Determination of Dissolved Oxygen of Wastewater by Winklet Method
Experiment IV
Solubility of Dissolved Oxygen
Purpose
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.
References
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.
Theory
The dissolution and evolution of dissolved oxygen (DO) in the air-water system can be represented by the equation: [pic] (1) which has the equilibrium constant: [pic] (2) 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: [pic] (3)
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: [pic] (4) 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
Solubility of Dissolved Oxygen
Purpose
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.
References
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.
Theory
The dissolution and evolution of dissolved oxygen (DO) in the air-water system can be represented by the equation: [pic] (1) which has the equilibrium constant: [pic] (2) 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: [pic] (3)
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: [pic] (4) 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
References: 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.