Transcript Water as a Brønsted acid or base

Chapter 7
Acids, bases and ions in aqueous
solution
TOPICS
 Properties of water
 Molarity, molality, standard state and activity
 Brønsted acids and bases
 Energetics of acid dissociation
 Aquated cations
 Amphoteric behaviour
 Coordination complexes: an introduction
 Solubility product constants
 Solubilities of ionic salts
 Common-ion effect
 Formation of coordination complexes
 Stability constants
7.1 Introduction Liquid water is approximately 55 molar H2O
7.2 Properties of water
Structure and hydrogen bonding
Fig. 7.2 The variation in the value
of the density of water between
283 and 373 K.
Fig. 7.1 Part of the structure of ordinary
ice; it consists of a 3-dimensional
network of hydrogen-bonded H2O
molecules.
The self-ionization of water
If a pure liquid partially dissociates into ions, it is self-ionizing
Water as a Brønsted acid or base
A Brønsted acid can act as a proton donor, and a Brønsted base can
function as a proton acceptor.
Worked example 7.2 Manipulating equilibrium constant data
Molarity: 1 M or 1 mol dm-3 contains 1 mol of solute dissolved in
sufficient volume of water to give 1 L (1 dm3) of solution.
Molality:1 mol of solute dissolved in 1 kg of water it is one molal
(1 mol kg-1).
Standard State :T = 298 K, 1 bar pressure (1 bar = 1.00 x105Pa)
Activity: When concentration greater than 0.1 M, interactions
between solute are significant.
7.4 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.
Organic Acids, Carboxylic acids: examples of mono-, di- and polybasic
acids
Inorganic acids
HCl, HBr, HI have negative pKa, where as HF is a weak acid
(pKa = 3.45)
Example of oxoacids include HOCl, HClO4, HNO3, H2SO4,H3PO4.
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.’
Note that:
. oxoacids may be mono-, di- or polybasic;
. not all the hydrogen atoms in an oxoacid
are necessarily ionizable.
Phosphinic acid
monobasic
Inorganic bases: hydroxides
NaOH, KOH, RbOH, CsOH are strong bases (LiOH is a weaker base)
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
7.5 The energetics of acid dissociation in aqueous solution
Hydrogen halides
Fig. 7.3 The energetics of the dissociation of a hydrogen halide, HX (X=F, Cl,
Br or I), in aqueous solution can be considered in terms of a cycle of steps. The
significance of each step is discussed in the text.
Each reaction is
exothermic, with DHo
values in the order
HF < HCl < HBr ~ HI.
H2S, H2Se and H2Te
7.6 Trends within a series of oxoacids EOn(OH)m
Bell’s rule Which relates the first acid dissociation constant to the number of
‘hydrogen-free’ O atoms in an acid of formula EOn(OH)m
7.7 Aquated cations: formation and acidic properties
Water as a Lewis base
A Lewis acid is an electron acceptor, and a Lewis base is an electron donor.
Hydration is the specific case of solvation when the solvent is water.
Fig. 7.5 (a) The first hydration shell of an Na+ ion; ion–dipole interactions operate
between the Na+ ion and the H2O molecules.
(b) If the metal–oxygen bond possesses significant covalent character, the first hydration
shell can be reasonably represented showing oxygen-to-metal ion coordinate bonds;
however, there is also an ionic contribution to the bonding interaction.
In practice, the character of the metal …. oxygen interaction varies with
the nature of the metal ion.
The configurations 7.6 and 7.7 have
been established in the first hydration
shell for dilute solutions of LiCl and NaCl
by detailed neutron diffraction studies. In
concentrated solutions, the plane of the
water molecule in 7.6 makes an angle of
up to 50o with the M+…O axis (Figure
7.6) implying interaction of the cation
with a lone pair of electrons rather
than an ion–dipole interaction.
Fig. 7.6 If the plane of each
water molecule in [M(OH2)6]+
makes an angle of ~ 50o with
the M+….O axis, it suggests that
the metal–oxygen interaction
involves the use of an oxygen
lone pair.
Aquated cations as Brønsted acids
In the aqueous chemistry of cations, hydrolysis refers to thereversible loss of
H+ from an aqua species. The term hydrolysis is, however, also used in a wider
context, e.g. the reaction:
PCl3 + 3H2O  H3PO4 + 3HCl
is a hydrolysis process.
Aquated cations can act as Brønsted acids by loss of H+ from
a coordinated water molecule
The position of the equilibrium (and thus, the strength of the acid) depends on
the degree to which the OH bonds are polarized , and this is affected by the
charge density of the cation.
Small cations such as Li+, Mg2+, Al3+, Fe3+, and Ti3+ possess high charge
densities.
7.8 Amphoteric oxides and hydroxides
Amphoteric behaviour
If an oxide or hydroxide is able to act as either an acid or a base, it is said
to be amphoteric .
The hexaaqua ion, 7.10 , may be isolated as, for example, the sulfate salt after
reaction with H2SO4. The ion [Al(OH)4], 7.11, can be isolated as, for example,
the Na+ salt if the source of hydroxide is NaOH.
Similarly, aluminium hydroxide is amphoteric (equations 7.41 and 7.42).
