Ionic products of water

1. Ionic products of water

2. Water molecules can function as both acids and bases. One water molecule (acting as a base) can accept a hydrogen ion from a second one (acting as an acid). This will be happening anywhere there is even a trace of water. • A hydroxonium ion and a hydroxide ion are formed. •  the hydroxonium ion is a very strong acid, and the hydroxide ion is a very strong base. As fast as they are formed, they react to produce water again. • 2H2O(l) H3O+(aq) + OH-(aq)

3. At 25℃ water molecules dissociate at a very small extent producing very little H+ cation and OH- anion • H2O (L) H+(aq) + OH-(aq) Kc = [H+] [OH-] [H2O] Kc [H2O] = [H+] [OH-] • Kc is a constant that represents the concentration of water • Kw is a constant that represents the ionic product of water this is formed when two constants are multiplied together. • The value of kw = [H+] [OH-] = (10-7 mol dm-3) x ( 10-7 mol dm-3) = 10-14 mol2 dm-6 Even though the constant varies with temperature

4. BH+ is the conjugate acid of a base and the equilibrium constant (Ka)for the conjugate acid is BH+ B + H Ka = [B] [H+] [BH+] = [B] [H+] [OH-] [BH+] [OH-] = [B] [H+] [OH-] [BH+] [OH-] = 1 Kw Kb therefore Ka Kb= Kw taking log10 pKa + pKb= pKw so if pKw = 14 then pKa + pKb = 14 this equation can be used to calculate the pH of weak base The relationship between Ka and Kb

5. Acid - Base indicators (also known as pH indicators) are substances which change color with pH . They are usually weak acids or bases, which when dissolved in water dissociate slightly and form ions • The acid and its conjugate base have different colors. when using litmus you can write the undissociated molecules as HL an the dissociated molecules as H+ and L- HL H+ + L- (red) (blue) At low pH values the concentration of H3O+ is high and so the equilibrium position lies to the left. The equilibrium solution has the color A. At high pH values, the concentration of H3O+ is low the equilibrium position thus lies to the right and the equilibrium solution has color B • HL H+ + L- this equation shows that when an acid is added (red) (blue) the concentration of H+ ion the equilibrium shift to the left and the color changes to red but when a base is added the equilibrium shifts to the right due to the H+ reacting with the OH- therefore the color changes to blue. Acid – Base Indicators

6. Kln is known as the indicator dissociation constant. The colour of the indicator turns from colour A to colour B or vice versa at its turning point. At this point. • The pH of the solution at its turning point is called the pKln and is the pH at which half of the indicator is in its acid form and the other half in the form of its conjugate base.

7. Indicators do not change colour at a specific concentration of hydrogen ion but rather at a narrow range of hydrogen ion. pH Range Of Indicators

8. A Universal Indicator is a mixture of indicators which give a gradual change in colour over a wide pH range - the pH of a solution can be approximately identified when a few drops of universal indicator are mixed with the solution. • Indicators are used to test the acidity and alkalinity of solutions and to determine the end point of a titration. • When an indicator is used in a titration the end point of the titration is when the colour of the indicator changes. The endpoint must coincide with the equivalence point in the titration . • The equivalence point is where the number of moles of acid is equal to the number of moles of the base .

9. pH Changes During Titration • The pH of a solution changes during titration as an acid is added to alkalis or vise versa. A plot of the changes in pH is known as a titration curve. • (A) (B) • Curve A shows that the pH changes when an alkali is added to an acid, also curve B shows that the pH changes when an acid is added to an alkali during titration. The pH changes sharply by several units at the equivalence point .

10. Any acid base indicator that changes colour at a pH in the vertical part of a grapy is suitable to detect a titration • Strong acid- Strong base titration: any indictor that changes colour between pH 3-11 is suited for this type of titration. Example methyl orange and phenolphthalein.

11. If you use phenolphthalein, you would titrate until it just becomes colourless (at pH 8.3) because that is as close as you can get to the equivalence point. • On the other hand, using methyl orange, you would titrate until there is the very first trace of orange in the solution. If the solution becomes red, you are getting further from the equivalence point.

12. An indicator which changes colour between the pH of 3-7 is suitable for this type of titration • Methyl orange is suitable for this type of titration because it changes colour at a pH of 3.7 where as phenolphthalein is not cause it changes at a pH of 9.1 Strong acid – Weak base titration

13. methyl orange starts to change from yellow towards orange very close to the equivalence point.

14. An indicator that changes colour between the pH of 7-11 would be suitable for this titration. • Example is phenolphthalein would work due to the pH being within the range in which it changes colour Weak Acid- Strong Base Titration

15. In this titration the change in pH is too gradual to show a distinct colour change (end point). • The curve is for a case where the acid and base are both equally weak. • Neither methyl orange nor phenolphthalein indicator can be used. Weak Acid- Weak Base Titration

16. Thank you.