exam 1 chem - patterns

Yes — this is probably the most useful study guide to memorize before you start grinding calculation problems. I’ll focus on the patterns, trends, cause-and-effect relationships, comparisons, and “if this increases, what happens to that?” questions from Chapters 11–13.

🧪 CHEM EXAM 1 — CONCEPTUAL PATTERNS & TRENDS STUDY GUIDE

⭐ 1. The BIGGEST Pattern: Intermolecular Forces

A lot of Chapter 11 can be reduced to one idea:

How strongly are the molecules attracted to each other?

As intermolecular forces get STRONGER:

As intermolecular forces get WEAKER:

🧠 Memorize this chain:

Strong IMF

→ molecules stick together more
→ harder to separate
→ harder to evaporate
→ LOW vapor pressure
→ LOW volatility
→ need more heat to boil
→ HIGH boiling point

At the same time:

→ molecules resist flowing
→ HIGH viscosity

→ molecules at the surface are pulled inward strongly
→ HIGH surface tension

This one chain answers a TON of conceptual questions.

2. Types of Intermolecular Forces

Your three major ones:

London dispersion

Present in ALL atoms and molecules.

Dipole-dipole

Present between polar molecules.

Hydrogen bonding

Special dipole-dipole interaction when:

H is bonded directly to N, O, or F.

Think:

H–N, H–O, H–F

Quick identification pattern

Every molecule: dispersion

Polar molecule: dispersion + dipole-dipole

H attached to N/O/F: dispersion + dipole-dipole + hydrogen bonding

3. Dispersion Force Patterns

Dispersion forces depend on how easily the electron cloud can be distorted — polarizability.

Your slides emphasize:

More electrons → greater polarizability → stronger dispersion forces

So, generally:

larger electron cloud / more electrons

→ stronger dispersion
→ stronger IMF
→ higher BP

Molecular shape matters too!

Your slides specifically show that molecular shape affects dispersion forces.

More opportunity for molecules to contact each other can produce stronger attractions than a more compact structure.

So when comparing similar molecules, don't ONLY look at mass.

Think:

electrons + molecular shape

4. Polarity Pattern

This is another big exam area.

A molecule having polar bonds does NOT necessarily mean the entire molecule is polar.

You must look at:

bond polarity + molecular geometry

Example: CO₂

O=C=O

The bonds are polar.

But the molecule is symmetrical.

The dipoles cancel.

➡ Nonpolar

BF₃

Symmetrical trigonal planar.

Dipoles cancel.

➡ Nonpolar

H₂O

Bent.

Dipoles do NOT cancel.

➡ Polar

🧠 Pattern:

Symmetrical molecule + identical surrounding atoms

→ dipoles may cancel
→ molecule can be nonpolar

Asymmetrical molecule

→ dipoles are less likely to cancel
→ molecule can be polar

Don't use symmetry alone blindly, though—look at the actual molecular geometry and dipoles.

5. Polarity → IMF → Physical Property

Expect questions that require multiple steps.

For example:

Which substance has the higher boiling point?

Don't immediately think about boiling point.

Work backward:

Step 1

Determine polarity.

Step 2

Determine intermolecular forces.

Step 3

Compare IMF strength.

Step 4

Translate IMF strength into the requested property.

Example:

CH₃OH

has O–H

→ hydrogen bonding

→ stronger IMF

→ harder to separate molecules

→ higher BP / lower VP / higher viscosity / higher surface tension

This reasoning pattern is more important than simply memorizing individual compounds.

6. Boiling Point vs. Vapor Pressure

These move in OPPOSITE directions.

Strong IMF:

BP ↑

but

VP ↓

Weak IMF:

BP ↓

but

VP ↑

🧠 Think:

If molecules easily escape, vapor pressure is high and you don't need as much heat to make them boil.

7. Temperature vs. Vapor Pressure

These move in the SAME direction.

Temperature ↑ → Vapor pressure ↑

Why?

Higher temperature gives molecules more energy.

More molecules escape into the gas phase.

Therefore vapor pressure increases.

So:

T ↑ → VP ↑

and

T ↓ → VP ↓

8. What Actually Causes Boiling?

This definition is worth memorizing exactly conceptually:

A liquid boils when its vapor pressure equals the external atmospheric pressure.

