11.1 Intermolecular Forces
PRE-HEALTH PREP
Intermolecular Forces
Hydrogen Bonding, London Dispersion & More
General Chemistry 11.1
Lesson Overview
Intermolecular forces—hydrogen bonding, dipole-dipole, London dispersion, and ion-dipole—are the attractive interactions between separate molecules that govern physical properties such as boiling point, viscosity, surface tension, and vapor pressure. Hydrogen bonding requires a hydrogen atom directly bonded to fluorine, oxygen, or nitrogen, while London dispersion forces are universal and scale with molecular size and surface area. When comparing molecules, the relative strength of these intermolecular forces depends on both the type of interaction present and the size difference between molecules, with larger nonpolar molecules sometimes exceeding smaller polar ones in overall intermolecular attraction.
What Are Intermolecular Forces?
Intermolecular forces are attractive forces that act between separate molecules. The prefix "inter-" means between, so intermolecular forces describe the interactions between distinct molecules — whether in a pure molecular substance or a mixture of molecular compounds. These forces are responsible for the physical properties of liquids and solids, including boiling point, viscosity, and vapor pressure.
Intermolecular forces do not apply to ionic compounds or network covalent compounds. In an ionic solid, all ions are held together by ionic bonds. In a network covalent solid like diamond, all atoms are linked by covalent bonds. But consider a molecular solid such as ice: the individual atoms within a single water molecule are joined by covalent bonds, yet what holds one water molecule to its neighbors is something far weaker — intermolecular forces.
Intermolecular Forces Are Much Weaker Than Chemical Bonds
Intermolecular forces are roughly 25 to 100 times weaker than ionic or covalent bonds. As a consequence, molecular compounds are generally held together much less strongly than ionic or network covalent solids, and they tend to have significantly lower melting and boiling points.
What Are the 4 Types of Intermolecular Forces?
There are four intermolecular forces you need to know: hydrogen bonding, dipole-dipole forces, London dispersion forces, and ion-dipole forces. Ion-dipole forces get their own discussion because they occupy a somewhat different category — they involve an ion and a molecule, not two molecules.
For the first three, a useful general ranking applies when all molecules under comparison are similar in size:
| Force | Relative Strength | Required Condition |
|---|---|---|
| London Dispersion | Weakest | All molecules (universal) |
| Dipole-Dipole | Intermediate | Polar molecules |
| Hydrogen Bonding | Strongest (of the three) | Polar molecule with F–H, O–H, or N–H bond |
Important Caveat: This Ranking Assumes Similar Molecular Sizes
The ranking above only holds reliably when the molecules being compared are close in size. When molecular sizes differ significantly, London dispersion forces can override the other forces — even hydrogen bonding, in extreme cases. The worked examples later in this lesson address exactly this situation.
| INTERMOLECULAR FORCES — Quick Reference | |
|---|---|
| Hydrogen Bonding | -a super strong dipole-dipole force -must have an F-H, O-H, or N-H bond as a pure liquid -must only have F, O, N to hydrogen bond with water (or other H-bond donor) |
| Dipole-Dipole Forces | -interaction between molecules having permanent dipole moments -the larger the dipole moment, the larger the force |
| London Dispersion Forces (van der Waals Forces) | -weak interactions due to a transient (temporary) dipole -all molecules have these but is the only force present for nonpolar molecules -higher molecular weight, larger surface area, and greater polarizability result in higher London forces |
| Ion-Dipole Forces | -interaction between ions and polar molecules -present when ions are dissolved in polar liquids (such as NaCl in H2O) |
Dipole-Dipole Forces
Dipole-dipole forces are intermolecular forces that exist between polar molecules. Recall that a polar molecule possesses a permanent dipole moment — an unequal distribution of electron density that creates partially positive (δ+) and partially negative (δ−) regions.
Example: HCl
Hydrogen chloride is a classic illustration. Chlorine is significantly more electronegative than hydrogen, so the H–Cl bond is highly polar: chlorine carries a partial negative charge (δ−) and hydrogen carries a partial positive charge (δ+).
Now consider two HCl molecules side by side. The partially negative chlorine of one molecule is attracted to the partially positive hydrogen of the adjacent molecule. That attraction — between two separate molecules, not within a single molecule — is a dipole-dipole force. It is far weaker than the covalent H–Cl bond itself, but it is real and measurable.
In general, the more polar the molecules, the stronger the dipole-dipole forces between them.
Dipole-dipole forces are most meaningful when comparing pure substances of similar size, because that is when you rank properties like boiling point, melting point, viscosity, surface tension, and vapor pressure. In a mixture of different polar molecules you can still identify which forces are present, but you would not ordinarily compare those bulk properties across the components of the mixture.
