PropertyConsequence Excellent solventTransport of nutrients and waste products, prerequisite of biogeochemical processes High dielectric constantSolubility of ionic compounds High surface tensionPhysiological 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 vaporizationControls the transfer of vapor between atmosphere and water High heat of meltingStabilization of temperature regime at freezing/melting High heat capacityStabilization of temperature Physical properties of water and their importance

Hydrogen bonds

Boiling points of structurally similar compound from the period Anomalous water properties: boiling point

Anomalous water properties: density Density maximum 4 0 C 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 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 KClO 4, 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. PbCl 2 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 Thermodynamic relations derived for chemical reactions can be applied to this process, e.g. the equilibrium constant K is the equilibrium constant of the reaction a i is the equilibrium activity of i compound, ν i 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° = 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.

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:

Dissolution of minerals - examples Calculate molar solubility of AgCl in water – dissolution reaction is AgCl(s) --> Ag + + Cl -. From strochiometric ballance [Ag + ] = [Cl - ]. K s = 1.76 x = [Ag + ][Cl - ] = [Ag + ] 2, [Ag + ] = 1.33 x and molar solubility of AgCl is 1.33 x mol/l. Concentration of Ca 2+ (aq) equal to 3.32 x mol/l was obtained from analysis of water in contact with fluorite (CaF 2 ). Calculate the ion product of CaF 2. Equilibrium reaction is CaF 2 (s) Ca 2+ (aq) + 2F - (aq) and K s = [Ca 2+ ][F - ] 2. 1 mol of CaF 2 leads to 1 mol of Ca 2+ and 2 moles of F - upon dissolution, [F - ] = 2[Ca 2+ ] K s = [Ca 2+ ](2[Ca 2+ ]) 2 ; K s = (3.32 x )(6.64 x ) 2 = 1.46 x

Dissolution of reactive gases: CO 2 in water Dissolution reaction is (1): for which we apply Henry’s law (H = mol/(l·bar) = 29.41·10 5 Pa·l/mol, atmospheric content of CO 2 is about 0.038%): Dissolved carbon dioxide is subject to hydrolysis leading to carbonic acid, reaction (2):

CO 2 in water II Carbonic acid dissociates to hydrogen carbonate, reaction (3), and further to carbonate, reaction (4): Water autoprotolysis also has to be considered, reaction (5): All values of equilibrium constants relate to 25°C.

CO 2 in water III Dissolution of CO 2 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: 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.

CO 2 in water – pressure dependence The amount of dissolved CO 2 in water depends only on partial pressure of CO 2 (and temperature). Examples: in deep waters, where hydrostatic pressure adds up to atmospheric pressure carbonated beverages

CO 2 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 CO 2 increases with pH.

Limestone solubility In contact with limestone, reaction (6) is added to the system of reactions (1)-(5): The reason for variation of ion product is the unknown mineralogical character of limestone. Solution of reaction system (1)-(6) is a function of CO 2 partial pressure and pH. Minimum solubility of limestone (expressed as concentration of Ca 2+ ions) in open water is about 0.3 mmol/l Ca 2+.

10 − × 10 −3 10 − × 10 −3 10 − × 10 −4 10 − × 10 −4 10 − × 10 − × 10 − × 10 −4 10 − × 10 −4 10 − × 10 −3 10 − × 10 − × 10 − × 10 −2 Limestone solubility II Dependence on partial pressure of CO 2 and pH. p(CO 2 ) pH c(Ca 2+ ) mol/l