CHEM

Yes. The material in those screenshots is way more wordy than it needs to be. The actual concepts fit together in a pretty logical chain:

electrons → electronegativity → bond polarity → molecular polarity → intermolecular forces → physical properties

Here are the notes I would actually put in your notebook.

Intramolecular Bonds, Molecular Polarity & Intermolecular Forces1. Start with the atom

An atom contains:

Protons: positive (+), in nucleus
Neutrons: neutral, in nucleus
Electrons: negative (−), outside nucleus

Most important for bonding:

Valence electrons = electrons in the outermost shell.

They are the electrons involved in:

forming bonds
breaking bonds
chemical reactions

2. Intramolecular vs. Intermolecular

This distinction is VERY important.

Intramolecular = INSIDE

Intramolecular bonds hold atoms together inside a substance.

Examples:

O—H bonds inside H₂O
C—H bonds inside CH₄
ionic attraction within NaCl

Main types:

Covalent: electrons are shared.

Ionic: electrons are transferred, creating oppositely charged ions.

Intermolecular = BETWEEN

Intermolecular forces (IMFs) are attractions between separate particles/molecules.

Example with water:

H—O—H ··· H—O—H

The O—H line is a covalent bond.

The ··· between the two water molecules is an intermolecular force.

Easy memory trick

INTRAmolecular = INside

INTERmolecular = between molecules

Intramolecular bonds are generally much stronger than intermolecular forces.

3. Covalent vs. Ionic BondsCovalent bond

Usually:

nonmetal + nonmetal

Electrons are shared.

Example:

H₂O

O and H share electrons.

Ionic bond

Usually:

metal + nonmetal

Electrons are transferred.

This forms:

cation = positive ion
anion = negative ion

Then:

opposite charges attract.

Example:

NaCl → Na⁺ + Cl⁻

NaCl does NOT normally exist as little individual NaCl molecules.

It forms a large ionic crystal lattice.

4. ElectronegativityElectronegativity (EN)

The ability of an atom to:

pull shared electrons toward itself in a covalent bond.

The more electronegative atom pulls the electrons closer.

Periodic trend

Electronegativity generally increases:

→ across a period

and

↑ up a group

So it increases toward the upper-right corner of the periodic table.

Most electronegative element:Fluorine (F)

Approximately:

Element	EN
F	4.0
O	3.5
N	3.0
Cl	3.0
C	2.5
H	2.1
Na	0.9

These are rounded values commonly used in general chemistry.

You don't necessarily need to memorize every electronegativity value unless your professor requires it.

Definitely remember:

F > O > N/Cl > C > H

5. Electronegativity Difference — ΔEN

To determine how polar a bond is:

FormulaΔEN=∣EN1−EN2∣

Always take the absolute value.

Example:

C—O

C = 2.5
O = 3.5

ΔEN=3.5−2.5=1.0

Therefore C—O is a polar covalent bond.

6. Bond Types Based on ΔEN

For your class, the notes are using approximately:

Nonpolar covalentΔEN≤0.4

Electrons are shared almost equally.

Examples:

C—C

H—H

C—H is generally treated as essentially nonpolar in general chemistry.

Polar covalent

Approximately:

0.5≤ΔEN≤1.9

Electrons are shared unequally.

Examples:

O—H
C—O
C—F
C—Cl

Ionic

Approximately:

ΔEN≥2.0

Very large difference in electronegativity.

Example:

Na—Cl

Important: These cutoffs are guidelines rather than absolute laws, so use the ranges your professor gives you.

7. Partial Charges: δ⁺ and δ⁻

In a polar covalent bond, electrons are NOT completely transferred.

Instead:

More electronegative atom:δ−Less electronegative atom:δ+

Example:

Hδ+−Oδ−

Oxygen is more electronegative than hydrogen, so oxygen pulls the electrons toward itself.

VERY IMPORTANT

δ⁺ and δ⁻ are partial charges.

They are NOT the same as:

Na⁺ or Cl⁻

Those are full ionic charges.

8. Bond Dipole

A polar bond creates a bond dipole.

Think:

electrons are being pulled toward one end of the bond.

The dipole arrow points toward the:

more electronegative atom

Example:

C → O

because O is more electronegative.

9. Dipole Moment

Dipole moment measures the amount of charge separation.

Symbol:

Unit:

Debye (D)

Generally:

larger dipole moment = greater polarity

But molecular polarity also depends on shape, not just individual bond polarity.

And that brings us to one of the MOST IMPORTANT concepts.

10. Polar Bond ≠ Polar Molecule

A molecule can contain polar bonds and still be:

NONPOLAR

Why?

Because the bond dipoles can cancel.

To determine whether an entire molecule is polar, ask TWO questions.

Question 1:

Does it have polar bonds?

Question 2:

Do those bond dipoles cancel because of molecular geometry?

