Chapter 13

Properties of Solutions

Warm - Up

• Why doesn’t salt dissolve in nonpolar solvents such as hexane?

• How does the orientation of water around Na+ differ from the orientation of water around Cl- during solvation?

• What is a hydrate?

Solutions

• One or more substances (solute)

disperse uniformly throughout the

solvent (the substance in greatest

amount).

Hold On a Sec!

Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.

Student, Beware!

• Dissolution is a physical change—you can get back the original solute by evaporating the solvent.

• If you can’t, the substance didn’t dissolve, it reacted.

Entropy and Dissolving

• Some ionic compounds dissolve EXOthermically (releasing energy (heat))

• These are SPONTANEOUS processes and are energetically “downhill”

• Most ionic compounds dissolve ENDOthermically (requiring energy) because those ionic bonds don’t want to break up. BUT they do happen because of ENTROPY!!!

Entropy

• Nature is happier with randomness and chaos. When ionic compounds break up, this is favorable ENTROPICALLY, so it still happens.

• If CaCl2 dissolves in water, what happens? Is this endo or exothermic? How do IMF’s come into play?

PRACTICE EXERCISE

Does the entropy of the system increase or decrease when the stopcock is opened to allow mixing of the two gases in this apparatus?

Solubility and Saturation

• Solubility: amount of solute needed to form a saturated solution in a given amount of solvent.

• Unsaturated: all solid dissolves and there’s room for more!

Supersaturation and Crystallization

• If you heat a solution you can get a supersaturated solution. Once cooling happens the substances are LESS soluble and form crystals.

How do we know if something will be soluble??

• Like dissolves like

• Substances with similar IMF’s will be soluble in one another.

• If something is polar it can interact with another polar molecule in a way that will solubilize it!

• Something nonpolar doesn’t have the IMF’s needed to dissolve in a polar solvent.

SAMPLE EXERCISE 13.2 Predicting Solubility Patterns

Predict whether each of the following substances is more likely to dissolve in carbon tetrachloride (CCl4) or in water: C7H16, Na2SO4, HCl, and I2.

Arrange the following substances in order of increasing solubility in water:

Vocab

• Miscible : pairs of liquids that dissolve in one another (EtOH and Water)

• Immiscible : liquids that don’t dissolve (oil and water)

Gases in Solution

• In general, the solubility of gases in water increases with increasing mass.

• Larger molecules have stronger dispersion forces.

Gases in Solution

• The solubility of liquids and solids does not change appreciably with pressure.

• The solubility of a gas in a liquid is directly proportional to its pressure.

Henry’s Law Sg = kPg

where

• Sg - solubility of the gas;

• K - Henry’s law constant

• Pg - partial pressure of the gas above the liquid.

Calculate the concentration of CO2 in a soft drink after the bottle is opened and equilibrates at 25°C under a CO2 partial pressure of 3.0 10–4 atm. The Henry’s law constant for CO2 in water at this temperature is 3.1 10–2 mol/L-atm.

Question

• Why do you only see the soda bubbles after you open the bottle???

Temperature

Generally, the solubility of solid solutes in liquid

solvents increases with increasing temperature.

Temperature

• The opposite is true of gases:

Carbonated soft drinks are more “bubbly” if stored in the refrigerator.

Warm lakes have less O2 dissolved in them than cool lakes.

Ways of Expressing Concentration Concentration Abv Definition Example

Mass Percent % (Mass of solute/mass of solution ) x 100

14% EtOH(aq) = 14g EtOH/100g water

Parts per million ppm 106 x mass of solute/ mass of solution

18 ppm Pb2+ in water = 18mg Pb2+ /1000 g

solution

Mole Fraction c Moles of one component/ total

moles

XNH3 = 0.10 One tenth of all the

moles is NH3

Molarity M Moles of solute / liters of solution

.15M HCl = .15 moles per liter of water.

Molality m Moles of solute/ kg of solvent

.20m NaCl = 0.2 moles per kg of water.

Warm Up (use table you copied Tuesday)

(a) A solution is made by dissolving 13.5 g of glucose (C6H12O6) in 0.100 kg of water. What is the mass percentage of solute in this solution? (b) A 2.5-g sample of groundwater was found to contain 5.4 g of Zn2+ What is the concentration of Zn2+ in parts per million?

