ACIDS AND BASES
Chapter 15
What are acids and bases
Arrhenius definition
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Arrhenius suggested that acids are compounds that
contain hydrogen and can dissolve in water to
release hydrogen ions into solution. For example,
hydrochloric acid (HCl) dissolves in water as follows:
HCl +H2OH+(aq) + Cl-(aq)
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Arrhenius defined bases as substances that dissolve
in water to release hydroxide ions (OH-) into
solution. For example, a typical base according to
the Arrhenius definition is sodium hydroxide
(NaOH):
NaOH +H2ONa+(aq) + OH-(aq)
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Acids release H+ into solution and bases release
OH-. If we were to mix an acid and base together,
the H+ ion would combine with the OH- ion to make
the molecule H2O, or plain water:
H+(aq) + OH-(aq) H2O
What are acids and Bases
Bronsted Lowry theory
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The Bronsted -Lowry definition of acids is very
similar to the Arrhenius definition, any substance
that can donate a hydrogen ion is an acid (under
the Bronsted definition, acids are often referred to
as proton donors because an H+ ion, hydrogen
minus its electron, is simply a proton).
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The Bronsted definition of bases is, however, quite
different from the Arrhenius definition. The Bronsted
base is defined as any substance that can accept a
hydrogen ion. A base is the opposite of an
acid. NaOH and KOH, would still be considered
bases because they can accept an H+ from an acid
to form water.
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The Bronsted-Lowry definition also explains why
substances that do not contain OH- can act like
bases. Baking soda (NaHCO3), for example, acts
like a base by accepting a hydrogen ion from an
acid as illustrated below:
Acid Base
Salt
HCl+ NaHCO₃ H2CO3 + NaCl
Neutralization reaction
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We have seen from the equations, acids release H+
into solution and bases release OH- If we were to mix
an acid and base together, the H+ ion would combine
with the OH- ion to make the molecule H2O, or plain
water:
H+(aq) + OH-(aq)  H2O. The neutralization reaction
of an acid with a base will always produce water and
a salt:
Acid Base Water Salt
HCl + NaOH  H2O + NaCl
HBr + KOH  H2O + KBr
Self ionization of water
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Water molecules collide with one another to cause
the self-ionization reaction represented by this
equation: 2H2O H3O+ + OHIt is a reversible reaction so the equation is usually
written with the arrows going in both directions:
2H2O(l) ↔H3O+ (aq) + OH- (aq)
The self ionization of pure water produces equal
amounts of H₃O+ ions and OH- ions. Therefore the
concentrations of these ions in pure water must be
equal. [OH- ]=[H3O+ ]
It has been found that at 25⁰C the concentration of
these two ions are equal to 1.00x10-7 mol/L
 Recall the equation for equilibrium constant
Keq which can be calculated from the equation.
Since self ionization is an equilibrium reaction, the
equilibrium constant called Kw, can be calculated as
follows.
Kw=[H3O+ ] [OH- ]=(1.00x10-7 )(1.00x10-7)=
1.00x10-14
Kw= 1.00x10-14 Kw is called as autoionisation
constant.
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Class Practice
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What is [OH- ] in a 3.00x10-5 M solution of HCl?
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Section review on page 549
Acids donate protons
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Any molecule or ion that can transfer a proton; a
hydrogen atom nucleus to another species.
Let us take the example of sulfuric acid
H2SO4(l) + H2O(l)- H 3O+ (aq) +HSO 4- (aq)
The hydrogen sulfate ion, is also a Bronsted Lowry
acid because it too can donate a proton to a water
molecule.
HSO 4- (aq) +H2O(l)↔ H3O+ (aq) +SO42- (aq)
Monoprotic , diprotic and triprotic
acids
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Monoprotic acids are acids that can donate one
hydrogen ion per molecule.
Example nitric acid HNO₃ or hydrochloric acid HCl.
