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