pH, Titration, and Indicators Acids and Bases pH, Titration, and Indicators

pH VIII. pH (power of hydrogen or hydronium) - measurement of hydronium concentration A. pH = -log [H3O+]; if [H3O+] = 10-7, then pH = 7 B. High [H3O+] gives low pH (more acidic with low pH C. pOH = -log [OH-] (power of hydroxide)

pH D. pH Tools to solve problems (MEMORIZE THESE!!) [H3O+][OH-] = 1 x 10-14 pH + pOH = 14 pH = -log [H3O+] pOH = -log [OH-] [H3O+] = 10-pH (antilog: put -pH in then use INV & log buttons on calculator) [OH-] = 10-pOH

pH Wheel pH Wheel [H3O+] (or [H+]) pOH [OH-] pH -log [H3O+] 14-pH -log [OH-] 10-pOH 1 x 10-14 [ ] [OH-]

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? [H3O+] = - log 2 x 10-4 M = 3.7  

pH Examples (cont) 3. If the hydroxide ion concentration is 3.5 x 10-6 M, what is the pH? [H3O+] = 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

E. pH Examples: (cont)  4. If the pH is 7.4, what are the hydronium and hydroxide ion concentrations? [H3O+] = 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

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

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 + H2O → H3O+ + In- Yellow Red  D. Look at pH of color change: called the transition interval See Figure 24 p. 662 (orange for indicator listed above)

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

Titration

X. Titration strong base/weak acid: pH 9 phenolphthalein 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

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

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. (Va)(Ma)(#H+) = (Vb)(Mb)(#OH-)

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? (Va )(Ma)(#H+) = (Vb)(Mb)(#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

Titration Examples 2. How many ml of .400 M NaOH are needed to neutralize 50.0 ml of .200 M HBr? (Va)(Ma)(#H+) = (Vb)(Mb)(#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