(APC 3.1) Intermolecular Forces Study Guide

Overview of Intermolecular Forces

Intermolecular forces (IMF) are the attractive forces that hold atoms, ions, or molecules together. These forces differ fundamentally in the arrangements and behaviors of their electrons.

Intermolecular forces explain macroscopic phase differences among chemical species with similar atomic compositions. For instance, carbon dioxide ($\text{CO}_2$), water ($\text{H}_2\text{O}$), and silicon dioxide ($\text{SiO}_2$) are each composed of exactly three atoms, yet at room temperature, $\text{CO}_2$ is a gas, $\text{H}_2\text{O}$ is a liquid, and $\text{SiO}_2$ is a solid due to differences in their interparticle interactions.

To determine the type of IMF a substance exhibits, examine the types of atoms present in the sample, classifying them as metals, nonmetals, or ions.

Classification and Relative Strengths of Forces

Van der Waals Forces (Weaker to Stronger Spectrum)

• London Dispersion Forces (LDF): Present in all substances. It is the only force present in nonpolar molecules or single nonmetal atoms (e.g., $\text{Ne}$, $\text{H}_2$, $\text{CO}_2$).

• Dipole-Dipole Interactions: Occurs in polar, asymmetrical molecules consisting exclusively of nonmetals (e.g., $\text{HCl}$, $\text{CH}_2\text{O}$, $\text{CH}_3\text{Cl}$).

• Hydrogen Bonding: A particularly strong type of dipole-dipole interaction occurring in polar molecules where hydrogen is directly bonded to nitrogen, oxygen, or fluorine ($\text{N}$, $\text{O}$, or $\text{F}$) (e.g., $\text{HF}$, $\text{H}_2\text{O}$, $\text{CH}_3\text{OH}$).

Strong Interparticle and Bonding Interactions (Difficult to Compare Directly)

• Ionic Bonding: Occurs between metals and nonmetals (e.g., $\text{NaCl}$, $\text{KBr}$, $\text{MgO}$).

• Metallic Bonding: Occurs exclusively among metal atoms (e.g., $\text{Au}$, $\text{Al}$, $\text{ZnCu}$).

• Covalent Network Bonding: Represents the strongest class of substance interactions. Specific examples include diamond ($\text{C(diamond)}$), graphite ($\text{C(graphite)}$), buckminsterfullerene ($\text{C}_{60}$), silicon dioxide ($\text{SiO}_2$), and silicon carbide ($\text{SiC}$).

London Dispersion Forces (LDF)

London dispersion forces are present in all chemical substances and must always be considered. They are the sole intermolecular force present in nonpolar covalent molecules (where nonpolar bond polarities cancel due to symmetry) and in noble gas atoms.

Origin and Mechanism

1. Electron Motion: LDF arises from the constant motion of electrons within an atom or molecule.

2. Temporary Dipole: At any given instant, electrons can become unevenly distributed. This creates a temporary dipole consisting of an excess of negative charge (partially negative, $\delta^-$) in one region and an electron-deficient region that is more positive (partially positive, $\delta^+$).

3. Induced Dipole: The temporary dipole in one atom or molecule alters the electron distribution of neighboring species. The negative end repels neighboring electrons, while the positive end attracts them.

4. Electrostatic Attraction: As neighboring electrons shift, an induced dipole forms that is electrostatically attracted to the temporary dipole of the first atom or molecule.

5. Transience: Because electrons continuously move, species return to an evenly distributed electron state. This short-lived occurrence makes LDF a relatively weak force.

Polarizability and Molecular Factors

Polarizability is the ability of an atom or molecule to form a temporary or induced dipole.

• Electron Count: London dispersion forces grow stronger as an atom or molecule possesses more electrons. Increased electron count enhances polarizability.

  • Halogen Trend Example: Halogens down Group 17 change physical states due to increasing LDF strength ($\text{F}_2 \rightarrow \text{Cl}_2 \rightarrow \text{Br}_2 \rightarrow \text{I}_2$):

    • $\text{F}_2$ has fewer electrons, resulting in lower polarizability and weaker attraction, allowing it to exist as a gas.

    • $\text{I}_2$ has significantly more electrons, higher polarizability, and stronger attractions that hold the molecules together into a solid.

• Molecular Shape and Surface Area:

  • Long, thin molecules possess greater surface area, making them more polarizable and resulting in stronger LDF.

  • Short, round molecules possess less surface area, making them less polarizable and resulting in weaker LDF.

