Chapter 8 Acids and Bases
Treatment of acid-base theory enhanced
Brønsted-Lowry Theory
Acid-Base Equilibrium Constants
Brønsted-Lowry Acids
Brønsted-Lowry Bases
*Trends in Acid-Base Behavior
Acid-Base Reactions of Oxides
Lewis Theory
*Pearson Hard-Soft Acid-Base Concepts
Applications of the HSAB Concept
Application of the HSAB Concept to Qualitative Analysis
The HSAB Concept in Geochemistry
Interpretation of the HSAB Concept
*Biological Aspects
Chapter 8. Acids and bases
Definitions of acids and bases: Arrhenius, Bronsted
acidity, Lewis
Arrhenius
acid: generates [H+] in solution
base: generates [OH-] in solution
normal Arrhenius equation: acid + base <---> salt
+ water
example: HCl + NaOH <---> NaCl + H2O
Bronsted-Lowery:The Brønsted/Lowry Defintions
specifies an acid as a proton donor and a base as a proton acceptor which
applies to aqueous systems.
The General Solvent System Definition is an extension
to any autoionizing solvent. An acid is defined as a substance increasing
the concentration of the characteristic cation of the solvent. One that
increases the concentration of the characteristic anion (or decreases the
concentrqtion of the cation) is a base.
acid: anything that donates a [H+]
(proton donor)
base: anything that accepts a [H+]
(proton acceptor)
normal Bronsted-Lowery equation: acid + base <--->
acid + base
example: HNO2 + H2O <---> NO2-
+ H3O+
Each acid has a conjugate
base and each base has a conjugate
acid. These conjugate pairs only differ by a proton. In this example:
HNO2 is the acid, H2O is the base, NO2-
is the conj. base, and H3O+ is the conj. acid.
2H2O 
H3O+ + OH-
HCl is an acid
NaOH and Na2O are bases
2H3N
NH4+ + NH2-
NH4Cl or urea, H2N(CO)NH2
are acids
NaNH2 and Na2N are bases
2NO2
NO+ + NO3-
NOCl is an acid
NaNO3is a base
2H2SO4
H3SO4+ + HSO4-
CH3COOH is an actually a base which is protonated to CH3COOH2+
HAsF6 might be an acid if it does not react.
NaHSO3 is a base
5.5 Anhydrous oxides
The Lux/Flood Definition
Covers things which would become acids or bases if dissolved in water.
CO2 + CaO 
CaCO3
Here CO2 is considered the acid - carbonic acid anhydride and
CaO is considered the base since it woud give Ca(OH)2 in water.
The Lux/Flood definition defines an acid as an oxide ion acceptor
and a base as an oxide ion donor and is mainly used for high temperature
anhydrous systems for example in steel-making (in acidic or basic "slags"):
CaO + SiO2
CaSiO3
2Na2O + P2O5
2Na3PO4
Aicd Bases in Water
5.1 Proton transfer equilibria in water
We typically talk about acid-base reactions in aqueous-phase
environments -- that is, in the presence of water. The most fundamental
acid-base reaction is the dissociation
of water:
In this reaction, water breaks apart to form a hydrogen ion
(H+) and a hydroxyl ion (OH-). In pure water, we
can define a special equilibrium constant (Kw) as follows:
Where Kw is the equilibrium constant for water
(unitless)
[H+] is the molar concentration of hydrogen
[OH- is the molar concentration of hydroxide
An equilibrium constant less than one (1) suggests that
the reaction prefers to stay on the side of the reactants -- in this case,
water likes to stay as water. Because water hardly ionizes,
it is a very poor conductor of electricity.
pH
What is of interest in this reading, however, is the
acid-base nature of a substance like water. Water actually behaves both
like an acid
and a base.
The acidity or basicity of a substance is defined most typically by the
pH
value, defined as below:
At equilibrium, the concentration of H+ is
10-7, so we can calculate the pH of water at equilbrium as:
pH = -log[H+]= -log[10-7] = 7
Solutions with a pH of seven (7) are said to be neutral,
while those with pH values below seven (7) are defined as acidic and those
above pH of seven (7) as being basic.
pOH
gives us another way to measure the acidity of a solution. It is just the
opposite of pH. A high pOH means the solution is acidic while a low pOH
means the solution is basic.
pOH = -log[OH-]
pH + pOH = 14.00
Salts
A salt
is formed when an acid and a base are mixed and the acid releases H+
ions while the base releases OH- ions. This process is called
hydrolysis.
The pH of the salt depends on the strengths of the original acids and bases:
| Acid |
Base |
Salt pH |
| strong |
strong |
pH = 7 |
| weak |
strong |
pH > 7 |
| strong |
weak |
pH < 7 |
| weak |
weak |
depends on which is stronger |
This is because the conjugate base of a strong acid is very weak and
cannot undergo hydrolysis. Similarily, the conjugate acid of a strong base
is very weak and likewise does not undergo hydrolysis.
