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Phase interactions are the basis of much of the interesting chemistry that shapes our planet. Dissolved oxygen in water makes it possible for fish and insects to live underwater. The oxygen in the atmosphere originated from microorganisms that lived in the sea. The transfer of oxygen from the sea to the atmosphere mad e it possible of life to develop on the land.
Phase interactions are responsible for the transport of nutrients (and pollutants) in soil, and the formation of all the sedimentary minerals.
The earth is unique among bodies of the solar system, in that it has oceans of liquid water. The phase interactions that result from the oceans stabilize the earth’s temperature and move water to the land where it condenses as fresh water to support both plant and animal lifeforms. An understanding of phase interactions is essential to understanding the chemical processes of the environment.
The two relationships that govern the number of gases that dissolve in water are
Henry’s Law
Dalton’s Law of Partial Pressures.
Dalton’s Law of Partial Pressures:
Dalton’s law of partial pressure states that the solubility of a gas in a liquid solution is a function of partial pressure of the gas.
PT = P1 + P2 + P3 + P4 + · · · ·
If we use the mole fraction of a gas in the solution as a measure of its solubility, then it can be said that the mole fraction of gas in the solution is proportional to the partial pressure of the gas over the solution.
For the dissolution of oxygen in water open to the air, the problem would be worked in this fashion:
PATM = PN2 + PO2 + PH2O + PCO2 + · · · ·
Since air is 20.95% O2, we can write:
PO2 = (0.2095)(Patm – PH2O)
And the solubility of O2 in water is: [O2(aq)] = KH PO2
Calculate the solubility of oxygen in water at 25 ° C, open to the atmosphere at 1.0 atm (760 mm/Hg) of pressure:
Vapor pressure of water at 25 ° C is 23.456 mm/Hg (from Handbook)
23.456 mm/Hg x 1.000 atm/760 mm/Hg = 0.031 atm
PO2 = 0.2095(1.000 – 0.031) = 0.2030 atm
[O2(aq)] = KHPO2 = (1.28 x 10-3 mol/L-atm)(0.2030 atm)
= 2.60 x 10-4 mol/L
Henry’s law
Henry’s law, named after the English chemist William Henry, describes the relationship between the concentration of a gas in a liquid and the pressure of that gas above the liquid. It states that the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid, assuming constant temperature. Mathematically, Henry’s law is expressed as:
C = k × P
Where:
- C represents the concentration of the gas in the liquid (usually expressed in terms of molarity or molality).
- k is the Henry’s law constant, which is specific to a particular gas-solvent system at a given temperature.
- P represents the partial pressure of the gas above the liquid.
According to Henry’s law, when the partial pressure of a gas above a liquid increases, more gas molecules dissolve in the liquid, resulting in a higher concentration of the gas. Conversely, if the partial pressure decreases, the concentration of the dissolved gas decreases as well.
Henry’s law is applicable to gases that do not chemically react with the solvent and where the concentration of the gas is relatively low. It is often used to describe the behavior of gases in solutions, such as the solubility of gases in water (e.g., the solubility of oxygen or carbon dioxide in water).
Henry’s law has various practical applications. For example, it explains phenomena like the absorption of gases in liquids (e.g., the absorption of carbon dioxide in soda) and the release of gases from liquids (e.g., the effervescence of carbonated beverages when the pressure is released). It also plays a significant role in fields such as environmental science, industrial processes, and the study of gas-liquid equilibria.
Clausius-Clapeyron equation
The Clausius-Clapeyron equation is a fundamental thermodynamic equation that relates the vapor pressure of a substance to its temperature and enthalpy of vaporization. It is named after the German physicist Rudolf Clausius and the French engineer Benoît Paul Émile Clapeyron, who independently derived the equation.
The Clausius-Clapeyron equation is given as:
ln(P2/P1) = -(ΔH_vap/R) × (1/T2 – 1/T1)
Where:
- P1 and P2 are the vapor pressures of the substance at temperatures T1 and T2, respectively.
- ΔH_vap is the enthalpy of vaporization, which represents the heat required to convert one mole of the substance from a liquid to a gas phase.
- R is the ideal gas constant.
- T1 and T2 are the corresponding temperatures in Kelvin.
The equation relates the change in vapor pressure to the change in temperature and the enthalpy of vaporization. It can be used to calculate the vapor pressure of a substance at a given temperature, or to determine the enthalpy of vaporization if the vapor pressures at two temperatures are known.
The Clausius-Clapeyron equation is particularly useful for studying phase transitions, such as boiling or evaporation. It provides insights into how the vapor pressure of a substance changes with temperature and can be used to predict the conditions under which a substance will boil or condense.
It is important to note that the Clausius-Clapeyron equation assumes ideal behavior and neglects factors such as intermolecular forces and non-ideal gas behavior. Therefore, it provides a reasonable approximation for many substances but may deviate for complex systems or under extreme conditions.