Types of Intermolecular Forces

London Dispersion Forces are the weakest type of intermolecular force, but they are present in ALL molecules and atoms — polar and nonpolar alike.

They arise from temporary (instantaneous) dipoles caused by the random motion of electrons. At any given moment, electrons may be unevenly distributed around a nucleus, creating a brief dipole that can induce a dipole in a neighboring molecule.

Key terminology:

• Instantaneous dipole: A temporary, fleeting charge separation in any atom or molecule
• Induced dipole: A dipole created in a neighboring particle by the electric field of the instantaneous dipole
• LDF are also called van der Waals forces or dispersion forces

London Dispersion Forces are unique among intermolecular forces because of one critical property.

Here's the step-by-step mechanism for London Dispersion Forces:

Step 1: Electrons in an atom or molecule are constantly moving. At any instant, there may be more electrons on one side than the other.

Step 2: This uneven distribution creates a temporary dipole (instantaneous dipole) with a partial negative side ((\delta^-)) and a partial positive side ((\delta^+)).

Step 3: The electric field of this temporary dipole distorts the electron cloud of a neighboring molecule, creating an induced dipole.

Step 4: The two dipoles attract each other — the (\delta^+) end of one molecule attracts the (\delta^-) end of the neighbor.

Step 5: These attractions are extremely short-lived, constantly forming and breaking, but their cumulative effect creates a net attractive force.

Fill in the blanks to describe how London Dispersion Forces form.

The strength of London Dispersion Forces depends on two main factors:

1. Molar Mass (Number of Electrons) More electrons (\rightarrow) larger electron cloud (\rightarrow) more easily polarized (\rightarrow) stronger LDF

Example: (\text{I}_2) (molar mass 254 g/mol) has much stronger LDF than (\text{F}_2) (molar mass 38 g/mol). That's why (\text{I}_2) is a solid at room temperature while (\text{F}_2) is a gas.

2. Surface Area (Molecular Shape) Greater surface area (\rightarrow) more points of contact between molecules (\rightarrow) stronger LDF

Example: n-pentane (straight chain) has a higher boiling point than neopentane (spherical/compact) even though they have the same molecular formula (\text{C}5\text{H}{12}). The elongated shape of n-pentane provides more surface area for LDF.

Consider the noble gases: He, Ne, Ar, Kr, Xe. They only experience London Dispersion Forces.

Molecular shape affects London Dispersion Forces through surface area contact.

Complete the following statements about London Dispersion Forces.

Dipole-dipole forces occur between polar molecules — molecules that have a permanent dipole moment.

A molecule is polar when:

• It contains polar bonds (bonds between atoms with different electronegativities)
• The molecular geometry does NOT cancel out the bond dipoles

In a dipole-dipole interaction, the partially positive end ((\delta^+)) of one molecule attracts the partially negative end ((\delta^-)) of another molecule.

Key points:

• Dipole-dipole forces are stronger than LDF for molecules of similar size
• They only occur between polar molecules
• Polar molecules ALSO experience LDF in addition to dipole-dipole forces

To experience dipole-dipole forces, a molecule must be polar. Let's review which molecules are polar.

Unlike LDF (which involve temporary dipoles), dipole-dipole forces involve permanent charge separations.

Alignment: Polar molecules tend to orient themselves so that the (\delta^+) end of one molecule is near the (\delta^-) end of a neighbor. This alignment maximizes attraction.

Example with HCl: $\underset{\delta^+}{\text{H}}\text{—}\underset{\delta^-}{\text{Cl}} \cdots \underset{\delta^+}{\text{H}}\text{—}\underset{\delta^-}{\text{Cl}}$

The positive H end of one HCl molecule is attracted to the negative Cl end of the next.

Important: Polar molecules experience BOTH dipole-dipole forces AND London Dispersion Forces. The total intermolecular attraction is the sum of all IMF types present.

Compare dipole-dipole forces and London Dispersion Forces by selecting the correct option for each statement.

Common molecules that experience dipole-dipole forces:

Molecules that are NOT polar (only LDF):

Determine which molecules experience dipole-dipole forces.

Dipole-dipole forces affect physical properties like boiling point.

Fill in the key facts about dipole-dipole forces.

Hydrogen bonding is a special, extra-strong type of dipole-dipole force. It occurs when hydrogen is covalently bonded to one of three highly electronegative atoms:

$\text{N, O, or F}$

The H atom in an N–H, O–H, or F–H bond carries a very large (\delta^+) charge because N, O, and F are so electronegative. This strongly positive H is then attracted to a lone pair on an N, O, or F atom of a neighboring molecule.

