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Buffers

Buffers. 4/13/2011. Definitions. Weak Acids : all proton donors that are in equilibrium Conjugate Bases : the ions that are left over after a weak acid loses it H + ions (protons) Buffer : a solution containing a mix of weak acid and the salt of its conjugate base

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Buffers

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  1. Buffers 4/13/2011

  2. Definitions • Weak Acids: all proton donors that are in equilibrium • Conjugate Bases: the ions that are left over after a weak acid loses it H+ ions (protons) • Buffer: • a solution containing a mix of weak acid and the salt of its conjugate base • A solution that can absorb added acids and bases w/o a large pH change

  3. The weak acid rxn • HA H+ + A- • In a pure acid, the [H+] and [A-] occur in equal amounts • In any solution, addition of H+ will decrease A- and vice versa • The conjugate base of HA is A- • Notice that if [A-] goes up, the H+ will go down. The solution will become more basic.

  4. Salt of a conjugate base To identify the salt: Take any weak acid Remove the H+ Replace with any other positive ion • A “salt” is defined as the product of an acid and a base • A salt of a conjugate base is the salt produced when a weak acid is neutralized - H A Na Salt of Conjugate Base Weak Acid Conjugate Base Examples: Weak acidSalt of Conjugate Base HFNaF HNO2 LiNO2 CH3COOH  KCH3CO3 H2SO3  NaHSO3 H2CO3 LiHCO3

  5. Non-Buffered Solutions 1 x 10-3HCl • If a strong acid is added to pure water, the pH changes radically • Original pH = 7 • Imagine 1 mL of 1 Molar HCl (1 x 10-3 moles) added to 1 Liter of water [H+]= 1x10-3 Cl- (spectator) pH 7 pH 3

  6. Buffered Solutions HA NaA • Add roughly equal amounts of weak acid and conjugate base salt • An equilibrium is set up  • And then altered HA H+ A- Na+ A- HA  H+ + A- H+ + A- (The addition of the A- lowers the H+)

  7. Buffers in Action The Buffer solution now has both an acid and a base in it. Addition of an acid will add more H+, which will be absorbed (neutralized) by the base. If a base is added, it will react with the H+ , shifting the equilibrium to the right, reducing HA and increasing A- Na+ A- HA H+ A- Acid Base OH- HA  H+ + A- H+ H2O

  8. Buffer Calculations HA  H+ + A- H+ + A- Imagine 1 liter of a buffer with 0.5 moles HA and 0.5 moles NaA Pretend (for now) that Ka = 1 x 10-7 Ka = [H+] [A-] = (X )[0.5] [HA] [0.5] So…with equal moles of weak acid and conjugate base, H+ = Ka = 1 x 10-7 And pH = 7 Like before, add 1 mL of 1 molar HCl. (you are adding .001 moles acid) The equilibrium will change… And so will the pH… Ka = [H+] [A-] = (X )[0.499] [HA] [0.501] The H+ has increased and = 1.004 x 10-7 And pH = 6.998 Original 0.5 mol 1 x 10-7 0.5 mol Add Acid + 0.001 mols Increase HA + 0.001 - 0.001Decrease A- New Values 0.501 mol X 0.499 mol pH 7 pH 6.998

  9. Practice calculations • What would be the pH of a solution of 0.2 M HF and 0.2 M NaF if the Ka = 1.6 x 10-5? • What would be the pH if the above solution was diluted by ½ (so each solute is 0.1 M)? • The Ka for Carbonic acid (H2CO3) is 4.5 x 10-7 (only the first H+ is ionzed). If the pH of a solution is 7.4, and the concentration of HCO31- is 0.22 M, what is the concentration of H2CO3?

  10. Buffers in the real world • Buffers are important in the biological world. • (imagine what could happen to a fish tank if the water was not buffered) • Your blood is strongly buffered to control the amount of CO2 in the blood. • CO2(aq) + H2O(l)H2CO3(aq)H+(aq) + HCO3-(aq) • Your blood measure pH instead of CO2(aq) • If your pH gets too low, your breathing rate goes up

  11. Summary of Buffers • They are made of weak acids and conjugate bases, and are in equilibrium with H+ • Because they contain both base and acid, they can absorb added bases and acids with only a very small change in pH • Calculations of acidity can be made if the Ka, acid and base concentrations are known.

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