1. Buffer

2. 11C Alyssa Ertheo Gana Kenneth Tamara

3. Buffer  An aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa.  Its pH changes very little when a small amount of strong acid or base is added to it and thus it is used to prevent changes in the pH of a solution.

4.  Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.  Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A -  HA ↔ H + + A -

5.  When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Chatelier's principle  If strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added

7. Making an acidic buffer To create an acidic buffer, we have to have a mixture of a weak acid and its salt. The sodium or potassium salt of the acid is used. 2CH 3 COOH + NaOH -> CH 3 COOH + CH 3 COONa + H 2 O As you can see, 1 mol of ethanoic acid is neutralized by the 1 mol of sodium hydroxide to form 1 mol of salt. There is 1 mol of ethanoic acid left over and that is also used in the buffer solution, as we will see later.

8. The salt is used to increase the concentration of the COO - ions in solutions, as naturally, as CH 3 COOH is a weak acid, it is only very slightly dissociated in solution, making it ineffective at absorbing H + ions.

9. Making an alkaline buffer To create an alkaline buffer, we have to have a mixture of a weak base and its chloride salt. 2NH 3 + HCl -> NH 3 + NH 4 Cl As you can see, 1 mol of ammonia is neutralized by the 1 mol of hydrochloric acid to form 1 mol of the ammonium chloride salt. There is 1 mol of ethanoic acid left over and that is also used in the buffer solution, as we will see later.

10. How buffer solutions work Acidic buffers CH 3 COOH CH 3 COO - + H + Adding an acid: An acid will always have H + ions. Any H+ added reacts with the CH 3 COO - ions to form CH 3 COOH, as seen in the equilibrium’s reverse reaction. It is thus removed from the solution.

11. How buffer solutions work CH 3 COOH CH 3 COO - + H + H + + OH - H 2 O Adding an alkali: An acid will always have OH - ions. Although OH - ions are notably absent from the equilibrium equation, they react with H + ions in the solution to form water via the 2 nd equilibrium.

12. How buffer solutions work Basic buffers NH 3 + H + NH 4 + Addition of an acid (H + ): the hydrogen ions added will react with NH 3 to form ammonium ions, thus the buffer solution “mops up” the added hydrogen ions.

13. How buffer solutions work NH 3 + H 2 O NH 4 + + OH - Addition of a base (OH - ): the ammonium ions react with the hydroxide ions to form ammonia and water (the reverse reaction); the extra ammonium ions from ammonium chloride help in increasing the concentration of ammonium ions in solution.

14. Exercise

17. Solubility Products, Ksp

20. Exercise

21. a. NiS(s) ↔ Ni2+(aq) + S2-(aq) Ksp = [Ni2+] [S2-], unit:mol2dm-6 b. Pb(OH)2 (s) ↔ Pb2+(aq) + 2OH-(aq) Ksp = [Pb2+] [OH-]2, unit:mol3dm-9 c. Ca3(PO4)2(s) ↔ 3Ca2+(aq) + 2 PO43-(aq) Ksp = [Ca2+]3 [PO43-]2, unit:mol5dm-15 d. Fe(OH)3(s) ↔ Fe3+(aq) + 3OH-(aq) Ksp = [Fe3+] [OH-]3, unit:mol4dm-12

22. Exercise

23. Solubility Product and Precipitation A mixture of two solutions will produce a precipitate depending on the ionic products of the ions present If the ionic product > Ksp, precipitation will occur If the ionic product < Ksp, no precipitation will occur

24. Exercise

26. BUFFER IN DAILY LIFE

27. Buffer in Daily Life In the Body Buffer solutions are everywhere – even in our own blood! Buffer solutions help keep the pH of our blood at a constant pH of around 7.4. If the pH is 0.5 above or below that value, unconsciousness and coma may result.

28. The buffer solution is produced by the formation of carbonic acid from carbon dioxide in respiration. CO 2 + H 2 O H + + HCO 3 - H + + HCO 3 - H 2 CO 3 The 2 nd equation shows how H+ ions from the addition of an acid are removed by the buffer solution, and thus pH remains the same. H 2 CO 3 acts the same as CO 2 in water, so the water will not become more acidic unless the amount of H + ions added exceed the “mopping” action of the buffer solution.

29. Buffer in Daily Life In the Body Another example of buffers within the human body is the "haemoglobin" complex, which binds to excess protons that the muscles release during exercise so that the body can use the oxygen they produce alongside the hydrogen.

30. In laboratory In the lab, buffer solutions are used to study the speed of reactions and in the preparation of specimens in biology. In the industry, they are used in paper manufacture, sugar manufacture, electroplating, fermentation (preventing solutions becoming too acidic and spoiling the batch) and maintaining the optimum pH for dying fabrics in the textile industry.

31. Buffer in Daily Life Commercial Products Buffers are also important in certain commercial household products. Shampoo, for example, contains buffers of citric acid and sodium hydroxide, which work to balance out the natural alkalinity of soap, which would otherwise burn the scalp.

32. Buffer in Daily Life Shampoos use buffer solutions to counter the basicity of soap, preventing irritation of the skin. Baby lotion uses buffer solutions to keep the pH constant at a value of 6, preventing the division and multiplication of bacteria. A baby, when applied baby lotion to his rash relieves it rather quickly. This is because baby lotion is buffered to a slightly acidic pH of six, which inhibits the growth of bacteria and other pathogens.

33. THE END