1. Chapter 16 Acid-Base Titration and pH

2. Aqueous Solutions and the Concept of pH Self-ionization of water – 2 water molecules produce a hydronium ion and a hydroxide ion by transferring a proton  H 2 0 + H 2 0  H 3 0 + + OH - Concentration of hydronium ion and the hydroxide ions are represented as [H 3 0 + ] and [OH - ]

3. For pure water the [H 3 0 + ] and [OH - ] are both 1.0 x 10 -7 M Ion product constant for water (K w ) is obtained by multiplying the [H 3 0 + ] and [OH - ]  K w for room temperature is 1.0 x 10 -14 M 2, but varies with temperature

4. Neutral, acidic, and basic solutions Neutral solutions have equal [H 3 0 + ] and [OH - ] Acidic solutions have a greater [H 3 0 + ] than [OH - ] Basic solutions have a lower [H 3 0 + ] than [OH - ]

5. Calculating [H 3 0 + ] and [OH - ] K w = [H 3 0 + ] x [OH - ] Remember use K w = 1.0 x 10 -14 M 2 Ex: A 1.0 x 10 -4 M solution of HNO 3 has been prepared. What is its [H 3 0 + ]? What is its [OH - ]?

6. The pH scale pH – the negative logarithm of the hydronium ion concentration  pH = - log [H 3 0 + ] pOH – the negative logarithm of the hydroxide ion concentration  pOH = -log [OH - ] pH + pOH = 14.0

7. pH less than 7 is acidic pH greater than 7 is basic pH equal to 7 is neutral

8. Calculations involving pH Sig figs for pH are different because of the logarithm There must be as many sig figs to the right of the decimal for the pH as what there were in the [H 3 0 + ]  Ex. [H 3 0 + ] is 1 x 10 -7 has one sig fig and his pH is 7.0

9. What is the pH of a solution if the [H 3 0 + ] is 3.4 x 10 -5 M? pH = - log [3.4 x 10 -5 M] pH = 4.47

10. Finding [H 3 0 + ] from pH pH = - log [H 3 0 + ] can be rearranged to solve for [H 3 0 + ] by using antilog [H 3 0 + ] = 10 -pH

11. Determine the hydronium ion concentration of an aqueous solution that has a pH of 4.0. pH = -log [H 3 0 + ] 4.0 = -log [H 3 0 + ] 10 -4.0 = [H 3 0 + ]

12. The pH of a solution is determined to be 7.52. What is the pOH, [H 3 0 + ], and [OH - ]?

13. The molarity of strong acids and bases can be used directly to calculate pH, but not weak acids and weak bases because they don’t ionize/dissociate completely; instead pH must be measured and then [H 3 0 + ] and [OH - ] calculated

14. Sect. 16-2: Determining pH and Titrations Acid-base indicators – compounds whose colors are sensitive to pH  Weak acids or weak bases  Different color in the ionized (In - ) vs. non- ionized form (HIn)  HIn  H + + In -

15. In an acid, accepts H + to form HIn In a base, OH - combines with H +, so more HIn ionizes to offset loss of H +, thus more In - is present

16. Transition interval – the pH range over which an indicator changes color  If a low pH, then the indicator is a stronger acid  If a higher pH, then the indicator is a weak acid pH meter – determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution

17. Titration – the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration  Used to determine equivalent volumes of acid and base  An example of a chemically equivalent amount would be 1liter of 0.1M HCl reacting with 0.1 mol solid NaOH

18. Equivalence point – the point at which the two solutions used in a titration are present in chemically equivalent amounts End point – the point in a titration at which an indicator changes color

19. When choosing indicators to be used in a titration remember:  Strong acid/strong base will have an equivalence point at 7  Strong acid/weak base will have an equivalence point below 7  Weak acid/strong base will have an equivalence point above 7

20. Molarity and titration Standard solution – the solution that contains the precisely known concentrations of a solute (“known” solution) Primary standard – a highly purified solid compound used to check the concentration of the known solution in a titration

21. Steps for determining molarity of unknown solution in titration 1. Use balanced neutralization reaction to determine chemically equivalent amounts of acid and base 2. Determine moles of known solution 3. Determine moles of solute of unknown solution (use stoichiometry) 4. Determine molarity of unknown

22. Example: In a titration, 27.4 mL of 0.0154M Ba(OH) 2 is added to a 20.0 mL sample of HCl solution of unknown concentration. What is the molarity of the acid solution? Ba(OH) 2 + 2HCl  BaCl 2 + 2H 2 O 1 mol 2 mol 1 mol 2 mol