Download Acids, bases and ions in aqueous solution

Chapter 5:
Acids, bases and ions in aqueous solution
5.1 Properties of water
Selected physical properties of water.
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Structure and hydrogen bonding
At atmospheric pressure, solid H2O can adopt one of two polymorphs, depending
upon the conditions of crystallization.
At higher pressures, five polymorphs exist which differ in their arrangement of the
oxygen atoms in the crystal lattice.
Part of the structure of ordinary ice; it
consists of a 3-dimensional network of
hydrogen-bonded H2O molecules
The hydrogen-bonded network may be
described in terms of a wurtzite lattice in
which the O atoms occupy the sites of
both the Zn and S centres; this places
each O atom in a tetrahedral environment
with respect to other O atoms
Each H atom lies slightly off the O……O
line, so that the intramolecular H–O–H
bond angle is 105°.
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The wurtzite lattice is very open, and
as a result, ice has a relatively low density
(0.92 g cm-3).
On melting (273 K), the lattice
partly collapses, allowing some of
the lattice cavities to be occupied
by H2O molecules. Consequently,
the density increases, reaching a
maximum at 277 K; between 277
and 373 K, thermal expansion is
the dominant effect, causing the
The variation in the value of the density of water
density to decrease.
between 283 and 373 K.
Note: Even at the boiling point (373 K), much of the hydrogen bonding
remains and is responsible for water having high values of the
enthalpy and entropy of vaporization
Molarity of water Show that pure water is approximately 55 molar.
Solution: Density of water =1 g cm-3
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Thus, 1000 cm3 (or 1dm3) has a mass of 1000 g
For H2O, M= 18 g.mol-1
1000
Number of moles in 1000 g =
= 55.5 mol.L-1
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Therefore, the concentration of pure water 55 mol.dm-3
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The self-ionization of water
If a pure liquid partially dissociates into ions, it is selfionizing.
Water itself is ionized to a very small extent (equation below) and the value of the
self-ionization constant, Kw , shows that the equilibrium lies well to the left-hand side
The self-ionization in equation 6.1 is also called autoprotolysis.
[H3O]+
2H2O(l)
Water
Kw = [H3
Oxonium ion
O+]
[OH-]
= 1.00
x10-14
+
[OH]Hydroxide ion
(at 298K)
Note:
In aqueous solution, protons are solvated and so it is more correct to write
[H3O]+(aq) than H+(aq). Even this is oversimplified because the oxonium ion is
further hydrated and species such as [H5O2]+, [H7O3]+ and [H9O4]+ are also present.
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Water as a Brønsted acid or base
Acids and Bases (definitions)
1. Arrhenius Acids and Bases
Acids are H+ donors
Bases are OH- donors
2. Arrhenius Broadened Definition
Acids increase H+ concentration or [H+] increases
Bases increase OH- concentration or [OH-] increases
3. Brønsted-Lowry Acids and Bases (1923)
Acids donate H+
Bases accept H+
Brønsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a substance that can donate a hydrogen ion H+, proton).
A Brønsted-Lowry base is a substance that can accept a hydrogen ion
Acids and bases occur as conjugate acid - base pairs.
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Water as a Brønsted acid or base
H3O+(aq)
2H2O(l)
+
OH-(aq)
The equation of self-ionization of water illustrates that water can function as
both a Brønsted acid and aBrønsted base.
In the presence of other Brønsted acids or bases, the role of water depends
on the relative strengths of the various species in solution.
In the presence of a Brønsted acid
HCl(aq)
H3O+(aq)
H2O(l)
+
Cl-(aq)
•
Hydrogen chloride is a much stronger acid than water. This means that:
HCl will donate a proton to H2O and the equilibrium lies well over to the
right-hand side, so much so that hydrochloric acid is regarded as being fully
dissociated, it is a strong acid.
•
Water accepts a proton to form H3O +, and thus behaves as a Brønsted base.
