Chemistry - lecture 3 - Molecular Shape, Molecular Forces and Non-Covalent Bonding

Learning Outcomes

  • Describe and explain the structure of simple molecules.

  • Apply MO theory to elucidate properties of molecular structures in biomolecules (e.g., peptide bonds, aromatic rings).

  • Use specific terminology when describing molecules and chemical bonding.

  • Catalogue different interactions in biomolecules.

Hybrid Orbitals

  • Definition: Atomic orbitals combine to form hybrid orbitals.

  • Hybrid orbitals offer better orientation for bonding than standard atomic orbitals.

  • Describes the orientation of bonding interactions.

  • Core Concept - Hybridisation:

    • Mixing of closely energy-aligned atomic orbitals (e.g., 2s and 2p) lowers overall energy.

    • Hybrid orbitals inherit characteristics of their constituent AOs based on mixing ratio.

    • Hybrid orbitals are directional, impacting molecular geometry.

Types of Hybridisation

sp3 Hybridisation

  • Occurs by mixing three p orbitals with one s orbital.

  • Forms four equivalent sp3 hybrid orbitals, arranged tetrahedrally (e.g., methane, alkyl carbons).

  • All hybrid orbitals possess the same energy level.

sp2 Hybridisation

  • Involves mixing two p orbitals with one s orbital.

  • Creates three sp2 hybrid orbitals arranged in a trigonal planar configuration, with the remaining p orbital perpendicular.

  • Key examples include C=C double bonds and C=O (carbonyl) groups.

sp Hybridisation

  • Results from mixing one p orbital with one s orbital.

  • Forms two sp hybrid orbitals arranged linearly, with two unhybridized p orbitals perpendicular to each other.

  • Important in triple bonds (e.g., carbon-carbon, cyanide (CN)).

VSEPR Theory

  • VSEPR = valence shell electron pair repulsion

  • Core Concept: Electron Pair Repulsion:

    • Analyzes arrangements of atoms around a central atom by minimizing repulsions among bonding pairs (σ bonds) and lone pairs.

    • Identification of hybridisation is based on the atom's valency.

  • Determining Geometry:

    1. Count valence electrons from the periodic table.

    2. Identify the number of bonds (single, double, triple).

    3. Consider formal charges.

    4. Example of methane (CH4):

      • 4 valence electrons, 4 single bonds (AX4 structure).

      • Example of ammonia (NH3):

      • 5 valence electrons, three bonds leave one lone pair (AX3E structure).

Relevant AXmEn Structures

  • Coordination number and Geometry:

    Coordination (m+n)

    No Lone Pair

    One Lone Pair

    Two Lone Pairs

    2

    Linear

    --

    --

    3

    Trigonal planar

    Bent

    --

    4

    Tetrahedral

    Trigonal pyramidal

    Bent (H2O)

Electron Pair Repulsion and Bond Angles

  • Influence of Lone Pairs:

    • Lone pairs exert stronger repulsion, altering bond angles (

    • e.g., tetrahedral 109.5°, NH3 at 106.8°, and H2O at 104.5°).

Carbonyl Group (sp2 Geometry)

  • Carbon and oxygen (O) are sp2 hybridised.

  • Bonding:

    • C-O double bond comprises a σ-bond and a π-bond, with C forming two additional σ-bonds.

    • Oxygen retains two lone pairs on its sp2 hybrid orbital.

    • Orbital energy mismatches lead to polarization in the bond.

Bond Rotations in Saturated Chains

  • Bond rotations: single bonds allow for free rotation where the bond connects.

  • Butane:

    • Demonstrates potential for "free" rotations about sigma bonds; can experience staggered or eclipsed conformations.

    • Staggered arrangements are lower in energy due to minimized steric clashes.

  • Restrictions to Rotations:

    • Steric hindrance, cyclic compounds, and adjacent π-systems limit rotational freedom.

Noteworthy Cases

Ethane

  • Preferred staggered configuration is attributed not just to sterics but also to hyperconjugation.

Bond Rotations – Multiple Bonds

  • Double/triple bonds restrict rotation due to presence of π-bonds, maintaining a nodal plane impeding rotation.

  • Rotation of two atoms and their orbitals will break the bond.

