CHE-102
Lecture Outline
Chapter 13
Info Not in Syllabus

Lecture Recordings: 3 Buttons, 1 for each instructor
 In the Dr. Gulde’s TR Button:
1. 9:30am LIVE lecture
2. Dr. Gulde’s Pre-Lecture (PL) Videos button
- Please watch prior to coming to lecture
- Allows time in lecture for students to work sample problems
- Embedded questions are not graded
 Location: Classroom Lecture Recordings → Recorded Lectures →Dr. Gulde’s TR lecture
→Dr. Gulde’s PL-Videos

Lecture Material: Each instructor as a folder
 In the Dr. Gulde’s Lecture Material Folder:
1. CHE102-Gulde Schedule-SP2024: more detailed than syllabus & has PL-Video due
dates
2. Dr. Gulde’s Skeleton Outlines: to take notes if you wish
3. Gulde-Additional Problems: Extra problems for you to work
4. YouTube Videos: videos to help explain material

Please bring to EVERY class
 Clicker device
 Outlines
 Calculator
 Periodic table
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1
CHE-102
Lecture Outline
Chapter 13
Properties of Solutions
Solutions

Solutions – are ____________geneous (___________________________) mixtures
 Solute – usually present in _______________________ amounts, _____________ dissolving
 Solvent – usually present in ______________________ amounts, _____________ dissolving
•
We are familiar w/ __________________ but doesn’t have to be
Solution Equilibrium

Dissolution (dissolves) – ______________ breaks apart (into ____________________________)

(re-)crystallize - when dissolved __________________ comes ________ of soln
(forms a __________)

Solution Equilibrium – when the _______________ of these processes are
________________
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2
CHE-102
Lecture Outline
Chapter 13
Types of Solutions

Terms for when a: _______________ or _______________ dissolve in a ________________
1. Saturated solutions – Soln is in __________________________ w/undissolved solute
 What this means: contains ___________ amount of solute for a solvent, at a specific temp
Sugar: 25°C = 211g/100g H2O
50 °C = 260g/100g H2O
 If add more solute it _____________ dissolve
 SOLUBILITY – amount of _______________________ needed to form a saturated solution
 The solubility of sugar at 25°C is _________________
2. Unsaturated solutions – contains _____________ than max amount of solute, at a specific temp
 If add more solute it _____________ dissolve, until reach __________________________
3. Supersaturated solutions – contains _____________ than the max amount of solute, at a specific
temp
 To make: _____________ saturated solution, add _____________ then cool to original temp
•
_______g @ 25°C
•
Heat to ________
•
Add until _____________g dissolved
•
Cool to ___________
•
Not _______________
 If add more solute seed crystal”, excess _______________________, until
__________________________ reached (Alum exp. From CHE101)
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3
CHE-102
Lecture Outline
Chapter 13
More Solution Terminology
 Terms for when a: _______________ mixes with a _____________ or __________