+
In the reverse direction, H3O+ acts as a weak acid and Cl- as a weak base; they
are, respectively, the conjugate acid and conjugate base of H2O and HCl.
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In the presence of a Brønsted base
NH3 (aq)
+
NH4+(aq)
H2O(l)
+
OH-(aq)
• In an aqueous solution of NH3, water behaves as a Brønsted acid, donating H+.
• NH4+ is the conjugate acid of NH3, while H2O is the conjugate acid of OH-.
• Conversely, NH3 is the conjugate base of NH4+, and OH- is the conjugate base of H2O.
• the value of K for equilibrium shows that NH3 acts as a weak base in aqueous solution
K=
[NH ][OH ] = 1.8 ×10
+
4
−
[NH 3 ]
−5
(at 298K)
Note: If an oxide or hydroxide is able to act as either an acid or a base,
it is said to be amphoteric.
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The equilibrium constants Ka, Kb and Kw
In dealing with acid–base equilibria in aqueous solution, three
equilibrium constants are of special significance:
• Ka is the acid dissociation constant.
• Kb is the base dissociation constant.
• Kw is the self-ionization constant of water.
For a general weak acid HA in aqueous solution:
H3O+(aq) + A-(aq)
HA(aq) + H2O(l)
[H O ][A ] = [H O ][A ]
+
−
3
Ka =
+
−
3
[HA][H 2O]
[HA]
By convenon, │H2O│ = 1; strictly, the activity of the solvent H2O is 1.
For a general weak base B in aqueous solution:
BH+(aq) + OH-(aq)
B(aq) + H2O(l)
[BH ][OH ] = [BH ][OH ]
+
Kb =
9
−
[B][H 2O]
pKa = - logKa
+
−
[B]
Ka = 10-pKa
pKb = - logKb
Kb = 10-pKb
Kw = [H3O+].[OH-] = 1.00 x 10-14
pKw = - logKw = 14:00
Kw = Ka . Kb
pH = - log [H3O+]
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5.2 Definitions and units in aqueous solution
Molarity and molality
A one molar aqueous
: Contains one mole of solute dissolved in a sufficient
solution (1M or 1 mol.dm-3)
volume of water to give 1dm3 (1L) of solution.
In contrast
A one molal aqueous : one mole of solute is dissolved in 1 kg of water,
solution ( 1 mol.kg-1)
the solution is said to be one molal (1 mol kg1).
Standard state
The standard state of a liquid or solid substance, whether pure or in a mixture,
or for a solvent is taken as the state of the pure substance at 298K and 1 bar
pressure (1 bar = 1.00 x 105 Pa)
The standard state of a gas is that of the pure gas at 298 K, 1 bar pressure and
exhibiting ideal gas behaviour.
For a solute in a solution, the definition of its standard state is referred to a
situation of infinite dilution: it is the state (a hypothetical one) at standard
molality (mo), 1bar pressure, and exhibiting infinitely diluted solution
behaviour.
In the standard state, interactions between solute molecules or ions are
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negligible.
Activity
When the concentration of a solute is greater than about 0.1 mol L-1. It becomes
necessary to define a new quantity called the activity, which is a measure of
concentration but takes into account the interactions between the solution species.
The relative activity, ai , of a component i is dimensionless and is defined by
the equation:
µi = µi0 + RTlnai
where µi : the chemical potential of component i
µi 0: is the standard chemical potential of i
R : the molar gas constant
T : the temperature in kelvin
The activity of any pure substance in its standard state is defined to be unity.
The relative activity of a solute is related to its molality by the
equation:
γm
ai = i 0 i
where γi : the activity coefficient of the solute
mi
mi : the molality
moi : the standard state molality,
Since moi is defined as being unity:
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ai = γ i mi
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5.3 Some Brønsted acids and bases
The larger the value of Ka, the stronger the acid.