11. Molecular Polarity

A molecule is usually polar when:

It contains polar bonds.
Its shape causes an uneven distribution of charge.

If the dipoles cancel:

molecule = NONPOLAR

If they do NOT cancel:

molecule = POLAR

12. Classic Example: CO₂

Structure:

O=C=O

Each C=O bond is polar.

Oxygen pulls electrons toward itself.

But CO₂ is:

linear

Bond angle:

180∘

So the two dipoles point in opposite directions:

← O = C = O →

They cancel.

Therefore:

CO₂ is NONPOLAR.

Even though it contains polar bonds.

This is extremely testable.

13. Compare CO₂ and H₂OCO₂

polar bonds ✅
linear ✅
dipoles cancel ✅
nonpolar molecule

H₂O

polar O—H bonds ✅
bent shape ✅
dipoles do NOT cancel ✅
polar molecule

So molecular geometry matters enormously.

14. VSEPR Geometry You Should Remember Here

For determining polarity:

2 electron groups

Linear

180∘3 electron groups

Trigonal planar

120∘4 electron groups

Tetrahedral

109.5∘

But remember:

electron geometry alone doesn't tell you whether a molecule is polar.

You must look at:

the atoms attached
lone pairs
symmetry
direction of bond dipoles

15. Hydrocarbons

Hydrocarbons contain only:

C + H

Examples:

CH₄
C₃H₈
C₄H₁₀
C₈H₁₈

In general chemistry, C—H bonds are usually treated as essentially nonpolar.

Therefore hydrocarbons are generally:

NONPOLAR

That will become important when identifying intermolecular forces.

16. Intermolecular Forces

These are forces BETWEEN particles.

The main ones you need are:

London dispersion forces
Dipole-dipole forces
Hydrogen bonding
Ion-dipole forces

17. London Dispersion Forces

Also called:

dispersion forces
London forces
instantaneous dipole-induced dipole forces

MOST IMPORTANT RULE:ALL atoms and molecules have London dispersion forces.

Polar? Yes.

Nonpolar? Yes.

Atoms? Yes.

Everything with electrons has them.

18. How London Dispersion Forces Happen

Electrons are constantly moving.

At one instant, more electrons might temporarily be on one side of an atom/molecule.

That creates an:

instantaneous dipole

One side becomes temporarily:

δ−

and the other:

δ+

That temporary dipole can distort the electron cloud of a nearby molecule.

Now the nearby molecule develops an:

induced dipole

The two molecules attract.

That attraction is:

London dispersion force.

19. PolarizabilityPolarizability

How easily the electron cloud can be distorted.

More easily distorted electron cloud:

↓

stronger temporary dipoles

↓

stronger London dispersion forces

20. What Increases Polarizability?

Generally:

More electrons → larger electron cloud → more polarizable

Therefore:

larger/heavier atoms and molecules generally have stronger dispersion forces.

Example:

F₂ < Cl₂ < Br₂ < I₂

Dispersion force strength increases going down.

Why?

I₂ has many more electrons than F₂.

Its electron cloud is much easier to distort.

21. Example: CCl₄ vs. CI₄

Both are nonpolar molecules.

But iodine is much larger and more polarizable than chlorine.

Therefore:

CI4

has stronger London dispersion forces than:

CCl4

22. Surface Area Matters Too

For molecules with similar molar masses:

Long/straight molecules

have more surface-to-surface contact.

→ stronger dispersion forces

Compact/branched molecules

have less surface contact.

→ weaker dispersion forces

Your professor's example is:

Pentane vs. Neopentane

Same molecular formula:

C5H12

But:

pentane = long/straight

neopentane = compact/branched

Therefore:

pentane has stronger dispersion forces

and a higher boiling point.

Remember:

More surface area → stronger LDF → higher boiling point

23. Dipole-Dipole Forces

Dipole-dipole attractions occur between:

POLAR molecules.

A polar molecule has permanent:

δ+δ−

The positive end of one molecule attracts the negative end of another.

Think:

δ+⋯δ−Rule

Polar molecule → dipole-dipole forces

And remember, it ALSO has London dispersion forces.

24. Acetone vs. PropaneAcetone

Contains:

C=O

The carbonyl group is strongly polar.

Therefore acetone is polar and has:

London dispersion
dipole-dipole forces

PropaneC3H8

Hydrocarbon.

Nonpolar.

Therefore its main IMF is:

London dispersion

This helps explain why acetone has stronger intermolecular attraction than a similar-sized nonpolar molecule.

25. Hydrogen Bonding

This one is VERY important.

Hydrogen bonding is a particularly strong form of dipole-dipole attraction.

For hydrogen bonding, look for:H directly bonded to:F, O, or N

Memorize:

FON

or:

H—FH—OH—N

If you see one of those bonds, the molecule can generally hydrogen bond with appropriate neighboring molecules.

