Ch. 13 Properties of Solutions
13.1 The Solution Process
the entropy of a system increases if its degree of disorder increases, or if its energy becomes dispersed over a greater number of particles. formation of solutions is favored by the increase in entropy that accompanies mixing
Like dissolves like
entropy of a system increases if its degree of disorder increases, or if its energy becomes dispersed over a greater number of particles.
The enthalpy of the system, however, changes little upon gas mixing because there are few intermolecular interactions between the gas particles
All of the intermolecular forces (dispersion forces, dipole–dipole interactions, hydrogen bonding) and their corresponding energies are part of the enthalpy of the system. Decreasing the enthalpy of the system would correspond to increasing favorable intermolecular interactions between gas particles. Furthermore, the balance between increasing the entropy and decreasing the enthalpy of a system is what determines whether a process is spontaneous. Thus, the formation of solutions is favored by the increase in entropy that accompanies mixing.
13.1.2 The Effect of Intermolecular Forces on Solution Formation
Below is intermolecular interactions involved in solutions
Partial negative charge goes on the more electronegative atom when polar
When solute and solvent forces are not greatly different the solute miscible in the solvent
3 kinds of intermolecular interactions
Solute–solute interactions between solute particles must be overcome to disperse the solute particles through the solvent.
Solvent–solvent interactions between solvent particles must be overcome to make room for the solute particles in the solvent.
Solvent–solute interactions between the solvent and solute particles occur as the particles mix.
When the solvent is water, the interactions are referred to as hydration.
Interactions between ions and nonpolar molecules tend to be very weak (weak london dispersion forces)
13.1.3 Energetics of Solution Formation
Equation for overall enthalpy change:
Endothermic: requires an input of energy to overcome their attractive interactions. Change in enthalpy will be positive (change in h solute and change in h solvent will always be endothermic)
Change in h mix will always be negative and exothermic
Endothermic vs exothermic results below
Hot packs are exothrmic and cold packs are endothermic
13.2 Saturated Solutions and Solubility
As a solid solute begins to dissolve in a solvent, the concentration of solute particles in solution increases, increasing the chances that some solute particles will collide with the surface of the solid and reattach. This process, which is the opposite of the solution process, is called crystallization
Equation for this
When the rates of these opposing processes become equal, a dynamic equilibrium is established
the solubility of a particular solute in a particular solvent is the maximum amount of the solute that can dissolve in a given amount of the solvent at a specified temperature, assuming that excess solute is present.
Keep in mind equilibrium graphs are not forever diagonal
13.3 Factors affecting solubility
the stronger the attractions between solute and solvent molecules, the greater the solubility of the solute in that solvent.
Liquids that mix in all proportions, such as acetone and water, are miscible, whereas those that do not dissolve in one another are immiscible
One way to enhance the solubility of a substance in water is to increase the number of polar groups the substance contains.
The solubilities of solids and liquids are not appreciably affected by pressure, whereas the solubility of a gas in any solvent is increased as the partial pressure of the gas above the solvent increases
The relationship between pressure and gas solubility is expressed by Henry’s law:
13.3.3 Temp effects
The solubility of most solid solutes in water increases as the solution temperature increases
the solubility of gases in water decreases with increasing temperature
As you heat an aqueous solution that contains a gas, you give enough energy to the gas particles for them to escape from the solution and go into the gas phase. Thus, increasing temperature reduces the solubility of the gas. On the other hand, if you heat up a mixture of water and an ionic solid, you give energy to the solid so it can break into its constituent ions that are then hydrated by water. Thus, increasing temperature increases the solubility of the ionic solid.
13.4 Expressing Solution Concentrations
One of the simplest quantitative expressions of concentration is the mass percentage
We often express the concentration of very dilute solutions in parts per million (ppm) or parts per billion (ppb).
mole fraction of a component of a solution is given by
The molality of a solution, denoted m, is a concentration unit that is also based on moles of solute. Molality equals the number of moles of solute per kilogram of solvent:
Molarity depends on the volume of solution.Molality depends on the mass of solvent.
To convert between molality and molarity, the density of the solution will be needed
Lowering of the freezing point and raising of the boiling point are physical properties of solutions that depend on the quantity (concentration) but not on the kind or identity of the solute particles. Such properties are called colligative properties.
depend on the collective effect of the number of solute particles
Other colligative properties: vapor–pressure lowering and osmotic pressure
Vapor pressure lowering
A liquid in a closed container establishes equilibrium with its vapor. The vapor pressure is the pressure exerted by the vapor when it is at equilibrium with the liquid (that is, when the rate of vaporization equals the rate of condensation).
A substance that has no measurable vapor pressure is nonvolatile, whereas one that exhibits a vapor pressure is volatile.
A solution consisting of a volatile liquid solvent and a nonvolatile solute forms spontaneously because of the increase in entropy that accompanies their mixing
Ideally, the vapor pressure of a volatile solvent above a solution containing a nonvolatile solute is proportional to the solvent’s concentration in the solution. This relationship is expressed quantitatively by Raoult’s law, which states that
an ideal solution is defined as one that obeys Raoult’s law. implies total uniformity of interaction. The molecules in an ideal solution all influence one another in the same way—in other words, solute–solute, solvent–solvent, and solute– solvent interactions are indistinguishable from one another. Real solutions best approximate ideal behavior when the solute concentration is low and solute and solvent have similar molecular sizes and take part in similar types of intermolecular attractions.