Critical AP Concept: Despite its name, a hydrogen "bond" is NOT a chemical bond — it is an intermolecular force (an attraction between molecules). It is much weaker than a covalent or ionic bond but much stronger than typical dipole-dipole forces or LDF.

Mnemonic: Hydrogen bonding requires "FON" — Fluorine, Oxygen, or Nitrogen bonded to H.

The name "hydrogen bond" is misleading. Let's clarify what it actually is.

For a molecule to participate in hydrogen bonding, it needs BOTH of the following:

1. A hydrogen donor: An H atom covalently bonded to N, O, or F

• This H carries a large (\delta^+) charge
• Examples: O–H in water, N–H in ammonia, F–H in HF

2. A hydrogen acceptor: A lone pair on an N, O, or F atom on a nearby molecule

• This lone pair attracts the (\delta^+) hydrogen
• The acceptor atom must have available lone pairs

Common hydrogen bonding molecules:

• (\text{H}_2\text{O}) — both a donor (O–H) and acceptor (lone pairs on O)
• (\text{NH}_3) — both a donor (N–H) and acceptor (lone pair on N)
• (\text{HF}) — both a donor (F–H) and acceptor (lone pairs on F)
• (\text{CH}_3\text{OH}) (methanol) — donor (O–H) and acceptor (lone pairs on O)
• (\text{CH}_3\text{COOH}) (acetic acid) — donor (O–H) and acceptor (lone pairs on O)

Identify which molecules can form hydrogen bonds.

For each molecule, select whether it can form hydrogen bonds with itself.

Water is the most famous example of hydrogen bonding, and it explains many of water's unusual properties:

1. Unusually high boiling point Water ((\text{H}_2\text{O}), MW = 18) boils at 100°C. Compare this to (\text{H}_2\text{S}) (MW = 34), which boils at -60°C. Despite being lighter, water has a MUCH higher boiling point because of hydrogen bonding.

2. Ice is less dense than liquid water Hydrogen bonds in ice form a rigid, open crystal structure with hexagonal symmetry. This creates more empty space than in liquid water, making ice less dense — which is why ice floats.

3. High surface tension Strong hydrogen bonds at the surface create a "skin" that resists breaking, allowing insects to walk on water.

4. High specific heat capacity Energy is needed to break hydrogen bonds before the temperature can rise, so water resists temperature changes.

Each water molecule can form up to 4 hydrogen bonds (2 as donor through its 2 O–H bonds, 2 as acceptor through its 2 lone pairs).

Hydrogen bonding explains many of water's anomalous properties.

Complete these key facts about hydrogen bonding.

Ion-dipole forces occur between an ion (a charged particle) and a polar molecule (a dipole).

These forces are extremely important in chemistry because they explain how ionic compounds dissolve in polar solvents like water.

How they work:

• A cation (positive ion) attracts the (\delta^-) end of a polar molecule
• An anion (negative ion) attracts the (\delta^+) end of a polar molecule

Ion-dipole forces are generally the strongest type of intermolecular force (stronger than hydrogen bonding, dipole-dipole, and LDF) because ions carry full charges rather than partial charges.

$\text{Strength: Ion-Dipole} > \text{H-bonding} > \text{Dipole-Dipole} > \text{LDF}$

(This ranking is for forces between molecules of similar size)

When NaCl dissolves in water, here's what happens at the molecular level:

Step 1: Water molecules orient around the ions on the crystal surface.

• The (\delta^-) oxygen end of water points toward (\text{Na}^+) cations
• The (\delta^+) hydrogen end of water points toward (\text{Cl}^-) anions

Step 2: Ion-dipole attractions between water and the ions on the surface become strong enough to overcome the ionic bonds holding the crystal together.

Step 3: Individual ions are pulled away from the crystal and become hydrated (surrounded by a shell of water molecules). This shell is called a hydration shell or solvation shell.

$\text{NaCl}(s) \xrightarrow{\text{H}_2\text{O}} \text{Na}^+(aq) + \text{Cl}^-(aq)$

The (aq) symbol means the ion is surrounded by water molecules — held by ion-dipole forces.

Ion-dipole forces are central to the dissolution of ionic compounds.

Select which end of the water molecule is attracted to each ion.

The strength of ion-dipole forces depends on two main factors:

1. Charge of the ion Higher charge (\rightarrow) stronger ion-dipole force

• (\text{Mg}^{2+}) creates stronger ion-dipole forces than (\text{Na}^+)
• (\text{Al}^{3+}) creates even stronger ion-dipole forces

2. Size of the ion Smaller ion (\rightarrow) higher charge density (\rightarrow) stronger ion-dipole force

• (\text{Li}^+) (small) creates stronger ion-dipole forces than (\text{K}^+) (larger)
• Both have +1 charge, but Li⁺ concentrates that charge in a smaller volume

Charge density = charge / volume

The combination of high charge and small size gives the highest charge density and the strongest ion-dipole interactions.