Peptide Bonds and Delocalisation

  • N typically displays sp3 hybridization but can also engage in sp2 hybridization, allowing resonance.

  • In peptide bonds, N, C, and O become sp2 centers, creating a planar structure.

Geometry of Peptide Bonds

  • All atoms connected to sp2 centers lie in the same plane, making the protein backbone a series of interconnected planes.

Ramachandran Plot

  • Dihedral angles illustrate limitations in configurations due to sterics and interactions.

  • Core Concept - φ and ψ angles account for the orientation of peptide planes in protein backbones, restricted by atom interactions.

    • Relative to peptide planes.

    • Limited by sterics, repulsion, and preferred interactions

Describing Molecular Geometry

  • Factors include:

    • Bond Length: Distance between bonded atoms.

    • Bond Angle: Angle between bonds at a central atom.

    • Dihedral Angles: Involves angles between planes formed by connected atoms.

  • Use of xyz coordinates offers insight into connectivity and molecular structure.

Bond Lengths

  • General trends:

    • Single > Double > Triple lengths.

  • Typically of order of one Angström (A) - 0.1 nm

  • Covalent radius influences the bond lengths so effective charge is significant.

Bond Angles

  • Defined by repulsions within bonding pairs and lone pairs (VSEPR).

    • General rules:

      • Bonds are symmetrically arranged around the central atom.

      • Lone pairs compress bond angles more than bonding pairs.

Sugar and Ring Conformations

  • Rings experience restricted motion; e.g., cyclohexane maintains distinct staggered conformations to minimize sterics.

  • Chair Conformation: Lowest energy state.

  • Boat Conformation: Higher energy state due to steric interactions and torsional strain, making it less favorable compared to the chair.

  • Methyl groups prefer equatorial positions to reduce steric repulsion.

Cyclic Sugars

  • Cyclic forms of sugars have distinct substituents and pucker conformations.

  • Examples: α-D-glucopyranose, β-D-ribofuranose, leading to characteristic structures due to spatial orientations.

Non-Covalent Interactions

  • Types of interactions include:

    • Hydrogen bonds, dipole interactions, dispersion forces, π-π interactions, and hydrophobic effects.

Van der Waals Forces

  • Collection of distance-dependent interactions between atoms and molecules that are independent of chemical bonding.

  • Induction (Debye) and dispersion (London) forces

  • Characteristics:

    • Weaker than covalent/ionic bonds.

    • Depend on distance and orientation of molecules.

    • Short ranged and usually non-directional.

  • Includes attractive and repulsive forces, some based on electrostatics.

Dipole Interactions

  • Dipole: two equal and opposite charges separated by some distance, leading to a directional field.

  • Molecular dipoles arise from geometric arrangements of electronegative atoms (e.g., water displays a net dipole).

  • Symmetric molecules (e.g., methane) may have canceled dipoles.

  • Multipoles: spatial distributions of charges that generate electric fields.

  • Monopole: a single point charge.

Permanent Dipole Interactions

  • Dipoles that exist due to the molecular geometry.

  • Molecules with permanent dipoles can interact, leading to dipole-dipole interactions (strength varies with distance).

Induced Dipoles

  • Neutral species can develop dipoles due to nearby charged species (permanent dipoles or ions), leading to Debye forces.

London Dispersion Forces

  • Caused by transient electron density fluctuations, leading to instantaneous dipoles.

  • Generally weak but significant in interactions among neutral molecules.

π-Interactions

  • π systems can interact attractively with other π systems through stacking (e.g., in biomolecule structures).

  • Aromatic compounds may exhibit quadrupole interactions with charged species.

Hydrogen Bonding

  • Formed between positively charged H and electronegative atoms, N or O.

  • Key to determining structural features in biomolecules (e.g., protein secondary structures).

Special Hydrogen Bond Interactions

  • C-H···O interactions:

    • Weak hydrogen bonds may form involving C-H groups if coupled with strongly electronegative atoms.

  • Salt Bridges:

    • Non-covalent, pH-dependent interactions in proteins between charged side chains.

Hydrophobicity

  • Hydrophobic regions tend to collapse in aqueous environments, preventing unfavorable interactions.

  • Hydrophobic behavior arises from minimized interactions with water; essential in biomolecular behavior.