A.P. Chemistry Chapter 16 Outline
Acids and Bases
II.
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Acids and bases conduct electricity because they are electrolytes; they release
ions into solutions. Hydronium ions, H3O+, are responsible for the properties of
acids and hydroxide ions, OH-, are responsible for those of bases.
I.
Properties of Acids and Bases
A. Acids
 Taste sour
 React with active metals
 Conduct electricity
 React with bases to produce salts and water
 Turns pH paper red
 Common acids: sulfuric, nitric, phosphoric, hydrochloric, and acetic
B. Bases
 Taste bitter
 Slippery
 Conduct electricity
 React with bases to produce salts and water
 Turns pH paper blue
 Common bases: sodium hydroxide, potassium hydroxide
Arrhenius Acids and Bases
A. An Arrhenius acid contains hydrogen and ionizes in aqueous solution to form
hydrogen ions.
B. An Arrhenius base produces hydroxide ions in aqueous solutions
C. A strong acid or base is one that ionizes completely in aqueous solution. A strong
acid is a strong electrolyte. Acid strength increases with increasing polarity and
decreasing bond energy. Strong acids and bases are product favored.
D. Weak acid and bases are weak electrolytes. Examples: HF, acetic acid, HCN,
and organic acids (acids which contain the carboxyl group –COOH)
III.
Bronsted-Lowry Acids and Bases
A. A Bronsted-Lowry acid is a molecule or ion that is a proton donor. Recall
that the hydrogen ion, H+, is a proton, all Arrhenius acids donate hydrogen
ions and are thus Bronsted-Lowry acids as well.
B. A Bronsted-Lowry base is a molecule or ion that is a proton acceptor.
Ammonia, NH3, is an example of a Bronsted-Lowry base; it can accept a
proton to form NH4+
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Acids are described as monoprotic, diprotic, or triprotic, depending on
whether they can donate one, two, or three protons respectively.
Example: HCl is monoprotic – in water it donates one H+
H2SO4 is diprotic – it can donate two hydrogen ions
H3PO4 is triprotic- it can donate three hydrogen ions
IV.
Conjugate Acids and Bases: a pair of molecules or ions related to each
other by the loss or gain of a single H+ ion. One member of the
conjugate acid-base pair is always a reactant and the other is always a
product
Example: in the reaction NH3(aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq) ammonia is a base
and accepts a base and accepts a proton (a hydrogen ion, this makes ammonium its
conjugate acid. Water acts as an acid, donating the proton; and this makes the hydroxide
ion its conjugate base.
i. The species that remains after a Bronsted-Lowry acid has given up
a proton is the conjugate base of that acid
ii. The species that is formed when a Bronsted-Lowry base gains a
proton is the conjugate acid of that base.
iii. The stronger an acid is, the weaker its conjugate base; the stronger
a base is, the weaker its conjugate acid
iv. Proton-transfer reactions favor the production of the weaker acid
and the weaker base
v. Any species that can react as either an acid or a base is described
as being amphoteric. Water is amphoteric.
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V.
Naming Acids and Bases
A. Binary acids: an acid that contains only two elements: hydrogen and one
of the more electronegative elements. The name of a binary acid begins
with the prefix hydro- and the root of the second name follows this prefix
followed by –ic. Example: HCl = hydrochloric acid
a. Oxyaxids: an acid that is a compound of hydrogen, oxygen, and a
third element, usually a nonmetal. Example: HNO3 = Nitric acid
B. Organic acids and Amines
The COOH group, the carboxylic acid group, is polar and the hydrogen atom
is acidic
Amines are compounds like ammonia, where the nitrogen atom has a lone pair
of electrons which can accept a hydrogen ion (so they act as weak bases)
VI.
Neutralization Reactions
i. Neutralization is the reaction of hydronium ions and hydroxide
ions to form water molecules.
ii. A salt is an ionic compound composed of a cation from a base and
an anion from an acid
VII.
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Lewis Acids and Bases
 A Lewis acid is an atom, ion, or molecule that accepts an
electron pair to form a covalent bond. A Lewis acid
formula does not necessarily include hydrogen. Any
compound in which the central atom has three valence
electrons and forms three covalent bonds can react as a
Lewis acid.
 A Lewis base is an atom, ion, or molecule that donates an
electron pair to form a covalent bond.
 A Lewis acid-base reaction is the formation of one or more
covalent bonds between an electron pair donor and an
electron pair acceptor.
VIII. The Auto-ionization of water
H2O(l) + H2O (l)↔H3O+ (aq) + OH- (aq) kw = [H3O+][OH-] = 1.0 x 10-14
This equation applies to water and all aqueous solutions
In a neutral solution [H3O+] = [OH-] = 1.0 x 10-7
In an acidic solution hydronium concentration is greater than hydroxide
concentration and vise a versa
IX.
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X.
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XI.
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The pH scale
The pH is defined as the negative log if the hydronium concentration
pH = -log[H3O+]
in a neutral solution the pH is 7. Acids have pH’s less than 7 and bases
have pH’s greater than 7
the pH range is from about 1 for strong acids to 14 for strong bases
Ionization constants of Acids and Bases
HA (aq) + H2O(l) ↔H3O+(aq) + A- (aq)
KA= [H3O+][A-]/[HA]
KA= the acid ionization constant; the stronger the acid the larger the
constant
B(aq) + + H2O(l) ↔BH+(aq) + OH- (aq)
KB= [BH+][OH-]/[B]
KB= the base ionization constant; the stronger the base the larger its
constant
KA x kB = kW = 1.0 x 10-14
The successive k values for polypro tic acids decrease by a factor of
104 or 105, indicating that each ionization step occurs to a lesser extent
than the one before it
Hydrated metal ions as Acids
Many metal ions form weakly acidic aqueous solutions