If the solvent–solute interactions in a solution are weaker than either the solvent–solvent or solute–solute interactions, the vapor pressure tends to be greater than that predicted by Raoult’s law.
When the solute–solvent interactions in a solution are exceptionally strong, as might be the case when hydrogen bonding exists, the vapor pressure is lower than that predicted by Raoult’s law.
Boiling point elevation
The change in boiling point for a solution compared to the pure solvent is:
Kb is the molal boiling- point-elevation constant for the solvent (which is a proportionality constant that is experimentally determined for each solvent). Aka base dissociation constant
Vant hoffman factor- i. For nonelectrolyte always =1, for electrolyte depends on how many ions dissociate. (The limiting value of i can be determined for a salt from the number of ions per formula unit. for NaCl is it 2 because accounting for Na and Cl)
Freezing point depression
The freezing point of a solution is the temperature at which the first crystals of pure solvent form in equilibrium with the solution.
The proportionality constant Kf is the molal freezing-point-depression constant, analogous to Kb for boiling point elevation
Change in t f should be negative because the solution freezes at a lower temp than the pure solvent
Example in real life is how people put CaCl2 on roads so they are not icy
Osmosis
Semipermeable: means When in contact with a solution, these materials allow only ions or small molecules—water molecules, for instance—to pass through their network of tiny pores.
osmosis, the net movement of solvent is always toward the solution with the lower solvent (higher solute) concentration, as if the solutions were driven to attain equal concentrations.
Osmosis is the process of a solvent moving from one compartment to another, across a semipermeable membrane, toward higher solute concentration
Osmotic pressure: The pressure that must be applied to a solution to stop osmosis from pure solvent into the solution
Obeys a law similar to the ideal gas laws
M is the molarity of the solution for above
Osmotic pressure is a colligative property
If two solutions of identical osmotic pressure are separated by a semipermeable membrane, no osmosis will occur. The two solutions are isotonic with respect to each other.
If one solution is of lower osmotic pressure, it is hypotonic with respect to the more concentrated solution. The more concentrated solution is hypertonic with respect to the dilute solution.
Placing a red blood cell in a solution that is hypertonic relative to the intracellular solution (the solution inside the cells) causes water to move out of the cell. This causes the cell to shrivel, a process called crenation.
Placing the cell in a solution that is hypotonic relative to the intracellular fluid causes water to move into the cell, which may cause the cell to rupture, a process called hemolysis.
13.6 Colloids
Colloids, colloidal dispersions: mixtures containing particles larger than normal solutes but small enough to remain suspended in the dispersing medium
Although colloid particles may be so small that the dispersion appears uniform even under a microscope, they are large enough to scatter light. Consequently, most colloids appear cloudy or opaque unless they are very dilute.(ex: chunky milk)
This scattering of light by colloidal particles, known as the Tyndall effect, makes it possible to see the light beam of an automobile on a dusty dirt road, or the sunlight streaming through trees or clouds.
Not all wavelengths are scattered to the same extent. Colors at the blue end of the visible spectrum are scattered more than those at the red end by the molecules and small dust particles in the atmosphere. As a result, our sky appears blue. At sunset, light from the sun travels through more of the atmosphere; blue light is scattered even more, allowing the reds and yellows to pass through and be seen.
The most important colloids are those in which the dispersing medium is water. These colloids may be hydrophilic (“water loving”) or hydrophobic (“water fearing”)
A hydrophilic molecule folds in such a way that its hydrophobic groups are away from the water molecules, on the inside of the folded molecule, while its hydrophilic, polar groups are on the surface, interacting with the water molecules.
The hydrophilic groups generally contain oxygen or nitrogen and often carry a charge
Hydrophobic colloids can be dispersed in water only if they are stabilized in some way. Otherwise, their natural lack of affinity for water causes them to separate from the water. One method of stabilization involves adsorbing ions on the surface of the hydrophobic particles. (Adsorption means to adhere to a surface It differs from absorption which means to pass into the interior, as when a sponge absorbs water). the electrostatic repulsion between adsorbed ions on neighboring colloid particles keeps the particles from sticking together rather than dispersing in the water.
Hydrophobic colloids can also be stabilized by hydrophilic groups on their surfaces. Oil drops are hydrophobic, for example, and they do not remain suspended in water. Instead, they aggregate, forming an oil slick on the water surface.
Sodium stearate, or any similar substance having one end that is hydrophilic (either polar or charged) and one end that is hydrophobic (nonpolar), will stabilize a suspension of oil in water. Stabilization results from the interaction of the hydrophobic ends of the stearate ions with the oil drops and the hydrophilic ends with the water.
Colloidal particles in a solution undergo random motion as a result of collisions with solvent molecules. Because the colloidal particles are massive in comparison with solvent molecules, their movements from any one collision are very tiny. However, there are many such collisions, and they cause a random motion of the entire colloidal particle, called Brownian motion.