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Buffers

Buffers. Chem 12A Mrs. Kay. Buffers help maintain a constant pH. They are able to accept small quantities of acids and bases without drastically changing their pH A buffer is composed of a weak acid molecule in equilibrium with its conjugate base and hydrogen ion (H+). 

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Buffers

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  1. Buffers Chem 12A Mrs. Kay

  2. Buffers help maintain a constant pH. • They are able to accept small quantities of acids and bases without drastically changing their pH • A buffer is composed of a weak acid molecule in equilibrium with its conjugate base and hydrogen ion (H+).  • CH3COOH    < --- >    CH3COO-    +   H+

  3. Acidic buffer solution: • An acidic buffer solution is simply one which has a pH less than 7. • Acidic buffer solutions are commonly made from a weak acid and one of its salts - often a sodium salt. • Example: acetic acid and sodium acetate CH3COOH and CH3COONa

  4. Alkaline buffer solutions • An alkaline buffer solution has a pH greater than 7. • Alkaline buffer solutions are commonly made from a weak base and one of its salts. • A frequently used example is a mixture of ammonia solution and ammonium chloride solution. NH3 and NH4Cl

  5. How do buffers work? • It has to contain things which will remove any hydrogen ions or hydroxide ions that you might add to it - otherwise the pH will change.

  6. Below is a weak acid, if we added CH3COONa, its like adding CH3COO- CH3COOH < -- > CH3COO- + H+ There will be a shift to the left • The solution will therefore contain these important things: • lots of un-ionised acetic acid; • lots of acetate ions from the sodium acetate; • enough hydrogen ions to make the solution acidic. Other things (like water and sodium ions) which are present aren't important to the argument.

  7. Look to below website for a flash on how buffers work: • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf

  8. Buffer capacity and pH • Buffer capacity is the amount of acid/base the buffer can neutralize before the pH begins to drastically change. • The pH of the buffer depends on the Ka for the acid and on the relative concentrations of the acid/base of the buffer

  9. Calculations What is the pH of a buffer that is 0.12 M in lactic acid (HC3H5O3) and 0.10 M of sodium lactate? For lactic acid, Ka= 1.4 x 10-4 • Determine the important species involved • Set up ice table • Set up equilibrium expression • Plug in concentrations at equilibrium

  10. 1. HC3H5O3 <-- > C3H5O3 -+ H+

  11. Ka = 1.4 x 10-4 =[C3H5O3 –][H+] [HC3H5O3] • Because of the small Ka, we expect x to be small relative to 0.12 and 0.10 M, so we can simplify our equation to… • Ka = 1.4 x 10-4 =[0.10][x] [0.12] Solve for x, x =1.7 x 10-4

  12. Finally, solve for the pH of the buffer pH = - log [1.7 x 10-4] = 3.77 Notice, when you put the value of x into the original equation it makes no difference for your final answer.

  13. Addition of strong acids/bases to buffers: • We assume that the strong acid or base added is consumed completely by the reaction with the buffer • HX < -- > H+ + X- • By adding a strong acid, the shift is to the left, therefore increasing [HX] and decreasing [X-] • By adding a strong base, the OH- is used up by HX to produce X-, so [X-] increases and [HX] decreases.

  14. Calculation • A buffer is made by adding 0.300 mol of HC2H3O2 and 0.300 mol of NaC2H3O2 to 1.0 L of water. The pH of the buffer is 4.74. Calculate the pH of this solution if 0.020 mol of NaOH is added (ignore volume changes)

  15. HC2H3O2 +H2O  H3O+ +C2H3O2-

  16. Ka = (x)(0.320) / (0.280) = 1.8 x 10-5 so x = [H3O+]= (0.280)(1.8 x 10-5 )/0.320 = 1.6 x 10-5 • pH = - log (1.6 x 10-5)= 4.80

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