Chemical Bonding: Intermolecular Forces and Reactivity of Covalent Bonds

Overview of Chemical Bonding and Intermolecular Forces

• Classification of Chemical Bonds:

  • Interatomic Bonds: Intramolecular forces that hold atoms together within a chemical entity. These include:
  • Ionic Bond
  • Covalent Bond
  • Co-ordinate Covalent or Dative Bond
  • Metallic Bond
  • Intermolecular Forces: Weak attractive forces existing between distinct molecules or molecular entities other than those formed due to direct chemical bond formation. These are broadly termed van der Waals' forces and include:
  • Instantaneous dipole–induced dipole ($\text{id-id}$) forces (also known as London dispersion forces or temporary dipole–induced dipole forces).
  • Permanent dipole–permanent dipole ($\text{pd-pd}$) forces.
  • Hydrogen bonding (a special, particularly strong category of permanent dipole–permanent dipole interaction).

• Definition of van der Waals' Forces:

  • The expression van der Waals' forces covers all types of intermolecular forces between molecular entities, distinct from primary chemical bond formation.
  • It describes the weak forces of attraction between molecules involving either instantaneous ($\text{id-id}$) or permanent dipole–permanent dipole ($\text{pd-pd}$) forces, including hydrogen bonding.

Instantaneous Dipole–Induced Dipole ($\text{id-id}$) Forces

• Mechanism and Definition:

  • Instantaneous dipole–induced dipole ($\text{id-id}$) forces (London dispersion forces) represent the weakest type of intermolecular force.
  • Dipole Origin: Electrons in an atom or molecule are in continuous motion. At any given instant, an asymmetric distribution of electrons can occur across the electron cloud, creating a temporary or instantaneous dipole with partial positive ($\delta^+$) and partial negative ($\delta^-$) ends.
  • Induction Process: The instantaneous dipole in Atom 1 distorts the electron cloud of a neighboring Atom 2, inducing a complementary dipole in Atom 2.
  • Attraction: The electrostatic attraction between the instantaneous dipole and the induced dipole creates the $\text{id-id}$ force.

• Factors Increasing $\text{id-id}$ Force Strength:

  • Increasing Number of Electrons: As the total number of electrons in a molecule increases, the electron cloud becomes larger and more easily polarizable, leading to stronger instantaneous and induced dipoles.
  • Increasing Number of Contact Points: Molecules with greater surface area can come into closer proximity over a larger area, increasing total intermolecular contact points.
  • Effect of Molecular Branching: Branching in organic molecules reduces the overall surface area and contact sites between adjacent molecules, thereby decreasing $\text{id-id}$ forces.

• Empirical Evidence and Data:

  • Noble Gas Group Trends: As the atomic number (and thus the total number of electrons) increases down the noble gases group, both the enthalpy change of vaporisation and the boiling point increase systematically. Stronger $\text{id-id}$ forces require greater thermal energy to overcome the intermolecular interactions to transition from liquid to gas phase.
  • Structural Isomer Comparison (Pentane vs. 2,2-dimethylpropane):
  • Both pentane ($\text{C}_5\text{H}_{12}$) and 2,2-dimethylpropane ($\text{C}_5\text{H}_{12}$) possess identical chemical formulas and an equal total number of electrons.
  • Pentane: Boiling point = $36\,^{\circ}\text{C}$. Unbranched straight-chain structure allows molecules to line up closely alongside one another, maximizing the available surface area and number of contact points, leading to higher van der Waals' forces.
  • 2,2-Dimethylpropane: Boiling point = $10\,^{\circ}\text{C}$. Compact, spherical geometry minimizes available surface area for neighboring contact, resulting in significantly weaker van der Waals' forces and a lower boiling point.

Permanent Dipole–Permanent Dipole ($\text{pd-pd}$) Forces

• Nature of $\text{pd-pd}$ Forces:

  • $\text{pd-pd}$ forces are intermolecular electrostatic forces of attraction operating between polar molecules possessing permanent dipoles.
  • These forces are intrinsically stronger than instantaneous dipole–induced dipole ($\text{id-id}$) forces for molecules of comparable size and electron count.

• Example: Propanone ($\text{CH}_3\text{COCH}_3$):

  • Propanone contains a polar carbonyl group ($\text{C}=\text{O}$).
  • Oxygen is more electronegative than carbon, drawing electron density away from the carbon atom to create a permanent bond dipole ($\text{C}^{\delta+}=\text{O}^{\delta-}$).
  • The permanent $\delta+$ charge on the carbon atom of one propanone molecule electrostatically attracts the permanent $\delta-$ charge on the oxygen atom of a neighboring propanone molecule.

