AP Chemistry 3.1-3.2: Intermolecular Forces and Properties of Solids

Dipole-dipole attraction

An electrostatic attraction between the partially positive end of one polar molecule and the partially negative end of another. It occurs because polar molecules have permanent separations of charge.

How do size and electron count affect London dispersion forces?

Larger, heavier particles generally have more electrons and more easily distorted electron clouds, so they experience stronger dispersion forces. Their valence electrons are also farther from the nucleus and less tightly held.

Why do the boiling points of the halogens generally increase from F₂ to I₂?

All of these nonpolar diatomic molecules rely primarily on dispersion forces. Increasing molecular size and electron count increase polarizability, strengthening attractions and raising the boiling point: $F_2 < Cl_2 < Br_2 < I_2$.

How does molecular shape affect boiling points when molecules have the same formula?

More elongated molecules generally have greater surface area for intermolecular contact and stronger dispersion forces. Thus, among pentane isomers, $n$-pentane has a higher boiling point than isopentane, which has a higher boiling point than the more compact neopentane.

Rank CH₄, SiH₄, GeH₄, and SnH₄ from lowest to highest boiling point.

The order is $CH_4 < SiH_4 < GeH_4 < SnH_4$. These molecules are essentially nonpolar, so the larger molecules have stronger dispersion forces and higher boiling points.

Rank C₂H₆, C₃H₈, and C₄H₁₀ from lowest to highest boiling point.

The order is $C_2H_6 < C_3H_8 < C_4H_{10}$. Their attractions are primarily dispersion forces, which increase with molecular size and polarizability.

Intermolecular forces (IMFs)

Electrostatic attractions between particles such as atoms, molecules, or ions. IMFs help hold particles together and strongly influence physical properties such as boiling point, melting point, and phase.

Intramolecular forces versus intermolecular forces

Intramolecular forces act within a particle, such as covalent bonds holding the atoms of HCl together. Intermolecular forces act between particles and are generally much weaker for small molecules.

How do kinetic energy and intermolecular forces together determine the phase of a substance?

IMFs favor particles staying close together, while particle kinetic energy favors movement apart. A solid, liquid, or gas results from the balance between these attractions and the particles' kinetic energies.

How does increasing temperature affect a substance's phase?

Increasing temperature increases the particles' average kinetic energy. If the kinetic energy becomes sufficient to overcome attractions, the substance may melt, vaporize, or undergo another transition to a less condensed phase.

How can increasing pressure liquefy a gas?

Compression brings gas particles closer together, increasing the opportunity for intermolecular attractions to hold them together. At a suitable temperature, the attractions can then become significant relative to the particles' kinetic energy, causing condensation.

How are particles arranged and how do they move in solids, liquids, and gases?

Solid particles are closely packed, often ordered, and vibrate about fixed positions. Liquid particles remain close but move past one another, whereas gas particles are far apart and move independently except during collisions.

London dispersion forces

Attractions caused by temporary instantaneous dipoles that induce dipoles in nearby particles. Dispersion forces occur in all atoms and molecules and are especially important for nonpolar substances.

How do instantaneous and induced dipoles produce a dispersion force?

Random electron motion can momentarily make one side of an atom or molecule more negative than the other. This temporary dipole distorts a neighboring electron cloud, creating an induced dipole whose opposite charge attracts the original dipole.

Which type of intermolecular force is present in every molecular substance?

London dispersion forces are present in every atom and molecule. Polar molecules also have dispersion forces in addition to their other possible attractions.

Polarizability

Polarizability is the ease with which a particle's electron cloud can be distorted by a nearby charge or dipole. Greater polarizability produces stronger London dispersion forces.

What structural feature allows a molecule to experience dipole-dipole attractions?

The molecule must have a permanent net dipole, meaning its bond dipoles do not cancel because of its bonding or molecular geometry. Polar molecules experience dipole-dipole attractions as well as dispersion forces.

Why does HCl have a higher boiling point than F₂ even though they have similar masses?

Both substances have dispersion forces of similar general magnitude, but HCl is polar and also has dipole-dipole attractions. Nonpolar F₂ lacks this additional attraction, so it boils at a lower temperature.

Which has the higher boiling point, CO or N₂, and why?

CO has the higher boiling point. CO and N₂ have similar masses and therefore similar dispersion forces, but CO is polar and experiences additional dipole-dipole attractions.

Which has the higher boiling point, ICl or Br₂, and why?

ICl has the higher boiling point. Their similar masses imply comparable dispersion forces, but polar ICl also has dipole-dipole attractions, whereas nonpolar Br₂ does not.

Hydrogen bonding

A particularly strong dipole-dipole attraction involving hydrogen covalently bonded to N, O, or F and an electronegative N, O, or F atom on a neighboring particle. Hydrogen bonding is intermolecular, not the covalent bond within the molecule.

What atoms must be involved for a molecule to donate a hydrogen bond?

The molecule must contain an $N-H$, $O-H$, or $F-H$ bond. The strongly polarized bond creates a concentrated partial positive charge on hydrogen that can attract a nearby electronegative atom.

Why are hydrogen bonds stronger than ordinary dipole-dipole attractions?

N, O, and F are very electronegative, and hydrogen is very small, so the partial charges are highly concentrated. This produces unusually strong electrostatic attractions between neighboring molecules.

Why is water a liquid at room temperature while the larger molecule ONF is a gas?

Water molecules form extensive hydrogen bonds through their O-H bonds. ONF has dispersion and dipole-dipole attractions but cannot form comparable hydrogen bonds, so its intermolecular attractions are much weaker despite its greater mass.

