Ch 14.1
Properties of Acids and Bases
Acids
 Are sour to taste
 React with bases to produce salts and
water.
 React with metals and release H2 gas
 Turn litmus paper red
 Conduct an electric current.
 Examples: citrus fruit, tomatoes, HCl,
and stomach acid
Acid Nomenclature
 Binary Acid: an acid that contains only
2 different elements: Hydrogen &one
of the more electronegative elements.
 Naming: begins with hydro- followed
by the root of the 2nd element & ends
in -ic.
 Oxyacid: an acid that is a compound of
hydrogen, oxygen, and third element,
usually a nonmetal.
Common Industrial Acids





Sulfuric Acid
Nitric Acid
Phosphoric Acid
Hydrochloric Acid
Acetic Acid
Bases
 Are bitter to taste
 Feel slippery
 React to neutralize acids, forming salt
and water
 Turn litmus paper blue
 Conduct an electric current.
 Examples: Baking Soda, Soaps,
Ammonia, and NaOH
Arrhenius Acid
 Produces H+ in an aqueous solution.
Once present, the H+ combines with
water to form H3O+ (called
hydronium).
 The dissociation of HCl looked like:
HCl  H+ + Cl But it was actually:
HCl + H2O  Cl- + H3O+
Arrhenius Base
 Produces OH- in solution.
 Examples:
NaOH  Na+ + OHNH3 + H2O  NH4+ + OH-
Dissociation
 The strength of an acid or base
depends on the amount of
dissociation that occurs. This
depends on the polarity of the bond
and the ease at which the bond can
break.
 Organic Acids containing
carboxyl/acid groups COOH, like
vinegar, are generally weak.
Ch 14.2
Acid-Base Theories
Bronsted-Lowry Acids & Bases
 Do not require aqueous solutions.
 Acids are defined as proton (H+) donors
and bases are proton (H+) acceptors.
 Examples:
HCl + NH3  NH4+ + ClA
B
H2O + NH3  NH4+ + OHA
B
Monoprotic and Polyprotic Acids
 Monoprotic: An acid that can donate
only one proton (H+) per molecule.
HCl + H2O  H3O+ + Cl Polyprotic: An acid that can donate more
than one proton (H+) per molecule.
 Diprotic: Can donate 2 protons. H2SO4
 Triprotic: Can donate 3 protons. H3PO4
Lewis Acids and Bases
 Lewis of electron dot structure fame
 Where both Arrhenius and BronstedLowry acids contain or produce H+,
Lewis acids don’t have to.
 An acid is defined as an atom, ion, or
molecule that accepts an electron pair
to form a covalent compound.
 A base donates an electron pair.
Lewis Acids and Bases
 They can exist in solid, liquid, or gas
phase.
 Examples
H+ + NH3  NH4+
H2O + HCl  Cl- + H3O+
BF3 + F-  BF4-
Ch 14.3
Acid-Base Reactions
Conjugate Acids and Bases
 When a Bronsted-Lowry acid gives up
a proton, it becomes a B-L Base
because it can then accept a proton.
Reactions usually involve two acidbase pairs.
HCl + NH3  NH4+ + ClA
B
CA
CB
H2O + NH3  NH4+ + OHA
B
CA
CB
 In general, strong acids have weak
conjugate bases and weak acids
have strong conjugate bases.
 Proton-transfer reactions favor the
production of the weaker acid and
weaker base.
Amphoteric Substances
 Substances which can be an acid or a
base.
 Whether it acts as one or the other
depends on the strength of what it reacts
with.
 Example: H2O:
H2O + NH3  NH4+ + OHH2SO4 + H2O  H3O++ HSO4-
Neutralization Reactions
 A base neutralizes an acid forming salt and
water. But things are not quite so simple.
The return of the SPECTATOR ION!
HCl + NaOH  NaCl + H2O
 Base Dissociation
NaOH  Na+ + OH Acid Dissociation
HCl +H2O  Cl- + H3O+
 Then write Overall and then Net Ionic.
Acid Rain
 Rain that is very acidic
 Industrial processes produce gases
such as NO, NO2, CO2, SO2, and SO3.
These can dissolve in water producing
acidic solutions that fall to the ground
in the form of rain. This Acid Rain can
erode statues and affect ecosystems.