• Distinction for Hydrogen Halides ($\text{HCl}$ and $\text{HBr}$):

  • The intermolecular forces present between $\text{HCl}$ or $\text{HBr}$ molecules are classified as permanent dipole–permanent dipole forces, not hydrogen bonds.
  • Although hydrogen is covalently bonded to halogen atoms, chlorine ($\text{Cl}$) and bromine ($\text{Br}$) atoms have larger atomic radii and lower electron densities compared to nitrogen, oxygen, and fluorine, preventing the formation of true hydrogen bonds.

Hydrogen Bonding

• Definition and Requirements:

  • A Hydrogen Bond is a specialized, unusually strong form of permanent dipole–permanent dipole force occurring specifically when a hydrogen atom is covalently bonded directly to a highly electronegative atom with small atomic radius: Fluorine ($\text{F}$), Oxygen ($\text{O}$), or Nitrogen ($\text{N}$).
  • Because $\text{F}$, $\text{O}$, and $\text{N}$ are extremely electronegative, they strongly pull bonding electron pairs toward themselves, generating a strong bond dipole ($\text{X}^{\delta-}-\text{H}^{\delta+}$) and leaving the small hydrogen nucleus almost completely exposed (deshielded).

• Relative Strength:

  • Hydrogen bonds represent the strongest type of intermolecular force.
  • While stronger than all other van der Waals' forces ($\text{id-id}$ and standard $\text{pd-pd}$), hydrogen bonds remain significantly weaker than primary intramolecular chemical bonds (such as covalent bonds).

• Structural Examples:

  • Water ($\text{H}_2\text{O}$): Hydrogen bonds form between the partial negative oxygen atom ($\text{O}^{\delta-}$) of one water molecule and the partial positive hydrogen atom ($\text{H}^{\delta+}$) of an adjacent water molecule.
  • Ammonia ($\text{NH}_3$): Hydrogen bonds form between the partial negative nitrogen atom ($\text{N}^{\delta-}$) lone pair of one ammonia molecule and the partial positive hydrogen atom ($\text{H}^{\delta+}$) of another.

• Determinants of Hydrogen Bonding Capacity:

  • The average number of hydrogen bonds formed per molecule depends on:
  • The total number of hydrogen atoms directly attached to $\text{F}$, $\text{O}$, or $\text{N}$ in the molecule.
  • The number of unshared electron lone pairs present on the electronegative $\text{F}$, $\text{O}$, or $\text{N}$ atom.

Physical Effects and Applications of Hydrogen Bonding

• Impact on Melting and Boiling Points:

  • Molecules capable of hydrogen bonding display anomalously high melting and boiling points compared to molecules with similar numbers of electrons that only experience $\text{id-id}$ or $\text{pd-pd}$ forces.
  • Group 15, 16, and 17 Hydrides Trend:
  • Comparing hydrides across increasing electron numbers ($\text{H}_2\text{O}$, $\text{H}_2\text{S}$, $\text{H}_2\text{Se}$, $\text{H}_2\text{Te}$ and $\text{HF}$, $\text{HCl}$, $\text{HBr}$, $\text{HI}$):
  • While boiling points generally increase from Period 3 to Period 5 hydrides due to increasing electron counts and $\text{id-id}$ forces, $\text{H}_2\text{O}$ and $\text{HF}$ exhibit exceptionally high boiling points despite having the smallest electron numbers in their respective series.
  • This anomaly occurs because significant thermal energy is required to break the strong hydrogen bonds during state transitions.

• Solubility in Water and Polar Solvents:

  • Covalent substances can dissolve readily in water if they possess the capability to form hydrogen bonds with water molecules.
  • Examples of Hydrogen-Bonding Soluble Systems:
  • Hydrogen Fluoride ($\text{HF}$): Highly soluble in water due to $\text{F-H}\cdots\text{O}$ hydrogen bonding.
  • Ammonia ($\text{NH}_3$): Dissolves extensively in water by forming $\text{N}\cdots\text{H-O}$ hydrogen bonds.
  • Ethanol ($\text{CH}_3\text{CH}_2\text{OH}$): Soluble in water via its polar hydroxyl ($\text{-OH}$) group forming hydrogen bonds with water molecules.
  • Hydrazine ($\text{NH}_2\text{NH}_2$): Soluble in ethanol ($\text{CH}_3\text{CH}_2\text{OH}$) because both molecules contain donor and acceptor sites ($\text{-NH}_2$ and $\text{-OH}$) for extensive hydrogen bonding.

• Peculiar Properties of Water:

  • High Melting and Boiling Points: Requires substantial thermal energy to disrupt the hydrogen bonding network relative to water's low molecular mass.
  • High Surface Tension: Cohesive intermolecular hydrogen bonding creates a strong surface film in liquid water.
  • Density Anomaly (Ice vs. Liquid Water):
  • In solid ice, hydrogen bonding forces water molecules to align in a fixed, open 3D crystalline lattice structure.
  • This open arrangement holds water molecules further apart on average than in liquid water, causing water to expand upon freezing.
  • Consequently, ice is less dense than liquid water, allowing ice to float on water.

