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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
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