Miscible – mixes _________________________

Immiscible – _______________________ mix significantly
 Ex: Oil & Water
WHY do Solutions form?
1. Entropy (∆S) – an increase in ________________________
(______________________________)
2. __________________ forces will ________________ each other
“_______________________________________________”!!
•
Polar solvents dissolve: ____________________________ solutes
•
Nonpolar solvents dissolve: ________________________ solutes
______ solution b/c forces
__________ similar
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4
CHE-102
Lecture Outline
Chapter 13
Type of Forces:
A. Intramolecular forces: forces b/t atoms _______________________ molecule
(aka _______________! ________________ lines)
 Ionic, Covalent, Metallic
B. Intermolecular forces (IMF): is b/t atoms ____________________________ molecules
(____________ lines)
 London dispersion, Dipole-Dipole, H-bonding, Ion-dipole
______molecular force
__________molecular force
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5
CHE-102
Lecture Outline
Chapter 13
Intermolecular Forces (IMF)
1. London Dispersion forces–movement of electrons w/in the electron cloud creating an
____________________________ dipole (_____________________________ partial charge)
 Dipole – distribution of electrons, one end has electron excess (______), the other electron
deficient (______)
 Can ___________________________ to _________dipole at any time
 Present between _____________ molecules
 London Forces Increase with:
a) increasing _________________________ (more ______________________ to distort)
b) _____________________________ shapes (more _____________________________)
2. Dipole-Dipole– attraction between neutral ______________________ molecules
 __________________________ dipole (
)
 __________________ than London
3. Hydrogen bonding – ________________ strong ___________________________ force
 Requires:
a) H atom bonded to ___________________ that is attracted to
b) lone pair of another __________________________ atom
4. Ion-dipole − between ions & partial charges (________________________) of
________________________ molecules
 Exist when _________________ compounds (ex. NaCl) ________________ in polar
substance (ex. water)
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6
CHE-102
Lecture Outline
Chapter 13
Sample: CH3Cl & H2O
1. List all the intermolecular forces associated with:
a) CH3Cl: (chloromethane)
b) H2O:
•
Dispersion?
•
Dispersion?
•
Dipole-dipole?
•
Dipole-dipole?
•
H-bonding?
•
H-bonding?
•
Ion-dipole?
•
Ion-dipole?
2. List all the IMFs between CH3Cl & H2O:
 Dispersion?
 Dipole-dipole?
 H-bonding?
 Ion-dipole?
3. Will a solution of CH3Cl and H2O form, are they miscible?
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7
CHE-102
Lecture Outline
Chapter 13
HOW do Solutions form?

__________________________________________ are broken & new ________________ created!

SOLVATION – when solute (ions or molecules) is _____________________________________
in a specific way
 HYDRATION – special case of solvation b/c __________________ = solvent
NaCl(s) in water

Water _________________
Water _____________
to Nacl
NaCl apart
There are 3 interactions involved in solution formation & each has an
_____________________ (∆H) change
+
−
Ex: NaCl(s) → Na (aq) + Cl (aq)
A. Solute−Solute interaction (∆Hsolute): must be ____________________
 ____________________ energy, ________thermic (____)
B. Solvent−Solvent interaction (∆Hsolvent): must be ____________________
 ____________________ energy, ________thermic (____)
C. Solute−Solvent interaction (∆Hmix): must be __________________
 ____________________ energy, ________thermic (____)

All 3 processes occur at the ______________________________
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8
CHE-102
Lecture Outline
Chapter 13

OVERALL enthalpy of Solution (∆Hsoln): _____________ all 3 processes

∆Hsoln can be endothermic (_____) or exothermic (______)
 Depends on size of __________

Solution formation is favored when ∆Hsoln is ___________thermic!
 Solute-Solvent particles are highly ______________________ to each other
 ___________ is large

Why does the disfavored _________thermic solution form?
(NH4NO3 ∆HSoln = +26kJ)
 ___________________, ∆S (_____________________) counteracts ______________
∆HSoln value
•
If ∆Hsoln is very __________________, _____________ soln forms
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9
CHE-102
Lecture Outline
Chapter 13
Solution Formation Review
1. Enthalpy (∆Hsoln): ____________________ (due to force interactions)
 Solutions usually form when ∆Hsoln ________________________
• Strong ___________________________________ interactions
2. Entropy (∆S): _________________________
 Solutions usually form when ∆S ______________________________
• Helps compensate for ∆Hsoln ________________________________________
Sample: Soln Forces CH3Cl & H2O
1. What processes are involved when CH3Cl dissolves in water?
Solute=
Solvent=
Question: Force Enthalpy
Label the following processes as exothermic or endothermic:
(a) Breaking solvent-solvent interactions to form separated particles
(b) Forming solvent-solute interactions from separated particles
(a)
(b)
A. Endothermic
Endothermic
B. Endothermic
Exothermic
C. Exothermic
Endothermic
D. Exothermic
Exothermic
Question: Force Enthalpy
From weakest to strongest, rank the following solutions
in terms of Solvent–Solute interactions:
NaCl in water,
butane (C4H10) in benzene (C6H6),
water in ethanol (C2H5OH)
A. NaCl in water < C4H10 in C6H6 < water in ethanol
B. Water in ethanol < NaCl in water < C4H10 in C6H6
C. C4H10 in C6H6 < water in ethanol < NaCl in water
D. C4H10 in C6H6 < NaCl in water < water in ethanol
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10
CHE-102
Lecture Outline
Chapter 13
Concentration of Solutions