Therefore:

VP = atmospheric pressure → BOILING

9. Atmospheric Pressure vs. Boiling Point

Atmospheric pressure ↓ → boiling point ↓

Atmospheric pressure ↑ → boiling point ↑

Why?

At lower external pressure, the liquid doesn't need to achieve as high a vapor pressure before:

VP = external pressure

So it can boil at a lower temperature.

High altitude pattern:

Altitude ↑

→ atmospheric pressure ↓

→ boiling point ↓

Your slides' Leadville, Colorado octane problem applies this relationship.

10. Volatility Pattern

Volatile = evaporates readily.

HIGH volatility:

weak IMF
easy evaporation
high vapor pressure
low boiling point

LOW volatility:

strong IMF
difficult evaporation
low vapor pressure
high boiling point

🧠 Memorize:

Volatility and vapor pressure go together.

Volatility and boiling point go opposite.

11. Viscosity Pattern

Viscosity = resistance to flow.

IMF ↑ → viscosity ↑

Stronger attractions make it more difficult for molecules to move past one another.

Therefore:

strong IMF → slow flow

weak IMF → easier flow

12. Surface Tension Pattern

IMF ↑ → surface tension ↑

Strongly attracted molecules resist spreading apart/increasing the liquid's surface area.

Therefore:

strong IMF → high surface tension

🔥 13. Phase Change Pattern

Memorize the direction:

SOLID → LIQUID → GAS

Particles become:

more separated
more mobile
more disordered

Going this direction requires energy.

Therefore:

Melting: solid → liquid

Vaporization: liquid → gas

Sublimation: solid → gas

are endothermic.

Going backward:

GAS → LIQUID → SOLID

releases energy.

Therefore:

Condensation: gas → liquid

Freezing: liquid → solid

Deposition: gas → solid

are exothermic.

14. Entropy Pattern During Phase Changes

Your slides connect phase changes with molecular randomness, ΔS.

Conceptually:

Solid → Liquid → Gas

Randomness increases.

Therefore:

entropy increases

Gas → Liquid → Solid

Randomness decreases.

Therefore:

entropy decreases

So:

15. Energy Needed for Phase Changes

Your slides define:

ΔHfus = energy needed for solid → liquid

ΔHvap = energy needed for liquid → gas

ΔHsub = energy needed for solid → gas

And using Hess's Law:

ΔHsub = ΔHfus + ΔHvap

Conceptually this makes sense because:

solid → liquid → gas

combines fusion and vaporization.

🧊 CHAPTER 12 — SOLID PATTERNS

16. Crystalline vs. Amorphous

Crystalline

organized + repeating + long-range order

Amorphous

no regular three-dimensional arrangement

Examples in your slides:

glass
rubber

🧠 Think:

Crystal = organized

Amorphous = irregular

17. Four Types of Crystalline Solids

Know what occupies/holds the solid together.

Ionic

ions

Held by electrostatic attractions.

Molecular

molecules

Held together by intermolecular forces.

Covalent network

atoms connected through an extensive network of covalent bonds

Metallic

metal atoms

with shared/delocalized electrons — the electron-sea idea.

This gives metals good:

conductivity + malleability.

18. Unit Cell Pattern

This table is worth straight-up memorizing.

Pattern:

SC → BCC → FCC

Coordination:

6 → 8 → 12

Packing:

52% → 68% → 74%

So increasing coordination corresponds here to tighter packing.

19. FCC and HCP Pattern

Both are closest-packed structures.

Therefore both have:

Coordination number = 12

and

Packing efficiency = 74%

Your slides identify:

FCC = cubic closest-packed (CCP)

20. Counting Atoms — Pattern

You should immediately recognize:

Corner

belongs to 8 unit cells

→ contributes 1/8

Face

belongs to 2 unit cells

→ contributes 1/2

Body center

completely inside

→ contributes 1

Therefore:

SC = 1 atom

BCC = 2 atoms

FCC = 4 atoms

🧪 CHAPTER 13 — SOLUTION PATTERNS

21. Solute vs. Solvent

Solvent

Component present to the greater extent

Solute

Component present to the lesser extent

A solution is:

homogeneous

meaning composition is uniform throughout.

22. Solubility & Intermolecular Forces

Your slides use CCl₄ and benzene as an example of two nonpolar liquids that are miscible.

The important conceptual connection is:

Similar intermolecular-force environments can favor mixing.

