1. Determining the pH and Titrations
Chapter 15.2

2. Acid-base indicators:
Acid-base indicators are chemical dyes whose colors are affected by acidic or basic solutions, and whose color changes as pH changes.
Gives an approximate value for pH of a solution.
Many indicators are weak acids with own pH range over which it changes colors – need to use indicator that changes color at equivalence point of titration.
pH meter
Determines exact pH by measuring voltage difference between 2 electrodes in a solution
Voltage changes as [H3O+] changes, used if exact pH value is needed

3. Acid-base indicators continued
Transition interval is pH range over which an indicator changes color
To titrate a strong acid/strong base – use bromothymol blue
To titrate a strong acid/weak base – use Methyl orange (pH change < 7)
To titrate a weak acid/strong base – use phenolphthalein or phenol red (pH change > 7)
Universal indicator mix of several indicators, used for pH paper

4. Titration Net equation for acid/base neutralization:
H3O+(aq) + OH-(aq) ↔ 2H2O(l)
Because acids & bases react; a progressive addition of acid to base (or base to acid) can be used to compare concentrations of acid and base
To analyze the acid or base content of a solution, a titration can be performed.
A titration is the controlled addition and measurement of the amount of a solution known concentration required to react completely with a measured amount of a solution of unknown concentration.

5. Basic steps of a titration
A measured volume of an acidic or basic solution of unknown concentration is placed in a beaker.
A buret is filled with the titration solution of known concentration. This called the standard solution, or titrant.
Measured volumes of the standard solution are added slowly and mixed into the solution in the beaker. This process is continued until it reaches its equivalence point.

7. Equivalence Point
Point at which the two solutions used in a titration are present in chemically equivalent amounts.
**Equivalence point is does not necessarily mean the solution has a pH = 7. **
Equivalence can be reached at different points for different types of solutions. Depending on the titration a different indicator is used:
Indicators with pH < 7 are used for strong-acid/weak-base titrations.
Ex: Methyl orange
Indicators with pH > 7 are used for weak-acid/strong-base titrations.
Ex: Phenolphthalein
There are no specific indicators used for weak-acid/weak-base titration b/c the equivalence point can be almost any value.

8. End Point­
Point in the titration where neutralization has occurred and the indicator changes color or there is a dramatic change in pH

9. Step for solving Titration: Combined Calculation
Write a balanced equation.
Determine number of [H+] and [OH-] from acid and base formulas
Use formula: MaVana = MbVbnb
(M = molarity, V = volume used, na = number of [H+] in acid or nb= number of [OH-] in base)
Rearrange equation to find information related to the unknown of unknown.

10. Practice Examples:
If 30.0 mL of 0.10 M NaOH neutralize 25.0mL of HCl, what is the molarity of HCl?
If 27.5 mL of M NaOH neutralize 25.0 mL of sulfuric acid. Determine the molarity of the acid.
__ NaOH + __ HCl → __ H2O + __ NaCl
MaVana = MbVbnb
(Ma)(25.0)(1) = (0.10)(30.0)(1)
(Ma) = 0.12M
__ NaOH + __ H2SO4 → __ H2O + __ Na2SO4 2 2
MaVana = MbVbnb
(Ma)(25.0)(2)= (27.5)(0.125)(1)
(Ma) = M → M