What is and Acid?
• Arrhenius Acid
• Defined as any chemical that increases
the concentration of hydrogen ions (H+)in
solution. (usually by a dissociation
reaction
Examples
• Hydrochloric acid HCl  H+ + Cl• Sulfuric acid H2SO4  2H+ + SO42• Phosphoric acid H3PO4  3H+ + PO43-
What is a Base?
• Arrhenius Base
• Defined as any chemical that increases
the hydroxide ions (OH-) concentration in
solution.
• Examples• NaOH  Na+ + OH• KOH  K+ + OH• Ca(OH)2  Ca2+ + 2OH• The Arrhenius definition was limited to
substances which were soluble in water
Not all Hydrogens are acidic
• The hydrogen must be part of a polar
bond in order to dissociate.
• For example:
• HF is acidic, but CH4 is not.
• In CH4, the hydrogen is part of a non-polar
covalent bond and does not dissociate in
solution!
• HF is a polar bond and HF  H+ + F-
Bronsted – Lowry Acid
• Defined as a molecule or ion that is a hydrogen
ion donor.
• Also known as a proton donor because H+ is a
proton.
• The acid will donate its H+ ion to a base in an
acid base reaction.
H+ + OH-  H2O
Acid + Base
Bronsted-Lowry Base
• Defined as a hydrogen ion acceptor.
• In an acid-base reaction the base
“accepts” the hydrogen ion from the acid.
NH3 + H+  NH4+
NH3 accepts the H+ from the acid.
Brønsted-Lowry Theory
Brønsted-Lowry Acid/Base Theory: deals with the exchange of hydrogen
ions, H+, between two species…slightly less specific than Arrhenius. All
Arrhenius are also Brønsted-Lowry, but not all Brønsted-Lowry are also
Arrhenius. The acid and base always appear in the same reaction for
this type.
Brønsted-Lowry Acid: a species that donates H+ to another
participant in a chemical reaction.
Brønsted-Lowry Base: a species that accepts H+ from another
participant in a chemical reaction.
Example: HNO2 + CN- → HCN + NO2Conjugate Acid / Conjugate Base: species resulting from the transfer of
the proton that are also capable of interacting.
Brønsted-Lowry Theory
Example: HNO2 + CN- → HCN + NO2Conjugate Acid / Conjugate Base: species resulting
from the transfer of the proton that are also capable
of interacting.
Lewis Acids and Bases
• Lewis Acid: An electron pair acceptor
• Lewis Base: An electron pair donor
• The Lewis definition includes all the Arrhenius
and Brönsted-Lowry acids and bases and
includes substances such as BF3 and boric acid,
which do not fit under the other definitions.
Acid Dissociations
• Monoprotic:
• HCl(aq)⇄ H+ (aq)+Cl-(aq)
• CH3COOH (aq) ⇄H+ (aq) + CH3COO(aq)
• HNO3 (aq) ⇄H+ (aq)+ NO3-(aq)
•
• Diprotic: (two step dissociation)
• H2SO4(aq) ⇄ H+(aq) + HSO4- (aq)
• HSO4-(aq) ⇄ H+(aq) + SO42- (aq)
• H+ may be written in aqueous media as
the hydronium ion: H3O+
The Autoionization of Water and
the Hydronium ion
• At 1 atm of pressure and 25 oC, water
undergoes the following equilibrium
dissociation:
•
2H2O(l) ⇄ H3O+(aq) + OH-(aq)
• The H3O+ is called the hydronium ion and
is used synonymously with the symbol for
a proton, H+.
• In reality, in an aqueous solution, the
protons are hydrated as the hydronium
ion, but the H+ is used for convenience.
• The water equilibrium can also be written:
H2O(l) ⇄ H+(aq) + OH-(aq)
•
• This indicates that water autoionizes to
produce both an acid (H+) and a base
(OH-). Any substance which can act as
both an acid and a base is termed
amphoteric or amphiprotic.
The Titration
• One of the most important lab procedures
involving acids and bases is the titration.
• A titration is an analytical procedure that allows
for the measurement of the amount of one
solution that is required to exactly react with the
contents of another solution.
• In acid-base terms, you add one solution to the
other until the equivalence point is reached.
• The use of a pH meter will produce a pH curve
(titration curve), so you can specifically calculate
at what pH your solutions have been
neutralized.
Acid-Base Titration Terms to
Know
•
•
•
•
•
•
Titrant: the standard solution of known molarity in the buret that is being
added to the solution in the flask. This is more often the acid than the
base.
Analyte: the solution in the flask of unknown concentration. Usually the
base.
Indicator: a compound that is added in small amounts (a few drops) in
acid-base titrations. It changes color over a certain pH range, and
indicates the end of the titration. This range should be matched with pH
at which you expect your solutions to reach the equivalence point.
Endpoint: the point at which the titration is stopped, when the indicator
permanently changes color. Traditionally, this is the point when the
titration is stopped, where the number of moles of titrant is equal to the
number of moles of analyte, or some multiple thereof (as in di- or triprotic acids)
Equivalence point (a.k.a. neutralization or endpoint): the point (in mL of
solution added) at which the number of moles of acid equal the number
of moles of base.
Half-equivalence point: the (in mL of solution added) at which the
number of number of moles of acid (or base) added is half the number
of moles of base (or acid) present in the solution.
Standardize a solution of KOH