Two liquids that are completely soluble in each other in all proportions are:

miscible

23. Saturated vs. Supersaturated

Saturated

Maximum equilibrium amount of dissolved solute.

At equilibrium:

dissolving ⇌ crystallizing

Supersaturated

Contains more than the equilibrium amount of dissolved solute.

⭐ 24. Gas Solubility vs. Temperature

This is a major pattern.

Temperature ↑ → gas solubility ↓

Temperature ↓ → gas solubility ↑

So gases generally stay dissolved better at lower temperatures.

🧠 Gases don't like being hot in solution.

25. Gas Solubility vs. Pressure

The opposite relationship:

Pressure ↑ → gas solubility ↑

Pressure ↓ → gas solubility ↓

This is the basis of Henry's Law in your slides.

Think about soda:

Closed soda

high pressure

→ more CO₂ remains dissolved

Open soda

pressure drops

→ CO₂ solubility decreases

→ CO₂ escapes.

⭐ 26. Temperature vs. Pressure — Don't Mix Them Up

For gas solubility:

🧠 Shortcut:

Cold + pressure = gas stays dissolved.

27. Colligative Properties

Colligative properties depend on:

NUMBER OF DISSOLVED PARTICLES

NOT their chemical identity.

Your four are:

1. Vapor-pressure lowering

2. Boiling-point elevation

3. Freezing-point depression

4. Osmotic pressure

This is one of the most important concepts in Chapter 13.

⭐ 28. THE MASTER COLLIGATIVE PROPERTY PATTERN

Add a nonvolatile solute to a solvent:

Vapor pressure ↓

Boiling point ↑

Freezing point ↓

This is VERY testable.

Write this on your scratch paper:

VP ↓ | BP ↑ | FP ↓

29. Why Vapor Pressure Decreases

Pure solvent:

Lots of solvent molecules available at the surface to escape.

Add nonvolatile solute:

The solvent's mole fraction decreases.

Therefore:

vapor pressure decreases.

This is represented in your slides by Raoult's Law.

30. Why Boiling Point Goes UP

Remember:

Boiling happens when:

vapor pressure = atmospheric pressure

Adding nonvolatile solute lowers vapor pressure.

So now the solution needs a higher temperature to reach the required vapor pressure.

Therefore:

Solute added → VP ↓ → BP ↑

This chain is worth understanding rather than memorizing separately.

31. Why Freezing Point Goes DOWN

Adding solute interferes with the solvent's ability to form its solid phase.

Therefore a lower temperature is needed for freezing.

Solute added → FP ↓

Hence:

freezing-point depression

⭐ 32. Number of Particles Matters

Colligative properties depend on particle number.

Your slides introduce the van't Hoff factor, i for this reason.

Idealized example:

NaCl

NaCl → Na⁺ + Cl⁻

One formula unit gives two particles.

So approximately:

i = 2

A nonelectrolyte that doesn't dissociate:

i = 1

Therefore, at the same appropriate concentration, a solute producing more dissolved particles has a larger colligative effect.

Meaning:

more particles

→ greater VP lowering
→ greater BP elevation
→ greater FP depression
→ greater osmotic pressure

Your slides also warn that real electrolyte solutions can deviate from ideal values because of ion pairs.

33. Molality vs. Molarity — Conceptual Difference

Molarity

mol solute / L solution

Depends on volume.

Molality

mol solute / kg solvent

Depends on mass.

Your slides emphasize:

Molality is temperature-independent.

Why?

Mass doesn't change simply because temperature changes.

Molarity can change with temperature because solution volume can change.

Percent by mass is also temperature-independent.

34. Osmosis Pattern

Osmosis is movement of:

SOLVENT

through a semipermeable membrane.

Direction:

LESS concentrated → MORE concentrated

So if:

Side A = low solute concentration

Side B = high solute concentration

solvent moves:

A → B

🧠 Think:

Water moves toward MORE solute.

35. Osmotic Pressure Pattern

Osmotic pressure is the pressure needed to stop osmosis.

Your slides give:

Π = MRT

So conceptually:

Concentration ↑ → osmotic pressure ↑

and, according to the equation:

Temperature ↑ → osmotic pressure ↑

For particle-count effects, the van't Hoff factor is relevant when dissociation is included.

🧠 MASTER PATTERN MAP

This is the section I'd memorize right before the exam.