Comparative Analysis of Intermolecular Interactions

• Evaluating Intermolecular Strengths Across Liquid Pairs:

  • Comparing mixture scenarios for total van der Waals' force strength:
  • Systems consisting exclusively of non-polar alkanes (e.g., octane and hexane) interact purely through $\text{id-id}$ forces.
  • Systems containing polar carbonyl groups (e.g., propanone $\text{H}_3\text{C-CO-CH}_3$ and butanone $\text{H}_3\text{C-CO-CH}_2\text{-CH}_3$) interact via $\text{pd-pd}$ forces.
  • Systems containing hydroxyl ($\text{-OH}$) and amine ($\text{-NH}_2$) groups (e.g., methanol $\text{H}_3\text{C-O-H}$ and propylamine $\text{H}_3\text{C-CH}_2\text{-CH}_2\text{-NH}_2$) form intermolecular hydrogen bonds, producing the strongest combined van der Waals' interactions.

• Comparative Boiling Point Analysis of Organic Compounds:

  • Ranking liquid organic compounds in decreasing order of boiling point:
  1. Butanol ($\text{H}_3\text{C-CH}_2\text{-CH}_2\text{-CH}_2\text{-O-H}$): Highest boiling point. Possesses an $\text{-OH}$ group capable of forming strong intermolecular hydrogen bonds requiring substantial energy to break.
  2. Pentanone ($\text{H}_3\text{C-CH}_2\text{-CH}_2\text{-CO-CH}_3$): Second highest. Contains a polar carbonyl ($\text{C}=\text{O}$) group capable of permanent dipole–permanent dipole ($\text{pd-pd}$) forces, which are stronger than pure dispersion forces but weaker than hydrogen bonding.
  3. Octane ($\text{H}_3\text{C-CH}_2\text{-CH}_2\text{-CH}_2\text{-CH}_2\text{-CH}_2\text{-CH}_2\text{-CH}_3$): Third. A non-polar, long straight-chain alkane. Lacks permanent dipoles or hydrogen bonding, but possesses a large electron count ($66$ electrons) and extensive surface area for high total $\text{id-id}$ forces.
  4. Methyl Hexane ($\text{H}_3\text{C-CH(CH}_3\text{)-CH}_2\text{-CH}_2\text{-CH}_2\text{-CH}_3$): Lowest boiling point. An alkane isomeric to heptane with molecular branching, which reduces molecular contact points and surface area, lowering total $\text{id-id}$ forces relative to straight-chain hydrocarbons of similar mass.

Factors Affecting the Reactivity of Covalent Bonds

• 1. Bond Energy:

  • Definition: The amount of energy required to break one mole of a specific covalent bond in its gaseous state.
  • Units: Kilojoules per mole ($\text{kJ\,mol}^{-1}$).
  • Effect on Reactivity:
  • A lower bond energy value indicates a weaker covalent bond that requires less energy to cleave.
  • Chemical compounds possessing weaker bonds exhibit higher chemical reactivity.
  • Example: The bond dissociation energy of the chlorine-chlorine bond ($\text{Cl-Cl}$) is lower than that of the hydrogen-hydrogen bond ($\text{H-H}$); consequently, chlorine gas reacts much more readily than hydrogen gas.

• 2. Bond Length:

  • Definition: The internuclear distance between the nuclei of two covalently bonded atoms.
  • Effect on Reactivity:
  • Longer bonds feature greater separation between bonding atomic nuclei, resulting in decreased electrostatic attraction between the nuclei and shared electron pair.
  • Consequently, longer bonds are generally weaker, break more easily during collisions, and increase chemical reactivity.
  • Example: Down Group 17 haloalkanes, the carbon-halogen bond length increases ($\text{C-F} < \text{C-Cl} < \text{C-Br} < \text{C-I}$). The $\text{C-I}$ bond is longer and weaker than the $\text{C-F}$ bond, rendering iodoalkanes significantly more reactive than fluoroalkanes.

• 3. Bond Polarity:

  • Definition: The unequal sharing of bonding electrons resulting from differences in electronegativity between two bonded atoms, producing partial positive ($\delta^+$) and partial negative ($\delta^-$) charges across the bond.
  • Effect on Reactivity:
  • Polar bonds establish localized partial charges that attract incoming ions or polar reagents.
  • Electron-rich species (nucleophiles) are attracted to $\delta^+$ centers, while electron-deficient species (electrophiles) are attracted to $\delta^-$ centers, increasing overall reaction rates.
  • Example: In bromoethane ($\text{CH}_3\text{CH}_2\text{Br}$), the bromine atom is more electronegative than carbon, polarizing the bond ($\text{C}^{\delta+}-\text{Br}^{\delta-}$). The electron-deficient $\delta+$ carbon atom readily attracts electron-rich nucleophilic reagents, enhancing chemical reactivity.