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The
Nature of
Acid-Base
Equilibria
Acids and Bases in Aqueous Solution
• Arrhenius Definition of Acids and Bases:
• An acid is a substance that gives hydrogen ions, H+,
when dissolved in water. In fact, H+ reacts with
water and produces H3O+.
• A base is a substance that gives hydroxide ions, OH-,
when dissolved in water.
• The neutralization reaction of an acid with a base
yields water plus a salt.
Chapter 8.1
Some Common Acids
• Some common strong acids:
•
•
•
•
Sulfuric acid, H2SO4
Hydrochloric acid, HCl
Nitric acid, HNO3
Perchloric acid HClO4
• Weak acids:
• Phosphoric acid, H3PO4
Some Common Bases
•Some common strong bases:
• Sodium hydroxide, NaOH;
• Calcium hydroxide, Ca(OH)2
• Magnesium hydroxide, Mg(OH)2
•Weak base:
• Ammonia, NH3
• Acetic acid, CH3CO2H
•The Bronsted-Lowry Definition of Acids and
Bases:
•Acid: Any substance that is able to give a hydrogen
ion, H+, to another ion or molecule. H+ ions are also
known as a proton. Therefore, acids are those
substances that can donate protons.
•Base: Any substance that is able to accept a
hydrogen ion, H+, from an acid. A base can be neutral
or negatively charged, for example, ammonia, NH3
and hydroxide ion, OH-.
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•Monoprotic acid : Acids with one proton to
donate, such as, hydrochloric acid, HCl;
nitric acid, HNO3
•Diprotic acid: Acids with two protons to
donate, such as sulfuric acid, H2SO4.
•Triprotic acid: Acids with three protons to
donate, such as phosphoric acid, H3PO4.
•Conjugate acid-base pair: Two substances
whose formula differ by only a hydrogen
ion, H+.
•Conjugate base: The substance formed by
loss of H+ from an acid.
•Conjugated acid: The substance formed by
addition of H+ to a base.
Acid Dissociated Constants
Water as Both an Acid and a Base
•According to Bronsted-Lowry acid-base theory,
water is both an acid and a base. Called
amphoteric.
•
NH3 + H2O  NH4+ + OH-
•
CH3COOH + H2O  CH3COO- + H3O+
• Strong acids have Ka value much greater than 1.
• Weak acids have Ka value much less than 1.
• Donation of each successive H+ from a
polyprotic acid is more difficult than the one
before it, so Ka value become successively
smaller.
• Most organic acids, which contains –CO2H
group, have Ka value near 10-5.
•The reaction of weak acid with water can be
described by an equilibrium equation.
Equilibrium
constant,
K,
and
water
concentration [H3O+] together makes the acid
dissociation constant Ka. Acid dissociation
constant is a measure of acid strength.
Dissociation of Water
•
H2O + H2O  H3O+ + OH-
•Ion product constant for water, Kw:
•Kw = [H3O+][OH-] = 1.00 x 10-14 at 25oC.
•
therefore, [H3O+]=[OH-] = 1.0 x 10-7 M
•Product of [H3O+] and [OH-] is a constant.
Therefore, in an acidic solution where [H3O+] is
large and [OH-] must be small.
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Measuring Acidity in Aqueous Solution: pH
•A pH value between 0 and 14 is used to indicate
concentration of H3O+ or OH- in solution.
Mathematically, the pH of a solution is defined as
the negative common logarithm of the H3O+
concentration:
•
pH = -log [H3O+ ] or
•
[H3O+ ] = 10-pH
•Acidic solution:
•Neutral solution:
• Basic solution:
pH < 7
pH = 7
pH > 7
[H3O+ ] > 1.00 x 10-7 M
[H3O+] = 1.00 x 10-7M
[H3O+ ] < 1.00 x 10-7 M
• “Acidity” depends on the concentration of
H3O+ ions
Acidic:
[H3O+] > [OH-]
Basic:
[H3O+] < [OH-]
Neutral:
[H3O+] = [OH-]
• Notice that neutral does NOT necessarily
mean pH 7
• pH is usually quoted with the same number of
significant digits as the concentration
• The p-scale can also be applied to ionization
constants
pKa = - log Ka
• The larger the value of Ka the smaller the value of
pKa and the stronger the acid
• From the acid-dissociation constant we can
calculate equilibrium concentrations as well as
pH
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