(b) A solution is made by dissolving 4.35 g glucose (C6H12O6) in 25.0 mL of water at 25°C. Calculate the molality of glucose in the solution.

Agenda

• How do you make the perfect airbag?

• Solutions problems

• Air bag lab

(c) An aqueous solution of hydrochloric acid contains 36% HCl by mass. (i) Calculate the mole fraction of HCl in the solution. (ii) Calculate the molality of HCl in the solution.

(d) A solution contains 5.0 g of toluene (C7H8) and 225 g of benzene and has a density of 0.876 g/mL. Calculate the molarity of the solution.

Now try these

Warm Up! • A solution is prepared by dissolving 16.2 g of

benzene (C6H6) in 282 g of carbon tetrachloride The concentration of benzene in this solution is __________ molal. The molar masses of and are 78.1 g/mol and 154 g/mol, respectively.

• A solution is prepared by adding 1.43 mol of solid KCl to 889 g of water. The concentration of KCl is __________ molar.

• The concentration of KCl in a solution prepared by adding 0.0660 mol of KCl to 1.00 mol of water is __________ % by mass.

Agenda

• What are colligative properties?

• How behavior of solutions differ from the behavior of the solvent

• Raoults Law

• Kf and Kb

• HW Set 2: 62 63-75 odds, 81, 85,89

Colligative Properties • Changes in colligative properties depend only

on the number of solute particles present, not on the identity of the solute particles.

• Volatile – substance with a measurable equilibrium VP. Non-volatile has no measurable VP.

• Among colligative properties are Vapor pressure lowering Boiling point elevation Melting point depression Osmotic pressure

Solution has LOWER VP

Solute inhibits the escape of solvent

Raoult’s Law

PA = XAPA where

• XA is the mole fraction of compound A

• PA is the normal vapor pressure of A at that temperature

NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

Raoult’s Law

Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25°C. Calculate the vapor pressure at 25°C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25°C is 23.8 torr (Appendix B).

PA = XAPA

The vapor pressure of pure water at 110°C is 1070 torr. A solution of ethylene glycol and water has a vapor pressure of 1.00 atm at 110°C. Assuming that Raoult’s law is obeyed, what is the mole fraction of ethylene glycol in the solution?

Boiling Point Elevation and Freezing Point Depression

Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

Boiling Point Elevation

The change in boiling point is proportional to the molality of the solution:

Tb = Kb m

where Kb is the molal boiling point elevation constant, a property of the solvent.

Tb is added to the normal boiling point of the solvent.

Boiling Point Elevation and Freezing Point Depression

Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved.

Tb = Kb m

Tf = Kf m

Practice

Automotive antifreeze consists of ethylene glycol (C2H6O2), a nonvolatile nonelectrolyte. Calculate the boiling point and freezing point of a 25.0 mass % solution of ethylene glycol in water.

Calculate the freezing point of a solution containing 0.600 kg of CHCl3 and 42.0 g of eucalyptol (C10H18O), a fragrant substance found in the leaves of eucalyptus trees. (FPCHCl3 = -63.5 °C)

Calculating Molar Mass

A solution of an unknown nonvolatile electrolyte

was prepared by dissolving 0.250 g of the

substance in 40.0 g of CCl4. The boiling point of

the resultant solution was 0.357°C higher than that

of the pure solvent. Calculate the molar mass of

the solute.

Colligative Properties of Electrolytes

Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

Colligative Properties of Electrolytes

However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M solution of methanol does.

Van’t Hoff Factor

• i, a measure of the extent to which electrolytes dissociate.

• Ideal value for i is the number of ions/formula unit

• NaCl, CaCl2, Al(NO3)3, Fe2(SO4)3 what’s i?

van’t Hoff Factor not always ideal

One mole of NaCl in water does not really give rise to two moles of ions.

van’t Hoff Factor

Some Na+ and Cl− reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

The van’t Hoff Factor

• Reassociation is more likely at higher concentration.

• Therefore, the number of particles present is concentration dependent.

The van’t Hoff Factor

We modify the previous equations by multiplying by the van’t Hoff factor, i

Tf = Kf m i

Osmosis

• Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles.

• In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

Osmosis

In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).

Osmotic Pressure

• The pressure required to prevent osmosis.

n

V = ( )RT = MRT

where M is the molarity of the solution

If the osmotic pressure is the same on both sides

of a membrane (i.e., the concentrations are the

same), the solutions are isotonic.