IMF

IMF ↑

→ BP ↑
→ viscosity ↑
→ surface tension ↑
→ VP ↓
→ volatility ↓

Temperature & vapor pressure

Temperature ↑ → VP ↑

Atmospheric pressure

Atmospheric pressure ↓ → BP ↓

Phase

Solid → Liquid → Gas

→ energy absorbed
→ entropy ↑

Gas → Liquid → Solid

→ energy released
→ entropy ↓

Gas solubility

Temperature ↑ → gas solubility ↓

Pressure ↑ → gas solubility ↑

Add nonvolatile solute

VP ↓

BP ↑

FP ↓

More dissolved particles

→ larger colligative effect

→ BP elevation ↑
→ FP depression ↑
→ osmotic pressure ↑

Osmosis

Solvent: low solute concentration → high solute concentration

Unit cells

SC → BCC → FCC

Atoms:

1 → 2 → 4

Coordination:

6 → 8 → 12

Packing:

52 → 68 → 74%

🚨 CONCEPTUAL QUESTIONS I WOULD EXPECT

Make sure you can answer questions like:

“Which compound has the highest boiling point?”

→ Identify IMF.

“Which has the highest vapor pressure?”

→ Look for weaker IMF.

“Which is most volatile?”

→ Weakest attractions / high VP.

“Which has the greatest surface tension?”

→ Strongest IMF.

“Which has the greatest viscosity?”

→ Strongest IMF.

“What happens to vapor pressure when temperature increases?”

→ Increases.

“What happens to boiling point when atmospheric pressure decreases?”

→ Decreases.

“What happens when a nonvolatile solute is added?”

→ VP ↓, BP ↑, FP ↓.

“What happens to gas solubility when temperature increases?”

→ Decreases.

“What happens to gas solubility when pressure increases?”

→ Increases.

“Which solution has the greatest colligative effect?”

→ Look for the greatest number of dissolved particles.

“Which direction does water move during osmosis?”

→ Toward the more concentrated solution.

“Which unit cell is most tightly packed?”

→ FCC among SC/BCC/FCC; FCC and HCP both have 74% packing in your slides.

“Which unit cell has coordination number 8?”

→ BCC.

“Which phase has the greatest molecular randomness?”

→ Gas.

⭐ Five lines I want you to KNOW cold

If you only had 60 seconds before walking into the exam, memorize:

Strong IMF → BP ↑, viscosity ↑, surface tension ↑, VP ↓, volatility ↓

Temperature ↑ → vapor pressure ↑

Atmospheric pressure ↓ → boiling point ↓

Gas solubility: temperature ↑ = solubility ↓; pressure ↑ = solubility ↑

Add nonvolatile solute → vapor pressure ↓, boiling point ↑, freezing point ↓

Those five patterns alone connect a huge portion of the conceptual material across your slides.

Yes — this is the version I would actually study from for the conceptual portion. Instead of long definitions, I'm organizing Chapters 11–13 into patterns you can memorize and then apply to multiple-choice questions. Everything below is based on your three exam slide decks.

🧪 CHEM EXAM 1 — MASTER CONCEPTUAL PATTERN SHEET

⭐ 1. Intermolecular Forces — THE MOST IMPORTANT PATTERN

Stronger IMF →

Boiling point ↑
Viscosity ↑
Surface tension ↑
Vapor pressure ↓
Volatility ↓
Evaporation becomes harder

Weaker IMF →

Boiling point ↓
Viscosity ↓
Surface tension ↓
Vapor pressure ↑
Volatility ↑
Evaporation becomes easier

Why?

Strong IMF → molecules hold onto each other → harder to pull them apart.

So if a question says:

Substance A has stronger intermolecular forces than substance B...

You should IMMEDIATELY know:

A has higher BP, viscosity, surface tension AND lower VP/volatility.

⭐ 2. IMF Types

Every molecule →

London dispersion forces

ALL atoms and molecules have dispersion forces.

Polar molecule →

London dispersion + dipole-dipole

H directly bonded to N, O, or F →

London dispersion + dipole-dipole + hydrogen bonding

Memorize:

H—N, H—O, H—F = hydrogen bonding

Simply having H somewhere in the molecule is NOT enough.

CH₄ ❌
HCl ❌
CH₃OH ✅
NH₃ ✅
H₂O ✅
HF ✅

⭐ 3. London Dispersion Pattern

Your slides emphasize two things that affect dispersion forces.