Acid-Base Character
For a molecule with a H-X bond to be an acid, the hydrogen must have a
positive oxidation
number so it can ionize to form a positive +1 ion. For instance, in
sodium hydride (NaH) the hydrogen has a -1 charge so it is not an acid
but it is actually a base. Molecules like CH4 with nonpolar
bonds also cannot be acids because the H does not ionize. Molecules with
strong bonds (large electronegativity
differences), are less likely to be strong acids because they do not ionize
very well. For a molecule with an X-O-H bond (also called an oxyacid)
to be an acid, the hydrogen must again ionize to form H+. To
be a base, the O-H must break off to form the hydroxide ion (OH-).
Both of these happen when dealing with oxyacids.
Strong Acids: These acids completely ionize in solution so they
are always represented in chemical equations in their ionized form. There
are only seven (7) strong acids:
HCl, HBr, HI, H2SO4, HNO3, HClO3,
HClO3, HClO4
To calculate a pH value, it is easiest to follow the standard "Start,
Change, Equilibrium" process.
Example Problem: Determine the pH of a 0.25 M solution of HBr.
Weak Acids: These are the most common type of acids. They
follow the equation:
HA(aq) <---> H+(aq) + A-(aq)
The equilibrium constant for the dissociation of an acid is known as
Ka. The larger the value of Ka, the stronger the
acid.
Example Problem: Determine the pH of .30 M acetic acid (HC2H3O2)
with the Ka of 1.8x10-5.
Strong Bases: Like strong acids, these bases completely
ionize in solution and are always represented in their ionized form in
chemical equations. There are only seven (7) strong bases:
LiOH, NaOH, KOH, RbOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Example Problem: Determine the pH of a 0.010 M solution of Ba(OH)2.
Weak Bases: They follow the equation:
Weak Base + H2O <---> conjugate acid + OH-
example: NH3 + H2O <---> NH4+
+ OH+
Kb is the base-dissociation constant:
Ka x Kb = Kw = 1.00x10-14
To calculate the pH of a weak base, we must follow a very similar "Start,
Change, Equilibrium" process as we did with the weak acid, however we must
add a few steps.
Example Problem: Determine the pH of 0.15 M ammonia (NH3)
with a Kb=1.8x10-5.
When dealing with weak acids and weak bases, you also might have
to deal with the "common ion effect". This is when you add a salt to a
weak acid or base which contains one of the ions present in the acid or
base. To be able to use the same process to solve for pH when this occurs,
all you need to change are your "start" numbers. Add the molarity
of the ion which comes from the salt and then solve the Ka or
Kb equation as you did earlier.
Example Problem: Find the pH of a solution formed by dissolving
0.100 mol of HC2H3O2 with a Ka of 1.8x10-8
and 0.200 mol of NaC2H3O2 in a total volume
of 1.00 L.
Acid-Base Titrations
An acid-base titration
is when you add a base to an acid until the equivalence point is reached
which is where the moles of acid equals the moles of base. For the titration
of a strong base and a strong acid, this equivalence point is reached when
the pH of the solution is seven (7) as seen on the following titration
curve:
For the titration of a strong base with a weak acid, the equivalence point
is reached when the pH is greater than seven (7). The half equivalence
point is when half of the total amount of base needed to neutralize the
acid has been added. It is at this point where the pH = pKa
of the weak acid.
In an acid-base titration, the base will react with the weak acid and form
a solution that contains the weak acid and its conjugate base until the
acid is completely gone. To solve these types of problems, we will use
the weak acid's Ka value and the molarities in a similar way
as we have before. Before demonstrating this way, let us first examine
a short cut, called the
Henderson-Hasselbalch Equation.
This can only be used when you have some acid and some conjugate base
in your solution. If you only have acid, then you must do a pure Ka
problem and if you only have base (like when the titration is complete)
then you must do a Kb problem.
Where:
pH is the log of the molar concentration of the hydrogen
pKa is the equilibrium dissociation constant for an acid
[base] is the molar concentration of a basic solution
[acid] is the molar concentration of an acidic solution
Example Problem: 25.0 mL of .400 M KOH is added to 100. mL of .150
M benzoic acid, HC7H5O2 (Ka=6.3x10-5).
Determine the pH of the solution.
This equation is used frequently when trying to find the pH of
buffer
solutions. A buffer solution is one which resists changes in pH upon
the addition of small amounts of an acid or a base. They are made up of
a conjugate acid-base pair such as HC2H3O2/C2H3O2-
or NH4+/NH3. They work because the acidic
species neutralize the OH- ions while the basic species neutralize
the H+ ions. The buffer capacity is the amount of acid or base the
buffer can neutralize before the pH begins to change to an appropriate
degree. This depends on the amount of acid or base in the buffer. High
buffering capacities come from solutions with high concentrations of the
acid and the base and where these concentrations are similar in value.