Hydrogen Bonding
Requirements for Hydrogen Bonding
Hydrogen bonding can be thought of as an exceptionally strong form of dipole-dipole interaction. However, it only arises under a specific structural condition: a hydrogen atom must be directly bonded to one of three small, highly electronegative atoms — fluorine (F), oxygen (O), or nitrogen (N).
Simply having hydrogen atoms and one of these electronegative atoms in the same molecule is not sufficient. The hydrogen must be bonded directly to the F, O, or N. In other words, the molecule must contain at least one of the following bonds:
- F–H (hydrogen fluoride, HF)
- O–H (water, alcohols, carboxylic acids)
- N–H (ammonia, amines, amides)
Memory Aid: "FON"
Hydrogen bonding requires a hydrogen directly bonded to Fluorine, Oxygen, or Nitrogen. These 3 elements are among the most electronegative in the periodic table, and their small atomic radii allow the partial charges to get close enough for an unusually strong interaction.
Why Hydrogen Bonds Are So Strong
When hydrogen is bonded to F, O, or N, the bond is intensely polar. The hydrogen carries an exceptionally large partial positive charge, and the electronegative atom carries a correspondingly large partial negative charge. As a result, the attraction between that hydrogen and a lone pair on a neighboring F, O, or N atom is far greater than an ordinary dipole-dipole interaction — strong enough to warrant its own name.
In this interaction, 2 roles are defined:
- The hydrogen bond donor is the molecule that contributes the F–H, O–H, or N–H hydrogen.
- The hydrogen bond acceptor is the molecule that contributes the lone pair on its F, O, or N atom.
Water: The Classic Hydrogen Bonding Example
Water is the most familiar molecule exhibiting hydrogen bonding. Each water molecule has two O–H bonds, making it a hydrogen bond donor up to twice. The oxygen also carries two lone pairs, making the same molecule a hydrogen bond acceptor up to twice. In liquid water, each molecule can therefore interact with up to four neighbors simultaneously.
Why Ice Is Less Dense Than Liquid Water
In ice, every water molecule participates in exactly four hydrogen bonds in a tetrahedral network. To achieve this perfectly ordered crystal structure, the molecules must space themselves out slightly more than they do in the looser arrangement of liquid water. This is why freezing water causes it to expand — the opposite of what almost every other substance does when it solidifies. This anomalous behavior will be revisited in the next lesson.
Formaldehyde: A Common Misconception About Hydrogen Bonding
Formaldehyde (CH₂O) contains both hydrogen atoms and an oxygen atom, which causes many students to assume it can hydrogen bond. It cannot. Examining the Lewis structure reveals that neither hydrogen is bonded to the oxygen — both hydrogens are bonded to the carbon. There is no O–H bond anywhere in the structure.
Common Mistake: Confusing Hydrogen Bond Acceptors with Donors
Formaldehyde does have an oxygen with lone pairs, so it qualifies as a hydrogen bond acceptor — it could interact with an O–H or N–H hydrogen from an adjacent molecule. However, it is not a hydrogen bond donor, because its own hydrogens are bonded to carbon. The C–H bond is effectively nonpolar, so those hydrogens carry no significant partial positive charge and cannot participate in hydrogen bonding. The result: formaldehyde exhibits only dipole-dipole forces (it is a polar molecule), not hydrogen bonding.
London Dispersion Forces
Temporary Dipoles and Induced Dipoles
Even perfectly nonpolar molecules attract one another — otherwise noble gases and nonpolar hydrocarbons could never condense into liquids. The explanation lies in the continuous motion of electrons. At any given instant, the electron density in a molecule is not perfectly symmetric. This creates a fleeting, temporary dipole. That temporary dipole then induces a dipole in a neighboring molecule, and the two experience a momentary attraction. For this reason, London dispersion forces are also called induced dipole forces.
These instantaneous attractions are short-lived and constantly shifting, but they are always present and always additive over a large surface area.
All Molecules Have London Dispersion Forces
London dispersion forces are universal — every molecule has them, because every molecule has electrons. Nonpolar molecules rely on London dispersion forces exclusively, since they lack a permanent dipole. Polar molecules have London dispersion forces in addition to dipole-dipole forces, and molecules with F–H, O–H, or N–H bonds have all three types of intermolecular forces.
Which Intermolecular Forces Does a Molecule Have?
Think of it as a hierarchy: start at London dispersion (every molecule has this), then ask whether the molecule is polar (add dipole-dipole), then ask whether it has an F–H, O–H, or N–H bond (add hydrogen bonding). Exam questions often ask either "list all intermolecular forces present" or "identify the strongest intermolecular force." For the strongest, simply identify the highest tier that applies.
