1. Water & pH Lecture 2 Ahmad Razali Ishak
Department of Environmental Health
Faculty of Health Sciences
UiTM Puncak Alam

2. Role of water in the life of organisms
Mammalian cells 70% water
Solvent for biological systems & for most chemical reactions that support life.
75% of the earth is covered with water
Has a very high specific heat-retains heat better than other materials

3. Properties of water Polarity Hydrogen bond Universal solvent
Hydrophobic interactions
Other non covalent interactions in biomolecules
Nucleophilic nature of water
Ionization of water

4. 1) Polarity
Water = H2O
2 H atoms are linked covalently to oxygen, each sharing an electron pair.
Non linear arrangement bond angle-104.5
Oxygen atom more electronegative than H atom-polar covalent bond
Creates a permanent dipole in H2O molecule

5. 2) Hydrogen bonds
Water molecules attract to each other due to polarity
H bonds: attraction of one slightly +ve H atom of one water molecule and one slightly –ve oxygen atom of another water molecule
Water molecule can form H- bond with 4 other water molecules
H bonds weaker than covalent bonds.
H bonds gives water a HIGH melting point, specific heat and heat of vaporization
The ability to form strong H bond is responsible for the many unique characteristics of water such as its high melting point and boilng point
3D structures of many important biomolecules including proteins (Hb) and nucleic acids (DNA) are stabilized by H bonds

6. 3) Universal Solvent
Water interact with and dissolve other polar and ionize (electrolytes) compounds
Water aligning around electrolytes to form solvation spheres
Solubility depend on polarity and ability to form H-bonds
Functional groups on molecules that confer solubility: carboxylates, protonated amines, amino, hydroxyl and carbonyl
As the number of polar groups increase in a molecule, so does its solubility in water.
Flavoring and CO2 gas dissolved in water to make soft drinks
Farmers use water to dissolve fertilizers
Medicines in water
Chlorines or fluorides added to water

7. 4)Hydrophobic interaction
Non polar molecules NOT soluble in water, hydrophobic
Amphiphatic molecule: Have both hydrophobic and hydrophilic portions
E.g. Detergents, and surfactants
Form micelles in aqueous solution
Used to trap grease and oils inside to remove them
Hydrophilic compounds interact (disslove) with water eg . Polar cpds (alcohols and ketones)& ionic cpds (KCl), amino acids
Hydrophobic compounds do not interact with water eg. Non polar cpds (hexane, fatty acids, cholesterol)

8. 5) Other non covalent interactions in biomolecules
Four major non covalent forces involved in structure and function of biomolecules:
H-bond
Hydrophobic interactions
Ionic bonds
Van der Waals forces

9. 6) Nucleophilic nature of water
Nucleophilic: electron rich
Electrophiles: electron deficient
Nucleophiles are negatively charged and have unshared electrons pairs: attack electrophiles during substitution or addition reactions.

10. 7)Self-Ionization of Water
H2O is amphoteric. i.e. as an acid it gives up H+ to become OH¯, and as a base it accepts H+ and becomes H3O+.
When water reacts with itself
H2O + H2O H3O+(aq) + OH¯(aq)
This reaction of water is called self-ionization of water.

11. Kw Keq = [H3O+][OH¯] [H2O]2 Concentration of pure water is constant
Keq [H2O]2 = [H3O+][OH¯]
Kw = [H3O+][OH¯]
Kw = ion-product constant of water

12. At 25 ˚C, both H3O+ and OH¯ ions are found at concentration of 1
At 25 ˚C, both H3O+ and OH¯ ions are found at concentration of 1.0 x 10-7 M
Kw = [1.0 x 10-7][1.0 x 10-7]
Kw = 1.0 x 10-14
If [H3O+] > 1.0 x 10-7, solution is acidic
If [OH¯] > 1.0 x 10-7, solution is basic
If [H3O+] = [OH¯] = 1.0 x 10-7, solution is neutral.

13. p. 48

14. Cont.. At equilibrium, pH = -log [H+] - log [1.0 x 10-7 ]
= 7 (neutral)
Low pH represent highest [H+] and lowest [OH]

15. Acids, Bases, and Buffers
Acids are proton donors and base are proton acceptors (Bronsted-lowry)
The strength of an acid is measured by its acid dissociation constant, K
The larger the K value, the stronger the acid and more H+ dissociates
The conc. H+ is expressed as pH, -ve log of H ion conc.

16. This tendency to ionize can be put in terms of an equation for the equilibrium:

17. Where [ ] = Molar concentration;
K = Ionization constant (acid dissociation constant)
Simplest example is water (H2O):

18. p. 47

19. p. 47

20. Table 2-6, p. 50

21. Buffers 1. Principle of Ionization of Weak Acids:
The Fundamental Concept of Buffers is: A Buffer Resists Change
pH buffers resist change in pH when either acid (H+) or base (OH-) is added to it.
Chemicals which are pH buffers are weak acids or bases
Acids = Proton (H+) donors
Bases = Proton Acceptors

22. Titration of a Weak Acid illustrating its Ionization and Buffering Property

23. Fig. 2-13b, p. 51

24. All weak acids have titration curves like this one
All weak acids have titration curves like this one. Bases (like ammonium, NH4+) are also weak acids and have similar titration curves.
The position where the Buffering zone is on the pH scale is related to the chemical nature of the weak acid:
Acetic acid ionizes in the Acidic portion of the pH scale

25. Fig. 2-16a, p. 58

26. Fig. 2-13b, p. 51

27. This relationship is known as the Henderson-Hasselbalch equation.
Useful in predicting the properties of buffer solutions used to control the pH of reaction mixtures.
The pK of a weak acid is the pH where [A-] = [HA]
At pH below the pK, [HA] > [A-]
At pH above the pK, [HA] < [A-]
Therefore the pK determines the buffering zone for a weak acid.

28. A similar expression pK can be used, pK=-log K
The pH of a solution of a weak acid and its conjugate base is related to the concentration of the acid and base- Henderson- Hasselbach equation.

29. Henderson Hasselbach Equation
p. 49

30. Example 1
Calculate the pH of a buffer solution made from 0.20 M HC2H3O2 and 0.50 M C2H3O2- that has an acid dissociation constant for HC2H3O2 of 1.8 x 10-5.

31. Answer pH = pKa + log ([A-]/[HA])
pH = pKa + log ([C2H3O2-] / [HC2H3O2])
pH = -log (1.8 x 10-5) + log (0.50 M / 0.20 M)
pH = -log (1.8 x 10-5) + log (2.5)
pH =
pH = 5.1

32. Example 2
Calculate the relative amounts of acetic acid and acetate ion present at the following points when 1 mol of acetic acid is titrated with NaOH. Use HH eqn. to calculate pH
0.1 mol NaOH added
0.3 mol NaOH added
0.5 mol NaOH added

33. Answer
Ratio 1:1, when 0. 1 mol of NaOH added, 0.1 mol acetic acid reacts with it to form 0.1 mol acetate ion, leaving 0.9 mol acetic acid
pH = pK + log 0.1/0.9 =
= 3.81

34. Buffer applications
In humans-in the intracellular and extra cellular fluid, there is a conjugate acid-base pairs that act as buffer
Major intracellular buffer is conjugate base-acid pair H2P04- / HPO42- (pK = 6.8)
Extracellular fluid- main buffer in blood and interstitial fluid is the bicarbonate H2CO3/HCO3- (pK = 6.2). Normal pH for blood is 7.4

35. Summary 2-3, p. 51

36. Summary 2-4, p. 53

37. Summary 2-5, p. 61