Transcript Chem 1310: Introduction to physical chemistry Part 0: Some

Chem 1310:
Introduction to physical chemistry
Part 4: Acids and bases in water
Peter H.M. Budzelaar
About acids and bases
Acids and bases play a key role in chemistry, in
particular in water (including in living things).
Acid-base reactions are fairly simple examples of
chemical equilibria. Because they are so important,
they still merit special attention.
Also, you will see a relation between chemical
structure and reactivity for the first time here.
What is an acid?
•Brønsted-Lowry:
– An acid is an H+ donor
– A base is an H+ acceptor
•Lewis:
– An acid is an electron-pair acceptor
– A base is an electron-pair donor
The Lewis definition is the more general one. In
water, the two are nearly equivalent, and talking
about H+ is easier than figuring out where the
electrons go, so we mostly use Brønsted-Lowry.
What is an acid? (2)
In water, we never have free H+ but rather H3O+.
This is what you will see in all equations. However,
acid-base reactions still involve transfer of a proton.
A "free proton" is so electron-poor and reactive it
will attach itself to anything it encounters.
Acids and bases in water
An acid donates a proton to water, forming H3O+.
H2S + H2O ⇋ HS- + H3O+
acid
conjugate base
conjugate base
acid
H2S is a neutral acid
H2O is a neutral base.
HS- is an anionic base.
H3O+ is a cationic acid.
The acid can donate a proton.
Its conjugate base can accept a proton.
Acids and bases in water (2)
NH4+ + H2O ⇋ NH3 + H3O+
acid
conjugate base
conjugate base
acid
HSO4- + H2O ⇋ SO42- + H3O+
acid
conjugate base
conjugate base
acid
NH4+ a cationic acid
HSO4- is an anionic acid.
Acids and bases in water (3)
The reaction doesn't have to be with water:
HSO4- + NH3 ⇋ SO42- + NH4+
acid
conjugate base
conjugate base
acid
But this is just a combination of the equations on the
previous slide.
Acids, bases and equilibria
For acid-base reactions (in water), we don't worry
about kinetics.
Equilibria are established instantly.
What about bases?
A base abstracts a proton from water, forming OH-.
H2O + NH3 ⇋ OH- + NH4+
conjugate acid
base
base
conjugate acid
It doesn't matter whether you are talking about an
acid and its conjugated base, or a base and its
conjugated acid...
What about bases?
H2O + H2O ⇋ OH- + H3O+
conjugate acid
base
base
conjugate acid
Water is an acid and a base!
Acid and base strength
In water, you cannot have acids stronger than H3O+
They simply protonate water, forming H3O+
quantitatively.
We call them "strong acids".
Nor can you have bases stronger than OH-.
They simply deprotonate water, forming OHquantitatively.
We call them "strong bases".
Water is a convenient reference
There will be a scale for acids:
strong
H3
O+
weak
H2O
very weak
and a similar one for bases:
very weak
H2O
weak
OH-
strong
Water as a reference
(H3O+/H2O/OH-)
H2SO4 + H2O
HSO4- + H3O+
H2SO4 is a very strong acid (>H3O+)
HSO4- is a very weak base, weaker than H2O
H2CO3 + H2O
HCO3- + H3O+
H2CO3 + OHHCO3- + H2O
H2CO3 is a weaker acid than H3O+
HCO3- is a weaker base than OHCH4 + OHCH3- + H2O
CH4 is a very weak acid, CH3- a very strong base.
Water as a reference
(H3O+/H2O/OH-)
H2O + H2O ⇋ OH- + H3O+
[H3O ][OH ]
14
Kw 
 10
2
[H2O]
KW depends on temperature
H2O is the
(factor of 10 over 30°).
"almost pure"
solvent
Always valid, also in presence
of added acids and bases.
If we know [H3O+], we also know [OH-] = KW/[H3O+].
No need to calculate separately.
Water as a reference
(H3O+/H2O/OH-)
In pure water, [H3O+] = [OH-], so:
x = [H3O+] = [OH-], x2 = 10-14, x = 10-7 mol/L.
We call a solution:
• neutral if [H3O+] = [OH-] = 10-7 mol/L.
• acidic if [H3O+] > [OH-] ([H3O+] > 10-7, [OH-] < 10-7)
• basic if [H3O+] < [OH-] ([H3O+] < 10-7, [OH-] > 10-7)
pH and acidity
pH is defined as
pH = - log [H3O+] ( [H3O+] = 10-pH )
base 10 logarithm, don't confuse with "natural logarithm" ln!
Defined in this way to have a convenient range,
normally 0-14:
pH = 7: [H3O+] = [OH-] = 10-7 mol/L, neutral
pH = 0: [H3O+] = 100 = 1 mol/L, very acidic (1M acid!)
pH = 14: [H3O+] = 10-14 mol/L, [OH-]= 1 mol/L,
very basic (1M base!)
pH and acidity
Since [H3O+][OH-] = 10-14,
log [H3O+] + log [OH-] = -14 = -pH + log [OH-]
pH = 14 + log [OH-] = 14 - pOH
this is
nearly always
negative!
Measuring pH
• Electronic ("pH meter")
Just put electrode in solution, read out the pH.
Needs to be calibrated, but is fairly accurate.
• Indicator (solution or paper)
Contains an organic base that changes colour on
protonation (or acid that changes colour on
deprotonation).
Convenient, easy to carry, limited accuracy
(ca 1 pH unit).
A scale for acid strength
We characterize the strength of an acid
by its Ka value:


