CHAPTER 16: (HOLT)ACID-BASE TITRATION AND pH

I. Concentration Units for Acids and Bases

• A. Chemical Equivalents

• 1. Definition: quantities of solutes that have equivalent combining capacity

• a. Acid: mass of one equivalent is numerically equal to the mass of one mole of the acid divided by the number of protons(H+ or H3O +) that one mole of the acid can provide

• Example:

• HCl 36 g/mol; 1 eq = 1H +; 36 g/mol H +

• H2SO4 98g/mol; 2 eq = 2H +; 49 g/mol H +

• B. Base: mass of one equivalent is numerically equal to the mass of one mole of the base divided by the number of protons(OH-) that one mole of the base can provide

• Example:

• NaOH 40 g/mol; 1 eq = 1 OH-; 40 g/mol OH-

• Ca(OH)2 74 g/mol; 2 eq = 2 OH-; 37 g/mol OH-

B. Normality

• Definition: number of equivalents of solute per liter of solution

• N = eq of solute

• L of solution

C. Relationship Between Normality and Molarity

• N = nM

• N: Normality

• n: number of equivalents (# of H+= or OH-)

• M: Molarity

• Example: • 1M HCl = 1N HCl 1M NaOH = 1N NaOH

• 1M H2SO4 = 2N H2SO4 1M Ca(OH)2 = 2N Ca(OH)2

II. Aqueous Solutions and the Concept of pH

• A. Self-Ionization of Water• 1. Definition: Two water molecules interact to produce

a hydronium ion and a hydroxide ion by proton transfer - forms a weak electrolyte

• 2. [ ] is symbol used to indicate concentration in moles per liter (Molarity)

• 3. H2O + H2O <---> H3O+ + OH- ;

• in pure water [H3O+ ] = [OH- ]

• 4. [H3O+ ][OH- ] = 10-14

• 5. If the [H3O+ ] increases then the [OH- ] decreases or

• If the [H3O+ ] decreases then the [OH- ] increases

B. The pH scale• 1. pH -- the negative of the common logarithm of

the hydronium ion concentration

• pH = -log[H3O+ ]

• 2. Acid: pH < 7• 3. Base: pH > 7• 4. Neutral: pH = 7

C. Calculations involving pH• pH = -log[H3O+ ]

• 0.001 M HCl = [H3O+ ] =1 x 10 -3

• pH = -log[1 x 10-3]• pH = 3 (acid)

• {Remember that [H3O+ ][OH- ] = 1 x 10-14 ; so if [H3O+ ] = 1 x 10-3; then [OH- ] = 1 x 10-11

• FYI: there is also pOH = - log[OH- ] and• pH + pOH = 14

III. Acid-Base Titrations• A. Indicators

• 1. Definitions:

• a. indicators - weak acid or base dyes whose colors are sensitive to pH, or hydronium, concentration

• b. transition interval - the pH range over which an indicator changes color

• 2. Types of indicators

• a. Change color at about pH 7

• b. Change color below pH 7

• c. Change color above pH 7

B. The Principle of Titration

• Definitions:

• 1. Titration - the controlled addition and measurement of the amount of a solution of known concentration that is required to react completely with a measured amount of a solution of unknown concentration

• 2.Standard solution - a solution that contains a precisely known concentration of a solute

• 3. Equivalence point - in a neutralization reaction, the point at which there are equivalent quantities of hydronium and hydroxide ions

• 4. End point - the point in a titration where an indicator changes color

• 5. Primary standard - a highly purified compound, when used in solution to check the concentration of the known solution in a titration

•

C. Molarity and Titration• 1. Determine the moles of acid (or base) from the

standard solution used during titration

• 2. From a balanced chemical equation, determine the ratio of moles of acid (base) to base (acid)

• 3. Determine the moles of solute of the unknown solution used during the titration

• 4. Determine the molarity of the unknown solution

D. Normality and Titrations

• Va x Na = Vb x Nb

• Va : volume of the acid

• Na : normality of the acid

• Vb: volume of the base

• Nb: normality of the base