3.1 Intermolecular Forces: Types, Relative Strengths, and Physical Properties

Fundamentals of Intermolecular vs. Intramolecular Forces

• Intramolecular forces are forces that hold atoms together within a single molecule (e.g., covalent bonds).
• Intermolecular forces (IMFs) are attractive or repulsive forces that occur between individual molecules or between atoms on different molecules.
• Formation of IMFs in polar substances:
  • In polar molecules such as water ($H_2O$), unequal sharing of electrons creates partial charges: a partial positive charge ($\delta+$) on hydrogen atoms and a partial negative charge ($\delta-$) on oxygen atoms.
  • The partial negative oxygen atom of one water molecule attracts the partial positive hydrogen atom of an adjacent water molecule.
  • This attraction is Coulombic (electrostatic) in nature.
• Relative strength comparison:
  • Like covalent and ionic bonds, IMFs are Coulombic electrostatic interactions.
  • However, IMFs are generally weaker than intramolecular covalent and ionic bonds.

Dipole-Dipole Interactions

• Dipole-dipole interactions occur specifically between two polar molecules.
• These interactions can be attractive or repulsive, but molecules naturally orient themselves in space to maximize attraction and minimize repulsion.
• The strength of dipole-dipole forces is directly related to the magnitude of the dipole moment within the molecule.
• Examples of polar molecules experiencing dipole-dipole interactions:
  • Carbon monoxide ($CO$)
  • Ammonia ($NH_3$)
  • Difluoromethane ($CH_2F_2$)

Dipole-Induced Dipole Interactions

• Dipole-induced dipole interactions occur between a polar molecule and a nonpolar molecule.
• Real-world application:
  • Nonpolar oxygen molecules ($O_2$) dissolve in polar water ($H_2O$), allowing aquatic organisms such as fish to take in oxygen underwater.
• Mechanism of attraction:
  • When the permanent dipole of a polar molecule ($H_2O$) approaches a nonpolar molecule ($O_2$), the negative partial charge of water repels the electrons in $O_2$.
  • This repulsion distorts the electron cloud of $O_2$, forcing the nonpolar molecule to temporarily adopt an induced dipole.
• Proximity and orientation dependencies:
  • The temporary dipole in $O_2$ depends on the orientation and proximity of the nearby water molecule.
  • If the water molecule moves away, the induced dipole in $O_2$ ceases to exist.
  • Dipole-induced dipole interactions are always attractive.

London Dispersion Forces (LDFs)

• London dispersion forces (also called induced dipole-induced dipole interactions or LDFs) occur between nonpolar molecules (such as two $Br_2$ molecules or two $O_2$ molecules).
• Mechanism of LDF formation:
  • Even in nonpolar molecules, electrons continuously move and undergo random fluctuations.
  • These random electron fluctuations create temporary, instantaneous dipoles.
  • When a temporary dipole in one nonpolar molecule approaches another nonpolar molecule, it induces a small temporary dipole in the second molecule.
  • Partial positive and negative regions form, resulting in an electrostatic attraction between the two molecules.
• Universal presence and polarizability:
  • All molecules exhibit London dispersion forces, including polar molecules.
  • LDFs are the primary type of intermolecular attraction present between nonpolar molecules.
  • The strength of LDFs depends on how easily the electron cloud can disperse or distort inside a molecule (polarizability).
  • Larger molecules possess larger electron clouds that are more polarizable, leading to greater LDF strength.
  • In large molecules, LDFs can be exceptionally strong.
• Halogen Trend Analysis ($F_2$, $Cl_2$, $Br_2$):
  • Fluorine ($F_2$), chlorine ($Cl_2$), and bromine ($Br_2$) are all nonpolar diatomic molecules.
  • Moving down the halogen group ($F_2 \nrightarrow Cl_2 \nrightarrow Br_2$), molecular size increases.
  • Larger molecules have larger, more polarizable electron clouds, resulting in stronger LDFs and higher boiling points.
• Isomer Comparison (Butane vs. 2-Methylpropane):
  • Butane and 2-methylpropane are nonpolar isomers sharing the same molecular formula ($C_4H_{10}$).
  • Butane possesses a linear shape, providing a larger surface area that allows for more intermolecular interaction and stronger LDFs.
  • 2-Methylpropane possesses a branched, spherical shape with a smaller surface area and lower LDF strength.
  • As a result, butane has a slightly higher boiling point than 2-methylpropane.