Practice weak acid problem:
C6H5COONa is a salt of a weak acid C6H5COOH.
A .10 M solution of C6H5COONa has a pH of 8.60.
-
calculate [OH-] of C6H5COONa
-
calculate K for: C6H5COO- + H2O
<---> C6H5COOH + OH-
-
calculate Ka for C6H5COOH
See the weak
acid solution.
Practice titration problem:
20.00 mL of 0.160 M HC2H3O2
(Ka=1.8x10-5) is titrated with .200 M NaOH.
-
What is the pH of the solution before the titration begins?
-
What is the pH after 8.00 mL of NaOH has been added?
-
What is the pH at the equivalence point?
-
What is the pH after 20.00 mL of NaOH has been added?
See the titration
solution.
The Lewis Definition :
Lewis acidity
5.7 Examples of Lewis acids and bases
acid: accepts an electron pair
base: donates an electron pair
The advantage of this theory is that many more reactions
can be considered acid-base reactions because they do not have to occur
in solution.
5.11 The fundamental types of reaction
The Lewis Definition considers an acid to be an electron pair acceptor
and a base as an electron pair donor. When they react, the resulting compound
is called an adduct and the bond a dative bond. It encompasses
the Brønsted/Lowry definition and many other reaction types:
H+ + :OH-
H2O
F3B + :NH3
F3BNH3
All the usual ligands can be viewed as Lewis bases and the metal ions as
lewis acids.
5.8 Boron and carbon group acids
5.9 Nitrogen and oxygen group acids
5.10 Halogen acids
Acid/Base Strength and Electonic Effects
(CH3)3N: > H3N:
> F3N:
(CH3)3B < "H3B"
< F3B
In the above series, F is the most electron withdrawing and CH3
the most electron releasing. However, there can be unexpected effects.
Consider the observed order:
BF3 < BCl3
< BBr3
The reduced strength of the fluoride compound is due to internal B-F p-bonding
which goes some way towards satisfying the electron deficiency of the boron.
A similar effect is observed in the B(OR)3 compounds (relative
to the S or Se analogues).
Systematics of Lewis acids and bases
Towards protons, pyridine is a weaker base than 2-methyl or 4 methyl
pyridine which are about equally stronger. However, towards B(CH3)3
the 4-methyl pyridine is the significantly stronger base. The methyl group
of the 2-methyl pyridine gets in the way. A similar effect is seen with
triethylamine and quinuclidine (CH(CH2CH2)3N.
Also, B(CH3)3 is a stronger acid than the
highly hindered B(C(CH3)3)3.
5.12 Hard and soft acids and bases
Hard and Soft Acids and Bases
Hard (Predominantly, having the electron concguration of inert
gases)
Acids: Li+; Na+;K+; Be2+;Mg2+;
Ca2+; Al3+; Cr3+; Fe3+;Co3+
Bases: NH3;RNH2;H2O;OH-;CO32-,
;SO42- ;F-
Boderline (Intermediate character)
Acids: Fe2+;Co2+; Ni2+;Cu2+;
Zn2+; Pb2+; Sn2+
Bases: Br-; Cl-
Soft (Electron confguration in the outer shell similar to Nio;
Pbo; Pto)
Acids: Cu+;Ag+;Cd2+;Hg2+;Au+
Bases: CN-;SCN-;RSH;RS-
Type (a) Hard Acids
-
Alkali and alkaline earth metals
-
Other lighter and highly charged cations e.g. Ti4+,
Fe3+, Co3+, Al3+
Type (b) Soft Acids
-
Heavier transition metal ions e.g. Hg22+, Hg2+,
Pt2+, Pt4+
-
Metals in lower oxidation states e.g. Ag+, Cu+ and
zero metals
Hardness is associated with low polarizability and a tendency towards
ionic bonding.
Softness is associated with greater polarizability and
more covalent bonding.
The hard and soft characters are not absolute but gradually vary. HSAB
RULES: Thermodynamic equilibrium: Hard acids pre-
fer to associate with hard bases and soft acids with soft bases. Kinetics:
Hard acids react readily with hard bases and
soft acids with soft bases. The basic principle is that like prefers
like:
Complexes of
Type (a) Metals |
Ligands |
Complexes of
Type (b) Metals |
| Strongest bonding |
R3N |
R2O |
F- |
Weakest bonding |
|
R3P |
R2S |
Cl- |
|
|
R3As |
R2Se |
Br- |
|
| Weakest bonding |
R3Sb |
R2Te |
I- |
Strongest bonding |
5.13 Thermodynamic acidity parametersAcid/Base Strength and Steric
Effects
The Drago-Wayland Equation
This empirical equation can be used to obtain a fairly quantitative value
for DHab, the enthalpy of formation
of and Lewis adduct:
-DHAB =
EA.EB + CA.CB
+ W
The factors E are a measure of the acid or base's tendency to form a bond
by electrostatic interaction which is roughly correlated with hardness.
