Today’s class

1. Today’s class • Any questions about the lab last class? • Look at unique properties of water • pH and pOH • Do some practice problems • Next day we’ll look at strong and weak acids and bases

3. 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: H2O(l) ↔H+(aq) + OH-(aq) • Above reaction called the autoionization of water • Can also be expressed as: H2O + H2O↔H3O++ OH- acid1 base2 acid2 base1 • The acid-base conjugate pairs are 1) H2O(acid) and OH-(base) and 2) H3O+(acid) and H2O(base)

4. 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: Kc= [H3O+][OH-] • We replace Kc with Kw to denote that the equilibrium constant refers to the autoionization of water so the expression is now: Kw= [H3O+][OH-] • Kw is called the ion-product constant

5. 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 1.0X10-7M • As you can see when you put these values into the equation Kw=1.0X10-14 • So when [OH-]=[H+] the soultion is said to be neutral. • Acidic solutions the [H+] > [OH-], basic solutions [H+] < [OH-]

6. 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.

7. 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 Kw=1.0X10-14 and [OH-] we can easily calculate [H+] by rearranging the equation [H+]= Kw = 1.0X10-14 =4.0X10-12M [OH-] 0.0025M

8. 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+]

9. 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+]>1.0x10-7M, then pH<7 Basic: [H+]<1.0x10-7M, then pH<7 Neutral: [H+]=1.0x10-7M, then pH=7

10. 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-]

11. pH and pOH • Recall that the ion-product constant for water at 25°C is: [H+][OH-]=Kw=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

12. 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

13. 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

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

15. 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

16. Let’s give it a try!