Download ACID BASE - Union City High School

  Acids
are substances that ionize in
aqueous solutions to form hydrogen
ions, thereby increasing the
concentration of H+ ions.
  Because hydrogen atom consists of a
proton and an electron, H+ is simply a
proton.
  Thus, acids are often called proton
donors.
  Molecules
of different acids can ionize to
form different numbers of H+ ions.
  Both
hydrochloric acid and nitric acid are
monoprotic acids, which yield one per
molecule of acid.
  Sulfuric
acid is a diprotic acid, one that
yields two H+ per molecule of acid.
  Bases
are substances that accept H+
ions.
  Bases
produce hydroxide ions when
they dissolve in water.
  When
dissolved in water, they
dissociate into their component ions,
introducing OH- ions into the solution.
ACID
BASE
Sour taste
Bitter taste
Neutralizes bases
Neutralizes acids
Turns litmus paper blue to red
Turns litmus paper red to blue
Soapy and slippery feeling
Indicator
Color in
strongly acidic
solution
pH at which
color changes
Color in
strongly alkaline
solution
Methyl orange
Red
4
Yellow
Litmus
Red
7
Blue
Phenolphthalein
Colorless
9
Red
Screened
Methyl Orange
Red
4
Green
  Acids
and bases that are strong
electrolytes (completely ionized in
solution) are called strong acids and
strong bases.
  Those
that are weak electrolytes (partly
ionized) are called weak acids and weak
bases.
STRONG ACIDS
STRONG BASES
Hydrochloric acid
Lithium hydroxide
Hydrobromic acid
Sodium hydroxide
Hydroiodic acid
Potassium hydroxide
Chloric acid
Rubidium hydroxide
Perchloric acid
Cesium hydroxide
Nitric acid
Calcium hydroxide
Sulfuric acid
Strontium hydroxide
Barium hydroxide
 
To classify a soluble substance as a strong
electrolyte, weak electrolyte, or nonelectrolyte,
we simply use the following table:
Strong
Electrolyte
Weak
Electrolyte
Nonelectrolyte
Ionic
All
None
None
Molecular
Strong acids
Weak acids
Weak bases
 
 
All other
compounds
If an acid is not listed, it is probably a weak electrolyte.
NH3 is only a weak base that we consider.
  When
a solution of an acid and that of
a base are mixed, a neutralization
reaction occurs.
  The
products of the reaction have
none of the characteristics properties
of either the acidic and the basic
solutions.
  By
analogy to this reaction, the term salt
has come to mean any ionic compound
whose cation comes from a base and
whose anion comes from an acid.
  A
neutralization reaction between an acid
and a metal hydroxide produces water
and salt.
1. 
Arrhenius Acids and Bases
2. 
Bronsted-Lowry Acids and Bases
3. 
Lewis Acids and Bases
 
Swedish chemist Svante Arrhenius (1859-1927)
proposed a revolutionary way of defining and
thinking about acids and bases.
 
He said that acids are hydrogen-containing
compounds that ionize to yield hydrogen ions (H+)
in aqueous solutions.
 
He also said that bases are compounds that ionize
to yield hydroxide ions (OH-) in aqueous solutions.
 
Acids that contain one ionizable hydrogen, such as
nitric acid, are called monoprotic acids.
 
Acids that contain two ionizable hydrogens, such
as sulfuric acid, are called diprotic acids.
 
Acids that contain three ionizable hydrogens, such
as phosphoric acid, are called triprotic acids.
 
In 1923, the Danish chemist Johannes Bronsted
and the English chemist Thomas Lowry
independently proposed a new definition.
 
Defines an acid as a hydrogen-ion donor.
 
Defines a base as a hydrogen-ion acceptor.
 
A conjugate base is the particle that remains when
an acid has donated a hydrogen ion.
 
A conjugate base is the particle that remains when
an acid has donated a hydrogen ion.
 
Conjugate acids and bases are always paired with
a base or an acid, respectively.
 
A conjugate acid-base pair consists of two
substances related by the loss or gain of a single
hydrogen ion.
 
The third theory of acids and bases was proposed
by Gilbert Lewis.
 
Lewis focused on the donation or acceptance of a
pair of electrons during a reaction.
 
This concept is more general than either the
Arrhenius theory or the Bronsted-Lowry theory.
 
A Lewis acid is a substance that can accept a pair
of electrons to form a covalent bond.
 
A Lewis base is a substance that can donate a
pair of electrons to form a covalent bond.
 
A hydrogen ion (Bronsted-Lowry acid) can accept
apair of electrons in forming a bond.
 
A hydrogen ion, therefore, is also a Lewis acid.
 
A Bronsted-Lowry base, or a substance that
accepts a hydrogen ion, must have a pair of
electrons available and is also a Lewis base.
 
A widely used system for expressing [H+] is the pH
scale, proposed in 1909 by the Danish scientist
Soren Sorenson.
 
It ranges from 0-14, neutral solutions have a pH of
7.
 
A pH of 10 is strongly basic.
 
The pH of a solution is the negative logarithm of
the hydrogen-ion concentration.
  The
pH may be represented
mathematically using the following
equation:
pH = - log [H+]
  Similarly,
the pOH of a solution equals
the negative logarithm of the
hydroxide-ion concentration.
pOH = - log [OH-]
 
A neutral solution has a pOH of 7.
 
A solution with a pOH less than 7 is basic.
 
A solution with a pOH greater than 7 is acidic.
 
A simple relationship between pH and pOH makes
it easy to find either one when the other is known.
pH + pOH = 14
pH = 14 – pOH
pOH = 14 - pH
2
3
4
5
6
Increasing Acidity
7
NEUTRAL
1
8
9
10
11
12
Increasing Basicity
10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-1110-12
10-1310-14
13
14
  What
is the pH of a solution with a
hydrogen-ion concentration of 1.0 x
10-10 M?
  Find the pH of each solution:
◦  [H+] = 1.0 x 10 -4 M
◦  [H+] = 0.0010 M
◦  [H+] = 1.0 x 10 -9 M
◦  [H+] = 1.0 x 10 -12 M
◦  [H+] = 0.010 M