Qualitatively: _______________________ term
 Insoluble
slightly soluble
soluble
very soluble
________/100gH2O
________/100gH2O
 Dilute vs. concentrated
•

The _______________ solute you have the ____________ concentrated the soln
Quantitatively: _______________________________ calculated
Concentration Units
1. Mass percent:
 Unit = ______________
 ____________________________ b/c also referred to as mass of the component in exactly
100g soln
• 7.2% by mass =
2. Mole fraction, X:
moles of A
 Unit = _______________
ΧA =
total moles of all components
 ________________________ indicates the component of interest
3. Molarity, M (molar soln):
moles of solute
 Unit = __________
M=
L of soln
 Most ________________ used
 ____________ to measure volume, but it changes with _________________
4. molality, m (molal soln):
 Unit = ___________
 ____________________________ of temp

m=
moles of solute
kg of solvent
Must know _____________________ of soln to convert b/t M & m.
 Memorize water density= __________________________________
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11
CHE-102
Lecture Outline
Chapter 13
Sample: Glycerol in Ethanol
1. Express the concentration of 50.0g glycerol (C3H8O3) in 60.0g of ethanol (C2H6O) if the density of
the solution is 1.10 g/mL, in a) %mass b) mole fraction
c) molality & d) molarity.
 Assume interested in ______________, unless specified
a) % mass:
b) mole fraction:
c) molality:
d) Molarity:
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12
CHE-102
Lecture Outline
Chapter 13
Sample: mass NaCl
2. If you have a 3.30% by mass NaCl solution, calculate the mass of NaCl in a 46.6g saline solution.
Sample: X NaCl
3. If you have a 33% by mass NaCl solution, calculate the mole fraction of NaCl.
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13
CHE-102
Lecture Outline
Chapter 13
Lecture Summary:
1. Mass Relations
2. Mole Fraction
3. Molarity, M
4. molality, m
Group: Glucose & H2O

What is the m & M concentration of a solution that has 49.8g glucose (C6H12O6) dissolved in
100.0kg of water?

NOTE:
 If water is solvent: M & m are ~ ________________ for _____________ solns
• Ex: 49.8g glucose (C6H12O6) in 100.0kg of water (previous problem)
 Because:
•
1 kg ______________________ ~ 1 kg __________________________
•
1 L ______________________ ~ 1 L __________________________
 So, the density of water can also be written ___________________________
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14
CHE-102
Lecture Outline
Chapter 13
Sample: m & X

Calculate the molality of methanol, CH3OH in water which methanol has a mol fraction of 0.133.
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15
CHE-102
Lecture Outline
Chapter 13
Factors Affecting Solubility (of solutes in water)
1. Structural Factors: Like Dissolves Like
a) Polarity: Solute solubility increases as _________________________________ increases
•
b/c more _______________________________
 _________________ dissolved better
b) Liquids: Solute solubility decreases as __________________________________________
increases
•
b/c solute becomes _____________________________________
•
Hexanol & Ethanol are both ___________________, but hexanol is ______________
c) Gases: main force is _________________ thus solubility
_____________________ with:
•
Increased _____________
•
Increased ______________________
Question: Vitamin