Similarly the factors C go with the tendency to form a covalent bond, and
is correlated with softness. Both factors are considered to be constant
regardless of the partner. The factor W is almost always zero (always in
the data given!) and is a fudge factor for use if it is suspected that
and acid or a base makes a constant contribution to DHAB
which is not dependent on the partner and so should not be included in
the product terms.
Table 7.2 lists the experimental E and C parameters for a variety
of acids and bases. It must be understood that every adduct of a "new"
acid or base whose DEAB is measured
allows a new entry to be added to the table. In the beginning four parameters,
had to be given arbitarily assigned values (remember the Pauling electronegativity
scale). They are marked b
As an example, consider:
(CH3)3B + :N(CH3)3
(CH3)3B:N(CH3)3
-DE = 5.79x1.19 + 1.57x11.20 = 24.47
kcal mol-1
(CH3)3B + :P(CH3)3
(CH3)3B:P(CH3)3
-DE = 5.79x1.11 + 1.57x6.51 = 14.80
kcal mol-1
"(CH3)3Al" +
:N(CH3)3
(CH3)3Al:N(CH3)3
-DE = 17.32x1.19 + 0.94x11.20 =
31.14 kcal mol-1
"(CH3)3Al" +
:P(CH3)3
(CH3)3Al:P(CH3)3
-DE = 17.32x1.11 + 0.94x6.51 = 25.35
kcal mol-1
The amine adducts are both stronger than the phosphine adducts, but the
difference is much more marked for the harder boron. Unfortunately the
table of acids does not contain any really soft ones where the order of
strength of the dative bond might be reversed.
5.2 Solvent leveling
Leveling Effect of Protonic Solvents
In any solution, the solvent is always the substance present in the largest
amount. If the solvent contains ionizable protons it is said to be protonic,
and if it is protonic, it will engage in acid-base reactions. Many solvents
are protonic, including water, liquid ammonia, and glacial acetic acid.
A protonic solvent is often amphiprotic, like water, and thus engages in
acid-base reactions. When an acid which is stronger than the conjugate
acid of water, such as hydrochloric acid, is added to water the stronger
acid will react with water to give the weaker acid which is the aquated
proton or hydronium ion,
H3O+: HCl(g) + H2O --> H3O+(aq)
+ Cl-(aq).
All acids which are stronger than the hydronium ion will react with
the essentially limitless supply of water to quantitatively produce hydronium
ion, and so their strength will be leveled to that of the hydronium ion.
This levelling effect is the reason why, in aqueous solution, the
strongest acid which can exist is the hydronium ion. All others will be
levelled to the hydronium ion by reaction. In aqueous solution hydrochloric
acid, sulfuric acid, perchloric acid, and nitric acid are all equally strong.
As with acids, so also with bases. When a base which is stronger than
the conjugate base of water, such as sodium oxide, is added to water the
stronger base will react with water to give the weaker base which is the
hydroxide ion,
OH-: Na2O(s) + H2O --> 2OH-(aq)
+ 2Na+(aq)
The actual base here is the oxide ion since the sodium ion is an extremely
weak acid or base.
All bases which are stronger than the hydroxide ion will react with
water to quantitatively produce hydroxide ion. Although oxide ion, amide
ion, ethoxide ion, and methoxide ion are all stronger bases than is hydroxide
ion, their strength in aqueous solution is leveled to that of hydroxide
ion. Addition of sodium oxide rather than sodium hydroxide to water will
not give a more basic solution. Sodium oxide, and the other oxides of alkali
metals and alkaline earths, are called basic oxides because their aqueous
solutions are basic. They are also referred to as anhydrides of hydroxides
(from anhydrous, meaning "without water") because they act like hydroxides
from which water has been removed.
The levelling effect operates in any protonic solvent. In liquid ammonia,
for example, all acids are levelled to the strength of the ammonium ion,
NH4+, and all bases are levelled to the strength
of the amide ion, NH2-. Many of the acids which are
weak in water act as strong acids in liquid ammonia because they are stronger
than ammonium ion. On the other hand, not all of the bases which are strong
in water are also strong in liquid ammonia.
Glacial acetic acid is another protonic solvent in which the levelling
effect takes place. Glacial acetic acid can be used to show that hydrogen
chloride is a weaker acid than is perchloric acid, since hydrogen chloride
behaves as a weak acid in glacial acetic acid. Methanol is also a protonic
solvent in which some of the acids which are strong (completely dissociated)
in water are found partially in molecular form.