Comparative Analysis: Chlorine vs. Hydrogen Chloride

• Physical properties and IMF identification:
  • Chlorine ($Cl_2$) is nonpolar and possesses only LDFs.
  • Hydrogen chloride ($HCl$) is polar and possesses both LDFs and dipole-dipole interactions.
• Heat of vaporization data comparison:
  • Chlorine ($Cl_2$) exhibits a higher heat of vaporization than hydrogen chloride ($HCl$).
• Explanation of relative IMF strengths:
  • The heat of vaporization data indicates that $Cl_2$ has stronger total intermolecular attractions than $HCl$.
  • Even though $HCl$ has two types of IMFs (LDFs and dipole-dipole interactions), $Cl_2$ has a larger, more polarizable electron cloud.
  • The LDFs in $Cl_2$ alone are greater in magnitude than the combined LDFs and dipole-dipole interactions in $HCl$.

Comparative Analysis: Ethane vs. Methanol

• Ethane ($C_2H_6$) and methanol ($CH_3OH$) have comparable molecular sizes and identical molar masses.
• Identification of IMFs:
  • Ethane ($C_2H_6$) is nonpolar and experiences only LDFs.
  • Methanol ($CH_3OH$) is polar and experiences both LDFs and dipole-dipole interactions.
• Physical state at room temperature:
  • Because their sizes and molar masses are comparable, their LDF strengths are similar.
  • Methanol exists as a liquid at room temperature due to its additional dipole-dipole forces, whereas ethane exists as a gas.

Hydrogen Bonding

• Group 16 Hydride Anomalies ($H_2S$, $H_2Se$, $H_2Te$, $H_2O$):
  • Group 16 hydrides all possess bent molecular geometries and are polar, experiencing both LDFs and dipole-dipole forces.
  • From $H_2S$ to $H_2Se$ to $H_2Te$, molar mass increases, electron clouds become more polarizable, LDFs strengthen, and boiling points increase steadily.
  • Water ($H_2O$) is the smallest molecule in the group but displays an abnormally high boiling point that breaks the group trend.
  • Water's high boiling point is caused by hydrogen bonding, which is significantly stronger than standard dipole-dipole forces or LDFs in small molecules.
• Nature and Criteria of Hydrogen Bonds:
  • A hydrogen bond is an unusually strong type of dipole-dipole intermolecular interaction, not an intramolecular chemical bond.
  • Mechanism:
  • An oxygen, nitrogen, or fluorine atom covalently bonded to a hydrogen atom pulls electron density strongly toward itself due to its high electronegativity.
  • Because a hydrogen atom has only one electron, pulling this electron away leaves an unshielded, bare proton.
  • This exposed partial positive charge strongly attracts an unshared lone pair on a highly electronegative atom ($N$, $O$, or $F$) of a neighboring molecule.
  • Required conditions:
  1. A hydrogen atom must be covalently bonded to a highly electronegative atom ($N$, $O$, or $F$) within the molecule.
  2. This hydrogen atom must interact with an $N$, $O$, or $F$ atom on a nearby molecule (or a distant region of the same molecule that can conform/bend around).
• Identification of Hydrogen Bonding in Diagrams (Positions 1–5):
  • Position 1: Represents a solid covalent bond between an $H$ atom and an $O$ atom inside the same molecule. Not a hydrogen bond.
  • Position 2: Interaction between an $H$ atom (covalently bonded to $O$) and an $N$ atom on a different molecule. Valid hydrogen bond.
  • Position 3: Interaction between an $H$ atom and an $O$ atom, but the $H$ atom is not covalently bonded to $N$, $O$, or $F$. Not a hydrogen bond.
  • Position 4: Interaction between an $H$ atom and an $F$ atom on another molecule, but the $H$ atom is not covalently bonded to $N$, $O$, or $F$. Not a hydrogen bond.
  • Position 5: Interaction between an $H$ atom (covalently bonded to $O$) and an $O$ atom on a different molecule. Valid hydrogen bond.

