Acids and Bases
Properties of Aqueous acids
• Taste sour (from the Latin acere, which
means 'sour' )
•
•
•
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Change blue litmus paper red
Conductive (electrolytes)
React with bases to form salt and water
Form hydrogen gas (H2) upon reaction
with active metals
Properties of Aqueous Bases
• Bitter
• Feel slippery or soapy
•
•
•
•
Change red litmus paper blue
Turn phenolphthalein red
Conductive (electrolytes)
React with acids to form salt and water
Definitions
• Arrhenius: aqueous solutions
– Acids release protons in solution
• HCl → H+ (aq) + Cl- (aq)
– Bases release hydroxide ions in solution
• NaOH → Na+(aq) + OH- (aq)
• Limited to protic acids (acids that have H+)
and hydroxide bases
Definitions
• Bronsted-Lowry
– Acids are proton donors
– Bases are proton receptors
• H2O + NH3 → OH- + NH4+
acid base
• Limited to protic acids but non-hydroxide
bases possible
• Solutions other than aqueous possible
Definitions
• Lewis acids/bases
– Acids: electron acceptor
– Bases: electron donor
– Example: NH3 + BF3
Lewis acids con’t
Metals often form Lewis
acids.
Transition metals have
empty “d” orbitals
that will accept an
electron pair
Naming Acids Review
• Acids form hydrogen ions when dissolved
in water. They are named according to the
anion in the compound
– Anions ending in “ide” : acid starts with
“hydro’ and anion ends with “ic”.
• Exp: HCl is hydrochloric acid
– Anions ending with “ite”: anion ends with
“ous”
• Exp: HNO2: nitrous acid
– Anions ending with “ate”: anion ends with
“ic”
• Exp: H2SO4: sulfuric acid
Hydronium ion
• In aqueous solutions, the proton donated
by an acid does not exist alone..it attaches
to a water molecule to form a new species,
the hydronium ion:
– HCl + H2O → H3O+ + Cl↑
hydronium ion
Conjugate acids and bases
• When an acid donates a proton and a base
accepts the proton, the remaining species in
solution are now in the position to behave as
acids and bases themselves
– Example:
base
conjugate acid
↑
has a proton to donate
may act as an acid
These are known as conjugate acids and bases
conjugate base
↑
has the ability to take a proton
may act as a base
Examples
• Determine the acid, base, conjugate acid
and conjugate base in the following
reactions:
– HNO3 + H2O
– HCO3 + H2O
– HCN + F– H3PO4 + H2O
Watch this!
• Acid Base summary with some
environmental info
Autoionization of water
• In any sample of water a very small
amount of ionization of the water
molecules does take place.
H2O(l) + H2O(l) → H3O+ (aq) + OH- (aq)
• In acid/base reactions that occur in water,
the acidity of the solution is determined by
the amount of H+ in solution
Autoionization of water
• The equilibrium expression for the
autoionization of water is:
Keq = [H3O+] [OH-] or [H+] [OH-]
= 1.0 x 10-14
This value is known as Kw or the equilibrium
constant for the ionization of water
Acidity
• When [H+] = [OH-] then each is equal to
1.0x 10-7 and the solution is neutral.
• If H+ > 1.0x 10-7 the solution is an acid
( H+ > OH-)
• If H+ < 1.0x 10-7 the solution is a base
(OH- > H+)
Calculating molarity of H+ and OH• Since [H+] [OH-] = 1.0 x 10-14 this equation may
be used to calculate changes in the amount of
either ion in solution.
• The addition of H+ or OH- will not change the
constant, but it will change the amount of each
ion individually.
• The amount added will always be far greater
than the initial concentration of H+ or OH- (1.0x
10-7) so the initial amount may be ignored.
Example
OH- is added to a solution until the final
concentration is equal to .0025M. What is the
concentration of H+?
[H+] [OH-] = 1.0 x 10-14
[H+] [.0025] = 1.0 x 10-14
[H+] [.0025] = 1.0 x 10-14
[.0025]
[.0025]
[H+] = 4.0 x 10-12
Is the solution an acid or a base?
• [H+] < 1.0 x 10-7 so it is a base
More examples
• Calculate the H+ and OH- in a solution
prepared by dissolving 0.0300 moles of HI
in enough water to form .500L of solution.
HI → H+ + I.0300mol/.500L = .0600M = H+
[.0600] [ OH-] = 1.00 x 10-14
[OH-] = 1.70 x 10-13
Acid or Base??
Practice
• Calculate the H+ and OH- in a solution that is
made by dissolving 40.0g HNO3 in water to
make 500. ml of solution.
40.0g/63.0g = .635 mol HNO3
.635 mol HNO3/.500L=1.27M HNO3
[H+] = 1.27
[1.27] [OH-] = 1.00 x 10 -14
[OH-]= 7.87 x 10-15 M
Strong acids vs weak acids
•
The only difference between a strong acid and a weak
acid is the amount of ionization that takes place.
•
Strong acids dissociate 100% (HCl, H2SO4 etc.)
•
Weak acids only dissociate a small amount
(CH3COOH, HF etc.). They will establish an
equilibrium with their ions in solution.
•
This has nothing to do with concentration!! Weak acids
may be very concentrated and still be dangerous
Strong Acids
HI
H+(aq) + I-(aq)
HBr
H+(aq) + Br-(aq)
HClO4 (perchloric acid)
H+(aq) + ClO4-(aq)
HCl
H+(aq) + Cl-(aq)
HClO3 (chloric acid)
H+(aq) + ClO3-(aq)
H2SO4
H+(aq) + HSO4-(aq) (HSO4- is a weak
acid that contributes additional protons)
H+(aq) + NO3-(aq)
HNO3
Strong Bases
• The hydroxide bases of the group I and
group II metals are strong bases even
though some (like Ca(OH)2) are not very
soluble.
• The conjugate bases of weak acids (Ka
10-13 or more) are considered strong
bases.
pH scale
• Because the values of H+ are so small, the
pH scale was developed to show the value
of H+ in an easier way.
pH = -log [H+]
Example: in an earlier problem H+ = .0600M
Therefore the pH would be: -log [.0600] = 1.22
***There is no significance to this except that it
makes the numbers easier to work with***
pH scale con’t
• The pH of a neutral solution:
when [H+] = [OH-] =1.0x 10-7
–log [1.0x 10-7] = 7
pH generally runs from 1.....7……14 where 1-6.99
is an acid, 7 is neutral and anything above 7 is a
base.
In reality, many solutions have pH higher than 14
and less than 1
Examples
• Calculate the pH of a solution prepared from
.00135 moles of NaOH in 100.ml of solution.
.00135/.100L = .0135M NaOH
[H+] [.0135] = 1.00 x 10-14
[H+] = 7.41 x 10-13
pH = ?
12.1
pOH
• pOH may also be calculated
• pOH = -log [OH-]
From the previous problem:
pOH= -log [.0135] = 1.87
Add the pH and the pOH. What do you notice??
pH + pOH = 14