Diprotic acids are acids that can donate two
hydrogen ions per molecule.
Examples sulfuric acids H₂SO₄.
Triprotic acids are acids that can donate three
hydrogen ions per molecule.
Example phosphoric acid H₃PO₄
Bases accepts protons
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Any atom, or molecule that receives a proton from
another species.
Ammonia is a typical base
NH₃(aq)+H₂O(l) ↔NH₄+ (aq)+ OH- (aq)
In the gas phase ammonia accepts a proton from
HCl and forms ammonium chloride which is
composed of ammonium and chloride ion.
Species that are both acids and bases.
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Water can act as an acid by donating a proton, and
can act as a base by accepting a proton.
H₂O(l)+H₂O(l)↔OH- (aq) +H₃O+ (aq)
A species that can act as either an acid or a base is
called as amphoteric.
Hydrogen carbonate ion is also an example of
amphoteric species.
HCO₃- (aq) + OH- (aq) ↔CO₃2- (aq)+H₂O(l)
It behaves as abase in the presence of an acid such as
formic acid
HCOOH (aq) + HCO₃- (aq) ↔HCOO- (aq) + H₂CO₃(aq)
Conjugate acids and bases
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Conjugate acids are formed when a base accepts a
proton.
Conjugate base are formed when an acid donates
a proton.
NH₃(aq)+H₂O(l)↔NH₄+ (aq) + OH- (aq)
Ammonium ion is the conjugate acid and OH ion is
the conjugate base.
Class Practice
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Page 554 concept check
Weak acids and bases
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Weak acids and bases are partially ionized in their
solutions, whereas strong acids and bases are
completely ionized when dissolve in water. Common
Weak Acids
Formic HCOOH
Acetic CH3COOH
Trichloro acetic CCl3COOH
Hydrofluoric HF
Hydrocyanic HCN
Hydrogen sulfide H2S
Water H2O
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Common Weak Bases
ammonia NH3
Trimethyl ammonia N(CH3)3
pyridine C5H5N
Ammonium hydroxide NH4OH
water H2O
Acid ionization constant
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The equilibrium constant of the ionization of a weak
acid in water is known as the acid ionization
constant Ka.
CH₃COOH(aq)+H₂O(l)↔CH₃COO- (aq) +H₃O+ (aq)
The equilibrium expression for this reaction is
written as follows
Ka=[H₃O+ ][CH₃COO- ]/[CH₃COOH] =1.75x10-5
Class practice
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A vinegar sample is found to be 0.837M acetic
acid. Its hydronium ion concentration is measured as
3.86x10-3 mol/L. Calculate Ka for CH₃COOH.
Home work
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Page 558
Total recall
pH
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What is pH?
The negative logarithm of the hydronium ion
concentration in a solution.
pH of a solution ranging from 1-7 are usually
acidic.
pH of a solution ranging from 7-14 are usually
basic.
pH of a solution which is 7 are neutral.
Class Practice
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Page 562
Practice problems all
pH and Kw
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Kw=[H₃O+][OH-]=1.00x10-14
What is the pH of a 0.0136M solution of Ba(OH)₂ a
strong base?
Class Practice
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Page 563
Practice problems
Buffers
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A buffer solution is one which resists changes in pH
when small quantities of an acid or an alkali are
added to it.
An acidic buffer solution is the one which has a pH
less than 7. Acidic buffer solutions are commonly
made from a weak acid and one of its salts often a
sodium salt.
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An alkaline buffer solution has a pH greater than 7.
Alkaline buffer solutions are commonly made from a
weak base and one of its salts.
Titration
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The operation of gradually adding one solution to
another to reach an equivalence point.
Equivalence point is the point in a titration when the
amount of added base or acid exactly equals the
amount of acid or base originally in solution.
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Titrant: The solution added to another solution in a
titration.
Home work
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Term Review all Page 580
15,17,29,41 ,47,66
Test prep all