Chapter 7: The Three States of Matter and Intermolecular Forces

Properties and Characteristics of the Three States of Matter

Existing as a gas, liquid, or solid depends on two primary factors: the balance between the kinetic energy of its particles and the strength of the interactions (intermolecular forces) between those particles.

• Gas

  • Shape and Volume: Expands to fill its entire container.

  • Arrangement of Particles: Randomly arranged, disorganized, and far apart.

  • Density: Low (< 0.01\,g/mL).

  • Particle Movement: Very fast.

  • Interaction Between Particles: None.

• Liquid

  • Shape and Volume: Features a fixed volume that takes the shape of the container it occupies.

  • Arrangement of Particles: Randomly arranged but close together.

  • Density: High ($1\text{ to }10\,g/mL$).

  • Particle Movement: Moderate.

  • Interaction Between Particles: Strong.

• Solid

  • Shape and Volume: Possesses a definite shape and volume.

  • Arrangement of Particles: Fixed arrangement of very close particles.

  • Density: High ($1\text{ to }10\,g/mL$).

  • Particle Movement: Slow.

  • Interaction Between Particles: Very strong.

Introduction to Intermolecular Forces

Intermolecular forces (IMFs) are the attractive forces that exist between molecules. The strength of these forces determines whether a compound has a high or low melting point and boiling point, as well as its physical state (solid, liquid, or gas) at a given temperature.

• Categories of IMFs (In order of increasing strength):

  1. London dispersion forces

  2. Dipole–dipole interactions

  3. Hydrogen bonding

London Dispersion Forces

London dispersion forces are very weak interactions caused by momentary changes in electron density within a molecule.

• Mechanism:

  • A change in electron density creates a temporary (instantaneous) dipole.

  • More electron density in one region creates a partial negative charge ($\delta^-$).

  • Less electron density in another region creates a partial positive charge ($\delta^+$).

  • The weak attraction between these temporary dipoles constitutes the London dispersion force.

• Occurrence: All covalent compounds exhibit London dispersion forces.

• Scaling: The larger the molecule, the larger the attractive force, and the stronger the resulting intermolecular forces.

Dipole–Dipole Interactions

Dipole–dipole interactions are the attractive forces between the permanent dipoles of two polar molecules.

• Characteristics of Polar Molecules:

  • They contain one or more polar bonds.

  • They have individual bond dipoles that do not cancel out due to the molecule's shape (symmetry matters).

• Identifying Nonpolar Molecules: A molecule is generally nonpolar if it meets three criteria:

  1. All outer atoms are identical.

  2. There are no lone pairs on the central atom.

  3. The overall shape is symmetric.

• Molecular Polarity Practice Identification:

  • $CO_2$: Non-polar

  • $H_2O$: Polar

  • $CCl_4$: Non-polar

  • $HCl$: Polar

  • $BCl_3$: Non-polar

  • $NH_3$: Polar

  • $N_2$: Non-polar

  • $CH_3OH$: Polar

  • $CH_3NH_2$: Polar

Hydrogen Bonding

Hydrogen bonding is the strongest of the three types of intermolecular forces.

• Definition: It occurs when a hydrogen atom bonded to an $O$, $N$, or $F$ atom is electrostatically attracted to an $O$, $N$, or $F$ atom in another molecule.

• Biological Application: In the DNA double helix, hydrogen bonding interactions (often visualized as dashed red lines) hold the structure together.

• Examples of Incapability: Formaldehyde and Trifluoromethane are not capable of hydrogen bonding because they do not contain a hydrogen atom directly bonded to an $O$, $N$, or $F$ atom.

Summary of Intermolecular Force Types

Comparing Strengths of Intermolecular Forces

To determine which species has stronger IMFs, one must evaluate the type of force present and molecular size:

• Comparison Exercises:

  • $NH_3$ vs. $CH_4$: $NH_3$ is stronger because it has hydrogen bonding, whereas $CH_4$ only has London dispersion forces.

  • $CH_3Cl$ vs. $CO_2$: $CH_3Cl$ is stronger because it has dipole-dipole interactions, while $CO_2$ only has London dispersion forces.

