Chapter 12

Remember that a solution is any homogeneous mixture. There are many types of solutions: SoluteSolvent Resulting Solution Examples gasgasgasair gasliquidliquid soda water gassolidsolid H 2 gas in palladium liquidliquidliquidwhiskey solidliquidliquid NaCl in water solidsolidsolidbrass

Some important terms to learn: Solvent – The substance in a solution that causes the solution to be made. Usually it is the most abundant substance in the mixture. Solute – The substance that is dissolved. There can be more than one solute in a solution. In the most common solution, the solute is a solid or a liquid and the solvent is a liquid.

Solubility - Usually expressed as g solute/ 100 g solvent (sometimes for aqueous solutions [any solution in which H 2 O is the solvent] g solute/ 100 mL solvent) Saturated Solution – A solution in which the maximum amount of solute that will dissolve is present. Each solute-solvent system is unique Unsaturated Solution – A solution in which less than the maximum amount of solute is dissolved.

Supersaturated Solution – Occasionally, during the dissolving process, more than the theoretical maximum amount of solute has dissolved. This is an unstable situation. In time, some of the solute crystallizes out of the solution. Frequent stirring during the solution process can help prevent this. Crystallization is basically the same as precipitation, except that the processes result in solids of different appearance. Precipitation usually results in small particles while crystallization frequently results in larger crystals.

In order for a solute to dissolve in a solvent, the solute particles have to fit in between the solvent particles. This usually can help us predict which solutes will dissolve in which solvents. This is usually accomplished with solute-solvent attractions. There is a general Rule of Thumb for solubility: Like Dissolves Like. We will try to explain each case briefly Some solutions have complete miscibility (gas in gas, many liquids in liquids). Miscibility means complete solubility in any proportion.

Let’s learn some more concentration terms: 1. Wt. % = Sometimes this is written as g solute/ 100 g solution 2. Mole Fraction of A = X A = 3. Molality (m) = The # of moles of solute dissolved in 1 kg of solvent

Let’s do examples 12.2 and 12.3 on pages

All the concentration terms (including molarity) can be converted into each other. Let’s do example 12.4 on page 510 in your text (we will also calculate weight %).

Solubility varies with temperature. For a gas dissolved in a liquid, the solubility always decreases as temperature increases. For solids dissolved in liquids, there is no general rule about the variation of solubility with temperature, except that most often solubility increases as temperature increases. See figure 12.3 on page 512 in your text:

External pressure has no effect on the solubilities of liquids and solids in a liquid solvent. It does effect greatly the solubility of gases in liquids. The effect can be calculated using Henry’s Law, which says that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas over the solution.

c A = kP A, where c A is the molar concentration of gas A, P A is the partial pressure of A over the solution and k is a constant that depends only on T. This law is dramatically illustrated with carbonated beverages. CO 2 is dissolved in water in a sealed container and CO 2 is added to the container, thus increasing the P CO2 and thus increasing the amount of CO 2 dissolved. When the container is opened, the CO 2 above the solution escapes, thus the solubility of the CO 2 decreases and CO 2 has to escape from the solution, hence the bubbling effect. If left open long enough all the CO 2 escapes and the beverage becomes flat.

Most gases obey Henry’s Law with some notable exceptions, especially if the gas reacts with water. In those cases much higher solubilities can be obtained. A good example is ammonia in water.

There are a whole class of physical properties of solutions that are interrelated. These are called Colligative Properties – Depend only on the number of solute particles in a solution but not on the nature of the solute particles. 1. Raoult’s Law – When a non-volatile solute (basically any solid), the partial pressure of a solvent, P S, over a solution is directly proportional to the mole fraction of the solvent, X S, and the VP of the pure solvent, P. P S = X S P

The net effect is that the vapor pressure of the solution is always lower than the vapor pressure of the pure solvent. Raout’s Law can be rewritten as:  P = X solute P

If both components are volatile, Raoult’s Law still works. It has to be applied for each component, solute and solvent: P A = X A P and P B = X B P Using Dalton’s Law of Partial Pressures P T = P A + P B or P T = X A P+ X B P

Raoult’s Law works for an ideal solution. Most solutions are not ideal, but we will consider all solutions to be ideal. Let’s do example 12.7 on page 517 in your text.

All the other colligative properties can be derived from Raoult’s Law: 2. A non-volatile solute will raise the B.P. of a solution above that of the pure solvent:  T b = k b m, where m is the molality and k b is a constant dependent on the solvent, not the solute, and is called the Molal Boiling- Point Elevation Constant

Similarly, a non-volatile solute will lower the freezing point of a solution below that of the pure solvent:  T f = k f m, where m is the molality and k f is a constant dependent on the solvent, not the solute, and is called the Molal Freezing-Point Depression Constant

Let’s do example 12.8 on page 522 of your text.

Osmosis – When 2 solutions are separated by a semipermeable membrane (one that allows solvent molecules to pass through but not solute particles [usually we are dealing with small solvent molecules like water and much larger solute particles]), solvent will flow from the lower concentration to the higher concentration. This is due to the attempt to achieve equilibrium (same concentration) between the 2 solutions. This process can be stopped if pressure is applied to the membrane from the high concentration side. The pressure needed to stop the flow is called the Osmotic Pressure (  )

We compare the concentrations of the 2 solutions with new terms: hypertonic – The solution of higher concentration hypotonic – The solution of lower concentration isotonic - 2 solutions of equal concentration

The flow of solvent is always from hypotonic to hypertonic until they are isotonic. Animal cell walls are semipermeable membranes and cells can be made to burst or shrivel if placed in hypotonic solutions or hypertonic solutions respectively.

 = MRT where M is the molarity (this is the only colligative property that uses molarity), R is the Universal Gas Law Constant and T is the Kelvin temperature. Note that this is very similar to the Universal Gas Law Equation. Helpful for remembering.

Colligative properties can be used to find the molecular weights of solutes in non-electrolyte solutions. You will do this in the last experiment in lab, making use of the Freezing Point Depression. Boiling Point Elevation can also be used. If we know K f and measure the change in freezing point or boiling point, then we can determine the molarity. If we know the mass of solvent and unknown solute, we can then determine the MW. Let’s do example and on pages 526 and 527 in your text.

What happens if the solute is an electrolyte? Remember that colligative properties depend on the number of particles only. NaCl, when dissolved in water produces Na +1 and Cl -1 ions. Thus one molecule of NaCl produces 2 particles. A 1.0 molal NaCl solution really has a 2.0 molal effect. Likewise, BaCl 2 produces 3 particles and a 1.0 molal solution has a 3.0 molal effect

For electrolytic solutions the colligative property equations all have the format of:  T f = i(k f m) where i is the actual number of particles produced per formula unit. This i is inserted in the same way in the other colligative property formulas. It is called the van’t Hoff factor. In reality the van’t Hoff factor is usually less than expected because even a strong electrolyte has some ions pairing up as one particle. Unless otherwise stated, we will ignore this and simply assume complete ionization.