chapter 12: solutions and their properties renee y. becker valencia community college

Chapter 12: Solutions and Their Properties

Renee Y. Becker

Valencia Community College

Introduction

1. A mixture is any intimate combination of two or more pure substances

2. Can be classified as heterogeneous or homogeneous

Heterogeneous

-The mixing of components is visually nonuniform and have regions of different

composition

Homogenous

-Mixing is uniform, same composition throughout

-Can be classified according to the size of their particles as either solutions or colloids

Solutions

Solution

1. Homogeneous mixtures

2. Contain particles with diameters in the range of 0.1–2 nm

3. Transparent but may be colored

4. Do not separate on standing

Colloids

1. Milk & fog

2. Diameters 2-500 nm

3. Do not separate on standing

Types of Solutions

Solution Formation

Solute • Dissolved substance, or smaller quantity substance

Solvent • Liquid dissolved in, larger quantity substance

Saturated solution• Contains the maximum amount of solute that will

dissolve in a given solvent.

Solution Formation

Unsaturated Contains less solute than a solvent has the

capacity to dissolve.

Supersaturated Contains more solute than would be present in a saturated solution.

Crystallization The process in which dissolved solute comes out

of the solution and forms crystals.

Energy Changes and the Solution Process Three Types of interactions

1. Solvent-solvent

2. Solvent-solute

3. Solute-solute

“Like dissolves like”

solutions will form when three types of interactions are similar in kind and magnitude

Energy Changes and the Solution Process

Example NaCl and water:Ionic solid NaCl dissolve in polar solvents like water

because the strong ion-dipole attractions between Na+ and Cl- ions and polar water molecules are similar in magnitude to the strong dipole-dipole attractions between water molecules and to the strong ion-ion attractions between Na+ and Cl- ions

Example Oil and waterOil does not dissolve in water because the two liquids

have different kinds of intermolecular forces. Oil is not polar or an ionic solvent


NaCl in H2O

1. Ions that are less tightly held because of their position at a corner or an edge of the crystal are exposed to water molecules

2. Water molecules will collide with the NaCl until an ion breaks free

3. More water molecules then cluster around the ion, stabilizing it by ion-dipole attractions

4. The water molecules attack the weak part of the crystal until it is dissolved

5. Ions in solution are said to be solvated they are surrounded and stabilized

by an ordered shell of solvent molecules


G, Free energy change

1. If G is negative the process is spontaneous, and the substance is dissolved

2. If G is positive the process is non-spontaneous, the substance is not dissolved

3. G = H -TS

H, enthalpy, heat flow in or out of the system, Hsoln heat of solution

S, entropy, disorder, Ssoln entropy of solution


Ssoln Entropy of Solution

Usually a positive number because when you dissolve something you are increasing disorder

Hsoln Heat of Solution

1. Harder to predict because it could be exothermic

(- Hsoln) or endothermic (+Hsoln)

2. The value of the heat of solution for a substance results from an interplay of the three kinds of interactions


1) Solvent-solvent interactions: Energy is required (endothermic) to overcome intermolecular forces between solvent molecules because the molecules must be separated and pushed apart to make room for solute particles

2) Solute-solute interactions: Energy is required (endothermic) to overcome interactions holding solute particles together in a crystal. For an ionic solid, this is the lattice energy. Substances with higher lattice energies therefore tend to be less soluble than substances with lower lattice energies.


3) Solvent-solute interactions: Energy is released (exothermic) when solvent molecules cluster

around solute particles and solvate them. For ionic substances in water, the amount of hydration

energy released is generally greater for smaller cations than for larger ones because water molecules can approach the positive nuclei of smaller ions more

closely and thus bind more tightly. Hydration energy generally increases as the charge on the

ion increases.


The solute–solvent interactions are stronger than solute–solute or solvent–solvent.

Favorable process

Exothermic rxn.

Exothermic -Hsoln


The solute–solvent interactions are weaker than solute–solute or solvent–solvent.

Unfavorable process.

Endothermic rxn

Endothermic +Hsoln


• Hydration– The attraction of ions for water molecules

• Hydration Energy– The energy associated with the attraction

between ions and water molecules

• Lattice energy– The energy holding ions together in a crystal

lattice

Example 1

Arrange the following in order of their expected increasing solubility in water:

Br2, KBr, C7H8

Example 2

1. Na+

2. Cs+

3. Li+

4. Rb+

1. Mg2+

2. Na+

3. Li+

Which would you expect to have the larger (more negative) hydration energy?