26. Example: Water

Water:

H−O−H

Contains O—H bonds.

Therefore water molecules form hydrogen bonds.

One molecule's:

Hδ+

is attracted to another molecule's oxygen lone pair / δ⁻ region.

This produces strong intermolecular attraction.

27. Why Water Has Unusual Properties

Hydrogen bonding helps explain water's:

relatively high boiling point
high surface tension
cohesion
ability to dissolve many polar/ionic substances

Despite H₂O's small molar mass, it remains liquid at room temperature because its molecules attract one another strongly through hydrogen bonding.

28. VERY COMMON TRICK: Acetone

Acetone contains oxygen.

But:

Does acetone hydrogen bond with itself?

NO.

Why?

Its H atoms are bonded to carbon, not oxygen.

It does NOT contain:

O—H

N—H

F—H.

So pure acetone does not hydrogen bond to itself.

BUT:

Acetone CAN hydrogen bond with water.

Why?

Water has O—H and can donate the H.

Acetone's oxygen has lone pairs and can accept the H-bond.

That distinction is testable.

29. Hydrogen Fluoride

HF contains:

H−F

Therefore HF forms hydrogen bonds.

Fluorine is extremely electronegative.

So:

Hδ+−Fδ−

has a strongly polarized bond.

30. More —OH Groups = More Hydrogen Bonding

Compare:

EthanolCH3CH2OH

one OH group.

Ethylene glycolHO−CH2−CH2−OH

two OH groups.

Ethylene glycol can form a more extensive hydrogen-bonding network.

Therefore it has much stronger intermolecular attraction and a much higher boiling point.

Approximate boiling points:

Ethanol:

78∘C

Ethylene glycol:

197∘C

Your screenshot has a typo where it mentions ethanol at 178°C. Ethanol is about 78°C.

31. Ion-Dipole Forces

Ion-dipole attractions occur between:

an ION

and

a POLAR molecule

Classic example:

NaCl dissolved in water

NaCl separates into:

Na++Cl−

Water is polar.

32. Water Around Na⁺

Water has:

oxygen = δ⁻
hydrogen = δ⁺

Around:

Na+

the oxygen δ⁻ side points toward Na⁺.

Think:

positive ion attracts negative side of water.

33. Water Around Cl⁻

Around:

Cl−

the hydrogen δ⁺ sides point toward Cl⁻.

Think:

negative ion attracts positive side of water.

This is:

ion-dipole attraction.

34. Solubility: "Like Dissolves Like"

A useful rule:

Like dissolves like.Polar substances tend to dissolve in:

polar solvents

Nonpolar substances tend to dissolve in:

nonpolar solvents

Ionic substances often dissolve in:

polar solvents

Example:

NaCl + H₂O ✅

because polar water stabilizes Na⁺ and Cl⁻ through ion-dipole forces.

NaCl + gasoline ❌

because gasoline is nonpolar and cannot effectively stabilize the ions.

35. IMF Strength — Basic Ranking

For the types you're studying, a useful general ranking is:

Ion-dipole>Hydrogen bonding>Dipole-dipole>London dispersion

BUT there is an important caveat:

A very large molecule can have extremely strong London dispersion forces, sometimes stronger overall than the dipole-dipole forces of a small polar molecule.

So don't blindly use the ranking when molecules are very different in size.

36. Do NOT Put Covalent Bonds in the IMF List

Your screenshots show a hierarchy that includes covalent/ionic bonds above the intermolecular forces.

That's okay if your professor is comparing all types of attractive forces, but remember:

Covalent and ionic bonds

are intramolecular/chemical bonding

while:

ion-dipole, H-bonding, dipole-dipole and dispersion

are intermolecular attractions.

So if asked:

"What is the strongest intermolecular force?"

Do NOT answer "covalent."

Covalent isn't an IMF.

37. Stronger IMF = Higher Boiling Point

This is one of the most important relationships.

To boil a liquid, molecules must separate from each other.

If molecules strongly attract each other:

more energy is required.

Therefore:

stronger IMF→higher boiling point

Usually also:

stronger IMF→higher melting point

although melting-point trends can be less straightforward because crystal packing matters.

38. Stronger IMF = Lower Vapor Pressure

If molecules strongly attract each other, fewer escape into the gas phase.

Therefore:

strong IMF→lower vapor pressure

And generally:

strong IMF→lower volatility

39. Quick Property Relationships

Memorize:

Stronger intermolecular forces:

⬆ boiling point

⬆ generally melting point

⬆ viscosity

⬆ surface tension

⬇ vapor pressure

⬇ volatility

40. Distance and Attraction

Electrostatic attraction becomes weaker as particles get farther apart.

So:

closer charges = stronger attractionfarther charges = weaker attraction

Also:

larger charges = stronger attraction

This principle applies throughout chemistry.

