Chem 2 - Chapter 12: Solutions

Front 1

Types of Solution

Back 1

gas in gas (air)
gas in liquid
gas in solid
liquid in liquid
liquid in solid
solid in liquid
solid in solid

** Attractive forces are an important factor in the formation of solutions (except for gas solutions)

Front 2

Solute

Back 2

the smaller portion of the solution which dissolves;

the dissolved substance

Front 3

Solvent

Back 3

represents the larger portion of the solution;

the liquid

Front 4

The Dissolution Process (the two perspectives)

Back 4

Macroscopic Perspective vs. Microscopic Perscpective

Front 5

Macroscopic Perspective

Back 5

1. Salts dissolve in a solvent
2. Become part of the solution

Front 6

Microscopic Perspective

Back 6

1. The solvent interacts with the solute
2. The ions are separate in the solution

Front 7

Entropy (H)

Back 7

The pervasive tendency for all kinds of energy to disperse (spread out) when it is not restrained from doing so;
The disorder or randomness in a system

There is no chance in the kinetic energy of the two gases

Front 8

Intramolecular Forces

Back 8

bonding forces;

exist WITHIN a molecule and influences the CHEMICAL properties of the substance

Front 9

Intermolecular Forces

Back 9

non-bonding forces;

exist BETWEEN molecules and influence the physical properties of the substance

Front 10

Effect of Intermolecular Forces:

Back 10

In the absence of intermolecular forces, two substances will mix spontaneously and form a solution

Intermolecular forces exist between solute molecules themselves, solvent molecules themselves, or solute and solvent Molecules
(Solute-Solute, Solvent-Solvent, or Solute-Solvent)

Front 11

Types of Intermolecular Forces:

Back 11

1. Ion-Dipole
2. Hydrogen Bond
3. Dipole-Dipole
4. Dipole-Induced Dipole
5. London-Dispersion

** Even though IMF's are diverse because of the different types that exist, they are relatively weak forces.

Front 12

Ion-Dipole

Back 12

The IMF that exists between an ion and a polar compound;

The Strongest of the IMF's

Ex:
NaCl(s) --> Na+(aq) + Cl- (aq)
-- with the addition of H2O

*The larger the charges, the stronger the force
(CaS is stronger than NaBr even though both are ion-dipole because it has an overall charge of 4 and NaBr only has a charge of 1)

Front 13

Hydrogen Bond

Back 13

the IMF that exists when hydrogen is directly connected to F, O, or N.

The 2nd Strongest IMF

Ex:
1. HF
2. H2O
3. NH3

Front 14

Dipole-Dipole

Back 14

the IMF that exists when two polar-covalent compounds interact

3rd Strongest IMF

Ex:
1. CH3F
2. CH2Br2
3. PH3

Front 15

Dipole/Induced-Dipole

Back 15

the IMF that exists when a non-polar covalent compound interacts with a polar covalent compound

4th Strongest IMF

Ex: P4 in PH3

Front 16

London-Dispersion

Back 16

the IMF that exists when two non-polar covalent compounds interact

Weakest IMF

Front 17

Cases when a Compound is Non-Polar:

Back 17

1. A central element is connected to the SAME elements
and has NO lone pairs
---Ex: CCl4 (see clutch notes for diagram)
CO2 (O=C=O)

2. If compound has ONLY Carbons and Hydrogens
---Ex: CH4, C6H6, or C5H12

3. If a non-metal is By Itself or connected to Copies of itself
---Ex: C(graphite), O2, Cl2, P4

Front 18

Rule of Thumb for Determining Solubility:

Back 18

"Like Dissolves Like"

Front 19

IMF effects of Melting Point, Boiling Point, and Surface Tension

Back 19

the stronger the IMF, then the higher the melting poing, boiling point, and Surface Tension

Front 20

Weights Affect on IMF

Back 20

The less a substance weighs, the weaker it is

Front 21

IMF Effect on Vapor Pressure

Back 21

The weaker the IMF, the higher the Vapor Pressure

Front 22

Energy Changes and Dissolution:

Back 22

Dissolution is either Exothermic or Endothermic

heat is either gained or lost during dissolution

Front 23

Energetics of Solution Formation (3 Steps):

Back 23

1. Separating the Solute into its constituent Particles

2. Separating the solvent particles to make room for the solute particles

3. Mixing the solute and solvent particles

Front 24

Energetics of Solution Formation (Equation):

Back 24

According to Hess's Law, the enthalpy of solutions is given by :

ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmixture
(endo, +) (endo, +) (exothermic, -)

Front 25

ΔHsolution = 0

Back 25

The sum of the endothermic terms is about equal to the magnitude of the exothermic term

Increasing entropy upon mixing drives the process.

Overall energy of the system remains constant

Front 26

ΔHsolution is Negative (exothermic)

Back 26

the sume of the endothermic terms is smaller in magnitude than the exothermic term

Exothermic

both tendency towards a lower energy and higher entropy drive the process

Front 27

ΔHsolution is Positive (endothermic)

Back 27

The sum of the endothermic terms is greater in magnitude than the exothermic term

As long as ΔHsolution is not too large, entropy will drive the process.
If ΔHsolution is too large, no solution is formed.

Front 28

Heat of Hydration

Back 28

the enthalpy change that occurs when 1 mol of gaseous solute ions are dissolved in water

Largely negative for ionic compounds

applicable ONLY to aqueous solutions

ΔHsolvent + ΔHmixture = ΔHhydration

Front 29

Solubility and Solution Equilibrium

Back 29

The solubility of a substance is the amount of the substance that will dissolve in a given amount of solvent.