Units of Concentration Molarity (M)

M = mole of solute / Liter of solution

Molality (m)m = moles of solute/mass of solvent (kg)

Mole Fraction (x)X = mole of component / total moles

Units of Concentration

Mass Percent (mass %)

Mass % =

(mass of component / total mass of sol’n) x 100%

Parts per million, ppm =

(mass of component / total mass of solution) x 106

Parts per billion, ppb =

(mass of component / total mass of solution) x 109

Example 3

What is the mass % concentration of a saline sol’n prepared by dissolving 1.00 mol of NaCl in 1.00 L of water? DensityH2O=1.00 g/mL MMNaCl =

58.443 g/mol

Example 4

Assuming that seawater is an aqueous solution of NaCl what is its molarity? The density of seawater is 1.025 g/mL at 20C and the NaCl concentration is 3.50 mass %

• Assume 1 L to make easier, 1000 mL

Example 5

What is the molality of a solution prepared by dissolving 0.385 g of cholesterol, C27H46O in 40.0

g of chloroform, CHCl3? What is the mole

fraction of cholesterol in the solution?

Example 6

What mass in grams of a 0.500 m solution of sodium acetate, CH3CO2Na, in water would you use to obtain 0.150 mol of

sodium acetate?• Assume 1 kg

Example 7

The density at 20°C of a 0.258 m solution of glucose in water is 1.0173 g/mL and the molar mass of glucose is 180.2 g/mol. What is the molarity of the solution?

• Assume 1 kg

Example 8

The density at 20°C of a 0.500 M solution of acetic acid in water is 1.0042 g/mL. What is the concentration of this solution in molality? The molar mass of acetic acid, CH3CO2H, is 60.05

g/mol.• Assume 1 L

Some Factors Affecting Solubility Solubility

The amount of solute per unit of solvent needed to form a saturated solution

MiscibleMutually soluble in all proportions

Effect of Temperature on Solubility

1. Most solid substances become more soluble as temperature rises2. Most gases become less soluble as temperature rises

Some Factors Affecting Solubility

Effect of Pressure on Solubility

1. No effect on liquids or solids

2. The solubility of a gas in a liquid at a given temperature is directly proportional to the partial pressure of the gas over the solution, @ 25°C

Henry’s Law solubility = k x P

k = constant characteristic of specific gas, mol/L atm

P = partial pressure of the gas over the sol’n

Some Factors Affecting Solubility

a) Equal numbers of gas molecules escaping liquid and returning to liquid

b) Increase pressure, increase # of gas molecules returning to liquid, solubility increases

c) A new equilibrium is reached, where the #’s of escaping = #

of returning

Example 9

Which of the following will become less soluble in water as the temperature is increased?

1) NaOH(s)

2) CO2(g)

Example 10

The solubility of CO2 in water is 3.2 x 10-2 M @

25°C and 1 atm pressure. What is the Henry’s-Law constant for CO2 in mol/L atm?

solubility = k x P

Physical Behavior of Solutions: Colligative Properties Colligative properties

Properties that depend on the amount of a dissolved solute but not its chemical identity

There are four main colligative properties:

1. Vapor pressure lowering

2. Freezing point depression

3. Boiling point elevation

4. Osmotic pressure

Physical Behavior of Solutions: Colligative Properties

In comparing the properties of a pure solvent with those of a solution…

1. Vapor pressure of sol’n is lower

2. Boiling point of sol’n is higher

3. Freezing point of sol’n is lower

4. Osmosis, the migration of solvent molecules through a semipermeable membrane, occurs when solvent and solution are separated by the membrane

Vapor-pressure Lowering of Solutions: Raoult’s Law 1. A liquid in a closed container is in equilibrium with its

vapor and that the amount of pressure exerted by the vapor is called the vapor pressure.

2. When you compare the vapor pressure of a pure solvent with that of a solution at the same temperature the two values are different.

3. If the solute is nonvolatile and has no appreciable vapor pressure of its own (solid dissolved) the vapor pressure of

the solution is always lower that that of the pure solvent.

4. If the solute is volatile and has a significant vapor pressure (2 liquids) the vapor pressure of the mixture is intermediate

between the vapor pressures of the two pure liquids.

Solutions with a Nonvolatile Solute When solute molecules displace solvent molecules at the

surface, the vapor pressure drops since fewer gas molecules are needed to equalize the escape rate and capture rates at the liquid surface.

Solutions with a Nonvolatile Solute!!!

Raoult’s Law Psoln = Psolv · Xsolv

Psoln = vapor pressure of the solution

Psolv = vapor pressure of the pure solvent

Xsolv = mole fraction of the solvent in the

solution

Raoult’s Law applies to only Ideal solutions

1. Law works best when solute concentrations are low and when solute and solvent particles have similar intermolecular forces.

2. If intermolecular forces between solute particles and solvent molecules are weaker than solvent molecules alone, solvent molecules are less tightly held, vapor pressure is higher than Raoult predicts

3. If intermolecular forces between solute and solvent are stronger than solvent alone, solvent molecules are more tightly held and the vapor pressure is lower than

predicted

4. No Van’t Hoff factor!!!

Example 11

What is the vapor pressure (in mm Hg) of a solution prepared by dissolving 5.00 g of benzoic acid (C7H6O2) in 100.00 g of ethyl alcohol (C2H6O) at

35°C? The vapor pressure of the pure ethyl alcohol at 35°C is 100.5 mm Hg

Example 11Psoln = Psolv · Xsolv

Solutions with a Nonvolatile Solute

Close-up view of part of the vapor pressure curve for a pure solvent and a solution of a nonvolatile solute. Which curve represents the pure solvent, and which the solution?