Which is more likely to be a water-soluble
vitamin?
A. Vitamin A
B. Vitamin C
C. They have the same solubility
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16
CHE-102
Lecture Outline
Chapter 13
2. Pressure: Strongly affects dissolved _________________, but __________ solids or liquids
 ____________ solubility increases as the partial pressure of the ______________________
the solvent increases
•
Increased pressure results in __________________________________ w/liquid
surface, thus more ___________________________________ by solvent
 Henry’s Law:
M = Solubility of gas (____________________)
P = Partial pressure of gas above solution (________)
k = Henry’s law constant (____________)
(varies w/temp & solute-solvent pair)
 Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 of
4.0atm over the liquid at 25°C. The Henry’s law constant for CO2 in water at this temp is
3.4x10-2 mol/L⋅atm.
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17
CHE-102
Lecture Outline
Chapter 13
3. Temperature: Depends on ________________:
 As temperature increases:
a) Solubility of most SOLID solutes in water ________________
(Solute=______________, Solvent==______________)
- b/c ______________ moves faster, able to form more
_______________________________________ forces
b) Solubility of most GAS solutes in water ________________
(Solute=______________, Solvent==______________)
- b/c ___________ moves faster, more particles able
to ___________ solvent
Question: Most Miscible
Which compound below is the MOST miscible with water?
A. CH3OH
B. CH3CH2OH
C. CH3CH2CH2OH
D. CH3CH2CH2CH2OH
E. They all dissolve the same
Question: Gas Solubility
In general, as the temperature of a solution increases, the solubility of a gas solute
A. increases
B. decreases
C. remains unchanged
Lecture Summary:
 Factors Affecting Solute Solubility in H2O:
1. Structure:
a) Polarity: Solubility ______ w/_______ polarity
b) Liquids: Solubility ______ w/________ carbon chain length
c) Gases: Solubility _______ w/_________ Mm & in polarity
2. Pressure: Affects _________ solutes, ________ liquids or solids
a) GAS solubility __________ w/__________ gas partial pressure
b) Henry’s Law:
3. Temperature: affects __________ of solute differently. If Temp increases…
a) Solid: solubility _______
b) Gas: solubility _______
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18
CHE-102
Lecture Outline
Chapter 13
Colligative Properties

Colligative property – depends on ____________________ but not the kind of particles
(_____________________ important)
1. Vapor Pressure lowering
2. Boiling point elevation
3. Freezing point depression
4. Osmotic pressure

van’t Hoff factor, i: compensates for ___________________________________ dissolved

Must consider what the _______________________ breaks up into?
 Non-Electrolytes: dissolve as entire ________________________________
• _______________________________ compounds, ex: Glucose:
• Estimated i = _________
 Typically ____________________________, but can be calculated
C6H12O6(s) water

→
 Strong Electrolytes: dissolve____________ into many ________________
• ___________________, Strong Acid or Strong Base
• Estimated i = _____________________
FeCl3(s)

→

water
Why is the calculated van’t Hoff factor (i) less than the estimated for _____________ compounds?
 Cations & anions can ____________________________ forming an ionic cluster
(aka an ___________________) which act as a __________________ particle
• Resulting in solute particle concentration being ______________________ the
estimated number
Ion pairing occurs ________________ as concentration of solute
_____________________
Ex:
NaCl:
FeCl3:
Estimated=
2
_____
Actual=
1.9
3.4
•
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19
CHE-102
Lecture Outline
Chapter 13
Question: Largest van’t Hoff

Which of the following has the largest van’t Hoff factor?
A. 1.00m Ethanol (CH3CH2OH)
B. 3.00m Glucose (C6H12O6)
C. 0.60m CaCl2
D. 0.50m K2SO4
E. 0.50m Fe2(SO4)3
Vaporization of PURE Liquid

