Introduction to pH

1. Introduction to pH BY G. Ram Kumar Department of Chemistry Pydah College P.G. Courses Gambheeram, Visakhapatnam

2. S.P.L Sorenson introduced the concept of pH in the year 1909 • Definition • Negative Logarithm of the hydrogen ion concentration pH = - log aH3O+ or pH = - log [H+] Where [H+] is in molarity

3. Battery Acid Normal Stomach Acid (1.0 - 3.0) Normal Rainwater (5.6) Milk Pure Water, Blood Seawater, Shampoo Household Ammonia (10.5 - 11.9) Oven Cleaner • pH < 7 then acidic • pH = 7 then neutral • pH > 7 then basic

4. Acid-base equlibria in water • H20 H+ + OH- • Ionic product of water (W) at 25 0C is 10-14 mol/l • [H+] = [OH-] = 10-7 mol/l • pH + pOH = 14

5. pH Indicators – using equilibria to detect changes in pH • Litmus - is a weak acid. It has a seriously complicated molecule which we will simplify to HLit. The "H" is the proton which can be given away to something else. The "Lit" is the rest of the weak acid molecule. The un-ionised litmus is red, whereas the ion is blue.

6. Adding hydroxide ions

7. Adding hydrogen ions

8. Methyl orange

9. Phenolphthalein The half-way stage happens at pH 8.3-10 A mixture of pink and colourless is simply a paler pink, The pH range of indicators The importance of pKind Think about a general indicator, HInd - where "Ind" is all the rest of the indicator apart from the hydrogen ion which is given away:

10. Think of what happens half-way through the colour change.

11. indicator pKind litmus 6.5 methyl orange 3.7 phenolphthalein 9.3 That means that the end point for the indicator depends entirely on what its pKind value is. For the indicators we've looked at above, these are:

12. indicator pKind pH range litmus 6.5 5 - 8 methyl orange 3.7 3.1 - 4.4 phenolphthalein 9.3 8.3 - 10.0 There is a gradual smooth change from one colour to the other, taking place over a range of pH. As a rough "rule of thumb", the visible change takes place about 1 pH unit either side of the pKind value.

14. Choosing indicators for titrations Strong acid vs strong base HCl Vs NaOH

15. Strong acid vs weak base HCl Vs NH4OH

16. Weak acid vs strong base CH3COOH Vs NaOH

17. Weak acid vs weak base CH3COOH Vs NH4OH

18. Na2CO3 Vs HCl

19. Buffer • Using equilibria to stabilize pH • Definition - a solution which resists the change in its pH values on its dilution or addition of small amounts of acid or base to it

20. Acidic Buffer • CH3COOH + CH3COONa Basic Buffer NH4Cl + NH4OH

21. Hendreson equations • Acidic buffers • pH = pKa + log (salt/acid) • Basic buffers • pOH = pKb + log (salt/base) • pH = 14 – pKb + log(base/salt)

22. pH electrodes GLASS ELECTRODE Glass electrode is a potentiometric sensor made from glass of a specific composition. silicate matrix based on molecular network of silicon dioxide (SiO2) with additions of other metal oxides, such as Na, K, Li, Al, B, Ca, etc.

24. 1 - a sensing part of electrode, a bulb made from a specific glass2 - sometimes electrode contain small amount of AgCl precipitate inside the glass electrode3 - internal solution, usually 0.1M HCl for pH electrodes 4 - internal electrode, usually silver chloride electrode or calomel electrode5 - body of electrode, made from non-conductive glass6 - reference electrode, usually the same type as 47 - junction with studied solution, usually made from ceramics or capillary with asbestos or quartz fiber.

25. Importance of pH • Biological process & Industrial process generally occur at specified pH values only • Crops require soils of specific pH values for optimum growth and better yields • Many chemical and analytical procedures require maintenance of specific pH

26. Blood has specific pH ( 7.35 ) and this must always be maintained • Many physiological processes influence the pH, but one of the largest contributors is the CO2 content of the blood. • CO2 + H2O = HCO3-1 + H+1

28. Thanks for your attention