1. At any one time, there are incredibly small numbers of hydronium ions and hydroxide ions present. H 2 O (l) H + (aq) OH - (aq) K w (ion-product constant for water) [H + ][OH - ] = 1.0 x 10 -14 1.0 x 10 -7 mol/L

2. Acids / bases dissolve in water - increase [H + ] / [OH - ] and cause an equilibrium shift. Acidic - [H + ] is greater than the [OH - ] Basic - [OH - ] is greater than [H + ] K W = 1.0 x 10 -14 K W = [H + ][OH¯] H 2 O (l) H + (aq) + OH¯ (aq)

4.  Definition pH and pOH.  Given pH, pOH, [H 3 O + ] or [OH¯], calculate the remaining values.  Calculate Ka/Kb, given the pH or pOH and the concentration of a weak acid solution.  Describe how an acid-base indicator works in terms of the colour shifts and Le Chatelier's Principle.

5. Acidic - [H 3 O + ] > [OH¯] Alkaline (basic) - [OH¯] > [H 3 O + ] Neutral - [OH¯] = [H 3 O + ] 1909 - Soren Sorensen developed a simplified system for the degree of acidity of a solution. pH - the potenz (power) of hydrogen - German potentia hydrogenii - Latin

6. Convenient way to express [H + ] pH = -log [H + ] Similarly, the concentration of hydroxide can be expressed as pOH : pOH = -log [OH - ] pH and pOH have no units.

7. [H + ][OH - ] = 1.0 x 10 -14 pH + pOH = 14 Special case - pH can be read straight from the value of the [H + ]. [H + ] = 1.0 x 10 -x then pH = x [OH - ] = 1.0 x 10 -x then pOH = x

8. Remember, in a neutral solution pH = pOH = 7 Acidic - pH < 7 Basic - pH > 7

10. Calculate the pH of an HCl solution whose concentration is 5.0 x 10 -6 mol/L. pH = -log[H + ] = -log(5.0 x 10 -6 M) = -(-5.30) = 5.30

11. The pH of a solution is 3.25. Calculate the hydrogen ion concentration in the solution. [H 3 O + ] = 10 -pH = 10 -3.25 = 5.6 x 10 -4 M ** also means: [OH - ] = 10 -pOH

12. The pH of a solution is 10.30, what is the hydroxide ion concentration? pOH = 14.00 - pH pOH = 14.00 - 10.30 = 3.70 [OH - ] = 10 -pOH = 10 -3.70 = 2.0 x 10 -4 mol/L [H + ] = 10 -pH K W = [H + ][OH¯] OR

13. What is the pH of 5.0 x 10 -5 M Mg(OH) 2 solution? Mg(OH) 2 (s) Mg 2+ (s) + 2 OH - (s) xx 2x [OH - ] = 2x = 2(5.0 x 10 -5 M) = 1.0 x 10 -4 M pOH = -log[OH - ] = 4.00 pH = 14.00 - 4.00 = 10.00

15. Measuring pH There are two ways to measure pH: pH Meters Indicators The [H + ] inside the probe (reference electrode) is compared to [H + ] outside the probe. Probe must be calibrated first. (inserted into known pH solution)

16. An indicator is a weak acid or base that undergoes a colour change when they gain or lose hydrogen ions. Natural pH indicators Beets, Blackcurrant juice, Blueberries, Carrots, Cherries, Curry Powder, Delphinium Petals, Geranium Petals, Grapes, Horse Chestnut Leaves, Hydrangea, Morning Glories, Onion, Pansy Petals, Petunia Petals, Poison Primrose, Poppy Petals, Purple Peonies, Rayhan Leaves, Red cabbage, Red Radish, Rhubarb, Rose Petals, Strawberries,Tea, Thyme, Turmeric, Tulip Petals, Violet Petals

17. Red Litmus stays red in the presense of an acid.

18. It turns blue in presense of a base.

19. Weak acid is colourless and its conjugate base ion is bright pink. Phenolphthalein is a commonly used indicator.

22.  The pH scale goes from 0 - 14.  pH = -log[H 3 O + ] and [H 3 O + ] = 10 -pH  pOH = -log[OH¯] and [OH¯] = 10 -pOH  pOH + pH =14  Indicators are weak acids or bases that change colour in response to changing hydronium ion concentrations.