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
Physical properties of water and their importance
Hydrogen bonds
Boiling points of structurally similar compound from the 4.-7. period
Anomalous water properties: boiling point
Anomalous water properties: density
Density maximum40C
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
)()( aqAgA
Dissolution can be described in terms of chemical reaction, e.g. for gas in water
Thermodynamic relations derived for chemical reactions can be applied to this process, e.g. the equilibrium constant
n
1i
νiiaK
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
p
p
f
fa i
i
ii
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
Standard states II
Standard state for (aqueous) solutions is solution at unit concentration:
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.
ii
i
i
ii c
c
c
f
fa
H
p
Ap
pAc
Aa
AaK
gi
aq
g
aqh
)(
)(
)(
)(
Equilibrium in dissolution reactions
Equilibrium constant for dissolution of A gas in water is:
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:
)()()( aqBaqAsAB
)()()(
)()(
aqaqs
aqaqs BcAc
ABa
BaAaK
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 F- upon
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):
)()( 22 aqCOgCO 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%):
lmolH
COpcaq /1029.1
1041.29
10108.3)( 55
542
3222 )( COHOHaqCO
Dissolved carbon dioxide is subject to hydrolysis leading to carbonic acid, reaction (2):
3322 1070.1
)( aqc
COHcK
CO2 in water II
Carbonic acid dissociates to hydrogen carbonate, reaction (3), and further to carbonate, reaction (4):
332 HCOHCOH4
32
33 1050.2
)(
)()(
COHc
HcHCOcK
233 COHHCO
11
3
23
4 1061.5)(
)()(
HCOc
HcCOcK
Water autoprotolysis also has to be considered, reaction (5):
OHHOH 214
5 10)()( HcOHcK
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:
lmolcKCOHc aq /10196.21029.11070.1)( 853232
lmol
COHcKHCOcHc
/1034.210193.2105.2
)()()(
684
3233
6.5]1034.2log[)](log[ 6 HcpH
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):
23
23 COCaCaCO
92236 10936.3)()( CacCOcK
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+.
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
15.9
66.58 ×
10−3
105.3
01.42 ×
10−2
Limestone solubility II
Dependence on partial pressure of CO2 and pH.
p(CO2) pH c(Ca2+) mol/l