More electrons →

Polarizability ↑ → dispersion forces ↑

So generally:

more electrons → stronger dispersion → BP tends to ↑

Shape also matters

Molecular shape affects how well molecules can make contact with each other.

More effective molecular contact can produce:

stronger dispersion forces

So when comparing similar nonpolar substances, think:

Number of electrons + molecular shape

⭐ 4. Molecular Polarity Pattern

Having polar bonds does NOT automatically make the molecule polar.

Think:

Bond polarity + molecular geometry → molecular polarity

Symmetrical arrangement where bond dipoles cancel →

Nonpolar molecule

Examples from your slides:

CO₂ → nonpolar

BF₃ → nonpolar

Dipoles don't cancel →

Polar molecule

Example:

H₂O → polar

Exam thought process

When you see:

"Which molecule is polar?"

Think:

Are the bonds polar? → What is the geometry? → Do the dipoles cancel?

⭐ 5. IMF → Boiling Point

IMF ↑ → BP ↑

Why?

Strong attractions mean:

more energy is required to separate molecules into the gas phase.

Therefore:

Strongest IMF = generally highest BP

⭐ 6. IMF → Vapor Pressure

These move in the OPPOSITE direction.

IMF ↑ → VP ↓

IMF ↓ → VP ↑

Why?

If molecules strongly attract one another, fewer escape into the vapor.

Memorize:

Strong IMF = High BP + Low VP

That's one of the most useful relationships on the entire exam.

⭐ 7. Vapor Pressure → Volatility

These move together:

VP ↑ → volatility ↑

VP ↓ → volatility ↓

A volatile substance escapes/evaporates easily.

Therefore:

Weak IMF → VP ↑ → volatility ↑ → BP ↓

And:

Strong IMF → VP ↓ → volatility ↓ → BP ↑

⭐ 8. Temperature → Vapor Pressure

Temperature ↑ → vapor pressure ↑

Temperature ↓ → vapor pressure ↓

Why?

Higher temperature gives molecules more energy.

More molecules can escape from liquid → gas.

Therefore:

T ↑ → evaporation easier → VP ↑

⭐ 9. When Does a Liquid Boil?

Know this definition:

Vapor pressure = external atmospheric pressure → BOILING

This explains several other patterns.

⭐ 10. Atmospheric Pressure → Boiling Point

Atmospheric pressure ↓ → BP ↓

Atmospheric pressure ↑ → BP ↑

At lower atmospheric pressure, the liquid doesn't need to develop as much vapor pressure to boil.

Therefore:

Altitude ↑ → atmospheric pressure ↓ → BP ↓

This is the conceptual idea behind the Leadville, Colorado problem in your slides.

⭐ 11. IMF → Viscosity

Viscosity = resistance to flow.

IMF ↑ → viscosity ↑

IMF ↓ → viscosity ↓

Think:

Stronger attractions → molecules have more difficulty sliding past each other → slower flow

⭐ 12. IMF → Surface Tension

IMF ↑ → surface tension ↑

IMF ↓ → surface tension ↓

Strongly attracted surface molecules are pulled inward more strongly.

Therefore:

Strong IMF = high surface tension

🔥 13. Phase Change Direction

Memorize this:

SOLID → LIQUID → GAS

Going right means:

Particle movement ↑
Particle separation ↑
Randomness/entropy ↑
Energy must be absorbed

Going left:

GAS → LIQUID → SOLID

means:

Particle movement ↓
Particle separation ↓
Randomness/entropy ↓
Energy is released

⭐ 14. Endothermic Phase Changes

Going toward gas requires energy.

Solid → liquid

Melting/Fusion

Liquid → gas

Vaporization

Solid → gas

Sublimation

Therefore:

Melting + Vaporization + Sublimation = ENDOTHERMIC

Think:

Solid → Liquid → Gas = absorb heat

⭐ 15. Exothermic Phase Changes

Going toward solid releases energy.

Gas → liquid

Condensation

Liquid → solid

Freezing

Gas → solid

Deposition

Therefore:

Condensation + Freezing + Deposition = EXOTHERMIC

Think:

Gas → Liquid → Solid = release heat

⭐ 16. Phase → Entropy

Entropy = molecular randomness/disorder.