This is why salts of small, highly charged ions (like (\text{MgCl}_2)) tend to have very exothermic heats of hydration.

Apply the factors that determine ion-dipole force strength.

Ion-dipole forces help explain the famous rule "like dissolves like":

• Ionic and polar solutes dissolve well in polar solvents (like water)

  • Ion-dipole forces stabilize the dissolved ions
  • Example: NaCl dissolves in water

• Nonpolar solutes dissolve well in nonpolar solvents

  • LDF between solute and solvent replace similar LDF in pure substances
  • Example: Oil (nonpolar) dissolves in hexane (nonpolar)

• Ionic/polar solutes do NOT dissolve well in nonpolar solvents

  • Nonpolar solvents cannot form ion-dipole forces
  • The ionic bonds in the crystal are too strong to break without ion-dipole stabilization
  • Example: NaCl does NOT dissolve in hexane

This principle is central to understanding solubility in AP Chemistry.

Complete the following statements about ion-dipole forces.

On the AP Chemistry exam, you will frequently need to rank molecules by the strength of their intermolecular forces. Here is the general hierarchy:

$\text{Ion-Dipole} > \text{Hydrogen Bonding} > \text{Dipole-Dipole} > \text{London Dispersion Forces}$

However, this ranking has an important caveat: it applies to molecules of similar size. A very large nonpolar molecule (with only LDF) can have stronger total IMF than a small polar molecule (with dipole-dipole forces).

Example: Hexane ((\text{C}6\text{H}{14}), nonpolar, only LDF) boils at 69°C, while formaldehyde ((\text{CH}_2\text{O}), polar, dipole-dipole) boils at -19°C. Hexane's much larger size gives it stronger LDF that outweigh formaldehyde's dipole-dipole advantage.

Strategy for ranking:

1. Identify ALL types of IMF each molecule experiences
2. The dominant IMF determines relative ranking
3. If molecules share the same dominant IMF, compare size (molar mass)

Use this flowchart to identify which IMFs a substance has:

Step 1: Is it an ion interacting with a polar molecule?

• YES (\rightarrow) Ion-dipole forces

Step 2: Does the molecule have H bonded to N, O, or F?

• YES (\rightarrow) Hydrogen bonding (+ dipole-dipole + LDF)

Step 3: Is the molecule polar? (Consider geometry and electronegativity)

• YES (\rightarrow) Dipole-dipole forces (+ LDF)

Step 4: Is the molecule nonpolar?

• YES (\rightarrow) London Dispersion Forces only

Remember: Each molecule experiences ALL of the forces at its level AND below:

• H-bonding molecules also have dipole-dipole AND LDF
• Dipole-dipole molecules also have LDF
• ALL molecules have LDF

For each molecule, identify the strongest IMF it experiences.

Select the strongest IMF for each molecule.

Problem: Rank the following in order of increasing boiling point: (\text{CH}_4), (\text{CH}_3\text{Cl}), (\text{CH}_3\text{OH}), (\text{CH}_3\text{CH}_3)

Solution:

Ranking (lowest to highest BP): $\text{CH}_4 < \text{CH}_3\text{CH}_3 < \text{CH}_3\text{Cl} < \text{CH}_3\text{OH}$

• CH₄ < CH₃CH₃: Both LDF only, but CH₃CH₃ has higher molar mass
• CH₃CH₃ < CH₃Cl: CH₃Cl has dipole-dipole forces in addition to LDF
• CH₃Cl < CH₃OH: CH₃OH has hydrogen bonding, the strongest IMF here (even though CH₃Cl has a higher molar mass)

Use your IMF knowledge to rank boiling points.

Remember: the IMF hierarchy assumes similar-sized molecules. Size can override the hierarchy.

Test your understanding of IMF ranking.

Intermolecular forces directly determine a substance's physical properties. Stronger IMFs mean molecules are held together more tightly, which affects how the substance behaves.

The key physical properties influenced by IMF strength:

Notice that vapor pressure is the opposite — stronger IMFs mean LOWER vapor pressure.

Boiling point is the temperature at which a liquid's vapor pressure equals atmospheric pressure. At this point, molecules throughout the liquid have enough kinetic energy to overcome intermolecular forces and enter the gas phase.