Periodic trends in Bronsted acidity
5.3 Periodic trends in aqua acid strength
Sulphuric acid (H2SO4)
-
One of the most important chemicals in industry. (About 40 megatons in
1981)
-
Made from sulphur by burning it to SO2 and then oxidizing the
SO2 to SO3 by either the "lead chamber" process or
the "contact" process. The older lead chamber process was a homogeneous
catalytic cycle using oxides of nitrogen as the intermediate oxidizing
agent. The concentration of sulphuric acid that could be produced was limited
to 78%. The contact process uses a V2O3 heterogeneous
catalyst. The process is sufficiently exothermic and the quantities are
such that a sulphuric acid plant can be a source of industrial or domestic
power.
-
Concentrated sulphuric acid is 18 M or 98%. Pure sulphuric acid is made
by adding the necessary amount of SO3. Addition of excess SO3
leads to "fuming suphuric acid" or "oleum" which contains polysulphuric
acids.
-
Sulphuric acid is a very effective (and dangerous) dehydrating agent: it
can produce carbon from sugar C6H12O6,
for example. Never ever dilute concentrated
sulphuric acid by adding water to it: always add the acid to water
Nitric Acid (HNO3)
-
Nitric acid is made by oxidizing ammonia with oxygen over heated Pt as
catalyst:
4NH3 + 5O2
4NO + 6H2O
6NO + 3O2
6NO2
6NO2 + 2H2O
4HNO3 + 2NO
---------------------------------------------------
4NH3 + 8O2
4HNO3 + 4H2O
-
Concentrated nitric acid is about 70% by mass in water or 16 M.
-
It is often slightly yellow due to th epresence of NO2.
-
There are autoionization equilibria in the pure colourless acid:
2HNO3
H2NO3+ + NO2-
H2NO3+
NO2+ + H2O
-
The acid below 2 M is not very oxidizing, but the concentrated acid will
dissolve almost all metals by its oxidizing action. Metals that are resistant
are Au, Pt, Rh and Ir plus Al, Fe and Cu by oxide film passivation under
certain conditions. The other product of oxidation to the metal ion is
either NO or NO2 depending on the acid concentration. Only Mg
will produce H2
Aqua Regia
This is a mixture of 3 volumes of concentrated hydrochloric acid to 1 volume
of concentrated nitric acid. It contains free Cl2 and ClNO and
has very powerful oxidizing properties as well as being a strong acid.
It is the acid of last resort to get metal into solution for analysis.
It will dissolve Au and Pt which form [AuCl4]- and
[PtCl6]2-. Mercuric sulphide will dissolve because
the sulphide is oxidized to SO2 and the mercuric ion is in the
form of [HgCl4]2-
Perchloric Acid
-
Normally comes as a 70 - 72% solution in water.
-
Pure perchloric acid is made by dehydrating the aqueous solution over Mg(ClO4)2.
It is not vey stable and decomposes to the anhydride, Cl2O7.
-
The aqueous acid and the pure compound both react explosively with organic
materials and certain inorganic compounds. Exercise
extreme caution.
-
Pure perchloric acid is a "superacid" with Ho = -13.8.
Hydrohalic Acids (except HF)
-
They are similar to oneanother except HF is anomalous.
-
They are all gases as pure substances which are very soluble in water in
which they are 100% dissociated.
-
Pure HCl (b.p. -85 oC) autoionizes to a small extent
like HF:
3HCl
H2Cl+ + HCl2
Compounds of both ions have been isolated too.
Because Br- and I- are reducing agents, the acids do not always
give the simple products expected in an acid/base reaction.
Hydrofluoric Acid
-
In aqueous solution, it is a weak acid, K = 7.2x10-5
because the H3O+ ion is stabilized by H-bonding
to the F- ions.
-
Hydrofluoric acid is difficult to work with because is attacks glass and
silica giving gaseous SiF4 or the [SiF6]2-.
It also causes burns which are very slow to heal or even get progressively
worse. Exercise extreme caution.
-
Pure HF is a very strong acid and autoionizes:
2HF
H2F+ + F-
F- + nHF
HF2-, H2F3-, H3F4-,
....
-
Very few compounds exist which are strong enough F- acceptors
to be considered acids. An example is SbF5 which leads to pure
HF having superacid properties.
-
Pure HF has a dielectric constant e/eo
= 84 and is a good solvent like water which is surprisingly gentle. It
can be used for example to remove Fe from metalloproteins without damaging
the apo-protein primary structure.
Binary Acids:
Compounds containing acidic protons bonded to a more electronegative
atom.
e.g. HF, HCl, HBr, HI, H2S
The acidity of the haloacid (HX; X = Cl, Br, I, F)
Series increase in the following order:
HF < HCl < HBr < HI
5.4 Simple oxoacids
Compounds containing acidic - OH groups in the molecule.
Acidity of H2SO4 is greater than H2SO3
because of the extra O (oxygens) attached to S pulling electrons away making
H+ to come off easier.
