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 +
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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
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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
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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
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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