Intramolecular and Intermolecular Forces Comprehensive Study Notes

Definitions and Core Distinctions of Atomic Forces

• Intramolecular vs. Intermolecular Forces
  • Forces can be classified based on whether they act within a single molecule or between multiple separate molecules.
  • Intramolecular Forces (Interatomic Forces): These are forces that hold atoms together within a molecule. They are characterized as stronger forces and are represented by chemical bonds. An example is the covalent $O-H$ bond within a single $H_2O$ molecule.
  • Intermolecular Forces (IMF): These are forces that exist between different molecules. They are generally weaker than intramolecular forces. An example is the attraction between separate $H_2O$ molecules. These forces are responsible for determining the physical properties of a substance, such as boiling point, melting point, and solubility.

The Five Primary Types of Intermolecular Forces

• 1. Mutually Induced Dipole Forces (London Forces)

  • These forces are present in all molecules, including non-polar covalent molecules.
  • Mechanism: Instantaneous dipoles are created by the random movement of electrons within a molecule. These instantaneous dipoles then induce dipoles in neighboring molecules, resulting in a mutual attraction.
  • Strength Factors: London forces become stronger as the number of electrons (and thus the molecular mass) increases and when the molecule has a larger surface area.
  • Energy Requirements: Breaking London dispersion forces does not require significant energy, which is why non-polar covalent compounds often freeze at very low temperatures.
  • Example: Attractions between $Cl_2$ molecules or within noble gases.

• 2. Dipole-Dipole Forces

  • These occur between polar covalent molecules.
  • Mechanism: An attraction exists between the positive pole ($\delta+$) of one polar molecule and the negative pole ($\delta-$) of another polar molecule. When two dipole molecules come into contact, they are held together by this attraction.
  • Characterization: These are considered strong intermolecular forces of attraction compared to London forces.
  • Example: Attractions between $HCl$ molecules.

• 3. Dipole-Induced Dipole Forces

  • These forces occur between a polar covalent molecule and a non-polar covalent molecule.
  • Mechanism: A polar molecule induces a dipole in a nearby non-polar molecule, leading to an attraction between the permanent dipole and the newly induced dipole.
  • Example: An $O_2$ molecule positioned next to an $HCl$ molecule.

• 4. Ion-Dipole Forces

  • These forces exist between ions and polar covalent molecules.
  • Mechanism: An ion is a charged atom; therefore, it will be attracted to one of the opposite poles of a polar molecule.
  • Example: When sodium chloride ($NaCl$) dissolves in water ($H_2O$), the positive sodium ions ($Na^+$) are attracted to the negative pole (the oxygen atom, $\delta-$) of the water molecules.

• 5. Ion-Induced-Dipole Forces

  • These forces exist between ions and non-polar covalent molecules.
  • Mechanism: An ion induces a dipole in a non-polar molecule, leading to a weak force that holds the compound together.
  • Example: Interaction between a ferrous ion ($Fe^{2+})$ and an oxygen molecule ($O_2)$.

Special Case: Hydrogen Bonding

• Definition: Hydrogen bonding is a special, exceptionally strong type of dipole-dipole interaction.
• Criteria for Formation: It occurs specifically when a hydrogen ($H$) atom is covalently bonded to a highly electronegative atom—specifically nitrogen ($N$), oxygen ($O$), or fluorine ($F$)—and that hydrogen atom is then attracted to a lone pair of electrons on an $N$, $O$, or $F$ atom of another molecule.
• Properties: Because hydrogen bonds are the strongest of the dipole-dipole attractions, they require considerable energy to break. Consequently, compounds containing these bonds have exceptionally high boiling and melting points.

Relationship Between Intermolecular Forces and Molecular Mass

• General Principle: As the molecular mass or total number of electrons in a molecule increases, the strength of the London forces increases. This leads to overall stronger intermolecular forces for similar types of molecules.
• Reasoning: Larger molecules possess more electrons and a greater surface area. This configuration allows for more frequent and stronger instantaneous dipoles to form, creating more robust attractions.
• Case Study: Alkanes
  • The increasing strength of intermolecular forces is clearly illustrated by the boiling points of the alkane series as the chain length (and thus mass) increases:
    • Methane ($CH_4$): Molecular Mass = $16\,g/mol$; Boiling Point = $-161\,^{\circ}C$
    • Ethane ($C_2H_6$): Molecular Mass = $30\,g/mol$; Boiling Point = $-89\,^{\circ}C$
    • Propane ($C_3H_8$): Molecular Mass = $44\,g/mol$; Boiling Point = $-42\,^{\circ}C$
    • Butane ($C_4H_{10}$): Molecular Mass = $58\,g/mol$; Boiling Point = $-1\,^{\circ}C$
    • Pentane ($C_5H_{12}$): Molecular Mass = $72\,g/mol$; Boiling Point = $36\,^{\circ}C$
    • Hexane ($C_6H_{14}$): Molecular Mass = $86\,g/mol$; Boiling Point = $69\,^{\circ}C$

Effects of Intermolecular Forces on Physical Properties

• 5.1 Boiling Point

  • Definition: The temperature at which the vapour pressure of a substance is equal to the atmospheric pressure.
  • Relationship: Stronger IMF $\rightarrow$ more energy needed to overcome attractions between molecules $\rightarrow$ higher boiling point.
  • Example: Water ($H_2O$, hydrogen bonding) has a higher boiling point than hydrogen sulfide ($H_2S$, dipole-dipole), which has a higher boiling point than methane ($CH_4$, London forces).

• 5.2 Melting Point

  • Definition: The temperature at which the solid and liquid phases of a substance are at equilibrium.
  • Relationship: Stronger IMF $\rightarrow$ more energy needed to break attractions in the solid phase $\rightarrow$ higher melting point.
  • Example: Iodine ($I_2$, London forces in large molecules) has a higher melting point than Bromine ($Br_2$), which has a higher melting point than Chlorine ($Cl_2$).

• 5.3 Vapour Pressure

  • Definition: The pressure exerted on a container when the vapour and liquid phases of a substance are at equilibrium.
  • Relationship: Stronger IMF $\rightarrow$ molecules are held more tightly in the liquid phase $\rightarrow$ fewer molecules escape into the gas phase $\rightarrow$ lower vapour pressure.
  • Example: Water has a lower vapour pressure than ethanol because water molecules are held more strongly by hydrogen bonding.

• 5.4 Solubility

  • Universal Rule: "Like dissolves like."
  • Mechanism: Solubility depends on the IMF between the solute and the solvent. High solubility occurs when the IMF between the solute and solvent are strong.
  • Polarity Rules:
    • Polar solutes are soluble in polar solvents (Strong IMF).
    • Non-polar solutes are soluble in non-polar solvents (Strong IMF).
    • Polar solutes are insoluble in non-polar solvents (Weak IMF).

Relative Strength Summary