The solubility of one substance in another is determined by entropy (nature's tendency towards mixing) and the IMF's present

Front 30

Saturated

Back 30

the maximum amount of dissolved solute is present in the solvent

When the rate of dissolution = the rate of deposition
(dynamic equilibrium has been reached)

Front 31

Unsaturated

Back 31

additional amounts of solute can be further dissolved in the solvent

As the solution becomes more concentrated, some of the sodium and chloride ions can begin to redeposit as solid NaCl

Front 32

Supersaturated

Back 32

more than the equilibrium concentration of solute has been dissolved

Front 33

Effect of Temperature on Solubility of Gases:

Back 33

1. Most salts increase in solubility as Temperature Increases (some decrease)

2. Most gases decrease in solubility as temperature increases

Front 34

Solubility of Gases in Water:

Back 34

The solubility of a gas depends on pressure

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

Gas + Solvent <=> solution
--Increasing the gas pressure stresses the equilibrium, and shifts it right
--More gas in solution

Front 35

Summary of Solubility of Gases in Water:

Back 35

1. Equilibrium: solvent solution <=> solvent gas phase
2. Reducing the volume increases the pressure
3. Equilibrium re-established after some of excess gas phase species enters solution phase

Front 36

Henry's Law:

Back 36

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

Sgas = Kh Pgas

Assumption: Low pressure, low concentrations, and a gas which does not react with the solvent

Front 37

Calculation Form of Henry's Law:

Back 37

C1/P1 = C2/P2

Front 38

Units of Concentration:

Back 38

1. Molarity
2. Mole Fraction
3. Mass Percent
4. Percent per Million/ per Billion
5. Parts by Volume
6. Molality

Front 39

Molarity (M):

Back 39

moles of solute dissolved per liters of solution

M = moles of solute/liters of solution

Problem:
It is Temperature Dependent

** Solubility = Molarity = Concentration

Front 40

Mole Fraction (X):

Back 40

Xcomponent = moles of component/ total moles making up the solution

Xsolute = moles of solute/total moles of solute + solvent

NOT temperature dependent
NO units
Dimensionless because its a ration

**Mol% = Mole fraction x 100%

Front 41

Mass Percent:

Back 41

Mass Percent = (mass of component/total mass of solution) x 100

Front 42

Percent Per Million (ppm):

Back 42

ppm = (mass of component/total mass of solution) x 10^6

Front 43

Parts Per Billion (ppb):

Back 43

ppb = (mass of component/total mass of solution) x 10^9

Front 44

Parts by Volume:

Back 44

parts by volume = (Volume of solute/volume of solution) x multiplication factor

Front 45

Molality (m):

Back 45

moles of solute dissolved per kg of solvent

molality = moles of solute/mass of solvent (kg)

Temperature INdependent

Front 46

Henry's Law:

Back 46

Explains the relationship between gas pressure and solubility:
--The solubility of a gas (Sgas) is directly proportional to the partial pressure of the gas (Pgas) above the solution.

Sgas = Kh x Pgas

Kh = Henry's Constant
Pgas = partial pressure of that gas

Front 47

Colligative Properties:

Back 47

Help to explain what happens to a pure solvent as we add solute to it

Front 48

Types of Colligative Properties:

Back 48

1. Vapor Pressure
2. Boiling Point
3. Freezing Point
4. Osmotic Pressure

Front 49

As Solute is Added to a pure Solvent:

Back 49

Boiling Point and Osmotic Pressure INCREASE

Freezing Point and Vapor Pressure DECREASE

Front 50

Vapor Pressure:

Back 50

The VP of a liquid is the pressure exerted by its vapor when the liquid and the vapor states are in dynamic equilibrium

Depends on the IMF's present and the temperature

Front 51

Dynamic Equilibrium:

Back 51

Rate of Evaporation = Rate of Condensation

Front 52

Vapor Pressure Lowering of Solutions:

Back 52

A solution always evaporates more slowly than a pure solvent does, because its VP is lower and therefore its molecules escape less easily

Front 53

Raoult's Law:

Back 53

The VP of a solution containing a non-volatile solute is = to the VP of the pure solvent times the mole fraction of the solvent

Psolution = Xsolvent Psolvent

Psolvent = the partial pressure of the PURE solvent

Front 54

Ptotal = ?

Back 54

Ptotal = PA + PB

PA = XA x PA (pure solvent)
PB = XB x PB (pure solvent)

so: Ptotal = XAPA (pure) + XBPB (pure)

Front 55

Colligative Properties:

Back 55

Depends on the amount of dissolved solute (i.e. number of solute particles) but NOT on its chemical identity.
The more concentrated the solution, the more the effect

Front 56

Freezing Point Depression:

Back 56

ΔTf = Kf x m x i

ΔTf always = a positive value

Kf = constant
m = molality
i = number of ions in the solute

Front 57

Boiling-Point Elevation:

Back 57

ΔTb = Kb x m x i

Front 58

Osmosis:

Back 58

the migration of solvent molecules through a semi-permeable membrane.

Occurs when the solvent and solution are separated by a membrane

The flow of solvent from a solution of lower solute concentration to one of higher solute concentration, through a semi-permeable membrane until equilibrium is reached

Front 59

Osmotic Pressure:

Back 59

Π = MRT*i

Π = Osmotic Pressure
M - molarity
R = gas constant (0.08206 L*atm/mol*K
T = Temperature in K

If external pressure is applied this osmosis can be stopped or even reversed
The pressure required to stop the osmotic flow is called osmotic pressure

Front 60

van't Hoff Factor (i):

Back 60

i = moles of particles in solution/moles of solute dissolved

Salts don't really dissociate at 100%

add i to all of the colligative properties equations for all ionic solutions