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