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