What Determines the Strength of London Dispersion Forces?
2 factors govern the magnitude of London dispersion forces: molecular size (closely related to molecular weight) and surface area. Larger molecules have more electrons, which creates a greater probability of significant temporary dipoles and stronger induced-dipole interactions. Greater surface area provides more sites for these momentary attractions to occur.
In general chemistry, molecular weight is the primary metric you will use to gauge London dispersion force strength. As a practical guideline: if two molecules differ in molecular weight by roughly 10–15% or less, that difference alone is unlikely to override other factors. If the difference is substantially larger than 15%, London dispersion forces will likely dominate regardless of polarity.
Ion-Dipole Forces
What Are Ion-Dipole Forces?
Ion-dipole forces are attractive interactions between an ion and a polar molecule. They are listed alongside intermolecular forces in most textbooks, but strictly speaking the term "intermolecular" (between molecules) is a misnomer here — ions are not molecules. Still, ion-dipole forces are important enough to discuss in this context.
Because ion-dipole forces require both an ion and a polar molecule, they can only exist in mixtures — never in a pure substance. You will never rank boiling points or viscosities using ion-dipole forces, since those discussions concern pure liquids.
The Classic Example: Dissolving NaCl in Water
Dissolving sodium chloride in water is the textbook illustration. When NaCl dissociates, the resulting Na⁺ and Cl⁻ ions are surrounded by water molecules. Water molecules orient themselves around Na⁺ with their partially negative oxygen atoms pointing inward, maximizing the attraction between the anion end of the water dipole and the positive ion. Around Cl⁻, water molecules flip orientation, presenting their partially positive hydrogen atoms toward the anion. This organized shell of water molecules — held in place by ion-dipole forces — is called a hydration shell, and the process is what makes ionic compounds soluble in polar solvents.
Strength of Ion-Dipole Forces
Ion-dipole forces span a wide range of strengths. The key variables are the charge density of the ion (smaller ions with higher charge produce stronger forces) and the polarity of the solvent molecule. On average, ion-dipole forces are stronger than hydrogen bonding, though the two ranges overlap at their extremes.
Ranking All 4 Intermolecular Forces
From weakest to strongest, the general ranking is: London dispersion < dipole-dipole < hydrogen bonding < ion-dipole. However, for pure substances, ion-dipole is irrelevant. When comparing pure liquids, hydrogen bonding is the strongest force available.
How Do Intermolecular Forces Affect Boiling Point and Viscosity?
Intermolecular forces govern several measurable properties of liquids. The stronger the forces holding molecules together, the more energy is required to disrupt them — and that shows up in multiple measurable ways. Understanding this relationship is essential for ranking boiling points, melting points, and other physical properties on exams.
Boiling Point and Melting Point
At the boiling point, molecules transition from the liquid phase (where they are in continuous contact and experiencing intermolecular forces) to the gas phase (where they are separated by large distances and those forces are negligible). Breaking these forces requires energy. The stronger the intermolecular forces, the more energy needed, and therefore the higher the boiling point.
The same logic applies to melting point — the transition from solid to liquid also requires breaking intermolecular forces, just fewer of them than the liquid-to-gas transition requires. The enthalpy of vaporization (the heat absorbed in boiling) and the enthalpy of fusion (the heat absorbed in melting) both increase with stronger intermolecular forces.
Viscosity
Viscosity is the resistance of a liquid to flow. Think of the difference between pouring honey and pouring water — honey is far more viscous. When a liquid flows, adjacent layers of molecules must slide past one another. Intermolecular forces create friction between these layers. Stronger intermolecular forces produce greater friction, which makes the liquid flow more slowly and thickly. Greater intermolecular forces therefore mean higher viscosity.
Surface Tension
Surface tension is the tendency of a liquid's surface to resist disruption. Molecules in the interior of a liquid are attracted equally in all directions by their neighbors. Molecules at the surface, however, are only attracted from the sides and below — there are no molecules above to balance the force. This net inward pull contracts the surface and creates a measurable tension.
Water Bugs and Overfilled Glasses
Two everyday manifestations of surface tension both trace back to water's hydrogen bonding network. A water strider (water bug) can walk on water because its body weight, distributed across its legs, is insufficient to rupture the strong hydrogen bonding network at the surface. Similarly, you can fill a glass of water slightly above its rim — water drops added one at a time will cause the surface to dome upward before it finally overflows. The stronger the intermolecular forces, the higher the surface tension.
