Solutions & Solubility PDF

Solutions & Solubility I. Terminology A. Solution – a homogeneous mixture that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent 1. Solutions do not have...

Solutions & Solubility I. Terminology A. Solution – a homogeneous mixture that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent 1. Solutions do not have to be liquids. Example: air is a solution of nitrogen, oxygen and trace gases B. Solute – material being dissolved (minority component by weight or concentration) C. Solvent – material doing the dissolving (majority component by weight or concentration) D. Molar concentration – moles per liter of solution E. Molal concentration – moles per kilogram of solvent (not kilograms of solution) F. Percent by weight to volume – grams of solute per 100 mL of solution G. Percent by weight to weight – grams of solute per 100 grams of solution H. Normality – equivalent per liter 1. Term largely abandoned in both chemistry and medicine I. Parts per million (ppm) – amount of solute in 1 million parts of solution J. Miscible – liquids that are soluble in each other in all proportions 1. Example: water and ethyl alcohol K. Immiscible – liquids that are poorly soluble in each other 1. Example: water and oil L. Colloids – not true solutions 1. Colloids consist of one phase uniformly dispersed in a second phase, but not actually dissolved II. Solubility A. The solubility of a solute is the amount of the solute that will dissolve in a given amount of solvent at a given temperature B. A saturated solution contains the maximum amount of a solute 1. Supersaturated solutions can exist, but are unstable and excess solute will come out of solution and recrystallize C. Polar substances are more soluble in polar solvents 1. Conversely, nonpolar substances are more soluble in nonpolar solvents III. Energy Changes & the Solution Process A. When a solute dissolves in a solution, there is an associated energy change 1. This energy change is called the ‘heat of solution’ or enthalpy of solution 2. Formally, enthalpy of solution is defined as the change that accompanies dissolving exactly one mole of solute in a solvent B. Enthalpy of solution can be either exothermic or endothermic 1. Enthalpy of solution from two processes a. The components of the ionic substance must be separated, which requires energy (endothermic) b. Ions are then attracted to the dipoles of the water molecule, lowering the energy state (exothermic) c. If the energy required to separate the ions exceeds the energy released by the combining of the ions with the water molecule, then the overall process is endothermic d. If the energy required to separate the ions is less than the energy released by the combining of the ions with the water molecule, then the overall process is exothermic IV. Factors Affecting Solubility A. Pressure has a dramatic effect on the solubility of gaseous solutes in liquid solvents 1. An increase in pressure causes an increase in the solubility of a gaseous solute 2. Henry’s Law describes the relationship between solubility and pressure S = KH x Pgas S= solubility, KH = Henry’s constant, Pgas = partial pressure B. Temperature increases can increase or decrease solubility 1. Solid and liquid solutes in liquid solvents experience and increase in solubility when the temperature is increased 2. Gaseous solutes in liquid solvents experience a decrease in solubility when temperature is increased V. Colligative Properties of Solutions A. The addition of a solute to a solution causes changes in the physical properties of the solution B. There are four colligative properties: vapor pressure, boiling point, freezing point, osmotic pressure 1. Vapor pressure of a liquid results from the most energetic molecules near the surface of the liquid escaping into the gas phase a. As we begin to introduce solute molecules into solution, some of escape points at the surface are now occupied by solute, thus reducing the vapor pressure i. Raoult’s Law can be used to quantify the change in vapor pressure from the introduction of a solute 2. Boiling point of a liquid is defined as the temperature at which the vapor pressure of the material is equal to the ambient pressure a. Solute molecules result in a fall in vapor pressure (see above) i. The resulting fall in vapor pressure means that more energy will have to be added to the system to raise the vapor pressure to above ambient pressure 3. Freezing point is defined as the temperature at which the liquid phase of the material is in equilibrium with the solid phase a. Freezing point is lowered by the addition of solute b. In order for a solution to enter the solid phase, it must organize into an orderly crystalline structure i. The presence of solute particles interferes with this process by getting in the way of the crystalline structure formation 4. Osmotic pressure is directly related to the number of solute particles in solution VI. Solubility Product A. Solubility product constants are used to describe saturated solutions of ionic compounds of relatively low solubility B. The solubility product can be expressed by the formula: Ksp = [M+]x [A-]y 1. In the formula above, M is the concentration of the cation and A is the concentration of the anion. 2. The exponents, x & y, are reflective of the ions dissociated in solution a. Example: PbCl2(s) --> Pb2+(aq) + 2 Cl-(aq) Ksp = [Pb2+][Cl-]2 3. Low values of the solubility product indicate that the substance is relatively insoluble in a given solvent