The order of acidity of oxyacids with a halogen (Cl, Br, or I) shows
a similar trend.
HClO4
HClO3
HClO2
HClO
Perchloric chloric
chlorus hyphochlorus
The Strengths of the Oxyacids
Consider compounds of type: XOn(OH)m e.g. PO(OH)3
Te(OH)6P
-
The successive K's differ by 10-4 to 10-5 or pKn-1
- pKn = 4.5±0.5.
-
The magnitude of K1 depends on the number of X=O groups:
| Number of X=O's |
K1 |
Acid Strength |
| 3 |
very very large |
very strong |
| 2 |
~102 |
strong |
| 1 |
~10-2 - 10-3 |
medium |
| 0 |
~10-7.5 - 10-9.5 |
weak |
-
The more X=O groups there are, the more canonical structures are available
to delocalize the charge and stabilize the anion. Perchloric acid is very
acidic and telluric acid is weak.
-
Exceptions have a deceiving formula, for example, phosphorous acid, H3PO3
is really HPO(OH)2 and H3PO2 is really
H2PO(OH) where only the OH protons are ionizable, but the acids
are stronger than one might guess (K1 = ~10-2 for
both). An the other hand, carbonic acid, (HO)2CO (K1
= 2 x 10-4) is really largely dissociated to a solution of CO2
in water and is weak by dilution, rather than inherently. Its apparent
K1 is ~10-6
5.6 Polyoxo compound information
Polyoxoanions of general composition [XaMbOc]d-
(X= P, Si; M =W, Mo) have applications in various disciplines including
medicine, catalysis, and separations science. In prospective applications
relating to nuclear waste remediation and processing, they generally function
as complexants to sequester and stabilize f ions, R. Remarkably, there
is a dearth of fundamental
information about the redox behavior and coordination of 4f (lanthanide)
and 5f (actinide) ions in polyoxoanion clusters with different X-M-O framework
structures. The interest with polyoxoanions concerns their interactions
with multivalent f elements. Two polyoxoanion systems—[P5W30O110]15- and
[P2W17O61]10-—have attracted our attention because of their remarkable
electrochemistry. Both are oxidants: multielectron reduction of the P-W-O
frameworks is reversible and nondestructive. By using synthetic methods,
spectroscopic, and electrochemical techniques, we are obtaining insights
that are leading to a predictive understanding of the mechanisms by
which the polyoxoanions and f ions are each affected by their mutual complexation.
Polyoxometalates constitute model systems for the study of the electron
and energy transfer in the infinite metaloxide
lattice and their simplicity allows to treat at the molecular scale
the coupling of electronic and nuclear movements, which is an inherent
problem for the mixed-valence systems. As is clear from such a variety
of both structure and reactivity of polyoxometalates, current research
on polyoxometalates are
1) structure/reactivity relationships with particular regard to the
mechanism of electron transfer reactions, 2) magnetic interaction and molecular
magnetic device, 3) energy-transfer mechanism and luminescence device (including
nonlinear optical
device), 4) encapsulation of templates in the photo-induced self-assembly
process, 5) template-exchange reaction
and topology, and 6) antibacterial effects on MRSA and VRE.
5.14 Solvents as acids and bases
Solvent Properties
Their usefulness is a function of:
-
Liquid temperature range
-
Dielectric constant
-
Its Lewis acid/base properties
-
Its Brønsted/Lowry acid base properties
-
Autodissociations properties
Liquid Range
The most useful are liquid at room temperature but we often need solvents
to go upto high or low temperatures. The range is large for DMF and propane-1,2-diol
carbonate. Ammonia is a good solvent for many reactions but it boils at
about -33oC. Hydrogen fluoride, which boils at about 20 oC
is also useful, but very corrosive.
Dielectric Constant
This factor is important if the sovent is to dissolve ionic compounds.
Water happens to have a very high dielectric constant (82) and is probably
the best all-round solvent for ionic compounds that we have. The key formula
here is:
F = q+q-/4per2
Lewis Acid/Base Properties
The extent to which the sovent can act as a ligand towards the ions of
an ionic solute, with the dielectric constant which is related but not
directly, determines how good a solvent it will be. In the crystal lattice
with simple ions, the anions effectively complex the cations and vice versa.
The solvent will have to do a better job.
Generally, because the cations are smaller than the anions, the
greatest gains are to be made by effective complexation of the cations,
so it is the Lewis basicity of the solvent that is more important. For
common solvents:
(CH3)2SO > H(CO)N(CH3)2»
H2O > (CH3)2CO >
(CH3CHCH2)O2CO > (CH3)2SO2
> CH3NO2 > C6H5NO2
» CH2Cl2
Protic Solvents
These solvents contain protons which can be inozed, in other words the
solvent is a Brønsted acid. Examples are H2O, HF, H2SO4,
HCN and even MH3.