Vapor Pressure
Vapor pressure is the one bulk property that moves in the opposite direction from the others. It measures the tendency of molecules to escape from the liquid surface into the vapor phase. For a molecule to escape, it must have enough kinetic energy to overcome the intermolecular forces holding it in the liquid. Stronger intermolecular forces raise that energy threshold, so a smaller fraction of molecules at any given temperature can escape. Stronger intermolecular forces therefore mean lower vapor pressure.
Summary: Direction of Each Effect
Stronger intermolecular forces → higher boiling point, melting point, enthalpy of vaporization, enthalpy of fusion, viscosity, and surface tension — but lower vapor pressure. Vapor pressure is the one property that runs contrary to all the others. Do not forget it.
Higher Intermolecular Forces Result in:
↑ Higher Boiling Pts
↑ Higher Enthalpy of Vaporization
↑ Greater Viscosity
↑ Greater Surface Tension
↑ Higher Critical Temperature/Pressure
↓ Lower Vapor Pressure
How Does Vapor Pressure Relate to Boiling Point?
The Vapor Pressure Curve
A vapor pressure curve plots vapor pressure (in torr or atm) on the y-axis against temperature (in °C) on the x-axis. As temperature increases, more molecules gain sufficient kinetic energy to escape the liquid, so vapor pressure rises — not linearly, but along a steep, logarithmic curve.
The key relationship: a substance boils when its vapor pressure equals the external (atmospheric) pressure. For water at sea level, that point is reached at 100 °C, where the vapor pressure equals 760 torr (1 atm). This is not a coincidence or definition — it is a physical consequence of the energy balance between intermolecular forces and kinetic energy at that temperature.
Boiling Point Changes with Altitude
At higher elevations, atmospheric pressure is lower. The vapor pressure threshold that triggers boiling is therefore reached at a lower temperature. On a tall mountain, water might boil at only 90 °C rather than 100 °C.
High-Altitude Cooking
This is not merely a chemistry curiosity — it has real culinary consequences. A recipe that instructs you to boil rice for 10 minutes assumes a boiling temperature of 100 °C. At high altitude, boiling occurs at a lower temperature, so the same rice takes longer to cook and requires adjustment. Cookbooks designed for high-altitude kitchens exist for precisely this reason.
Comparing Intermolecular Forces: 4 Worked Examples
The following 4 examples each present a set of 3 molecules. The goal in each case is to identify which molecule has the greatest overall intermolecular forces — and to reason through why. These intermolecular forces examples cover the most common scenarios you will encounter on a general chemistry exam.
Example 1: All Nonpolar Hydrocarbons
Consider a set of molecules containing only carbon and hydrogen (hydrocarbons). All C–H and C–C bonds are nonpolar, so each molecule is nonpolar overall. None of them have dipole-dipole forces or hydrogen bonding — London dispersion forces are the only game in town.
When comparing molecules that rely solely on London dispersion forces, size is the deciding factor. The largest molecule has the most electrons, the strongest London dispersion forces, and therefore the greatest overall intermolecular forces. Consequently, it has the highest boiling point, highest melting point, highest viscosity, highest surface tension — and the lowest vapor pressure.
Vapor Pressure Goes the Other Way
When asked which molecule in a nonpolar series has the highest vapor pressure, pick the smallest one. Weaker intermolecular forces make it easier for molecules to escape into the vapor phase.
Example 2: Similar-Sized Molecules, Different Force Types
Now consider 3 molecules with very similar molecular weights: CH₃CH₃ (MW ≈ 30), CH₃F (MW ≈ 34), and CH₃OH (MW ≈ 32). Because the sizes are so close, the type of force becomes the deciding factor.
- CH₃CH₃ (ethane): no polar bonds, nonpolar molecule — London dispersion forces only.
- CH₃F (fluoromethane): polar molecule, but the H atoms are bonded to carbon, not to fluorine — so there is no H–F bond. No hydrogen bonding. Dipole-dipole forces apply.
- CH₃OH (methanol): contains an O–H bond. Hydrogen bonding applies.
Common Mistake: Assuming CH₃F Can Hydrogen Bond
Students frequently assume that because CH₃F contains fluorine — the most electronegative element — it must hydrogen bond. It does not. Draw the Lewis structure: the 3 hydrogens are bonded to carbon, not to fluorine. There is no H–F bond. Hydrogen bonding requires the hydrogen to be directly attached to F, O, or N.
Since all 3 molecules are similar in size, the IMF type determines the ranking. CH₃OH has hydrogen bonding and therefore the greatest intermolecular forces — the highest boiling point, melting point, viscosity, and surface tension, but the lowest vapor pressure.
