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

2. Where [H+] is in molarity
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. pH < 7 then acidic pH = 7 then neutral pH > 7 then basic
Battery Acid
Normal Stomach Acid
( )
Normal Rainwater (5.6)
Milk
Pure Water, Blood
Seawater, Shampoo
Household Ammonia
( )
Oven Cleaner
pH < 7 then acidic
pH = 7 then neutral
pH > 7 then basic

4. Acid-base equlibria in water
H 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 pH range of indicators The importance of pKind
The half-way stage happens at pH 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. litmus methyl orange phenolphthalein
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:
indicator
pKind
litmus
6.5
methyl orange
3.7
phenolphthalein
9.3

12. 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.
indicator
pKind
pH range
litmus
6.5
5 - 8
methyl orange
3.7
phenolphthalein
9.3

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 glass 2 - sometimes electrode contain small amount of AgCl precipitate inside the glass electrode 3 - internal solution, usually 0.1M HCl for pH electrodes 4 - internal electrode, usually silver chloride electrode or calomel electrode 5 - body of electrode, made from non-conductive glass 6 - reference electrode, usually the same type as 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