General pattern:

Solid < Liquid < Gas

Therefore:

Solid → liquid

ΔS positive

Liquid → gas

ΔS positive

Solid → gas

ΔS positive

Going backward:

Gas → liquid

ΔS negative

Liquid → solid

ΔS negative

Gas → solid

ΔS negative

Memorize:

Toward GAS → ΔS +

Toward SOLID → ΔS −

⭐ 17. Enthalpy of Phase Changes

Know what each represents:

ΔHfus

solid → liquid

ΔHvap

liquid → gas

ΔHsub

solid → gas

And because:

solid → liquid → gas

your slides use Hess's Law:

ΔHsub = ΔHfus + ΔHvap

⭐ 18. Phase Equilibrium Pattern

Your slides give:

ΔG = ΔH − TΔS

At phase equilibrium:

ΔG = 0

Therefore:

ΔH = TΔS

This is why you can determine entropy changes from the temperature and enthalpy of a phase transition.

🧊 CHAPTER 12 — SOLIDS

⭐ 19. Crystalline vs. Amorphous

Crystalline →

Rigid + repeating + long-range order

Particles occupy specific positions in a lattice.

Amorphous →

No regular three-dimensional arrangement

Examples in your slides:

glass and rubber

Memory:

Crystal = organized

⭐ 20. Four Types of Crystalline Solids

Know what is present and what holds it together.

IONIC

ions + electrostatic attractions

MOLECULAR

molecules + intermolecular forces

Those IMFs can include:

dispersion
dipole-dipole
hydrogen bonding

COVALENT NETWORK

atoms + huge network of covalent bonds

METALLIC

metal atoms + shared/delocalized electron "sea"

This contributes to:

conductivity + malleability

⭐ 21. Unit Cell Pattern

A unit cell is:

the basic repeating structural unit of a crystalline solid

The major structures in your slides:

Simple Cubic (SC)
Body-Centered Cubic (BCC)
Face-Centered Cubic (FCC)
plus HCP

⭐ 22. SC → BCC → FCC Pattern

MEMORIZE THIS WHOLE THING:

Simple Cubic

1 atom/cell → CN 6 → 52% packed

BCC

2 atoms/cell → CN 8 → 68% packed

FCC

4 atoms/cell → CN 12 → 74% packed

So:

SC → BCC → FCC

Atoms:

1 → 2 → 4

Coordination:

6 → 8 → 12

Packing:

52 → 68 → 74%

That's a beautiful pattern to memorize instead of three separate sets of numbers.

⭐ 23. Coordination Number Pattern

Coordination number = number of nearest-neighbor atoms.

SC = 6

BCC = 8

FCC = 12

HCP = 12

In the structures shown in your slides:

Higher coordination number → tighter packing

⭐ 24. Closest Packing

Your slides give:

FCC/CCP

CN = 12
packing = 74%

HCP

CN = 12
packing = 74%

So:

FCC and HCP = closest packed = 74%

FCC corresponds to:

cubic closest packing (CCP)

⭐ 25. Atom Contribution Pattern

This is partly numerical but also very conceptual.

Corner atom →

shared by 8 cells

so contributes:

1/8

Face atom →

shared by 2 cells

so contributes:

1/2

Body-center atom →

not shared

so contributes:

1

Therefore:

Corner = 1/8

Face = 1/2

Center = 1

🧪 CHAPTER 13 — SOLUTIONS

⭐ 26. Solution Pattern

A solution is:

homogeneous

meaning composition is the same throughout.

Greater amount →

solvent

Lesser amount →

solute

There can be multiple solutes.

⭐ 27. Miscible Pattern

Two liquids are miscible when:

they are completely soluble in each other in all proportions.

Your slides use nonpolar CCl₄ and nonpolar benzene as an example.

The broader pattern supported by this section is that intermolecular forces help determine whether substances will mix.

⭐ 28. Dissolving Can Be Endothermic OR Exothermic

Don't assume dissolving always releases heat.

Your slides give:

NH₄Cl dissolving

Endothermic → solution gets cold

AlCl₃ dissolving

Exothermic → solution gets hot

Solution formation depends on breaking existing attractions and forming new attractions.

⭐ 29. Entropy Can Favor Dissolution

A negative/exothermic ΔH can favor dissolution, BUT dissolution doesn't have to be exothermic.

Your slides specifically explain that entropy matters too.

Particles in solution can have greater mobility than particles locked into a crystalline solid.

Therefore:

Dissolution can occur even when ΔHsolution is positive.