[A
][H
O
]
3
+
K

HA + H2O ⇋ A + H3O
a
[HA]
The strength of its conjugate base
is given by its Kb value:

[HA][OH
]
K

A + H2O ⇋ HA + OH
b
[A ]
[A ][H3O ] [HA][OH ]


14
K a Kb 

[H
O
][OH
]

K

10
3
w
[HA]
[A ]
A scale for acid strength
You do not need separate values for Ka and Kb!
(for the same acid/base pair)
Whether a Ka for an acid is tabulated, or a Kb for its
conjugated base, is a matter of convention and
convenience.
Often, you will find tables of pKa (= - log Ka) and
pKb (= - log Kb).
pH calculations
for strong acids and bases
Strong acids and bases are always fully dissociated,
so we know immediately how much H3O+ or OH- is
generated.
0.02 M CsOH solution.
OH- is a strong base. [OH-] = 2·10-2 mol/L,
[H3O+] = 5·10-13 mol/L, pH = 12.3
0.13 M HClO4 solution.
HClO4 is a strong acid. [H3O+] = 0.13 mol/L,
pH = 0.9
pH calculations
for strong acids and bases
Take care! We were neglecting the contribution of
water itself, assuming all H3O+/OH- came from the
added acid/base.
For very low concentrations this is no longer true,
and we need to use the "x method".
Low concentrations of
strong acids and bases
6·10-7 mol/L HNO3. Originally, we had
[H3O+] = [OH-] = 10-7 mol/L. We add 6·10-7 mol/L
[H3O+], but some will be consumed by reaction
with [OH-].
H3O+
OH-
initial
10-7+6·10-7
10-7
change
-x
-x
equilibrium
7·10-7-x
10-7-x
Low concentrations of
strong acids and bases
Calculation: KW  [H3O  ][OH ]  (7 107  x)(107  x)
 x 2  8 107 x  7 1014  1014
 x 2  8 107 x  6 1014  0
8 107  (8 107 ) 2  4 * 6 1014
x
2
8 107  6.3 107

 0.8 107
2
Final [H3O+] = 6.2·10-7, pH = 6.2.
pH calculations
for weak acids and bases
Cannot assume all acid has dissociated.
Need to use the equilibrium expression for Ka.
0.022 M acetic acid (Ka = 1.8·10-5).
HOAc
OAc-
H3O+
initial
0.022
0
(10-7)
change
-x
+x
+x
equilibrium
0.022-x
x
x
pH calculations
for weak acids and bases
Calculation:
[H3O ][OAc ]
x2
Ka 

 1.8 105
[HOAc]
0.022 x
Simplified version, according to book (MSJ p791),
assuming x « 0.022:
x2
x2