The smaller the value of pKa, the stronger the acid.
The larger the value of Kb, the stronger the base.
The smaller the value of pKb, the stronger the base.
Carboxylic acids: examples of mono-, di- and polybasic acids
In organic compounds, acidity is quite often associated with
the presence of a carboxylic acid group (CO2H) and it is
relatively easy to determine the number of ionizable
hydrogen atoms in the system.
Acetic acid
is a monobasic acid since it can donate only one proton
H3O+(aq) + MeCOO-(aq)
CH3COOH(aq) + H2O(l)
[H O ][MeCOO ]
+
Ka =
3
[MeCOOH]
−
= 1.75 x 10-5
at 298 K
In aqueous solution; it is a weak acid
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Ethanedioic acid (oxalic acid)
can donate two protons and so is a dibasic acid.
oxalic acid undergo stepwise dissociation in aqueous solution, the
following equations describe the steps :
Each dissociation step has an associated equilibrium constant (acid dissociation
constant), and it is general for polybasic acids that Ka(1) > Ka(2) and so on; it is
more difficult to remove H+ from an anion than from a neutral species.
Values of equilibrium constants may be temperaturedependent, and the
inclusion of the temperature to which the stated value applies is important.
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H4EDTA
The tetrabasic acid, and the anions derived from
it are commonly encountered in coordination
chemistry; the trivial name for this acid is
N,N,N’,N’-ethylenediaminetetraacetic acid and is
generally abbreviated to H4EDTA.
Inorganic acids: examples of mono-, di- and polybasic acids
hydrogen halides
In inorganic chemistry, hydrogen halides and oxoacids are of particular
significance in terms of acidic behaviour in aqueous solution.
Each of the hydrogen halides is monobasic:
for X =Cl, Br and I, the equilibrium lies far to the right-hand side, making
these strong acids ( Ka > 1 ↔ pKa = - log(Ka) < 0 ).
pKa(HCl) ≈ -7
pKa(HBr) ≈ -9
pKa(HI) ≈ -11
Hydrogen fluoride on the other hand is a weak acid:
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pKa(HF) = 3.45
oxoacids
The IUPAC definition of an oxoacid is ‘a compound which contains
oxygen, at least one other element, at least one hydrogen bound to
oxygen, and which produces a conjugate base by proton loss.
oxoacids may be mono-, di- or polybasic
not all the hydrogen atoms in an oxoacid are necessarily ionizable
Nitric acid : strong acid
pKa = -1.64
Nitrous acid : weak acid
pKa = 3.37 (285K)
Hypochloric acid : weak acid
pKa = 4.53
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Sulfuric acid : dibasic acid
in aqueous solution, the first dissociation step lies well over to the right-hand side:
pKa ≈ -2
but [HSO4]- is a weaker acid.
pKa = 1.92
Two series of salts can be isolated, e.g. sodium hydrogensulfate (NaHSO4)
and sodium sulfate (Na2SO4).
NOTE : In the oxoacids above, each hydrogen atom is attached to oxygen in the free
acid, and the number of H atoms corresponds to the basicity of the acid.
However, this is not always the case:
e.g. although phosphinic acid has the formula H3PO2, there is only one OH
bond and H3PO2 is monobasic :
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Inorganic bases:
hydroxides
The group 1 hydroxides NaOH, KOH, RbOH and CsOH are strong bases, being
essentially fully ionized in aqueous solution
LiOH is weaker (pKb = 0.2)
Inorganic bases:
nitrogen bases
The term ‘nitrogen bases’ tends to suggest ammonia and organic amines (RNH2),
but there are a number of important inorganic nitrogen bases related to NH3
Ammonia dissolves in water, and functions as a weak base, accepting
H+ to form the ammonium ion :
pKb = 4.75
Thus, a value of pKa for ‘ammonia’ of 9.25 is really that of the ammonium
ion and refers to equilibrium
pKa = 9.25
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