Don't use the rule:

"Endothermic substances can't dissolve." ❌

⭐ 30. Saturation Pattern

Unsaturated

Can dissolve more solute.

Saturated

Contains the maximum equilibrium amount.

At saturation:

Rate dissolving = rate crystallizing

Supersaturated

Contains:

more than the equilibrium amount of solute

⭐ 31. Gas Solubility → Temperature

SUPER IMPORTANT.

Temperature ↑ → gas solubility ↓

Temperature ↓ → gas solubility ↑

Memorize:

HOT → gas comes OUT

COLD → gas stays IN

⭐ 32. Gas Solubility → Pressure

Henry's Law:

Pressure ↑ → gas solubility ↑

Pressure ↓ → gas solubility ↓

Think:

PRESSURE PUSHES GAS INTO SOLUTION

This is exactly what your slides emphasize.

⭐ 33. Master Gas-Solubility Pattern

Put the last two together:

COLD + HIGH PRESSURE = MORE GAS DISSOLVED

and:

HOT + LOW PRESSURE = LESS GAS DISSOLVED

This is a great way to remember both simultaneously.

⭐ 34. Colligative Properties

Definition:

Properties that depend on the amount/number of dissolved particles, not their chemical identity.

Four from your slides:

1. Vapor-pressure lowering

2. Boiling-point elevation

3. Freezing-point depression

4. Osmotic pressure

🚨 35. THE BIGGEST CHAPTER 13 PATTERN

Add a nonvolatile solute:

VP ↓ | BP ↑ | FP ↓

Say it until it's automatic:

Vapor pressure DOWN

Boiling point UP

Freezing point DOWN

⭐ 36. Why Adding Solute Lowers Vapor Pressure

Add nonvolatile solute:

→ solvent mole fraction ↓

→ fewer solvent molecules contribute to vapor

→ vapor pressure ↓

This is the idea represented by Raoult's Law in your slides:

Psolution = P°solvent × Xsolvent

So conceptually:

Xsolvent ↓ → Psolution ↓

⭐ 37. Why Adding Solute Raises Boiling Point

Remember:

boiling happens when VP = atmospheric pressure

Add nonvolatile solute:

→ VP ↓

→ need more heat to raise VP enough

→ BP ↑

So memorize the causal chain:

Solute ↑ → VP ↓ → BP ↑

⭐ 38. Why Adding Solute Lowers Freezing Point

Adding solute interferes with formation of the solvent's solid phase.

Therefore:

lower temperature is required to freeze

So:

Solute ↑ → FP ↓

⭐ 39. Amount of Solute → Colligative Effect

The slides say colligative properties depend on the number of dissolved particles.

Therefore, conceptually:

More dissolved particles →

greater vapor-pressure lowering

greater boiling-point elevation

greater freezing-point depression

greater osmotic pressure

So if two otherwise comparable solutions contain different numbers of dissolved particles:

more particles = bigger effect

⭐ 40. van't Hoff Factor Pattern

The van't Hoff factor tells you about particles produced in solution.

Your slides define:

i = moles particles in solution / moles solute dissolved

For idealized:

Nonelectrolyte

Doesn't break apart.

i = 1

NaCl

NaCl → Na⁺ + Cl⁻

approximately i = 2

Therefore, more particles can produce a larger colligative effect.

BUT your slides give an important exception:

Actual electrolyte solutions can have smaller colligative effects than ideal predictions, especially at higher concentrations.

Why?

Ion pairing

Some ions remain associated, reducing the effective number of independent particles.

⭐ 41. Molarity vs. Molality Pattern

Molarity

moles solute / L solution

Volume can change with temperature.

Therefore:

Molarity can be temperature dependent

Molality

moles solute / kg solvent

Mass doesn't change because of temperature.

Therefore:

Molality is temperature independent

Percent by mass

Also based on mass.

Therefore:

Percent by mass is temperature independent

Your slides specifically point this out.

⭐ 42. Osmosis Pattern

Osmosis = movement of SOLVENT through a semipermeable membrane.

Direction:

LESS concentrated → MORE concentrated

So water/solvent moves:

low solute concentration → high solute concentration

Memory:

💧 WATER MOVES TOWARD MORE SOLUTE

Be careful: the question may try to trick you into following the solute instead.

⭐ 43. Osmotic Pressure Pattern

Osmotic pressure is:

pressure required to stop osmosis.