• $\pi$ Bonding: The presence of $\pi$ bonding (double or triple bonds) enhances polarizability and strengthens LDF.

Dipole-Dipole Interactions

Dipole-dipole interactions occur exclusively between polar molecules possessing a permanent dipole (permanent positive and negative charge centers).

Key Characteristics

• Permanent vs. Temporary: Unlike LDF, dipole-dipole attractions involve permanent dipoles created by asymmetrical electron distributions across nonmetal bonds.

• Electrostatic Alignment: The partially positive end ($\delta^+$) of one polar molecule is electrostatically attracted to the partially negative end ($\delta^-$) of a neighboring molecule throughout the system (e.g., $\delta^+\text{H-Cl}\delta^- \cdots \delta^+\text{H-Cl}\delta^-$).

• Influence of Polarity: Interactions are stronger when the substance is more polar, corresponding to a more unequal electron distribution. Electronegativity differences between bonded elements dictate the magnitude of the dipole.

Hydrogen Bonding

Hydrogen bonding is a specialized, strong type of dipole-dipole attraction. Despite the word "bonding" in its name, it is an intermolecular attraction rather than a covalent chemical bond.

Structural Requirements and Strength

• Direct Attachment Required: Hydrogen bonding occurs only in molecules containing hydrogen directly bonded to nitrogen, oxygen, or fluorine ($\text{N}$, $\text{O}$, or $\text{F}$).

• Electronegativity Origin: The large electronegativity difference between $\text{N}$, $\text{O}$, or $\text{F}$ and hydrogen generates a large molecular dipole, leading to strong attractions between neighboring molecules.

• Relative Strength Trend: Among hydrogen bonds, strength follows the electronegativity of the host atom: $\text{H-F} > \text{H-O} > \text{H-N}$. Fluorine is the most electronegative element, while nitrogen is the least electronegative of the three.

• Multiple Sites: Hydrogen bonding strength increases when a molecule presents multiple sites capable of forming hydrogen bonds.

Important Counter-Examples

• A $\text{C-H}$ bond (such as in hydrocarbons like ethane) does not engage in hydrogen bonding because carbon lacks sufficient electronegativity, and $\text{C-H}$ bonds are nonpolar.

• An intramolecular covalent bond within a single molecule is not a hydrogen bond; hydrogen bonding strictly describes the interaction between separate molecules.

Induced Dipole Interactions

Dipole-Induced Dipole Interactions

• Definition: Occurs when a polar molecule with a permanent dipole interacts with a nonpolar molecule that ordinarily exhibits only LDF.

• Mechanism: The permanent dipole charge of the polar molecule distorts the electron cloud of the nonpolar molecule, inducing a temporary dipole.

• Examples:

  • Polar hydrogen chloride ($\text{HCl}$) inducing a dipole in nonpolar Argon ($\text{Ar}$).

  • Polar water ($\text{H}_2\text{O}$) inducing a dipole in nonpolar oxygen gas ($\text{O}_2$).

• Factors Affecting Strength: Forces strengthen with a larger permanent dipole on the polar molecule and higher polarizability of the nonpolar molecule.

Ion-Induced Dipole Interactions

• Definition: Occurs when a fully charged ion (from an ionic compound) interacts with a nonpolar molecule.

• Mechanism: A nonpolar atom initially presents a spherical electron cloud surrounding its nucleus with no dipole. Upon the approach of a charged ion, the electrons respond to the ion's charge, causing the atom to develop an induced dipole ($\delta^+$, $\delta^-$).

• Factors Affecting Strength: Forces strengthen with larger ionic charge magnitude, smaller ion size, and higher polarizability of the nonpolar species.

Ion-Dipole Interactions

• Definition: Occurs when a fully charged ion interacts with a polar molecule possessing a permanent dipole.

• Mechanism Example (Dissolution of $\text{NaCl}$ in Water): When ionic $\text{NaCl}$ dissolves in water, it dissociates into $\text{Na}^+$ and $\text{Cl}^-$ ions. The polar water molecules align around the ions based on charge:

  • For $\text{Na}^+$ cations, the partially negative oxygen ends ($\delta^-$) of water face inward toward the ion, while the partially positive hydrogen ends ($\delta^+$) face away.

• Relative Strength: Ion-dipole interactions tend to be stronger than dipole-dipole forces because they involve full ionic charges interacting with permanent dipoles.

• Factors Affecting Strength: Interactions strengthen with larger ionic charges, smaller ion sizes, and larger permanent dipoles on the polar molecules.

Summary of Interparticle Interactions