Autodissociation
Also known as autoionization, so examples in protic solvents are:
2H2O
H3O+ + OH-
3HF
H2F+ + HF2-
2H2SO4
H3SO4+ + HSO4-
2H2O
H3O+ + OH-
2H3N
H4N+ + NH2-
The cations and anions produced can interact with the solutes helping to
explain certain reaction products. Also, the autodissociation reactions
are not really this simple - consider water: The more general equation
is:
(n+m+1)H2O
[H(H2O)n]+ + [OH(H2O)m]-
The ions [H3O]+, [H(H2O)2]+
and [H(H2O)3]+ have been characterized,
the first two in crystal structures.
K'25oC = [H+][OH-]/[H2O]
= (1.0x10-14)/55.56 and K25oC = [H+][OH-]
= 1.0x10-14
Aprotic Solvents
Such solvents fall into three broad classes:
-
Non-polar, or weakly polar solvents which do not dissociate and are not
strongly coordinating. Examples are hydrocarbons and CCl4. They
are poor solvents for everything except like substances, i.e. other nonpolar
molecules.
-
Strongly solvating (usually polar) solvents which do not dissociate. Examples
are acetonitrile ( CH3CN), N,N-dimethyformamide (DMF, H(CO)N(CH3)2),
dimethyl sulphoxide (DMSO (CH3)2SO), tetrahydrofuran
(THF C4H8O) and liquid sulphur dioxide (SO2).
They are similar in being very strongly coordinating towards
cations although SO2 forms an adduct with acetonitrile. The
dielectric constants can very quite a bit (DMSO = 45, THF = 7.6) which
will govern their ability to dissolve ionic materialas.
-
Highly polar and autoionizing solvents. Examples are:
2BrF3
BRF2+ + BrF4-
2NO2
NO+ + NO3-
2Cl3PO
Cl2PO+ + Cl4PO-
Molten Salts
These are the extreme case of autoionizing solvents. Generally they will
be very high temperature systems but lower temperatures can be attained
by the right eutectic mixtures, e.g.
LiNO3/NaNO2/KNO3 mixtures
can reach as low as 130 oC
(C2H5)2NH2Cl melts
at 215 oC
[N,N-RR'N2C3H3]+Cl-
with AlCl3 can be liquid at room temperature.
Uses:
Aluminum is made by electrolyis of Al2O3
in a cryolite {Na3AlF6]
Re3Cl9 + Et2NH2Cl
[Et2NH2][Re2Cl8] with its quadruple
bond.
Solvents for Electrochemistry
Solvent which are useful for electrochemistry must have two characteristics:
a highish dielectric constant so that they are good solvents for ionic
compounds, and they must be redox resistant. Water is actually not ideal
because, although its dielectric constant is very high, it is susceptible
to both oxidation and reduction at relatively small potentials:
H2O
O2 + 4H+(10-7
M) + 4e-
E = -0.82 V
i.e. ½X2 + e-
X- E > +0.82 V might generate
oxygen from water
H+(10-7 M) + e-
½H2 E = -0.41 V
i.e. M
M+ + e-
E > +0.41 V might generate H2 from water
Acetonitrile or DMF are often used as an electrochemistry solvent for example
in cyclic voltammetry for organometallic substances usually with a redox
inert supporting electrolyte such as t-butylammonium perchlorate.
Purity of Solvents
Water and oxygen are th most common contaminants and they can be very difficult
to remove. Many research labs have several continuously operating stills
from which the solvents can be removed under cover of nitrogen or argon.
Hydrocarbon solvents and ethers can be distilled over sodium (or potassium)
in the presence of benzoquinone as an intense blue indicator (of dryness).
Various hydrides can also be used.
More reactive solvents can be dried effectively over molcular
sieves and vacuum distilled to remove dissolved oxygen.
It is important to remember that even traces of contaminents will
mess up sensitive reactions because the solvent is present in such excess.
Heterogeneous acid-base reactions
Heterogeneous equilibria.
Heterogeneous equilibria have more than one phase present in the reaction.
We have discussed how to write equilibria expressions for solutions (use
K ) and gases (use K ). What about:
CaCO3(s)---> CaO(s) + CO2(g)
Recall that in equilibria expressons, [A] actually means the ratio
of the concentration in M to the concentration of the standard state. The
standard state concentrations are (a) that of the pure solid, (b) that
of a pure liquid. Hence, [A] for a pure solid or a pure liquid is 1 by
definition.
Consequently, terms for pure solids and pure liquids do NOT appear
in equilibria expressions.