Vaporization – lose _________ moving molecules (__________)
 Closed container: molecules trapped
• ____________ liquid converted to a gas
• ____________ gas converted to a liquid
- Reversible
 Establishes Equilibrium
• Equilibrium - when 2 reversible processes occur at the same
rate

Vaporization/Condensation rates are ________________
 Molecular level: molecules ___________________________ entering & leaving liquid
surface
• Amount of gas present is _________________________
- ___ always in gas
- ___ always in liquid
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20
CHE-102
Lecture Outline
Chapter 13
Vapor Pressure

Vapor pressure, (VP) – pressure exerted by a ___________ on a ___________ at equilibrium
 A ______________ vapor pressure means ________________________ molecules present
Vapor Pressure LOWERING

Terms to describe a SOLUTE
 Volatile – ______________ evaporates, ____________ Vapor Pressure (acetone)
 Nonvolatile – ______________ evaporates, ____________ Vapor Pressure (glycerin)

Adding a _____________________________ solute to a solvent ALWAYS
___________________ the vapor pressure of the SOLUTION
 ___________ b/t solvent & solute cause the ___________________ to be _____________
stable in the liquid state
•
_______________ solvent _______ particles produced

_______________ solute, the _______________ the SOLUTION VP
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21
CHE-102
Lecture Outline
Chapter 13
Raoult’s Law:
 ___________ solutions obey Raoult’s Law:
 Solute conc. is _____________
 Solute & solvent have similar ____________ & ______________
Psoln = VP of soln (w/___________________________ solute)
XSolvent = mole fraction
(
°
Psoln = (X solvent ) Psolvent
)
P°solvent = VP of __________________ solvent
nsolute = mols __________________________ dissolved
Example:

Which will lower the VP more when dissolved in water?
A. 3.0 mol NaCl(s)
B. 2.0 mol AlCl3(s)
 How many moles of solute would 3.0 moles of NaCl have?
 How many moles of solute would 2.0 moles of AlCl3 have?
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22
CHE-102
Lecture Outline
Chapter 13
Sample: Psoln NaCl & AlCl3

Assume a pure solvent has the following Vapor Pressure (P°solvent) & mole amounts (nsolvent), what is
the Psoln when NaCl &AlCl3 are added?
P°solvent = 14.0 atm
nsolvent = 6.0 mols
Psoln = ?
 Adding nonvolatile solute
____________________
the VP of the SOLUTION
 _________ solute particles
__________ VP more
Vapor Pressure Lowering

Can calculate the amount a solvent’s VP is _______________________:
 Calculate the Vapor Pressure Lowering for NaCl & AlCl3 in the previous problems.
•
NaCl: ∆P = 14.0atm – ______________ = ______________
•
AlCl3: ∆P = 14.0atm – ______________ = ______________
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23
CHE-102
Lecture Outline
Chapter 13
Boiling & Freezing Points

Phase diagram of Vapor Pressure Curves of a Pure Solvent (Black curve)
 Normal Boiling point – temp when liquid VP equals ____________ pressure (__________)
 Freezing point – temp when crystals of ________________________ form (___________)
•

________________________
Add nonvolatile solute, Vapor Pressure is ________________________!
 Blue line=Solution

(Pure Solvent = Black curve)
What happens to:
 Normal Boiling point – ________________ temp needed to reach 1atm
•
Temp _______________________!
 Freezing point – ____________________________ temp needed to form crystals of solvent
•
Temp _______________________________!
Boiling point ______________________
Freezing point ______________________
∆Tb = (i)(Kb)(m)
∆Tf = (i)(Kf)(m)
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24
CHE-102
Lecture Outline

i = van’t Hoff factor (# ___________________________________)

m = ___________________________ (mol _______________/kg_______________)

∆T (of ______________________) = always __________________
Chapter 13
 ∆Tb = _______________________
 ∆Tf = _______________________


Kb = Solvent molal boiling point elevation constant
Kf = Solvent molal freezing point depression constant