Your slides give:

Π = MRT

Therefore conceptually:

M ↑ → Π ↑

T ↑ → Π ↑

And particle-number effects are important for electrolyte solutions.

⭐ 44. Raoult's Law When BOTH Liquids Are Volatile

If only the solvent is volatile:

only that volatile component contributes to vapor pressure.

But if both components are volatile:

Both contribute.

Your slides show:

Ptotal = PA + PB

Therefore:

Total vapor pressure = sum of individual partial pressures

🧠 THE ENTIRE EXAM AS PATTERNS

This is the section I'd screenshot and memorize.

💪 INTERMOLECULAR FORCES

IMF ↑ → BP ↑ | viscosity ↑ | surface tension ↑ | VP ↓ | volatility ↓

🌡 TEMPERATURE & VAPOR PRESSURE

Temperature ↑ → VP ↑

🏔 ATMOSPHERIC PRESSURE

Atmospheric pressure ↓ → BP ↓

Altitude ↑ → atmospheric pressure ↓ → BP ↓

💨 VOLATILITY

VP ↑ → volatility ↑ → BP ↓

🔥 PHASE CHANGES

Solid → Liquid → Gas = energy absorbed + entropy ↑

Gas → Liquid → Solid = energy released + entropy ↓

🧊 SOLID STRUCTURES

SC → BCC → FCC

Atoms: 1 → 2 → 4

CN: 6 → 8 → 12

Packing: 52% → 68% → 74%

🧊 CLOSEST PACKING

FCC + HCP → CN 12 → 74% packed

💨 GAS SOLUBILITY

Temperature ↑ → gas solubility ↓

Pressure ↑ → gas solubility ↑

or simply:

COLD + PRESSURE = GAS STAYS DISSOLVED

🧂 ADD NONVOLATILE SOLUTE

VP ↓ | BP ↑ | FP ↓

🔢 COLLIGATIVE PROPERTIES

More dissolved particles → bigger colligative effect

Meaning:

VP lowering ↑

BP elevation ↑

FP depression ↑

osmotic pressure ↑

💧 OSMOSIS

Solvent: less concentrated → more concentrated

Water moves toward MORE solute

🌡 CONCENTRATION UNITS

Molality + mass % = temperature independent

Molarity = volume based → can change with temperature

🚨 OPPOSITE VS. SAME-DIRECTION PAIRS

This is especially useful for multiple choice.

Move in SAME direction:

IMF ↑ → BP ↑

IMF ↑ → viscosity ↑

IMF ↑ → surface tension ↑

Temperature ↑ → VP ↑

VP ↑ → volatility ↑

Atmospheric pressure ↑ → BP ↑

Gas pressure ↑ → gas solubility ↑

Particles ↑ → colligative effect ↑

Concentration ↑ → osmotic pressure ↑

Move in OPPOSITE directions:

IMF ↑ → VP ↓

IMF ↑ → volatility ↓

VP ↑ → BP ↓ (when comparing substances based on their IMF/vapor-pressure behavior)

Temperature ↑ → gas solubility ↓

Altitude ↑ → atmospheric pressure ↓

Altitude ↑ → BP ↓

Nonvolatile solute ↑ → VP ↓

Nonvolatile solute ↑ → BP ↑

Nonvolatile solute ↑ → FP ↓

🎯 THE 10 PATTERNS I WOULD KNOW WITHOUT THINKING

If your professor asks a conceptual question and you know these automatically, you'll be in great shape:

1. IMF ↑ → BP ↑, viscosity ↑, surface tension ↑, VP ↓, volatility ↓

2. More electrons → polarizability ↑ → dispersion forces ↑

3. Temperature ↑ → vapor pressure ↑

4. Atmospheric pressure ↓ → boiling point ↓

5. Solid → liquid → gas = heat absorbed + entropy ↑

6. SC → BCC → FCC = atoms 1→2→4, CN 6→8→12, packing 52→68→74%

7. Temperature ↑ → gas solubility ↓

8. Pressure ↑ → gas solubility ↑

9. Add nonvolatile solute → VP ↓, BP ↑, FP ↓

10. More dissolved particles → larger colligative effect

And one bonus because professors LOVE reversing it:

💧 OSMOSIS: solvent moves from LESS solute → MORE solute

Those are the cause-and-effect relationships running through essentially all of the conceptual material emphasized in your Chapter 11, 12, and 13 slides.