  • $HNO_3$ vs. $NO$: $HNO_3$ is stronger because it has hydrogen bonding, compared to the dipole-dipole interactions in $NO$.

  • $I_2$ vs. $F_2$: $I_2$ has stronger London dispersion forces because the molecules are larger.

  • $H_2O$ vs. $NO_2$: $H_2O$ is stronger because it has hydrogen bonding, while $NO_2$ has dipole-dipole interactions.

Effects of IMFs on Boiling Point and Melting Point

• Boiling Point (BP): The temperature at which a liquid is converted to the gas phase.

• Melting Point (MP): The temperature at which a solid is converted to the liquid phase.

• Relationship to IMFs: The stronger the intermolecular forces, the higher the boiling point and melting point.

  • Example 1: Methane ($CH_4$) has London forces only ($bp = -162\,^\circ C$, $mp = -183\,^\circ C$). Water ($H_2O$) has hydrogen bonding ($bp = 100\,^\circ C$, $mp = 0\,^\circ C$). Water has stronger forces, thus higher $bp$ and $mp$.

  • Example 2 (Molecular Weight): Propane ($bp = -42\,^\circ C$, $mp = -190\,^\circ C$) vs. Butane ($bp = -0.5\,^\circ C$, $mp = -138\,^\circ C$). Butane is a larger molecule with stronger forces and higher $bp/mp$.

• Molecular Weight Principle: When comparing two compounds with similar IMFs, the higher-molecular-weight compound generally has more surface area, a larger force of attraction, and therefore a higher boiling and melting point.

The Liquid State: Vapor Pressure and Phase Changes

• Evaporation: The conversion of liquids into the gas phase. This is an endothermic process, meaning it absorbs heat from the surroundings.

• Condensation: The conversion of gases into the liquid phase. This is an exothermic process, meaning it gives off heat to the surroundings.

• Vapor Pressure: The pressure exerted by gas molecules in equilibrium with the liquid phase.

  • Vapor pressure increases as temperature increases.

  • The boiling point of a liquid is defined as the temperature at which its vapor pressure equals $760\,mmHg$.

• IMF Relationship to Vapor Pressure: The stronger the intermolecular forces, the lower the vapor pressure at a given temperature. Stronger forces keep molecules in the liquid state; weaker forces allow more molecules to escape into the gas phase.

Questions and Discussion

Q: What types of intermolecular forces are present in the following molecules?

• Ar (Argon): London dispersion forces only. (Atoms exhibit London forces).

• HCl: London dispersion forces and dipole-dipole interactions. (Polar molecule).

• $CH_3Cl$: London dispersion forces and dipole-dipole interactions.

• $CH_3NH_2$: London dispersion forces, dipole-dipole interactions, and hydrogen bonding. (Contains an $N-H$ bond).

• $H_2O$: London dispersion forces, dipole-dipole interactions, and hydrogen bonding. (Contains $O-H$ bonds).

• $CO_2$: London dispersion forces only. ($CO_2$ is non-polar).

• $BCl_3$: London dispersion forces only. ($BCl_3$ is non-polar).

• $Br_2$: London dispersion forces only. (Non-polar molecule).

• $H_2$: London dispersion forces only. ($H_2$ has no net dipole).

• $CH_2O$ (Formaldehyde): London dispersion forces and dipole-dipole forces. (Polar molecule; no $H$ bonded to $O$).

• $CH_2F_2$, $CH_2Cl_2$, $CH_3I$: London dispersion forces and dipole-dipole forces. (All are polar).

• $HNO_3$, $HF$, $NH_3$: London dispersion forces, dipole-dipole forces, and hydrogen bonding. (All have $H$ bonded to $F$, $O$, or $N$).

Q: Comparing Butanol and Dimethyl ether (Same molecular formula):

• IMFs in Butanol: London dispersion, dipole-dipole, and hydrogen bonding (due to its alcohol group).

• IMFs in Dimethyl ether: London dispersion and dipole-dipole interactions.

• Higher Boiling Point: Butanol (due to stronger hydrogen bonding).

• Higher Vapor Pressure at a given temperature: Dimethyl ether (due to weaker IMFs).