Periodic trends in amphoteric properties
The character of the oxides of the elements across a row of the periodic
table ( s- and p-blocks) changes from basic to acidic, consistent with
a change from metallic to non-metallic character of the element.
Elements that lie close to the so-called ‘diagonal line’ (Figure 7.8)
possess amphoteric oxides and hydroxides.
In group 2, Be(OH)2 and BeO are amphoteric but M(OH)2 and MO
(M=Mg, Ca, Sr or Ba) are basic.
Among the oxides of the p -block, Al2O3, Ga2O3, In2O3, GeO, GeO2,
SnO, SnO2, PbO, PbO2, As2O3, Sb2O3 and Bi2O3 are amphoteric.
Fig. 7.8 The so-called ‘diagonal line’ divides metals from non-metals, although
some elements that lie next to the line (e.g. Si) are semi-metals.
7.9 Solubilities of ionic salts
Solubility and saturated solutions
The solubility of a solid at a specified temperature is the amount of solid
( solute ) that dissolves in a specified amount of solvent when
equilibrium is reached in the presence of excess solid.
The solubility may be expressed in several ways, for example:
•. mass of solute in a given mass of solvent (g of solute per 100 g of water);
•. moles of solute in a given mass of solvent;
•. concentration (mol dm3);
•. molality (mol kg 1);
•. mole fraction.
Fig. 7.9 The temperature-dependence of
the solubilities in water of potassium
iodide and sodium nitrate. The solubility
of sodium chloride is essentially
temperature independent in the range
273–373 K.
Tabulated values of solubilities of ionic salts refer to the maximum amount of
solid that will dissolve in a given mass of water to give a saturated solution.
Solubilities may also be expressed in concentrations, molalities or mole
fractions.
Sparingly soluble salts and solubility products
The solubility product,or
solubility constant, Ksp
The energetics of the dissolution of an ionic salt: DsolGo
where F = Faraday constant = 96 485 C mol1
In solubility
7.10 Common-ion effect
If a salt MX is added to an aqueous solution containing the solute MY (the
ion Mn+ is common to both salts), the presence of the dissolved Mn+ ions
suppresses the dissolution of MX compared with that in pure water; this is
the common-ion effect.
7.11 Coordination complexes: an introduction
In a coordination complex , a central atom or ion is coordinated by one or
more molecules or ions ( ligands ) which act as Lewis bases, forming
coordinate bonds with the central atom or ion; the latter acts as a Lewis
acid. Atoms in the ligands that are directly bonded to the central atom or
ion are donor atoms.
In a complex:
. a line is used to denote the interaction between an anionic ligand and the
acceptor;
. anarrow is used to show the donation of an electron pair from a neutral
ligand to an acceptor.
When a Lewis base donates a pair of electrons to
a Lewis acid, a coordinate bond is formed and the
resulting species is an adduct. The centred dot in,
for example, H3B THF indicates the formation of an
adduct.
Fig. 7.10 (a) The structure of pentane-2,4-dione (acetylacetone), Hacac (see Table 7.7 ); (b) Fe(III)
forms an octahedral complex with [acac] ; (c) the structure of the coordination complex [Fe(acac)3],
determined by X-ray diffraction
7.12 Stability constants of coordination complexes
In the formation of a complex [ML6]z+ from [M(OH2)6] z+ , each displacement of
a coordinated water molecule by ligand L has a characteristic stepwise stability
constant , k1, k2, k3, k4, k5, or k6
Thermodynamic considerations of complex formation: an introduction
For a given metal ion, the thermodynamic stability of a chelated complex
involving bidentate or polydentate ligands is greater than that of a complex
containing a corresponding number of comparable monodentate ligands. This
is called the chelate effect .
7.13 Factors affecting the stabilities of complexes containing only
monodentate ligands
Stabilities of complexes of non d-block metal ions of a given charge
normally decrease with increasing cation size: Ca2+ < Sr2+ < Ba2+
For ions of a similar size, the stability of a complex with a specific
ligand increases substantially as the ionic charge increases:
Li+ < Mg2+ < Al3+
Hard and soft metal centres and ligands
Pearson’s classification of hard and soft acids comes from a consideration
of a series of donor atoms placed in order of electronegativity:
Hard acids (hard metal cations) form more stable complexes with hard bases
(hard ligands) while soft acids (soft metal cations) show a preference for soft
bases (soft ligands)
In aqueous solution, complexes formed between class (a), or hard,
metal ions and ligands containing particular donor atoms exhibit trends
in stabilities as follows:
In contrast, trends in stabilities for complexes formed between class (b),
or soft, metal ions and ligands containing these donor atoms are:
Complex formation usually involves ligand substitution. In aqueous
solution, for example, ligands displace H2O and this is a competitive
rather than simple combination reaction
Suppose M2+ is a hard acid. It is already associated with hard H2O
ligands, i.e. there is a favourable hard–hard interaction. If L is a soft
base, ligand substitution will not be favourable. If L is a hard base, there
are several competing interactions to consider:
Overall, it is observed that such reactions lead to only moderately stable
complexes, and values of DHo for complex formation are close to zero.
Now consider the case where M2+ in equation
is a soft acid and L is a soft base. The competing interactions will be: