Chapter 13

Explain how the relative strengths of solute-solute interactions, solvent-solvent interactions, and solvent-solute interactions affect solution formation.

A solution always forms if the solvent-solute interactions are comparable to, or stronger than, the solvent-solvent interactions and the solute-solute interactions.

Define colligative property.

Colligative properties are properties that depend on the amount of solute and not the type of solute. Examples of colligative properties are vapor pressure lowering, freezing point depression, boiling point elevation, and osmotic pressure.

State the kind of intermolecular forces that would occur between the solute and solvent in lard (nonpolar) solution.

Dispersion

State the kind of intermolecular forces that would occur between the solute and solvent in motor oil (nonpolar) solution. dispersion hydrogen bonding ion-dipole dipole-dipole

Dispersion

State the kind of intermolecular forces that would occur between the solute and solvent in ethylene glycol (polar, contains OH groups) solution.

Dispersion, hipole-hipole, and hydrogen bond

What are the common units for expressing solution concentration?

Parts by volume, mole percent, molality, parts by mass, Molarity, and mole fraction

What are three steps involved in evaluating the enthalpy changes associated with solution formation?

Step 1: Separate the solute into its constituent particles. This step is always endothermic (positive H) because energy is required to overcome the forces that hold the solute together. Step 2: Separate the solvent particles from each other to make room for the solute particles. This step is also endothermic because energy is required to overcome the intermolecular forces among the solvent particles. Step 3: Mix the solute particles with the solvent particles. This step is exothermic because energy is released as the solute particles interact with the solvent particles through the various types of intermolecular forces.

How does this temperature dependence affect the amount of oxygen available for fish and other aquatic animals?

The decreasing solubility of gases with increasing temperature results in a lower oxygen concentration available for fish and other aquatic life in warm waters.

What is the heat of hydration (ΔHhydration)?

The heat of hydration is the enthalpy change that occurs when 1 mol of gaseous solute ions are dissolved in water. In aqueous solutions, ΔHsolvent and ΔHmix can be combined into a single term called the heat of hydration (ΔHhydration).

How does the solubility of a gas in a liquid depend on pressure?

The higher the pressure of a gas above a liquid, the more soluble the gas is in the liquid.

How does the solubility of a gas in a liquid depend on temperature?

The solubility of gases in liquids decreases with increasing temperature.

Explain why the lower vapor pressure for a solution containing a nonvolatile solute results in a higher boiling point and lower melting point compared to the pure solvent.

A nonvolatile solute lowers the vapor pressure of a solution relative to that of the pure solvent. The vapor pressure lowering occurs at all temperatures, which shifts the vaporization curve in the phase diagram. This means that the temperature must be raised above the pure solvent normal boiling point in order for the solution vapor pressures to be raised to 1 atm. This shift also results in a lowering of where the vapor pressure curve intersects the solid-gas curve. The net effect is that the solution has a lower melting point and a higher boiling point than the pure solvent.

Define saturated, unsaturated, and supersaturated solutions.

A saturated solution is a solution in which the dissolved solute is in dynamic equilibrium with the solid (or undissolved) solute. If you add additional solute to a saturated solution, it will not dissolve. An unsaturated solution is a solution containing less than the equilibrium amount of solute. If you add additional solute to an unsaturated solution, it will dissolve. A supersaturated solution is a solution containing more than the equilibrium amount of solute. Such solutions are unstable and the excess solute normally precipitates out of the solution. However, in some cases, if left undisturbed, a supersaturated solution can exist for an extended period of time.

What does it mean to say that a substance is soluble in another substance?

A substance is soluble in another substance if they can form a homogeneous mixture.

How does the solubility of a solid in a liquid depend on temperature?

Although there are exceptions, the solubility of most solids in water increases with increasing temperature.

Explain the difference between an ideal and a nonideal solution.

An ideal solution is a solution that follows Raoult's law at all concentrations for both the solute and the solvent. A nonideal solution will exhibit deviations from Raoult's law in the vapor pressure of a component as mole fraction of this component decreases from 1 (i.e. the pure component).

How does the enthalpy of solution depend on the relative magnitudes of ΔHsolute and ΔHhydration ?

Because the ion-dipole interactions that occur between a dissolved ion and the surrounding water molecules are much stronger than the hydrogen bonds in water, ΔHhydration is always largely negative (exothermic) for ionic compounds. Using the heat of hydration, we can write the enthalpy of solution as a sum of just two terms, one endothermic and one exothermic: ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmix = ΔHsolute + ΔHhydration = endothermic(+) term + exothermic(-) term. For ionic compounds, ΔHsolute, the energy required to separate the solute into its constituent particles, is simply the negative of the solute's lattice energy (ΔHsolute = - ΔHlattice) For ionic aqueous solutions, then, the overall enthalpy of solution depends on the relative magnitudes of ΔHsolute and ΔHhydration, with three possible scenarios (in each case we refer to the magnitude or absolute value of ΔH): (1) If ΔHsolute < ΔHhydration the amount of energy required to separate the solute into its constituent ions is less than the energy given off when the ions are hydrated. ΔHsoln is therefore negative and the solution process is exothermic and the solution feels warm to the touch. (2) If ΔHsolute > ΔHhydration the amount of energy required to separate the solute into its constituent ions is greater than the energy given off when the ions are hydrated. ΔHsoln is therefore positive and the solution process is endothermic (if a solution forms at all) and the resulting solution feels cool to the touch. (3) If ΔHsolute = ΔHhydration the amount of energy required to separate the solute into its constituent ions is about equal to the energy given off when the ions are hydrated. ΔHsoln is therefore approximately zero and the solution process is neither appreciably exothermic nor appreciably endothermic and there is no noticeable change in temperature.

What is Henry's law? For what kinds of calculations is Henry's law useful?

Henry's law quantifies the solubility of gases with increasing pressure as follows: Sgas = kH*Pgas, where Sgas is the solubility of the gas; kH is a constant of proportionality (called the Henry's law constant) that depends on the specific solute, solvent, and temperature; and Pgas is the partial pressure of the gas. The equation simply shows that the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid. If the solubility of a gas is known at a certain temperature, the solubility at another pressure at this temperature can be calculated.

Pick an appropriate solvent from the table below to dissolve each substance. State the kind of intermolecular forces that would occur between the solute and solvent in each case. Pick an appropriate solvent to dissolve motor oil (nonpolar). Water Acetone Methanol Ethanol Hexane Toluene Carbon tetrachloride

Hexane, Toluene, and Carbon tetrachloride

Why do two ideal gases thoroughly mix when combined? What drives the mixing?

Ideal gases do not interact with each other in any way (that is, there are no significant forces between their constituent particles). When the two gases mix their potential energy remains unchanged, so this does not drive the mixing. The tendency to mix is related, rather, to a concept called entropy. Entropy is a measure of energy randomization or energy dispersal in a system. Recall that a gas at any temperature above 0 K has kinetic energy due to the motion of its atoms. When the gases are separated, their kinetic energies are also confined to those regions. However, when the gases mix the kinetic energy of each gas becomes spread out or dispersed over a larger volume. Therefore, the mixture of the two gases has greater energy dispersal, or greater entropy, than the separated components. The pervasive tendency for all kinds of energy to spread out, or disperse, whenever they are not restrained from doing so is the reason that two ideal gases mix.

What is the effect on vapor pressure of a solution with particularly strong solute-solvent interactions? With particularly weak solute-solvent interactions?

If the solute-solvent interactions are particularly strong (stronger than solvent-solvent interactions), then the solute tends to prevent the solvent from vaporizing as easily as it would otherwise and the vapor pressure of the solution will be less than that predicted by Raoult's law. If the solute-solvent interactions are particularly weak (weaker than solvent-solvent interactions), then the solute tends to allow more vaporization than would occur with just the solvent and the vapor pressure of the solution will be greater than predicted by Raoult's law.

How does this pressure dependence account for the bubbling that occurs upon opening a can of soda?

In a sealed can of soda pop, for example, the carbon dioxide is maintained in solution by a high pressure of carbon dioxide within the can. When the can is opened, this pressure is released and the solubility of carbon dioxide decreases, resulting in bubbling.

Explain dynamic equilibrium with respect to solution formation.

In any solution formation, the initial rate of dissolution far exceeds the rate of deposition. But as the concentration of dissolved solute increases, the rate of deposition also increases. Eventually the rate of dissolution and deposition become equal—dynamic equilibrium has been reached.

State the kind of intermolecular forces that would occur between the solute and solvent in potassium chloride (ionic) solution.

Ion-dipole and dispersion

What is osmosis? What is osmotic pressure?

Osmosis is defined as the flow of solvent from a solution of lower solute concentration to one of higher solute concentration through a semipermeable membrane—a membrane that selectively allows some substances to pass through but not others. The osmotic pressure is the pressure required to stop the osmotic flow and is given by the following equation: П = MRT.

How are parts by mass and parts by volume used in calculations?