Memorize H2O Temps
Sample: MgCl2
1. Calculate the change in boiling point (in °C) when a solution of 275g of MgCl2 is added to 2.0kg of
water. The Kb of water is 0.51°C/m
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25
CHE-102
Lecture Outline
Chapter 13
Sample: Ethylene glycol
2. Ethylene glycol [EG, CH2(OH)CH2(OH)] is a common car antifreeze. It is water soluble and is a
nonvolatile nonelectrolyte (bp=197°C, d=1.11kg/L). Calculate the freezing point of a 30.0% (by
volume) solution.
3. The boiling point of pure ethanol (C2H5OH) is 78.4°C & it has a boiling point elevation constant of
1.22°C/m. If 38.96g of pentanol (a non-electrolyte) are added to 200.0g of ethanol to produce a
solution boiling temp of 81.1°C, what is the molar mass of pentanol?
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26
CHE-102
Lecture Outline
Chapter 13
Osmosis

Osmosis – net movement of ____________________ molecules across a
semipermeable membrane
 Semipermeable membrane – allows passage of certain _______________
molecules (water) but not large solute (molecules or ions) due to
_______________________________

Movement always occurs from less ______________________ (____________________)
→ more ________________________ (______________________________)
 As if to _____________________ the concentrations
•
Note: Bio uses high solvent → low solvent

As solvent moves, liquid levels become __________________ resulting in different
__________________________ on each end of the tube

Osmotic pressure (π) – pressure required to ____________ osmosis
 Similar to ____________________________ law
n
π = i  RT = (i )(M )(R )(T )
V 
π = Osmotic pressure (______)
i = _______________ particles in solution
V = volume (_______)
R = ideal gas constant (0.08206 L⋅atm/mol⋅K)
T = temperature (_____)
M = __________________ of particles (molecules or ions)
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27
CHE-102
Lecture Outline
Chapter 13
Osmosis Terms
A. Isotonic: soln of ____________________ concentration
 Solvent flow into cell ________________________ by flow out
• Cell size ___________________________
B. Hypertonic – soln of ______________ concentration
 If place red blood cells in Hypertonic soln:
•
Crenation – water _______________ the cell
− Cell size _______________________________
C. Hypotonic – soln of __________________ concentration
 If place red blood cells in Hypotonic soln:
•
Hemolysis – water _________________ the cell
− Cell ______________________________
Sample: Highest Freezing
1. Which of the following aqueous solutions will have the highest and lowest freezing point?
A. Pure water
B. 0.10m sucrose (C12H22O11)
C. 0.10m NaCl
D. 0.10m Na2SO4
Question: Highest Boiling
Which of the previous aqueous solutions will have the highest boiling point?
A.
B.
C.
D.
Pure water
0.10m sucrose (C12H22O11)
0.10m NaCl
0.10m Na2SO4
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28
CHE-102
Lecture Outline
Chapter 13
Group: Hemoglobin

A solution is prepared by dissolving 50.0g of hemoglobin, in enough water to make 500.0mL of
solution. The osmotic pressure of the solution is measured & found to be 28.6mmHg at 25°C.
Calculate the molar mass of hemoglobin (assume no change in volume when hemoglobin is added to
water).
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29
CHE-102
Lecture Outline
Chapter 13
Solutions Summary

Solution formation
 Entropy – increased disorder
 Enthalpy – favorable intermolecular interactions
•

Intermolecular forces of solute-solute, solvent-solvent & solute-solvent
Solubility
1. Influence of pressure – Henry’s Law
2. Influence of molecular structure – “like dissolves like”
•
Factors affecting solubility in water
3. Influence of temperature

Concentration
1. Mass %, ppm, ppb
2. Mole fraction
3. Molarity
4. Molality

Colligative properties
 van’t Hoff factor, i
1. Vapor Pressure – Raoult’s Law
2. Boiling pt. elevation
3. Freezing pt. depression
4. Osmosis
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30