Rank propane, dimethyl ether, and ethanol from lowest to highest boiling point.

The order is $CH_3CH_2CH_3 < CH_3OCH_3 < CH_3CH_2OH$. Propane is nonpolar, dimethyl ether is polar, and ethanol forms hydrogen bonds; their similar sizes make the differences in IMFs especially important.

Why does methylamine have a much higher boiling point than ethane despite their similar masses?

Methylamine contains N-H bonds and can form hydrogen bonds, while ethane is nonpolar and has only dispersion forces. The stronger attractions in methylamine substantially raise its boiling point.

Ion-dipole attraction

An electrostatic attraction between an ion and the oppositely charged end of a polar molecule. For example, water molecules orient their partially negative oxygen atoms toward cations and their partially positive hydrogen atoms toward anions.

Why do ionic compounds often dissolve in water?

Ion-dipole attractions between water molecules and the ions can stabilize separated ions. Dissolution is favored when these attractions, along with other changes in energy, compensate for disrupting the ionic lattice and separating water molecules.

How can boiling point, melting point, or enthalpy of vaporization be used to compare intermolecular forces?

For substances of comparable composition and structure, stronger IMFs generally require more energy to separate particles. Therefore, higher boiling or melting points and larger enthalpies of vaporization or fusion often indicate stronger attractions.

How do intermolecular and intramolecular energies compare in liquid HCl?

Separating liquid HCl molecules into a gas requires much less energy than breaking the H-Cl covalent bonds within the molecules. This illustrates that intermolecular attractions are typically far weaker than intramolecular bonds.

Why can geckos adhere to both polar and nonpolar surfaces?

Millions of microscopic spatulae on a gecko's toes create enormous contact area. Dispersion forces between the temporary charge distributions of the spatulae and the surface add together to produce substantial adhesion.

How can geckos switch between adhesion and release?

Changing the angle at which their spatulae contact a surface changes the effective contact and shear forces. Curling and uncurling their toes allows them to engage or disengage the many weak dispersion-force contacts.

What is a crystalline solid?

A solid whose atoms, ions, or molecules occupy a definite, repeating, ordered arrangement. Most metals and ionic compounds typically form crystalline solids.

What is an amorphous solid?

A noncrystalline solid whose particles lack a long-range repeating arrangement and are more randomly organized. Rapid cooling or restricted molecular motion can prevent a liquid from forming an ordered crystal.

What determines whether a substance forms a crystalline or amorphous solid during cooling?

The particles must have enough time and mobility to arrange into an ordered pattern for crystallization. If cooling is sufficiently rapid or molecular motion is restricted, the substance may solidify in a disordered amorphous state.

How can an amorphous solid become crystalline?

Under appropriate conditions, particles in an amorphous solid can gain enough mobility to reorganize into a repeating structure. This produces a transition from the noncrystalline to the crystalline state.

How do hydrogen bonds contribute to the structure of DNA?

Hydrogen bonds between complementary nitrogenous bases hold the two DNA strands together. Adenine pairs with thymine through two hydrogen bonds, while cytosine pairs with guanine through three.

Why can the two strands of DNA unzip relatively easily even though many attractions hold them together?

The cumulative hydrogen bonds stabilize the double helix, but each hydrogen bond is much weaker than the covalent bonds within each strand. The strands can therefore separate without breaking the covalent backbone.

What are the four major types of crystalline solids?

The four major types are molecular solids, ionic solids, metallic solids, and covalent network solids. They differ in the particles occupying lattice sites and in the forces or bonds holding those particles together.

What are molecular solids, and what properties do they typically have?

Molecular solids consist of discrete molecules held together by intermolecular forces. They generally have relatively low melting points, are soft or brittle, and do not conduct electricity because their particles are neutral and not freely mobile.

What are ionic solids, and what properties do they typically have?

Ionic solids consist of cations and anions arranged in a crystal lattice and held together by strong electrostatic attractions. They are typically hard, brittle, and high-melting; they do not conduct as solids but conduct when molten or dissolved because the ions can move.

Why are ionic solids brittle?

When layers of an ionic crystal shift, ions of like charge can become adjacent. Their strong repulsion causes the lattice to fracture rather than deform smoothly.

What are metallic solids, and what properties do they typically have?

Metallic solids consist of metal atoms or cations in a lattice surrounded by delocalized valence electrons. They generally conduct heat and electricity, and their nondirectional metallic bonding makes them malleable and ductile.

What are covalent network solids, and what properties do they typically have?

Covalent network solids contain atoms connected throughout an extended lattice by covalent bonds. They are typically very hard, have very high melting points, and are poor electrical conductors, although graphite conducts because it has delocalized electrons.

How do the bonding and particles in a solid determine its melting point and hardness?

Stronger attractions or bonds between particles generally require more energy to melt and resist deformation more effectively. Thus, ionic and covalent network solids tend to have high melting points, while molecular solids usually melt more easily; hardness also depends on how readily the structure can deform.

How does the electrical conductivity of solid types compare?

Metallic solids conduct as solids because they contain mobile delocalized electrons. Ionic solids conduct when molten or dissolved but not as solids because their ions are fixed. Molecular and most covalent network solids are poor conductors, with graphite as an important exception.

How do the properties of molecular, ionic, metallic, and covalent network solids compare?

Molecular solids are generally soft, low-melting, and nonconductive. Ionic solids are hard, brittle, high-melting, and conductive only when ions can move. Metallic solids are conductive, malleable, and ductile. Covalent network solids are very hard and high-melting, and are usually nonconductive.