Parts by mass and parts by volume are ratios of masses and volume, respectively. A parts by mass concentration is the ratio of the mass of the solute to the mass of the solution, all multiplied by a multiplication factor, where percent by mass (%) is the desired unit, the factor = 100; where parts per million by mass (ppm) is the desired unit, the factor = 10^6; and for parts per billion by mass (ppb), the factor = 10^9. The size of the multiplication factor depends on the concentration of the solution. For example, in percent by mass, the multiplication factor is 100%, so percent by mass = (mass solute)/(mass solution)*100%. A solution with a concentration of 28% by mass contains 28 g of solute per 100 g of solution.

What is Raoult's law? For what kind of calculations is Raoult's law useful?

Raoult's law quantifies the relationship between the vapor pressure of a solution and its concentration as P(solution) = x(solvent)*P^o(solvent) where Psolution is the vapor pressure of the solution, x(solvent)is the mole fraction of the solvent, and P^o(solvent) is the vapor pressure of the pure solvent. This equation allows you to calculate the vapor pressure of a solution or to calculate the concentration of a solution, given the vapor pressure of the solution.

How is this temperature dependence exploited to purify solids through recrystallization?

Recrystallization is a common technique to purify a solid. In this technique, the solid is put into water (or some other solvent) at an elevated temperature. Enough solid is added to the solvent to create a saturated solution at the elevated temperature. As the solution cools, it becomes supersaturated and the excess solid begins to come out of solution. If the solution cools slowly, the solid forms crystals as it comes out of solution. The crystalline structure tends to reject impurities, resulting in a purer solid. The solvent is chosen such that the solubility of the impurities is high at the lower temperature reducing the tendency to co-precipitate.

Define solution, solute, and solvent.

Solution: Mixture where one substance dissolves another. Solute: The substance that dissolves (minor component). Solvent: The substance that does the dissolving (major component).

What does the statement, like dissolves like, mean with respect to solution formation?

The statement "like dissolves like" means that similar kinds of solvents dissolve similar kinds of solutes. Polar solvents, such as water, dissolve many polar or ionic solutes, and nonpolar solvents, such as hexane, dissolve many nonpolar solutes.

Explain the role and meaning of the van't Hoff factor in determining the colligative properties of solutions containing ionic solutes.

The van't Hoff factor (i) is the ratio of moles of particles in solution to moles of formula units dissolved: i = (moles of particles)/(moles of formula units dissolved). The van't Hoff factor often does not match its theoretical value, due to the fact that the ionic solute is not completely dissolved into the expected number of ions, leaving ion pairs in solution. The result is that the number of particles in the solution is not as high as theoretically expected.

What is the effect of a nonvolatile solute on the vapor pressure of a liquid? Why is the vapor pressure of a solution different from the vapor pressure of the pure liquid solvent?

The vapor pressure of the solution is lower than the vapor pressure of the pure solvent. The simplest explanation for the lowering of the vapor pressure of a solution relative to that of the pure solvent is related to the concept of dynamic equilibrium itself. In dynamic equilibrium the rate of vaporization is equal to the rate of condensation. When a nonvolatile solute is added, however, the solute particles interfere with the ability of the solvent particles to vaporize, simply because they occupy some of the surface area formerly occupied by the solvent. The rate of vaporization is therefore diminished compared to that of the pure solvent. The change in the rate of vaporization creates an imbalance in the rates; the rate of condensation is now greater than the rate of vaporization. The net effect is that some of the molecules that were in the gas phase condense into the liquid. As they condense, the reduced number of molecules in the gas phase causes the rate of condensation to decrease. Eventually the two rates become equal again, but only after the concentration of molecules in the gas phase has decreased, which means a lower vapor pressure for the solution compared to the pure solvent.

Pick an appropriate solvent to dissolve ethylene glycol (polar, contains OH groups). Water Acetone Methanol Hexane Toluene Carbon tetrachloride

Water, Acetone, and Methanol

Pick an appropriate solvent to dissolve potassium chloride (ionic). Water Methanol Ethanol Hexane Toluene Carbon tetrachloride

Water, methanol, and ethanol

Why is entropy important in discussing the formation of solutions?

When two substances mix to form a solution there is an increase in randomness, due to the fact that the components are no longer segregated to separate regions. This makes the formation of a solution energetically favorable, even when it is endothermic.

What kinds of intermolecular forces are involved in solution formation?

Whether two substances will spontaneously mix to form a homogeneous solution is dependent on a number of different types of intermolecular forces including dispersion forces, dipole-dipole forces, hydrogen bonding, and ion-dipole forces.

Define entropy.

a measure of the disorder of a system