Title: pH pKa2 pH pKa1
1Acids and Bases in aqueous solution.
pKa2 2.33
pKa1 9.45
pH pKa2 pH pKa1
2The properties of water
Water is held together by H-bonding, which gives
it its surprisingly high B.Pt. and M.Pt. for such
a small molecule. Note that the heavier H2S
molecule is a gas at room temperature. Water is
also weakly ionized into H and OH- ions, as
discussed next.
2.75 Ã…
H-bonds
3The self-ionization of water
- Water ionizes weakly to produce OH- and H ions,
where Kw is the ionization constant of water,
according to - H2O(l) ? H (aq) OH-(aq) Kw
10-14 - This means that the product of the concentration
(molarity) of H and OH- must always be
10-14. - Kw 10-14 H OH- 1
- Thus, for pairs of H and OH- concentrations
in solution we have -
- H 1.0 M, then OH- 10-14/1.0 10-14
M - H 0.01 M, then OH- 10-14/0.01 10-12
M - H 10-7 M, then OH- 10-14/10-7 10-7
M - H 10-12 M, then OH- 10-14/10-12
10-2 M - H 10-15 M, then OH- 10-14/10-15 10
M
4pH and pOH
- The concentration of H and OH- can be
expressed as the pH and pOH (p in pH or pOH
means potency, and means that any constant K
or concentration with a p in front of it is
represented as pK -log K , or pH -log H) -
- pH - log H pOH - log OH-
- We can show that
-
- pKw 14.0 pH pOH
5- Thus, for any solution where the pH is known,
- pH pOH 14.0, or pOH 14 pH.
- H pH pOH OH-
- 10 M -1.0 15.0 10-15 M 1 M
0.0 14.0 10-14 M 0.01 M 2.0 12.0 10-12
M 10-5.7 M 5.7 8.3 10-8.3 M 10-7 M 7.0
7.0 10-7.0 M - 10-9.3 M 9.3 4.7 10-4.7 M
- 10-12 M 12.0 2.0 10-2 M
- 10-14 M 14.0 0.0 1 M
- 10-15.09 M 15.09 -1.09 12 M
6Acid dissociation constants (Ka)
- A weak Brønsted acid such as acetic acid
(CH3COOH) will dissociate to give off protons in
aqueous solution according to - CH3COOH (aq) ? CH3COO- (aq) H (aq)
- The extent of such dissociation is controlled by
the acid dissociation constant, Ka. -
- Ka CH3COO- H
- CH3COOH
acetic acid acetate ion
conjugate acid conjugate base
proton
7The acid dissociation constant of acetic acid
- The value of Ka for acetic acid is 10-4.84, so
the pKa is 4.84. This gives us the expression - 10-4.84 CH3COO- H 2
- CH3COOH
-
- This expression can be used to solve problems
relating to the ionization of acetic and other
acids. For example, what is the pH of a 0.1 M
solution of acetic acid? Equation 2 must be
obeyed at all times, so we have - 10-4.84 CH3COO-H/0.1 H2/0.1
- so H (10-4.84 x 0.1)0.5 10-2.92 M, pH
2.92.
8- Notice that in solving the above problem,
ionization of acetic acid produces equal
concentrations of H and CH3COO- ions. One
assumes that the amount of ionization is small,
so that CH3COOH is not corrected for the small
amount that it decreases, and we assume that
CH3COOH is still 0.1 M. - b) What is the pH of a 1 M solution of ammonia
if the pKa of NH4 is 9.2? - Note pKa of an acid plus pKb of its conjugate
base pKw pKa(NH4) pKb(NH3) pKw 14.0. - pKb NH3 14.0 9.2 4.8, so Kb 10-4.8.
- 10-4.8 NH4 OH- / NH3
- so again we have OH- (10-4.8 x 1)0.5,
10-2.4. - If OH- 10-2.4 M, then pOH 2.4, and pH
14-2.4 11.6
9The species distribution diagrams of acids and
bases
- We can calculate the percentage of an acid (e.g.
