2. Unique properties of Water Can act as either an acid or a base Weak electrolyte so it is a poor conductor of electricity but does undergo ionization to a small extent: H 2 O (l) ↔ H + (aq) + OH - (aq) Above reaction called the autoionization of water Can also be expressed as: H 2 O + H 2 O ↔ H 3 O + + OH - acid 1 acid 2 base 1 base 2 The acid-base conjugate pairs are 1) H 2 O(acid) and OH - (base) and 2) H 3 O + (acid) and H 2 O(base)The acid-base conjugate pairs are 1) H 2 O(acid) and OH - (base) and 2) H 3 O + (acid) and H 2 O(base)

3. Ion product of Water In the autoionization equation the concentration of water remains unchanged because only a small fraction of molecules are ionized. This means that the equilibrium constant for the reaction is: K c = [H 3 O + ][OH - ] We replace K c with K w to denote that the equilibrium constant refers to the autoionization of water so the expression is now: K w = [H 3 O + ][OH - ] K w is called the ion-product constant

4. Ion product of Water Continued In pure water at 25°C the concentrations of H + and OH - ions are equal and are found to be 1.0X10 -7 M As you can see when you put these values into the equation K w =1.0X10 -14 So when [OH - ]=[H + ] the soultion is said to be neutral. Acidic solutions the [H+] > [OH - ], basic solutions [H+] < [OH - ]

5. Example The concentration of OH - ions in a household ammonia clean solution is 0.0025M. Calculate the concentration of H +. State whether the solution is acidic or basic.

6. Soultion Since we know the concentration of OH -, we can use the relationship between [H + ] and [OH - ] in water or aqueous solutions given by the ion-product of water equation Knowing the K w =1.0X10 -14 and [OH - ] we can easily calculate [H + ] by rearranging the equation [H + ]= K w = 1.0X10 -14 =4.0X10 -12 M [OH - ] 0.0025M

7. pH In 1909 Soren Sorensen, a Danish chemist, introduced the concept of pH, a scale for measuring acidity. Since concentrations of [H + ] and [OH - ] ions in aqueous solutions were so small they were hard to work with. Sorensen decided to use a more practical measure called pH. It is defined as: pH=-log[H+] or pH=-log[H3O+]

8. pH Since pH is a way to express the hydrogen ion concentration we can distinguish between acidic and basic conditions using pH in the following way: Acidic: [H + ]>1.0x10 -7 M, then pH<7 Basic: [H + ]<1.0x10 -7 M, then pH<7 Neutral: [H + ]=1.0x10 -7 M, then pH=7

9. pOH Much the same as pH scale, pOH can be used as a more practical measure to work with when stating the concentration of hydroxide ion We define pOH as: pOH=-log[OH - ]

10. pH and pOH Recall that the ion-product constant for water at 25°C is: [H + ][OH - ]=K w =1.0x10 -14 If we take the negative logarithm of both sides we will see: -log[H + ]-log[OH - ]=-log(1.0x10 -14) Which is the same as: pH+pOH=14 Which gives us another way to show the relationship between [H + ] and [OH - ] concentrations

11. Example The concentration of H + ions in a solution is 0.004M, find the pH of the solution. State if the solution is acidic or basic

12. Solution So we have the hydrogen ion concentration, we can easily find the pH by using the folllowing equation: pH=-log[H+] pH=-log(0.004) pH=2.40 Since the pH is less than 7 we know the solution is acidic

13. Example Calculate the [H + ] in a soultion with a pH of 8.5.

14. Solution When given the pH of a soultion we must rearrange the equation to calculate [H + ] so [H + ]=10 (-pH) Using the equation we get [H + ]=10 (-8.5) [H + ]=3.16x10 -9