Physical Pharmacy Solubility of Gases PDF

# **Physical Pharmacy (II)(Cont.)** ## **Solubility of Gases** **By:** Prof. Mokhtar M. El-Baseir ### **Solubility of Gases in Liquids** **Examples of pharmaceutical solutions of gases:** 1. Hydrochloric acid. 2. Ammonia water. 3. Effervescent preparations containing carbon dioxide (CO2) that a...

# **Physical Pharmacy (II)(Cont.)** ## **Solubility of Gases** **By:** Prof. Mokhtar M. El-Baseir ### **Solubility of Gases in Liquids** **Examples of pharmaceutical solutions of gases:** 1. Hydrochloric acid. 2. Ammonia water. 3. Effervescent preparations containing carbon dioxide (CO2) that are dissolved and maintained in solution under positive pressure. 4. Aerosol products in which the propellant is either CO2 or nitrogen. ### **Factors Affecting Solubility of Gases in Liquids** 1. **Temperature** Basically, solubility increases with temperature. However, gases become less soluble in each other and in water, but more soluble in organic solvents as the temperature increases. 2. **Pressure** For the majority of solid and liquid solutes, pressure does not affect the solubility of solids or liquids. Solubility of gas increases with increased pressure. The effect of pressure on the solubility of a gas is expressed by **Henry's Law**: "In a very dilute solution at constant temperature, the concentration of dissolved gas is proportional to the partial pressure of the gas above the solution (at equilibrium)". **Henry's Law may be written as**: $C2 = σP$ Where C2 is the concentration of the dissolved gas in g/L, σ is the proportionality constant for a particular solution being considered (solubility coefficient) and P is the partial pressure of the undissolved gas above the solution. **Note**: Henry's law holds only when there is no chemical reaction between the solute and the solvent. The amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution. **Example 1.** What is the solubility of oxygen in water at 25°C and a partial pressure of about 610 mm Hg if the Henr'ys law constant is expressed as σ = concentration (g/liter H₂O)/pressure (mm Hg) = 5.33 x 10-5? $C2 = σP$ $C2 = 5.33 * 10^{-5} * 610 = 0.0325 g/liter$ **Example 2.** If 0.0160 g of oxygen dissolves in 1 liter of water at a temperature of 25°C and at an oxygen pressure of 300 mm Hg, calculate σ. $C2 = σP$ $σ = C2/P$ $σ = 0.0160/300 = 5.33 * 10^{-5}$ ## **Other Factors Affecting the Solubility of Gases in Liquids** 1. **Salting-out** Gases are often liberated from solutions in which they are dissolved by the introduction of electrolytes such as sodium chloride and sometimes by nonelectrolytes such as sucrose. This phenomenon is known as "salting out." The resultant escape of gas is due to the attraction of salt ions or the highly polar nonelectrolyte for the water molecules, which reduces the density of the aqueous environment adjacent to the gas molecules. Salting out may also occur in solutions of liquids in liquids and solids in liquids. 2. **Chemical Reaction** Gases such as hydrogen chloride, ammonia, and carbon dioxide show an increase in solubility due to a chemical reaction between the gas and solvent, usually with resultant increase in solubility. **Example:** Hydrogen chloride is about 10,000 times more soluble in water than is oxygen. ### **Solubility Calculations** The solubility of gas in a liquid may be expressed either by Henry's law constant (σ) or by Bunsen absorption coefficient (α). #### **Bunsen Coefficient (α)** The volume of gas in liters (reduced to standard condition of 0°C and 760 mm pressure) that dissolves in 1 liter of solvent and a partial pressure of 1 atmosphere of the gas at definite temperature. $α = \frac{V_{gas}}{V_{solution} * P}$ Where * $V_{gas}$ is the volume of gas at STP dissolved in a volume $V_{solution}$ of solution at a partial gas pressure of P. **Example** If 0.0160 g of oxygen dissolves in 1 liter of water at a temperature of 25°C and at an oxygen pressure of 300 mm Hg calculate the Bunsen coefficient. In order to compute the Bunsen coefficient one must first reduce the volume of gas to STP using the ideal gas equation $V=nRT/P$ $V_{gasSTP} = 0.0160/32 * 0.08205 * 273.16 = 0.0112 at STP$ $α = \frac{V_{gas}}{V_{solution} * P} = \frac{0.0112}{1 * 300/760} = 0.0254$