For CaCO3(s)---> CaO(s) + CO2(g), Kp
= [CO ],
Acid-base catalysis is of great importance in large scale chemicals
processing, e.g. petroleum refining. Mineral acids are often used in these
industrial processes; however, these liquids are corrosive, dangerous,
and difficult to dispose of. Efforts have been made to replace liquid acids
with solid catalysts such as acidic resins, supported Lewis-acids, zeolites,
phosphates, and modified metal oxide based systems; demanding that the
acidity of these materials be characterized. The numerous concepts which
have been developed to describe acidity (e.g. Brønsted, Lewis, Pearson,
Arrhenius, Bjerrum), very often for selected materials/systems, reflect
the difficulty in finding a general approach to the subject and a widely
applicable technique of measurement. Measurements of acidity are usually
indirect, i.e. they evaluate the interaction of an acid with a "probe".
The pH scale is a precise definition and measure of Brønsted acidity
but it is actually a very specific case: the interaction of a proton donor
type acid with the probe molecule water in the solvent water. The equilibrium
acid/probe is influenced by the proton donor capability of the acid, the
proton acceptor capability of the probe, and the solvation of all species
involved.
Transferring these ideas to solid surfaces with Brønsted sites
it becomes evident that again proton donor and acceptor
capabilities of acid site and probe molecule are important but also
the coordination of the probe molecule – particularly a large
molecule – to atoms surrounding the proton, thus the geometry of the
entire "site". With this background, it becomes
understandable that acidity scales of solid surfaces can successfully
be established when a "family" of samples is investigated,
i.e. structurally related compounds (e.g. zeolites with different Si/Al
ratios), while comparing results from different probes or
different materials often reveals inconsistencies. This also explains
why acidity measurements may not help to predict reactivity,
i.e. if the reactant and the probe are chemically non related. Understanding
the reaction mechanism and the role of acid sites in
the initiation or the propagation of a reaction, often a chain reaction
in acid catalysis, is essential to allow a correlation of
acid-base-properties and reactivity.
Super Acids
A superacid is defined as any acid stronger than 100% sulfuric acid.
In organic chemistry, superacids have been used extensively to stabilize
carbocations (R3C+, protonated arenes, etc.).
These cations are typically characterized by NMR spectroscopy at low temperatures.
In inorganic chemistry, superacids have been used to observe a wide variety
of reactive cations (S82+, H3O2+,
Xe2+, HCO+, etc.), some of which have
been isolated for structural characterization. Nevertheless, many
compounds decompose in typical superacid media and desired reactive cations
cannot be observed (e.g. R3Si+, C60+,
HC60+, P4+ etc.). This
identifies two of the major limitations of presently known superacids:
the nucleophilicity of their anions and the oxidizing power of the added
Lewis acid. For example, the silylium ion (R3Si+)
cannot be closely approached in a typical superacid such as HF/SbF5
because F- and SbF6- anions coordinate
to silicon. C60+ is unstable in HF/SbF5
because of oxidative destruction. Even the supposedly non-oxidizing
triflic acid, CF3SO3H, is reported to decompose
C60. In addition, superacids are typically viscous, corrosive
liquids -- difficult to handle stoichiometrically and limited in their
commercial applications.
The strongest
acid in aqueous solution is H3O+ since all the strong acids
protonate water to 100% of their capacity. In non-aqueous solution it is
possible to generate more powerful proton donors.
The strength of such strong acids is measured by the Hammett acidity
function:
Ho = pKBH+ - log([BH+]/[B])
Here B and BH+ usually refer to a colorimetric indicator, that
is either B or (more likely) BH+ is strongly coloured, and pKAB+
is small or negative. Typical indicators are:
| Indicator |
pKAB+
(in H2SO4) |
| m-nitroaniline |
+2.5* |
| p-nitroaniline |
+0.99* |
| o-nitroaniline |
-0.29 |
| 2,4-dinitroaniline |
-4.53 |
| 3-methyl-2,4,6-trinitroaniline |
-8.22 |
| 2,4,6-trinitroaniline |
-10.10 |
*The first two provide overlap with the normal range of the pH
scale (~0 - 14).
Notice that the Hammett scale is actually equivalent to the pH scale
in the case of aqueous solutions:
A table of selected Ho values is given below. In each
case, the equilibrium leading to the proton donor species is shown.
Ho = -log([B][H+]/[BH+] - log([BH+]/[B])
= -log[H+]
| Acid |
Ho |
| 2H2SO4
H3SO4+ + HSO4- |
-12 |
| H2SO4 + H2S2O7H3SO4+
+ HS2O7- (in oleum) |
-15 |
| 3HF
H2F+ + HF2- |
-11 |
| 2HF + SbF5
H2F+ + HSBF6 |
-12 |
| 2HSO3F
H2SO3F+ + SO3F- |
-15 |
| 2HSO3F + SbF5H2SO3F+
+ SbF5SO3F- |
-19 |
The super acids can be used in a variety of reactions leading to unusual
cations:
(CH3)3COH + super
acids
(CH3)3C+ + H2O
I2 + super acids
I2+ or I3+