Stronger IMFs (\rightarrow) Higher boiling point

Boiling requires breaking IMFs (not covalent bonds). The stronger the IMFs, the more energy (higher temperature) is needed.

Examples showing IMF effect on boiling point:

The trend clearly shows: as IMF strength increases, boiling point increases.

Use IMF analysis to predict relative boiling points.

Vapor pressure is the pressure exerted by the vapor above a liquid at a given temperature. It measures how easily molecules escape from the liquid phase into the gas phase.

Stronger IMFs (\rightarrow) LOWER vapor pressure

This is the inverse relationship! When IMFs are strong, molecules are held tightly in the liquid phase and fewer can escape to become vapor.

Key concept: Vapor pressure and boiling point are inversely related:

• High vapor pressure (\rightarrow) low boiling point (volatile substance)
• Low vapor pressure (\rightarrow) high boiling point (less volatile)

Examples:

• Diethyl ether (weak IMFs): high vapor pressure, evaporates quickly, low BP
• Water (hydrogen bonding): moderate vapor pressure, BP = 100°C
• Glycerol (extensive H-bonding): very low vapor pressure, very high BP

A substance with high vapor pressure is described as volatile.

Test your understanding of the relationship between IMFs and vapor pressure.

Surface Tension is the energy required to increase the surface area of a liquid. It arises because molecules at the surface only experience IMF attractions from the side and below (not above), creating a net inward force.

Stronger IMFs (\rightarrow) Higher surface tension

Water has exceptionally high surface tension due to its extensive hydrogen bonding network. This allows small insects to walk on water and enables water to form droplets.

Viscosity is a liquid's resistance to flow. More viscous liquids flow slowly (like honey), while less viscous liquids flow freely (like gasoline).

Stronger IMFs (\rightarrow) Higher viscosity

Glycerol ((\text{C}_3\text{H}_8\text{O}_3)) is very viscous because it has three O–H groups, allowing extensive hydrogen bonding. Each glycerol molecule can form multiple hydrogen bonds with its neighbors, creating a tangled network that resists flow.

Temperature effect: Increasing temperature decreases viscosity because molecules have more kinetic energy to overcome IMFs.

For each property, select how it changes with STRONGER intermolecular forces.

Complete the following statements about IMFs and physical properties.

Let's bring everything together for the AP exam. IMF questions are among the most commonly tested topics. Here's your complete toolkit:

The IMF Hierarchy (similar-sized molecules): $\text{Ion-Dipole} > \text{H-bonding} > \text{Dipole-Dipole} > \text{LDF}$

Quick Identification Guide:

1. Ion + polar molecule? (\rightarrow) Ion-dipole
2. H bonded to N, O, or F? (\rightarrow) Hydrogen bonding
3. Polar molecule? (\rightarrow) Dipole-dipole (+ LDF)
4. Nonpolar molecule? (\rightarrow) LDF only

Property Trends (stronger IMFs (\rightarrow)):

• Boiling point: (\uparrow)
• Melting point: (\uparrow)
• Surface tension: (\uparrow)
• Viscosity: (\uparrow)
• Vapor pressure: (\downarrow) (inverse!)

Remember: ALL molecules have LDF. Polar molecules have DD + LDF. H-bonding molecules have H-bond + DD + LDF.

Mistake 1: Saying hydrogen bonds are chemical bonds Hydrogen bonds are IMFs (between molecules), NOT chemical bonds (within molecules). Never say a molecule "has hydrogen bonds" — say it "exhibits hydrogen bonding" or "participates in hydrogen bonding."

Mistake 2: Thinking HCl has hydrogen bonding HCl has an H–Cl bond, but Cl is NOT N, O, or F. HCl only has dipole-dipole forces + LDF. The "hydrogen" in hydrogen bonding doesn't mean any bond to hydrogen — it specifically means H bonded to FON.

Mistake 3: Confusing IMFs with intramolecular bonds IMFs are forces BETWEEN molecules. Covalent, ionic, and metallic bonds are forces WITHIN a compound. Boiling and melting break IMFs, not covalent bonds.

Mistake 4: Ignoring molar mass when comparing LDF A large nonpolar molecule can have stronger IMFs than a small polar one. Always consider size.

Mistake 5: Forgetting that LDF are in ALL molecules Polar molecules have dipole-dipole AND LDF. Hydrogen bonding molecules have H-bonding AND DD AND LDF.

This question tests your ability to identify IMFs and predict properties.

Analyze intermolecular forces in a set of molecules.

Dissolution and ion-dipole forces.

For each substance, select the STRONGEST intermolecular force it experiences.

Watch out for this common AP trap!

Complete these essential AP Chemistry IMF facts.