1. Acids and Bases pH, Titration, and Indicators

2. pH VIII. pH (power of hydrogen or hydronium) - measurement of hydronium concentration A. pH = -log [H 3 O + ]; if [H 3 O + ] = 10 -7, then pH = 7 B. High [H 3 O + ] gives low pH (more acidic with low pH C. pOH = -log [OH - ] (power of hydroxide)

3. pH D. pH Tools to solve problems (MEMORIZE THESE!!) [H 3 O + ][OH - ] = 1 x 10 -14 pH + pOH = 14 pH = -log [H 3 O + ] pOH = -log [OH - ] [H 3 O + ] = 10 -pH (antilog: put -pH in then use INV & log buttons on calculator) [OH - ] = 10 -pOH

4. pH Wheel pH pOH [OH - ] [H 3 O + ] (or [H + ]) 14-pH 14-pOH 10 -pH -log [H 3 O + ] 10 -pOH -log [OH - ] 1 x 10 -14 [ ]

5. E. pH Examples: 1. If the pH is 2.3, what is the pOH? pOH = 14 – 2.3 = 11.7 2. If the hydronium ion concentration is 2 x 10 -4 M, what is the pH? [H 3 O + ] = - log 2 x 10 -4 M = 3.7

6. pH Examples (cont) 3. If the hydroxide ion concentration is 3.5 x 10 -6 M, what is the pH? [H 3 O + ] = 1 x 10 -14 = 2.9 x 10 -9 M 3.5 x 10 -6 pH = - log 2.9 x 10 -9 = 8.54 OR: pOH = - log [OH - ] = - log 3.5 x 10 -6 = 5.46 pH = 14 – 5.46 = 8.54

7. E. pH Examples: (cont) 4. If the pH is 7.4, what are the hydronium and hydroxide ion concentrations? [H 3 O + ] = 10 -pH = 10 -7.4 = 4 x 10 -8 M [OH - ] = 1 x 10 -14 = 2.5 x 10 -7 M 4 x 10 -8 OR: pOH = 14-7.4 = 6.6 [OH - ] = 10 -pOH = 10 -6.6 = 2.5 x 10 -7 M

8. pH Scale F. pH scale goes from 0-14 0-2 = strong acid 2-7 = weak acid 7 = neutral 7-12 = weak base 12-14 = strong base

9. Indicators Indicators: compounds whose colors are sensitive to pH. A. Color changes as pH changes. B. Weak organic acids whose colors differ from their conjugate base C. HIn + H 2 O → H 3 O + + In - Yellow Red D. Look at pH of color change: called the transition interval See Figure 24 p. 662 (orange for indicator listed above)

10. Indicators E. Limitations: 1. Solutions must be colorless (or close) 2. Not very precise – relies on eyesight 3. Only good for very narrow pH range

11. Titration

12. X. Titration A. Measuring the amount of standard solution (known concentration) that reacts completely with a measured amount of solution of unknown concentration. B. Equivalence point - the point where the two solutions are present in chemically equivalent amounts (H + = OH - ) C. End point - the point where the indicator used changes color D. Indicators: strong acid/strong base: pH 7 bromothymol blue strong acid/weak base: pH 4 methyl red strong base/weak acid: pH 9 phenolphthalein

13. Titration Titration curves: p. 499 (Draw in notes ) Strong acid/strong base Weak Base/strong acid Weak Acid/strong base

14. Titration F. Steps: p. 500-501 G. Calculations: Remember that at the equivalence point the moles of H + = moles of OH - (times by #H or OH - ) millimoles of H + = millimoles of OH - 1. Find mmoles of H + by multiplying the following for the acid: vol (ml) x concentration (M) x # of H + in the acid’s formula 2. Find mmoles of OH - by multiplying the following for the base: vol (ml) x concentration (M) x # of OH - in the base’s formula 3. Set them equal to each other and solve for the unknown. (V a )(M a )(#H + ) = (V b )(M b )(#OH - )

15. Titration Examples H. Examples: 1. If 22.6 ml of Mg (OH) 2 are used to neutralize 30.4 ml of.100 M HCl, what is the concentration of the base? (V a )(M a )(#H + ) = (V b )(M b )(#OH - ) (30.4 ml)(.100 M)(1) = (22.6ml)(X)(2) X = (30.4 ml)(.100 M)(1) (22.6 ml)(2) =.0673 M Mg(OH) 2

16. Titration Examples 2. How many ml of.400 M NaOH are needed to neutralize 50.0 ml of.200 M HBr? (V a )(M a )(#H + ) = (V b )(M b )(#OH - ) (50.0 ml)(.200 M)(1) = (X)(.400 M)(1) X = (50.0 ml)(.200 M)(1) (.400 M)(1) = 25.0 ml NaOH