CH3COOH) or its conjugate base (CH3COO-) that is
present as the acid or the conjugate base at any
given pH value if the pKa is known (4.84). Thus,
for acetic acid at pH 3.7 we have - 10-4.84 CH3COO- H / CH3COOH
- CH3COO- 10-3.7 / CH3COOH
- CH3COO- / CH3COOH 10-4.84 / 10-3.7
10-1.14 0.072 - so percent of CH3COO- at pH 3.7
- 100 x (0.072/(10.072)) 7.2 .
- of CH3COOH 100 7.2 92.8 .
10CH3COOH
CH3COO-
pH50 pKa
11The relationship between pH50 and the pKa
- The pH at which the concentrations of CH3COOH
and CH3COO- in solution are equal, i.e. both are
50, is known as the pH50. On the previous slide
it was indicated that the pKa of the acid equals
pH50. This arises simply as follows - Ka 10-4.84 CH3COO- H
- CH3COOH
- 10-4.84 H
- or pKa pH50
-
- Thus, we saw that for CH3COOH the crossover
point for CH3COOH and CH3COO- occurred at pH
4.84, which is the pKa. Similarly, on the next
slide, we see that the crossover point for NH4
and NH3 occurs at pH 9.22, which is the pKa of
NH4.
CH3COO- CH3COOH at pH50, so these cancel
12The species distribution diagram of
ammonia/ammonium ion
NH4
NH3
pH50 pKa
13Multiprotic acids and bases
- Many acids have more than one ionizable proton,
and many bases have more than one proton acceptor
site. A familiar example of this is triprotic
phosphoric acid - H3PO4 ? H2PO4- H pKa3 2.15
- H2PO4- ? HPO42- H pKa2 7.20
- HPO42- ? PO43- H pKa1 12.38
-
- It can be shown that for a monoprotic acid, when
the concentration of the acid and its conjugate
base are equal (both 50), then pKa pH. For
polyprotic acids the average extent of
protonation at any pH is given by the function
nbar. For the phosphate ion -
- nbar average number of protons bound per
phosphate
14The relationship between nbar, pH, and the pKa
values of multiprotic acids
- On a previous slide we showed that for a
monoprotic acid such as CH3COOH, the pKa pH50.
At pH50 for a monoprotic acid, in fact, nbar
0.5. This is the point at which half of the
acetate ions have a proton on them, while half do
not. For polyprotic acids, similar considerations
apply. At nbar 0.5 for phosphate, we have
HPO42- PO43-, and the pH at nbar 0.5
pKa1. In fact the half-values of nbar give us the
pKa values from the corresponding pH values - nbar pH pKan
- 0.5 pH pKa1
- 1.5 pH pKa2
- 2.5 pH pKa3 etc.
- Thus from a curve of nbar versus pH for
phosphoric acid, we can estimate the pKa values
from the pH values corresponding to nbar - 0.5, 1.5, and 2.5
15Plot of nbar versus pH for phosphate
H3PO4
pKa3 2.15
H2PO4-
HPO42-
pKa2 7.20
pKa1 12.38
PO43-
pH
pH pKa3 pH pKa2
pH pKa1
16Species distribution diagram for phosphoric acid
H3PO4 H2PO4-
HPO42- PO43-
pH
17pKa2 2.33
pKa1 9.45
pH
pH pKa2 pH pKa1
18NTA (nitrilo- triacetate)
pKa2 2.52 pKa1 9.46
EDTA (ethylene- diamine tetraacetate)
pKa1 9.52
pKa2 6.13
pKa3 2.69
19Lewis acids and bases
- You will recall Brønsted acids and bases, where
a Brønsted acid is a proton donor and a Brønsted
base is a proton acceptor. A broader definition
is that of Lewis Acids and Bases, where a Lewis
Acid is an electron acceptor, and a Lewis Base is
an electron donor.
BF3, a Lewis Acid
F-, a Lewis Base
Gilbert Newton Lewis (1875-1946)
Lewis Acid accepts electrons from the Lewis Base
to form a complex.