Chemistry - Intermolecular Forces and Properties

Condensed phase

Solid and liquid phases. Particles in a solid or liquid are fairly close together compared with those of a gas.

Gas

Particles far apart; possess complete freedom of motion.

Liquid

Particle are closely packed but randomly oriented; retain freedom of motion; rapidly change neighbors.

Solid

Particles are closely packed in an ordered array; positions are essentially fixed.

Intermolecular force

Attraction between molecules; generally not as strong as intramolecular forces.

Intramolecular force

Attraction between atoms within a molecule. (Ionic, metallic, or covalent bonds)

Dispersion force "London dispersion force"

The motions of electrons in one atom influence the motions of electrons in its neighbors. The instantaneous dipole on one atom can induce an instantaneous dipole on an adjacent atom, causing the atoms to be attracted to each other. The strength of the dispersion force depends on the polarizability of a molecule.

Polarizability

The molecule's ease with which the charge distribution is distorted to induce an instantaneous dipole. The greater the polarizability, the more easily the electron cloud can be distorted to give an instantaneous dipole. Therefore, more polarizable molecules have larger dispersion forces.

Molecular weight vs. Dispersion

The polarizability increases as the number of electrons in an atom or molecules increase. The strength of dispersion forces therefore tends to increase with increasing atomic or molecular size. Because molecular size and mass generally parallel each other, dispersion forces tend to increase in strength with increasing molecular weight.

Molecular shape vs. Dispersion

Intermolecular attraction is greater for a linear molecule because the molecules can come in contact over the entire length of the long, somewhat cylindrical molecules. Less contact is possible between the more compact and nearly spherical molecules.

Linear molecule

Larger surface area enhances intermolecular contact and increases dispersion force.

Spherical molecule

Smaller surface area diminishes intermolecular contact and decreases dispersion force.

Instantaneous dipole

Temporary dipole that occurs for a brief moment in time when the electrons of an atom or molecule are distributed asymmetrically.

Induced dipole

A temporarily uneven distribution of electrons in an otherwise nonpolar atom or molecule.

Permanent dipole

These occur when two atoms in a molecule have substantially different electronegativity.

Dipole-Dipole force

Result of a permanent dipole moment in polar molecules. Electrostatic attractions between the partially positive end of one molecule an d the partially negative end of a neighboring molecule. Repulsion can also occur when the positive (or negative) ends of two molecules are in close proximity.

Polarity vs Dipole-Dipole

For molecules of approximately equal mass and size, the strength of intermolecular attractions increases with increasing polarity. (Boiling point increases as the dipole moment increases)

Hydrogen bond

Attraction between a hydrogen atom attached to a highly electronegative atom (usually F, O, or N) and a nearby small electronegative atom in another molecule or chemical group. Because the electron-poor hydrogen is so small, it can approach an electronegative atom very closely and, thus, interact strongly with it.

Ion-Dipole force

The attraction between an ion and a polar molecule. Cations are attracted to the negative end of a dipole, and anions are attracted to the positive end. The strength of the attraction increases as either the ionic charge or the magnitude of the dipole moment increases.

Comparing intermolecular forces

-When the molecules of two substances have comparable molecular weights and shapes, dispersion forces are approximately equal in the two substances. (generally use dipole-dipole forces)
-When the molecules of two substances differ widely in molecular weights, and there is no hydrogen bonding. (generally use dispersion forces)

Viscosity

A measure of the resistance of fluids to flow. The ease with which the molecules of the liquid can move relative to one another depend on the attractive forces between molecules and on whether the shapes and flexibility of the molecules are such that they tend to become entangled. (long molecules can become tangled like spaghetti)

Surface tension

The intermolecular, cohesive attraction that cause a liquid to minimize its surface area. The energy required to increase the surface area of a liquid by a unit amount.

Adhesive force

Intermolecular forces that bind a substance to a surface.

Cohesive force

Intermolecular forces that bind similar molecules to one another, such as the hydrogen bonding in water.

Capillary action

The process by which a liquid rises in a tube because of a combination of adhesion to the walls of the tube and cohesion between liquid particles.

Phase changes

The conversion of a substance from one state of matter to another. The phase changes we consider are melting and freezing (solid ⇌ liquid), sublimation and deposition, and vaporization and condensation (liquid ⇌ gas).

