not lead to a significant change in the concentrations of these species.

  • Taking logs and rearranging gives:


Buffer Capacity

Buffer capacity is related to the amount of strong acid or base that can be added without causing significant pH change.

Depends on amount of acid & conjugate base in solution:

highest when [HA] and [A–] are large.

highest when [HA]  [A–]

Buffer Preparation and Capacity

  • Buffer Preparation

    • If the pH of a required buffer is pKa of available acid then use equimolar amounts of acid and conjugate base

    • If the required pH differs from the pKa then use the Henderson-Hasselbalch equation.


Buffer Preparation and Capacity

Most effective buffers have acid/base ratio less than 10 and more than 0.1  pH range is ±1


Biological systems, e.g. blood, contain buffers: pH control essential because biochemical reactions are very sensitive to pH.

Human blood is slightly basic, pH  7.39 – 7.45.

In a healthy person, blood pH is never more than 0.2 pH units from its average value.

pH < 7.2, “acidosis”; pH > 7.6, “alkalosis”.

Death occurs if pH < 6.8 or > 7.8.

Buffers in Natural Systems


“Extracellular” buffer (outside cell)

Buffer System in Blood

H+(aq) + HCO3–(aq)H2CO3(aq)

H2CO3(aq) H2O(l) + CO2 (g)

  • Removal of CO2 shifts equilibria to right, reducing [H+], i.e., raising the pH

  • The pH can be reduced by:

H2CO3(aq) + OH–(aq) HCO3–(aq) + H2O(l)


Example

In the H3PO4 / NaH2PO4 / Na2HPO4 / Na3PO4 system, how could you make up a buffer with a pH of 7.40?

DATA: Ka1 = 7.2 x 10-3, Ka2 = 6.3 x 10-8, Ka3 = 4.2 x 10-13

  • To make up a buffer, we need pH near pKa

    pKa1 = 2.14, pKa2 = 7.20, pKa3 = 12.38

     must use mixture of H2PO4- and HPO4-

  • Could go through whole procedure...or simply use Henderson-Hasselbalch equation ...


original

amounts

Example (Continued)

  • Require buffer with a pH of 7.40

  • the required ratio of Na2HPO4 to NaH2PO4 = 1.58:1


Practice Examples

  • 1What is the pH of a 0.045 M solution of KOBr?

  • The pKa of HOBr is 8.63.

  • (a) 4.74b) 4.99c) 8.25d) 9.01e) 10.64

  • A buffered solution is 0.0500 M CH3COOH and 0.0400 M NaCH3CO2. If 0.0100 mol of gaseous HCl is added to 1.00 L of the buffered solution, wahat is the final pH of the solution?For acetic acid, pKa = 4.76

  • (a) 4.76(b) 4.46(c) 4.66(d) 4.86(e) 4.54


  • Summary: Acids & Bases 3

    • Learning Outcomes - you should now be able to:

    • Complete the worksheet

    • Understand acid and base equilibria

    • Identify conjugate acid/base pairs

    • Perform calculations with strong acids/bases

    • Use the buffer concept and construct buffers

    • Apply the Henderson-Hasselbalch equation

    • Answer Review Problems 11.36-11.37 and 11.96-106 in Blackman

    • Next lecture:

    • Titrations


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