Ion-Dipole Interactions

• Occur when ionic compounds dissolve in polar solvents (e.g., aqueous solutions).
• Mechanism of dissolution:
  • Dipoles of water molecules interact with charged ions in an ionic lattice, causing the cations and anions to separate.
  • These ion-dipole interactions are stronger than hydrogen bonds.
• Orientation of water molecules around ions:
  • Chloride anion ($Cl^-$): Negatively charged; surrounded by water molecules oriented with partially positive hydrogen atoms ($\delta+$) pointing toward the $Cl^-$ ion.
  • Sodium cation ($Na^+$): Positively charged; surrounded by water molecules oriented with partially negative oxygen atoms ($\delta-$) pointing toward the $Na^+$ ion.
  • These orientations follow Coulombic principles where opposite charges attract.

Isomeric Comparison: Ethanol vs. Dimethyl Ether

• Structural Comparison:
  • Ethanol ($CH_3CH_2OH$) and dimethyl ether ($CH_3OCH_3$) are structural isomers sharing the formula $C_2H_6O$.
  • Both molecules are polar and have identical molar masses and similar electron cloud sizes, giving them similar LDF contributions.
• Intermolecular Force Comparison:
  • Ethanol contains an $H$ atom directly bonded to an $O$ atom, allowing ethanol molecules to form hydrogen bonds with one another.
  • Dimethyl ether has its $O$ atom situated in the center of the molecule between two carbons ($C-O-C$), meaning no hydrogen is bonded directly to oxygen; hydrogen bonding is impossible.
• Impact on Boiling Point:
  • Ethanol has a significantly higher boiling point than dimethyl ether because hydrogen bonding produces stronger total IMFs, making ethanol molecules harder to separate into the gas phase.

Physical Properties Influenced by Intermolecular Forces

• Properties directly related to IMF strength (as IMFs INCREASE, property INCREASES):
  • Melting point
  • Boiling point
  • Surface tension (the ability of a liquid surface to resist external forces)
  • Viscosity (resistance of a fluid to flow)
  • Heat of vaporization (energy required to convert a liquid to a gas)
• Properties inversely related to IMF strength (as IMFs INCREASE, property DECREASES):
  • Vapor pressure (pressure exerted by a gas at equilibrium with its liquid in a closed container)
  • Volatility (ease of evaporation)

Rules for Comparing Relative Intermolecular Force Strengths

• Comparing molecules of similar size and polarizability:
  • LDFs are approximately equal between the substances.
  • The substance capable of additional IMFs (e.g., dipole-dipole or hydrogen bonding) will possess stronger overall IMFs.
  • Example: Propane ($C_3H_8$) vs. Ethanol ($C_2H_5OH$). Both have similar molar masses and LDFs. Propane is nonpolar (LDFs only), while ethanol exhibits hydrogen bonding. Ethanol has stronger total IMFs.
• Comparing molecules with significantly different numbers of electrons but identical IMF types:
  • The larger molecule possesses a larger, more polarizable electron cloud, producing stronger LDFs and stronger overall IMFs.
  • Example: Methane ($CH_4$) vs. Octane ($C_8H_{18}$). Both are nonpolar with LDFs only, but octane is much larger, has stronger LDFs, and has a higher boiling point.
• Comparing molecules with different sizes and different types of IMFs:
  • Relative IMF strength cannot always be predicted purely from chemical formulas, as large electron clouds can produce LDFs that outweigh dipole-dipole or hydrogen bonding interactions.
  • Experimental data (such as boiling point) must be used to establish relative IMF strength.

AP Practice Example: Carbon Disulfide ($CS_2$) vs. Carbonyl Sulfide ($COS$)

• Structural and IMF Breakdown:
  • Carbon disulfide ($CS_2$): Linear and symmetrically structured; nonpolar. Possesses only London dispersion forces (LDFs).
  • Carbonyl sulfide ($COS$): Linear but unsymmetrically arranged ($S=C=O$); polar. Possesses both LDFs and dipole-dipole interactions.
• Boiling Point Analysis:
  • Experimental data shows the boiling point of $CS_2$ is higher than that of $COS$.
• Explanation:
  • Although $COS$ exhibits two distinct types of IMFs (LDFs and dipole-dipole), $CS_2$ has a significantly larger and more polarizable electron cloud than COS$.\n - The London dispersion forces in CS_2$are greater in magnitude than the combined LDFs and dipole-dipole interactions in$COS$, resulting in stronger overall intermolecular attractions and a higher boiling point for$CS_2$$.