Solubility of Dissolved Oxygen
To demonstrate the effect of partial pressure, temperature and salinity on the solubility of dissolved oxygen and to demonstrate the interference of nitrite in dissolved oxygen analysis by the Winkler Method. To demonstrate the use of the oxygen electrode and the difference between activity and concentration.
1. Mancy, K. H., Jaffe, T., "Analysis of Dissolved Oxygen in Natural and Waste Water," USDHEW Public Health Service, Pub. 999-WP-37, Cincinnati, Ohio, 1966. 2. Standard Methods, 17th ed., American Public Health Association, 1989. 3. Stumm, W., Morgan, J. J., Aquatic Chemistry, 3rd ed., Wiley Interscience, 1996. 4. Sawyer, C. N., McCarty, P. L., and Parkin, G. F. Chemistry for Environmental Engineering, 5th ed., McGraw Hill, 2003.
The dissolution and evolution of dissolved oxygen (DO) in the air-water system can be represented by the equation:
which has the equilibrium constant:
where aO2(aq) is the activity of oxygen in the water and fO2(g) is the fugacity, or activity, of oxygen in the gas phase. When the fugacity of oxygen is approximated by its partial pressure (PO2) and the activity of oxygen in water by its concentration, [O2(aq)], then Henry's Law can be expressed as:
KH approaches K in dilute solutions and for perfect gases. Henry's Law is generally a valid approximation for natural, fresh waters in equilibrium with the atmosphere. At constant temperature, oxygen partial pressure is given by the equation:
where VFO2 is the volume fraction of oxygen in dry air (generally 0.208), P is the total pressure of the gas phase, and PH2O is the vapor pressure of the water. Under normal atmospheric conditions, the vapor pressure correction is negligible.
As ionic strength increases, it is necessary to take into account the difference between activity and concentration of the dissolved oxygen (DO). These two terms are related by the activity coefficient, as:
For non-electrolytes, t is related to the ionic strength, I, by the empirical equation:
where Ks is the "salting out" coefficient and
with Ci being the concentration in mole/liter of an ion with charge Zi.
The activity coefficient is greater than one for DO, as it is for most non-electrolytes. As the ionic strength is increased, more water molecules are required to hydrate the ionic species and fewer are available to hydrate, or dissolve, oxygen. It is important to note that the DO activity at equilibrium remains constant with increasing ionic strength (see Equation 2) because the fugacity of oxygen in the gas phase is not a function of the ionic strength of the solution. Thus, as ionic strength increases, the DO concentration at equilibrium must decrease.
The effect of temperature on oxygen solubility is given by the Van't Hoff equation:
where K1 and K2 are the equilibrium constants at temperatures T1 and T2, respectively, ∆H is the enthalpy of reaction in cal/mole and R is the ideal gas constant, 1.99 cal/(mole.K). This equation is based on the assumption that ∆H does not vary with temperature over the range T1 to T2.
Measurement of Dissolved Oxygen
One of the more common procedures for measuring DO concentrations is the Winkler method. In this procedure O2 is reduced by Mn2+ at high pH.
4 e- + 4 H+ + O2 ( 2 H2O
2 Mn2+ + 4 OH- ( 2 MnO2(s) + 4 H+ + 4 e-
2 Mn2+ + 4 OH- + O2 ( 2MnO2(s) + 2 H2O
Upon addition of H2SO4, MnO2 oxidizes I-,
4 e- + 2 MnO2(s) + 8 H+(2 Mn2+ + 4 H2O
4 I- ( 2 I2 + 4 e-
MnO2(s) + 4 H+ + 2 I- ( I2 + Mn2+ + 2 H2O(10)
The I2, which is elly present as the triiodide ion, I3-, is then titrated...