Physical properties of water and their importance
Property
Consequence
Excellent solvent
Transport of nutrients and waste
products, prerequisite of
biogeochemical processes
High dielectric constant
Solubility of ionic compounds
High surface tension
Physiological control factor;
droplets and surfaces
Transparent for visible and
(partially) for UV radiation
Allows photosynthesis in aqueous
media
Highest density in liquid
state at 4 °C
Floating ice, stratification, isolation
of water biota from freezing
High heat of vaporization
Controls the transfer of vapor
between atmosphere and water
High heat of melting
Stabilization of temperature regime
at freezing/melting
High heat capacity
Stabilization of temperature
Hydrogen bonds
Anomalous water properties: boiling point
Boiling points of structurally similar compound from the 4.7. period
Anomalous water properties: density
Density maximum
40C
Consequence: density of ice
is lower than density of liquid water
Solubilty of liquids and solids
Water as a solvent
Water is the most common polar solvent. Some
solutes remain in aqueous solution in molecular
form, other – electrolytes – dissociate to ions.
Ionic crystals are usually well soluble (i.e.
solubility at least 0.1-1 mol/l). Solubility of salts
generally increases with temperature, in contrary
to gas solubility.
Some rules for solubility of solids with ionic structure
• Most sodium, potassium and ammonium salts are well soluble.
Exception is KClO4, which is often used for precipitation of potassium
ion from aqueous solutions.
• Nitrates are usually well soluble.
• Carbonates and phosphates are usually insoluble or sparingly
soluble, exceptions are sodium, potassium and ammonium salts.
Potassium-magnesium phosphate is used for precipitation of
magnesium ion from aqueous solutions.
• Halides are usually well soluble, exceptions are silver, lead and
mercury (I) halides. PbCl2 is sparingly soluble, silver and mercury (I)
chlorides are essentially insoluble.
• Sulfates are usually well soluble, exceptions are calcium, barium
strontium, lead and mercury (I) sulfates. Silver sulfate is sparingly
soluble.
• Sulfides are usually insoluble in water.
Solubility of nonelectrolytes
Solubility in the form of
molar concentration in
aqueous solution can be
estimated also from Henry’s
law constant and vapor
pressure.
Dissolution as a chemical reaction
Dissolution can be described in terms of chemical
reaction, e.g. for gas in water
A( g )  A(aq)
Thermodynamic relations derived for chemical reactions
can be applied to this process, e.g. the equilibrium
constant
n
K   aiν i
i 1
K is the equilibrium constant of the reaction ai is the
equilibrium activity of i compound, νi is the stoichiometric
coefficient of i compound
Activity and standard states
Activity is defined as the ratio of actual fugacity of a
compound to its fugacity in a standard state. Standard
states are chosen differently for compounds in different
phases. E.g. for gases the standard state is ideal gas at
standard pressure p° = 101325 Pa. Corresponding
activity is
fi
pi
ai    
fi
p
Standard states II
Standard state for (aqueous) solutions is solution at unit
concentration:
fi
ci
ai      ci
fi
ci
Standard state for pure solid or liquid compounds is
chosen as pure solid or liquid, leading to unit activity at all
conditions. The same standard state is used for solvents
in solutions.
Equilibrium in dissolution reactions
Equilibrium constant for dissolution of A gas in water is:
c( Aaq ) p 
a( Aaq )
p
Kh 


a( Ag )
pi ( Ag )
H
Henry‘s law constant is apparently a certain form of equilibrium
constant.
Solubility of solid ionic compound that is (partially) dissolved in water is
described by the ion product:


AB( s)  A (aq)  B (aq)
Ks 


a( Aaq
)a( Baq
)
a ( ABs )


 c( Aaq
)c( Baq
)
Dissolution of minerals - examples
• Calculate molar solubility of AgCl in water – dissolution reaction is
AgCl(s) --> Ag+ + Cl-.
From strochiometric ballance [Ag+] = [Cl-]. Ks = 1.76 x 10-10 = [Ag+][Cl-]
= [Ag+]2, [Ag+] = 1.33 x 10-5 and molar solubility of AgCl is 1.33 x 10-5
mol/l.
• Concentration of Ca2+(aq) equal to 3.32 x 10-4 mol/l was obtained
from analysis of water in contact with fluorite (CaF2). Calculate the
ion product of CaF2 .
Equilibrium reaction is CaF2(s) <--> Ca2+(aq) + 2F-(aq) and Ks =
[Ca2+][F-]2. 1 mol of CaF2 leads to 1 mol of Ca2+ and 2 moles of Fupon dissolution, [F-] = 2[Ca2+]
Ks = [Ca2+](2[Ca2+])2; Ks = (3.32 x 10-4)(6.64 x 10-4)2 = 1.46 x 10-10.
Dissolution of reactive gases: CO2 in water
Dissolution reaction is (1):
CO2 ( g )  CO2 (aq)
for which we apply Henry’s law (H = 0.034 mol/(l·bar) = 29.41·105
Pa·l/mol, atmospheric content of CO2 is about 0.038%):
p(CO2 ) 3.8 104 105
5
caq 


1
.
29

10
mol / l
5
H
29.4110
Dissolved carbon dioxide is subject to hydrolysis leading to carbonic
acid, reaction (2):
CO2 (aq)  H 2O  H 2CO3
c( H 2CO3 )
K2 
 1.70 103
caq
CO2 in water II
Carbonic acid dissociates to hydrogen carbonate, reaction (3), and
further to carbonate, reaction (4):
H 2CO3  H   HCO3
HCO3  H   CO32
c( HCO3 )  c( H  )
K3 
 2.50 10 4
c( H 2CO3 )
c(CO32 )  c( H  )
11
K4 

5
.
61

10
c( HCO3 )
Water autoprotolysis also has to be considered, reaction (5):

H 2O  H  OH

K5  c(OH  )  c( H  )  1014
All values of equilibrium constants relate to 25°C.
CO2 in water III
Dissolution of CO2 in water is described by the system of reactions
(1)-(5). Reactions (4) and (5) may be neglected for an open system
(in equilibrium with the atmosphere), allowing a simplified solution:
c( H 2CO3 )  K2  caq  1.70 103 1.29 105  2.196 108 mol / l
c( H  )  c( HCO3 )  K 3  c( H 2CO3 ) 
2.5 10 4  2.193 10 8  2.34 10 6 mol / l
pH   log[ c( H  )]   log[ 2.34 106 ]  5.6
pH of water in equilibrium with the atmosphere (open water not in
contact with buffering minerals such as calcite, atmospheric water) is
about 5.6. In reality pH of rain droplets is slightly higher (about 6) due
to non-equilibrium conditions.
CO2 in water – pressure dependence
The amount of dissolved CO2 in water depends
only on partial pressure of CO2 (and temperature).
Examples:
• in deep waters, where hydrostatic pressure adds
up to atmospheric pressure
• carbonated beverages
CO2 in water – pH dependence
In buffered waters where pH is fixed, only the concentrations of other
species are calculated. Their relative abundance is shown in the
graph vs. pH:
Total amount of dissolved CO2 increases with pH.
Limestone solubility
In contact with limestone, reaction (6) is added to the
system of reactions (1)-(5):
CaCO3  Ca 2  CO32
K 6  c(CO32 )  c(Ca 2 )  3.36  9 10 9
The reason for variation of ion product is the unknown
mineralogical character of limestone. Solution of reaction
system (1)-(6) is a function of CO2 partial pressure and
pH. Minimum solubility of limestone (expressed as
concentration of Ca2+ ions) in open water is about 0.3
mmol/l Ca2+.
p(CO2)
pH
c(Ca2+) mol/l
10−12
12.0
5.19 × 10−3
10−10
11.3
1.12 × 10−3
10−8
10.7
2.55 × 10−4
10−6
9.83
1.20 × 10−4
10−4
8.62
3.16 × 10−4
3.8 × 10−4
8.27
4.70 × 10−4
10−3
7.96
6.62 × 10−4
10−2
7.30
1.42 × 10−3
10−1
6.63
3.05 × 10−3
1
5.96
6.58 × 10−3
10
5.30
1.42 × 10−2
Limestone solubility II
Dependence on partial pressure
of CO2 and pH.