 1.8 105  x 2  4.0 107  x  6.3 104
0.022 x 0.022
Final pH: 3.2
Check your answer: 0.022-6.3·10-4  0.021,
reasonable (but not ideal) approximation.
pH calculations
for weak acids and bases
Full calculation:
x2
 1.8 105
0.022 x
 x 2  4.0 107  1.8 105 x
 x 2  1.8 105 x  4.0 107  0
 1.8 105  (1.8 105 ) 2  4 * 4.0 107
x
2
 1.8 105  1.27 103

 6.2 10 4
2
Final pH: 3.2
pH calculations
for weak acids and bases
For very weak acids (Ka < 10-10) or very small
concentrations of weak acids ([HA] < 10-5)
we would need to take explicitly into account:
• partial dissociation of the acid
• auto-ionization of water
This gets too complicated for solving without a
computer.
Types of acids
Acids have a proton that is easily lost (transferred).
It will be attached to an electronegative atom X
(typically, at least as electronegative as nitrogen).
Examples:
H-F is acidic
H-OH is somewhat acidic
H-OClO3 is very acidic
Why are some acids weak, others strong?
There are trends, but there is not a simple, single rule.
H-X acids (X not oxygen)
CH4
NH3
H2O
HF
SiH4
PH3
H2S
HCl
GeH4
AsH3
H2Se
HBr
SnH4
SbH3
H2Te
HI
increasing
X electronegativity
(MSJ p355)
more stable X-
decreasing
X-H bond strength
(MSJ p352)
easier loss of H+
very weak acids
weak acids
strong acids
H-X acids (X not oxygen)
Second (and third) dissociation is always much more
difficult than first.
It is harder to remove H+ from an anion than from a
neutral molecule!
So HS- is a much weaker acid than H2S.
H-O-X acids
H2CO3 HNO3
HNO2
H2O
H2O2
H4SiO4 H3PO4
H3PO3
H2SO4
H2SO3
H3AsO4
H3AsO3
HOCl
HClO2
HClO3
HClO4
H2SeO4 HOBr
H2SeO3 HBrO2
HBrO3
Stronger if:
• X more
electronegative
• More oxygen atoms
attached to X
Both stabilize negative
charge on anion
Also transition-metal acids: H2CrO4, HMnO4, etc
Very strong oxo-acids:
So what about HNO2, HClO2, H2SO3 ?
Hydrated metal ions
Cr(OH2)63+ + H2O ⇋ Cr(OH2)5(OH)2+ + H3O+
Cr3+ is a Lewis acid, complexes
with Lewis base H2O.
The complexation makes water more acidic.
Higher charge, smaller radius of metal ion
 stronger Brønsted-acidity of hydrated ion.
Organic acids:
carboxylic acids
O
The C=O group
O
is electronC + H2O
withdrawing
OH
and allows resonance
stabilization. Without it,
the OH group is hardly acidic.
C
O
O
C
O
+ H3O+
Common organic acids
O
O
H C
CH3
OH
O
C
CH3CH2CH2
OH
Formic acid
(the stuff that hurts
when an ant bites)
OH
Acetic acid
(vinegar,
wine gone bad)
O
O
CH3(CH2)16
C
Butyric acid
("unwashed" smell)
O
CH3
O
C
C
OH
O
OH
Stearic acid
Benzoic acid
(the Na salt is
household soap)
(common food
preservative)
C
C
OH
Acetylsalicylic acid
(Aspirin)
The strength of an organic acid
(see MSJ p782)
O
Formic (our reference)
H C
OH
O
CH3
C
OH
O
CH2F
C
OH
Fluoroacetic: stronger
(F withdrawing)
O
CF3
Trifluoroacetic: strong!
C
OH
O
Acetic: weaker
(CH3 donating)
C
OH
O
O
C C
OH
CH3
Benzoic: stronger
(resonance)
Pyruvic: stronger
(withdrawing)
Curiosities
HCN: Hydrocyanic acid, Prussic acid.
Negative charge stabilized by N atom.
Toxic! (not because of acidity)
C N
HN3: Hydrazoic acid. Nitrogens are
N N N
less effective than oxygens at stabilizing negative
charge, but hey can do so.
Toxic (like HCN).
Salts used in explosives and airbags.
Curiosities (2)
HO
Ascorbic acid (vitamin C).
Essential food component
and preservative.
The anion has a
delocalized negative
charge and several
electron-withdrawing
oxygens.
OH
H
HCOH
O
O
H2COH
O
OH
H
HCOH
H2COH
O
O
O
OH
H
HCOH
H2COH
O
O
Types of bases
• Anions of weak acids.
Weaker acid  stronger conjugate base (Ka*Kb = Kw!).
• Hydroxides
• Ammonia, amines
Phosphines are much weaker bases
Bases are often used for cleaning purposes. They
break down many organic compounds ("liquid
plummr"). Don't get strong base on your skin.
Amines cause the "fishy" smell of fish.
A special case:
amino-acids, zwitterions
Amino-acids combine an acidic and a basic group in
the same molecule. In the solid state and in neutral
polar solvents, they exist as zwitterions:
H2N
CH2
COOH
H3N
CH2
COO
This is not resonance but tautomerism!
Amino acids
• In acidic solution, they exist in the protonated
cationic form:
H3N
CH2
COOH
• In basic solution, they exist in the deprotonated
anionic form: H N
COO
2
CH2
Peptides and proteins
Peptides (including proteins) are formed by coupling
amino acids head-to-tail, via peptide linkages.
These linkages are neither acidic nor basic.
Only the head, tail and some side groups (see MSJ
p592) are acidic or basic.
H2N
CH2
COOH
+
H2N
CH
COOH
- H2O
CH3
O
H2N
CH2
CH3
NH
CH
COOH
O
H3N
CH2
C
CH3
NH
CH
COO
Strength of bases
Bases must be able to accept a proton, so they must have an
available electron pair:
H
H2O is a weak base.
+
+H O
+H O
O
O
H
H
H
H
H
N
H
3
H
H
+ H 3O +
O + H3O+
O + H2O
H
H
H
O
O
2
H
H
N
H
H
+ H2O
+ H2O
+ OH-
NH3 is a stronger base.
N less electronegative 
electron pair more available
OH- is a strong base,
O2- is even stronger
(negative charge!).
Strength of bases
In general, provided an atom has an available
electron pair, basicity goes up:
• going up in the periodic table (NH3>PH3)
• going left in the periodic table (NH3>H2O)
• with increasing negative charge
(PO43->HPO42->H2PO4-)
• with more electron-donating substituents
(CH3OH>H2O)
Solutions of salts
• A salt consists of an anion (conjugate base of an
acid) and a cation (often an acid itself, or
conjugate acid of a base). It is formed by
neutralization of the acid with the base (or vice
versa).
• If both the original acid and the original base are
strong, the solution will be neutral (the conjugate
base and acid are very weak).
NaCl = salt of HCl (strong) and NaOH (strong): neutral.
Solutions of salts
• If the original acid is strong but the base is weak,
the solution is acidic.
NH4Cl: from HCl + NH3. NH4+ is (weakly) acidic.
• If the original acid is weak but the base is strong,
the solution is basic.
NaOAc: from HOAc + NaOH. OAc- is (weakly) basic.
• If both acid and base were weak:
– If Ka(acid) > Kb(base), solution is acidic.
– If Ka(acid) < Kb(base), solution is basic.
pH calculations for salts
• For a salt of strong acid and weak base: need only
consider conjugate acid of weak base. See
calculations for weak acids.
• For a salt of weak acid and strong base: need only
consider conjugate base of weak acid. See
calculations for weak bases.
• For salts of weak acid and weak base: need to
include both, calculation can be complicated.
Lewis acidity that is not
Brønsted-Lowry acidity
Metal ions are strong Lewis acids. They bind H2O,
but also other Lewis bases like NH3 or anions of
various acids.
Cu(H2O)62+ + 4 NH3
light blue
HgI2 + 2 Ired,
insoluble
AgCl + 2 S2O32white,
insoluble
HgI42-
Cu(NH3)42+ + 6 H2O
deep dark blue
colourless solution
Ag(S2O3)23- + Cl-
Lewis acidity that is not
Brønsted-Lowry acidity
Other elements can act like Lewis acids
(but generally the electron pairs move around then)
O
O C O
+ OH
O C
O
O
O
O S
+ OH
O
O S
O
OH
H