41. Saturated vs. Unsaturated Fats — What the Chemistry Point Is

The important chemistry concept here is molecular shape and dispersion forces.

Saturated fatty acids

Contain mainly C—C single bonds.

Their hydrocarbon chains can often pack together relatively well.

Better packing:

→ greater intermolecular contact

→ stronger dispersion forces

→ generally higher melting point

Unsaturated fatty acids

Contain one or more C=C double bonds.

A cis double bond produces a bend/kink in the chain.

The kink prevents molecules from packing together as tightly.

Therefore:

→ less contact

→ weaker dispersion forces

→ lower melting point

Many unsaturated fats are liquids at room temperature.

Important correction

Don't memorize the statement from the generated guide that saturated fats are simply "difficult for human enzymes to break down" or that unsaturated fats are inherently "easily processed." That's an oversimplification and isn't the intermolecular-forces concept you need here.

Focus on:

shape → packing → dispersion forces → melting point

42. Real Gases and IMF

Ideal gas law:

PV=nRT

Ideal gases are imagined to have:

no intermolecular attractions
negligible particle volume

Real gases don't behave perfectly ideally.

43. When Do Real Gases Deviate Most?LOW temperature

and

HIGH pressure

Why?

At low temperature:

particles move more slowly.

At high pressure:

particles are pushed closer together.

Therefore intermolecular attractions become more important.

Remember:Low T + High P = most nonideal

44. The Master IMF Decision Tree

When they give you a molecule, do this.

STEP 1 — Is it ionic + mixed with a polar molecule?

Example Na⁺ in H₂O.

YES →

Ion-dipole

If no, continue.

STEP 2 — Does the molecule contain H directly bonded to N, O, or F?

YES →

Hydrogen bonding

It ALSO has:

dipole-dipole
London dispersion

Continue if no.

STEP 3 — Is the molecule polar?

YES →

Dipole-dipole

plus London dispersion.

If no →

London dispersion only

for a typical nonpolar molecular substance.

45. Example PracticeCH₄

Nonpolar.

IMF:

London dispersion

HCl

Polar.

No H—F, H—O or H—N.

IMFs:

Dipole-dipole + London dispersion

Strongest:

dipole-dipole

H₂O

Polar.

Contains O—H.

IMFs:

London dispersion
dipole-dipole
hydrogen bonding

Strongest:

hydrogen bonding

NH₃

Contains N—H.

Polar.

Strongest IMF:

hydrogen bonding

Contains H—F.

Strongest IMF:

hydrogen bonding

CO₂

C=O bonds are polar.

But molecule is linear and symmetrical.

Dipoles cancel.

Molecule = nonpolar.

Strongest IMF:

London dispersion

CCl₄

Each C—Cl bond is polar.

But molecule is symmetrical tetrahedral.

Dipoles cancel.

Molecule = nonpolar.

Strongest IMF:

London dispersion

CH₃OH

Contains O—H.

Polar.

Strongest IMF:

hydrogen bonding

CH₃OCH₃

Contains oxygen but NO O—H.

Polar molecule.

Strongest IMF:

dipole-dipole

NOT hydrogen bonding with itself.

46. Biggest Mistake to Avoid

Do NOT do this:

"It has oxygen, therefore it hydrogen bonds."

Wrong.

You need:

H−N, H−O, or H−F

The H must be directly bonded to N, O, or F for the molecule to donate a hydrogen bond.

47. Another Huge Mistake

Don't confuse:

Bond polarity

with

Molecular polarity

Example:

CO₂ has polar bonds

but is a nonpolar molecule.

Why?

Symmetry cancels the dipoles.

48. What Determines London Dispersion Strength?

Memorize these TWO things:

1. Polarizability / number of electrons

More electrons → stronger LDF.

2. Surface area

More surface contact → stronger LDF.

So:

bigger electron cloud→stronger LDF

and

greater surface area→stronger LDF

49. The Entire Chapter in About 10 Lines

If you remember nothing else, memorize this:

Intramolecular = bonds inside a substance.
Intermolecular = attractions between particles.
Higher electronegativity = stronger pull on bonding electrons.
Unequal electron sharing creates a polar bond.
Polar bonds do NOT automatically mean a polar molecule.
Shape determines whether bond dipoles cancel.
All molecules have London dispersion forces.
Polar molecules also have dipole-dipole forces.
H directly bonded to F, O, or N → hydrogen bonding.
Ion + polar molecule → ion-dipole.
Larger/more polarizable molecules → stronger dispersion.
More surface area → stronger dispersion.
Stronger IMFs generally → higher boiling point.
Low temperature + high pressure → real gases deviate most from ideal behavior.

And memorize this ranking:Ion-dipole>H-bond>Dipole-dipole>LDF

This is the version I would study rather than trying to memorize every sentence in those screenshots.