Example 3: Polar vs. Nonpolar with a Large Size Difference
Consider CH₄ (MW = 16), CH₃CH₃ (MW = 30), and CH₃F (MW = 34). Two nonpolar molecules and one polar molecule — no hydrogen bonding in any of them.
The instinct to pick the polar molecule is understandable, but check the sizes first. CH₄ is dramatically smaller than the other two (MW = 16 vs. 30–34), so it almost certainly has the weakest forces. Set it aside. Now compare CH₃CH₃ (MW = 30) and CH₃F (MW = 34): these differ by about 13%, well within the "similar size" threshold. With similar sizes, the polar molecule wins — CH₃F adds dipole-dipole forces on top of London dispersion forces and takes the top spot.
Result: CH₃F Has the Greatest IMFs
CH₃F wins because (1) it is slightly larger than CH₃CH₃, giving it slightly stronger London dispersion forces, and (2) it is polar, adding dipole-dipole forces that CH₃CH₃ lacks. Both effects favor CH₃F.
Example 4: When a Larger Nonpolar Molecule Beats a Polar One
Consider CH₃CH₃ (MW = 30), CH₃F (MW = 34), and C₃H₈ (propane: MW = 44). Again, no hydrogen bonding present.
Now compare the size of the largest molecule to the polar one: 44 vs. 34 is a difference of roughly 29% — far beyond the 10–15% threshold. At this size differential, the extra London dispersion forces in C₃H₈ outweigh the dipole-dipole advantage of CH₃F. The larger nonpolar molecule wins.
Result: C₃H₈ Has the Greatest IMFs
Despite being nonpolar, C₃H₈ is large enough (MW = 44) that its London dispersion forces exceed the combined London dispersion + dipole-dipole forces of CH₃F (MW = 34). Size wins when the difference is large enough.
This pair of examples illustrates the crucial nuance: in Example 3 the polar molecule won because the size difference was small (~13%); in Example 4 the nonpolar molecule won because the size difference was large (~29%). The lesson is that polarity does not automatically trump size.
How to Compare Intermolecular Forces: 3 Decision Rules
Applying intermolecular forces concepts on an exam comes down to a clear decision process. Here are the 3 rules that will serve you well:
- First, check for hydrogen bonding. If any molecule in the comparison has an F–H, O–H, or N–H bond, it almost certainly has the greatest intermolecular forces — unless another molecule is dramatically larger (roughly double the molecular weight or more).
- If no hydrogen bonding is present, compare sizes. If the molecules are similar in size (within ~10–15% by molecular weight), the polar molecule wins over the nonpolar one. If they differ substantially in size, the larger molecule wins regardless of polarity.
- Remember that vapor pressure goes the opposite direction. The molecule with the greatest intermolecular forces has the lowest vapor pressure, not the highest.
Can a Giant Nonpolar Molecule Beat Hydrogen Bonding?
In principle, yes. A very large nonpolar hydrocarbon — think of a wax molecule like C₃₀H₆₂ — accumulates so many London dispersion interactions across its enormous surface area that it can have higher overall intermolecular forces than a small molecule with hydrogen bonding. This is why waxes are solid at room temperature despite being nonpolar. However, you would need roughly double the molecular weight (or more) before a nonpolar molecule could overcome hydrogen bonding, and general chemistry exams rarely require you to identify the exact crossover point.
Master these 3 decision rules, practice identifying which intermolecular forces are present in a given molecule, and you will be well equipped to rank boiling points, melting points, viscosities, surface tensions, and vapor pressures with confidence.
Key Takeaways
- Intermolecular forces act between separate molecules and are 25–100× weaker than covalent or ionic bonds; they do not apply to ionic or network covalent solids.
- Dipole-dipole forces exist between polar molecules; the more polar the molecules, the stronger the force.
- Hydrogen bonding — the strongest of the 3 molecular forces — requires a hydrogen atom bonded directly to F, O, or N. Having these atoms in the molecule is not enough; the H must be bonded to one of them.
- London dispersion forces are present in all molecules and increase with molecular size (molecular weight) and surface area.
- Ion-dipole forces operate between an ion and a polar molecule; they are generally stronger than hydrogen bonding but only exist in mixtures, never in pure substances.
- Stronger intermolecular forces produce higher boiling points, melting points, enthalpies of vaporization/fusion, viscosity, and surface tension — but lower vapor pressure.
- A substance boils when its vapor pressure equals the external pressure; at higher altitudes the lower atmospheric pressure means a lower boiling point.
- When comparing IMF strength: if sizes are similar (<15% difference in MW), polarity and force type are decisive; if sizes differ substantially, the larger molecule wins regardless of polarity.