Why?

• The lower vapor pressure of a sol’n relative to that of a pure solvent is due to the difference in their entropies of vaporization, Svap. Because the entropy of the solvent

in a sol’n is higher to begin with, Svap is smaller for the

sol’n than for the pure solvent. As a result vaporization of the solvent from the sol’n is less favored (less negative Gvap), and the vapor pressure of the solution is

lower.

Solutions with a Volatile Solute!! Ptotal = PA + PB

Ptotal = (P°A · XA) + (P°B · XB)

P°A = vapor pressure of pure A

XA = mole fraction of A

P°B = vapor pressure of pure B

XB = mole fraction of B

Ptotal should be intermediate to A & B

Close-up view of part of the vapor pressure curves for two pure liquids and a mixture of the two. Which curves represent the mixture?

1) Red

2) Green

3) blue

Example 12

What is the vapor pressure ( in mm Hg) of a sol’n prepared by dissolving 25.0 g of ethyl alcohol (C2H5OH) in 100.0 g of water at 25°C? The vapor

pressure of pure water is 23.8 mm Hg and the vapor pressure of ethyl alcohol is 61.2 mm Hg at 25°C

Example 13

a) Is the boiling point of the second pure liquid higher or lower than that of the first liquid?

b) On the diagram where is the approximate position of the

second pure liquid?

The following phase diagram shows part of the vapor pressure curves for a pure liquid (green curve) and a solution (red curve) of the first liquid with a second volatile liquid (not shown)

Boiling Point Elevation and Freezing Point Depression of Solutions


1. Red line is pure solvent

2. Green line solution of nonvolatile solute

3. Vapor pressure of sol’n is lower

4. Temp at which vapor pressure = 1 atm for sol’n is higher

5. Boiling point of sol’n is higher by Tb

6. Liquid/vapor phase transition line is lower for sol’n

7. Triple point temp is lower for sol’n

8. Solid/liquid phase transition has shifted to a lower temp.

9. The freezing point of the sol’n is lower by Tf


Tb = Kb · m

Tf = Kf · m

Kb = molal boiling-point elevation constant

Kf = molal freezing-point depression constant

m = molality

NO Van’t Hoff Factor!!


The higher boiling point of a solution relative to that of a pure solvent is due to a difference in their entropies of vaporization, Svap. Because the solvent in a solution has a

higher entropy to begin with, Svap is smaller for the solution

than for the pure solvent. As a result, the boiling point of the solution Tb is higher than that of the pure solvent.


The lower freezing point of a solution relative to that of a pure solvent is due to a difference in their entropies of fusion, Sfusion. Because the solvent in a solution has a

higher entropy level to begin with, Sfusion is larger for the

solution than for the pure solvent. As a result the freezing point of the solution Tf is lower than that of the pure

solvent.

Example 14

What is the normal boiling point in °C of a solution prepared by dissolving 1.50 g of aspirin (C9H8O4)

in 75.00 g of chloroform (CHCl3)? The normal

boiling point of chloroform is 61.7 °C and Kb of

chloroform is 3.63 °C kg/mol

Osmosis and Osmotic Pressure 1. Semipermeable membranes allow water or other

small molecules to pass through, but they block the passage of large solute molecules or ions.

2. When a solution and a pure solvent are separated by the right kind of semipermeable membrane, solvent molecules pass through the membrane in a process known as osmosis.

3. Passage of solvent through the membrane takes place in both directions

Osmosis and Osmotic Pressure

4. Passage from the pure solvent side to the solution side is more favored and faster.

5. The amount of liquid on the pure solvent side decreases

6. The amount of liquid on the solution side increases

7. The concentration of the solution decreases


Osmotic Pressure

1. The amount of pressure necessary to achieve equilibrium

2. = MRT

= osmotic pressure

M = molarity

R = gas constant, .08206 L atm/K mol

T = temperature in kelvins

Example 15

What osmotic pressure in atm would you expect for a solution of 0.125 M C6H12O6 that is separated

from pure water by a semipermeable membrane at 310 K?

= MRT

Example 16

A solution of unknown substance in water at 300 K gives rise to an osmotic pressure of 3.85 atm. What is the molarity of the solution?

= MRT

M = /RT

Some uses of Colligative Properties

1. The most important use of colligative properties in the laboratory is for determining the molecular mass of an unknown substance.

2. Any of the four colligative properties can be used but using osmotic pressure gives the most accurate results

Example 17

• What is the molar mass of sucrose if a solution prepared by dissolving 0.822 g of sucrose in water and diluting to a volume of 300.0 mL has an osmotic pressure of 149 mm Hg at 298 K?

If you have to calculate the molar mass of a compound which of the following will give you the most accurate results?

1) Osmotic pressure

2) Vapor pressure lowering

3) Boiling point elevation

4) Freezing point depression

chapter 12: solutions and their properties renee y. becker valencia community college

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