Dynamic equilibrium

A state of balance in which opposing processes occur at the same rate.

heat of fusion

The enthalpy change, ∆Hfus for melting a solid. (~6.01Kj/mol - water)

heat of vaporization

The enthalpy change, ∆Hvap for vaporization of a liquid. (~40.7Kj/mol - water)

heat of sublimation

The enthalpy change, ∆Hsub for vaporization of a solid. The sum of heat of fusion and heat of vaporization. (~47Kj/mol - water)

Normal melting point

The melting point at 1 atm pressure. (0 °C - water)

Boiling point

The temperature at which its vapor pressure equals the external pressure, acting on the liquid surface. At this temperature, the thermal energy of the molecules is great enough for the molecules in the interior of the liquid to break free from their neighbors and enter the gas phase. The boiling point increases as the external pressure increases.

Normal boiling point

The boiling point of a liquid at 1 atm (760 torr) pressure.

Intermolecular vs Heat of vap/fus

Stronger intermolecular forces means higher heat of vap/fus.

Supercooling

Heat is removed so rapidly that the molecules have no time to assume the ordered structure of a sold. A supercooled liquid is unstable; particles of dust entering the solution or gently stirring is often sufficient to cause the substances to solidify quickly.

Calculating ∆H for temperature and phase changes

Specific heat: (moles)(grams/mol)(J/g-K)(∆K) = Joules/1000 = Kilojoules
Heat of phase change: (moles)(Kj/mol) = Kilojoules

Critical pressure

The pressure at which a gas at its critical temperature is converted to a liquid state.

Critical temperature

The highest temperature at which it is possible to convert the gaseous form of a substance to a liquid. Increases with an increase in the magnitude of intermolecular forces.

1 atm is equal to

101,325 Pascals
760 Torr
14.6959 PSI

Clausius-Clapeyron Equation

The natural log of the vapor pressure of a liquid is inversely proportional to its temperature.
ln P = (-∆Hvap)/(R) * 1/(T) + C
y = m x + B
P - pressure
T - temperature in kelvin
R - ideal gas constant
∆Hvap - molar enthalpy of vaporization
C - constant

Ideal gas constant

62.36L⋅mmHg/K⋅mol
0.08206L⋅atm/K⋅mol
8.314J/K⋅mol

Enthalpy of vaporization of a substance

∆Hvap = -slope * R
slope = -∆Hvap/R

Vapor pressure

The pressure of air when evaporation and condensation are at equilibrium. At any temperature some liquid molecules have enough energy to escape the surface and become a gas. As temperature rises the fraction of molecules that have enough energy to break free increases.

Ionic solid

Held together by the mutual electrostatic attraction between cations and anions. Differences between ionic and metallic bonding make the electrical and mechanical properties of ionic solids very different from those of metals: Ionic solids do not conduct electricity well and are brittle.
- Soluble in water
- High melting point
- High lattice energy

Metallic solid

Held together by a delocalized “sea” of collectively shared valence electrons. This form of bonding allows metals to conduct electricity. It is also responsible for the fact that most metals are relatively strong without being brittle
- Malleability
- Ductility

Alloy

Material that contains more than one element and has the characteristic
properties of a metal.

Substitutional alloy

When atoms of the solute in a solid solution occupy positions normally occupied by a solvent atom. Atoms of solute around the same size of solvent atoms. Formed when the two metallic components have similar
atomic radii and chemical-bonding characteristics.

Interstitial alloy

When the solute atoms occupy interstitial positions in the “holes” between solvent atoms. The solute atoms must have a much smaller bonding atomic radius than the solvent atoms

Hetergeneous alloy

Components not dispersed uniformity.

Molecular solid

Molecular solids are held together by the intermolecular forces: dispersion forces, dipole–dipole interactions, and hydrogen bonds. Because these forces are relatively weak, molecular solids tend to be soft and have low melting points.

Network covalent solid

Held together by an extended network of covalent bonds. This type of bonding can result in materials that are extremely hard, like diamond, and it is also responsible for the unique properties of semiconductors
- Tend to have higher melting and boiling points

Heating curve

Plot of heat vs. temperature. Within a phase, heat is the product of specific heat, sample mass, and temp change. The temperature of a substance does not change during a phase change.