Acids and bases, salts and
solutions
Chapter 10-1 – 10-9, 11-1 – 11-4
Key concepts
• Compare and contrast the Arrhenius and
Brønsted-Lowry theories of acids and bases
• Describe hydrated protons
• Properties of acid and base solutions
• Arrange acids according to acid strength
• Balance acid-base equations
• Amphoterism
• Lewis acid-base theory
• Molarity calculations in titrations
• Equivalents
Various properties of acids and
bases
Acids
Bases
Taste
Sour (e.g., vinegar, citrus)
Bitter (e.g., baking soda)
Indicators
Litmus paper: red
Phenophthalein: clear
Litmus paper: blue
Phenophthalein: pink/purple
Reactivity w/
metals
Reacts w/ metals above H in
activity series to produce H2 (g)
Normally non-reactive
Reactivity
w/MO, MOH
React to form salt + H2O
Metal hydroxides are bases
Other rxns
Strong acid + weak acid salt →
weak acid + strong acid salt
Reacts with lipids (soap +
grease)
Acid + base → salt + water
conductivity
Electrolytic solution
Electrolytic solution
Arrhenius theory of acids/bases
• Developed in 1884.
• An acid is a substance containing
hydrogen that produces _______ in
aqueous solution.
• A base is a substance containing the OH
group that produces _______ in aqueous
solution.
Protons are not alone….
• Protons combine with water molecules to
form __________ ___________.
• We commonly represent this as the
hydronium ion, H3O+(aq), but writing
H+(aq) means the same thing.
Brønsted-Lowry theory (1923)
• An acid is a _______ _________.
• A base is a ________ __________.
• Bases are no longer restricted to
compounds that release OH- in solution.
For instance, NH3 is a base. What does it
look like after reacting with a proton?
Ionization of weak acids/bases
• While strong acids dissociate completely,
not all reactions are complete and
irreversible (in fact, most are not).
• Rxns with weak acids/bases are
reversible. Example: HF + H2O
• What is the acid? What is the base?
(depends on which side of rxn you look at)
Conjugate acid-base pairs
• Conjugate acid-base pairs differ in
structure by ___ _________.
• Some examples of conjugate pairs:
Acid/base strength
• The strength of an acid is _______ proportional
to the strength of its base.
• Strong acids have ________________.
• Weaker acids have ________________. As the
acid gets weaker and weaker, what happens to
the conjugate base? What does this tell you
about the amount of ionization taking place?
Amphoterism
• Some substances can both give and
accept protons. This process is called
amphoterism.
• Water is the prime example of amphiprotic
behavior.
2H2O → H3O+ + OH-
Acid strength
• The hydrohalic acids: HF, HCl, HBr, HI
• What are the sizes of the halogens? How
will this affect the H-X bond?
• HF bond is very strong vs. the other
halogens.
• F- causes ordering of the H2O molecules
(how does that happen?)
Leveling solvents
• In aqueous solution, no acid is stronger
than H3O+(aq). All other acids completely
dissolve in water to form H3O+.
• Because of this, all strong acids are of
equal strength in water.
• A similar effect is observed for strong
bases, which completely dissolve to form
OH-.
Ternary acids and bases
• What is a ternary
acid?
• Ternary acids are
hydroxyl compounds
of a
______________.
• Ionize to produce H+.
• Compare to other
hydroxyl
compounds…
• Metal hydroxides—
ionize to produce
________________
and are ________ in
aqueous solution.
Ternary acid strength
• H2SO4 vs H2SO3. What’s the difference in acid
strength?
– Compare oxidation number of sulfur in each.
• Acid strength increases with _______ oxidation
number of the central atom.
• Order the following acids from weakest to
strongest:
– HBrO3, HBrO, HBrO4, HBrO2
important!
• When comparing ternary acid strength,
make sure the compounds have similar
structure.
• Where are the hydrogens located?
• (H3PO3 vs H3PO4)
Neutralization of Brønsted-Lowry
acids/bases
HA + MB → HB + MA
• In many cases, HB ends up being ______.
• Classic example: strong acid + strong base.
• What happens in the reaction of hydrochloric
acid and sodium hydroxide?
• What is the net ionic equation?
Weak acid + strong base
• General reaction:
HxA (aq) + x OH- (aq) → A- (aq) + x H2O (l)
– When does x vary?
– Examples:
Acid salts
• Acid salts are salts of ______ acids that
still contain __________ ___________.
Lewis theory
• The most general of all acid-base theories
– Discards the proton acceptor/donator all together.
• A Lewis acid _______ a share in an electron
pair.
• A Lewis base _________ a share in an electron
pair.
• Lewis acids and bases are neutralized when a
________ _______ forms.
Arrhenius or B-L theory
a better description
for most aqueous sol’ns
When is Lewis
theory used?
LEWIS THEORY
Bronsted-
Arrhenius
Lowry
Lewis theory a good
descriptor for
nonaqueous
solvents or transition
metals
Acid-Base calculations
• Molarity calculations play an important part
in acid-base reaction stoichiometry
• Much of what we will learned in Chapter 3
will be used here.
Molarity
• M = mol/L or
• M = mmol/mL
• we can use moles and liters, or millimoles
and milliliters, and the molarity is still the
same.
Similarities between acid-base and
other reaction calculations
• We still compare moles to moles, not
volumes to volumes or molarities to
molarities.
• Additionally, knowing the limiting reactant
is very important (i.e, what will run out
first—acid or base?)
Some examples
1. what volume of 0.800 M NH3 is required
to neutralize 22.0 mL of 12.0 M HCl?
2. 25.0 mL of 0.0500 M Ca(OH)2 added to
10.0 mL of HNO3.
– Is the solution now acidic or basic?
– how many moles excess acid or base are in
the solution?
– how much additional Ca(OH)2 or HNO3 sol’n
required to neutralize solution?
TITRATIONS
• Combining a known concentration with an
unknown concentration solution.
• Titrant: The solution of one reactant
(usually of unknown concentration) that is
carefully added to the solution of the other
reactant until the resulting solution is just
neutralized (no excess acid or base).
• How do we know when to stop?
Titrations (con’t)
•
•
•
•
indicators:
How to measure the volume of titrant?
Buret:
equivalence point: The point where
_____________ _______________ amounts
of acid and base have reacted.
• end point: The point where the indicator
____________ ___________.
• For accurate work, one wants the end point and
equivalence point to coincide with each other.
Primary and secondary standards
• Reading on standardization: your text
goes over the requirements of a primary
standard. You should be familiar with
these requirements.
• Primary standards are used to determine
the concentration of solutions, which
become secondary standards.
• Example: KHP and NaOH.
EQUIVALENT WEIGHTS AND
NORMALITY
• One mole of acid is
_____________________________
• But, one equivalent of acid contains
______________________________.
• The equivalent weight, then, corresponds
to molar mass/(# of equ./mol)
normality
• number of equivalents per liter, or N = eq/L
= meq/mL
• N = M  eq/mol
• Let’s do a couple of examples…
EQUIVALENTS in acid/base
reactions
• 1 eq acid always reacts with 1 eq base.
• Va Na = Vb Nb