Lecture 3: Molecular Shape, Molecular Forces, and Non-Covalent Interactions

Lecture 3: Molecular Shape, Molecular Forces, and Non-Covalent Interactions

Overview

  • Lecturer: Dr. Konstantin Roeder

  • Institution: Randall Centre for Cell & Molecular Biophysics

  • Date: January 2026

Recap of Previous Lecture

  • Electrons are located in orbitals characterized by four quantum numbers:

    • Principal quantum number ($n$)

    • Azimuthal quantum number ($l$)

    • Magnetic quantum number ($m_l$)

    • Spin quantum number ($m_s$)

  • Electrons occupy available orbitals from low to high energy in accordance with the n+l rule.

  • Fully filled shells act as core electrons and do not participate in chemical bonding.

  • Valence electrons are found in partially filled shells and subshells, critical for bonding.

  • Molecular orbitals (MOs) arise from the combination of atomic orbitals (AOs) in various phases, resulting in:

    • Bonding MOs

    • Non-bonding MOs

    • Antibonding MOs

  • Molecular orbital theory aids in understanding bond polarization and conjugated systems such as aromatic rings.

Synopsis of Current Lecture

  1. Hybridization of atomic orbitals

  2. Valence Shell Electron Pair Repulsion (VSEPR) theory

  3. Chemical bonding and molecular geometries

  4. Discussion on molecular geometry terminology

  5. Non-covalent interactions III

  6. Summary of interactions in biomolecules

Learning Outcomes

  1. Describe and explain the structure of simple molecules.

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

  3. Integrate specific terminology regarding molecular descriptions and chemical bonding.

  4. Catalogue the various interactions among atoms in biomolecules.

Hybrid Orbitals

  • Atomic Orbital Mixing: Atomic orbitals can combine to form hybrid orbitals, improving predictions of bonding orientations within molecules.

  • Hybridization: The process whereby atomic orbitals of close energies in the same atom mix to form hybrid orbitals.

    • Derived from Hydrogen-like orbitals; specifically mixing 2s and 2p orbitals can lead to lower energy configurations.

Character of Hybrid Orbitals

  • Hybrid orbitals exhibit directional properties reflecting characteristics of the atomic orbitals from which they were derived.

Hybridization Types

sp3 Hybridization

  • Formation: Mixing of three p orbitals with one s orbital, resulting in four equivalent sp3 hybrid orbitals.

  • Geometric Arrangement: Tetrahedral.

  • Example: Methane (CH4) and similar alkyl carbons.

sp2 Hybridization

  • Formation: Mixing of two p orbitals with one s orbital, yielding three sp2 hybrid orbitals.

  • Geometric Arrangement: Trigonal planar with the remaining unhybridized p orbital perpendicular.

  • Examples: C=C double bonds and carbonyl groups (C=O).

sp Hybridization

  • Formation: Mixing of one p orbital with one s orbital, producing two sp hybrid orbitals.

  • Geometric Arrangement: Linear orientation with two unhybridized p orbitals, placed perpendicularly.

  • Examples: Carbon-carbon triple bonds (e.g., in cyanide, CN).

VSEPR Theory (Valence Shell Electron Pair Repulsion)

  • Core Concept: Electron pairs in hybridized orbitals repel each other, dictating molecular geometry by optimizing distances between pairs (σ bonds and lone pairs).

  • Procedure for Geometry Determination:

    1. Count valence electrons from the periodic table.

    2. Identify the number of single, double, and triple bonds.

    3. Account for charges.

  • Example Structures:

    • Methane (CH4): Central atom C (4 valence electrons), four single bonds, no lone pairs, shape is AX4.

    • Ammonia (NH3): Central atom N (5 valence electrons), three bonds to H, one lone pair, shape is AX3E.

Geometry Summary for AXmEn Configurations

  • Coordination number (m+n)

    • 2: Linear

    • 3: Trigonal planar, Bent (when lone pairs are present)

    • 4: Tetrahedral, Trigonal pyramidal, Bent (e.g., H2O)

Electron Pair Repulsion and Bond Angles

  • Stronger Repulsion: Lone pairs cause more significant repulsion, altering bond angles.

Carbonyl Group and Geometry

  • Hybridization: Both carbon and oxygen atoms are sp2 hybridized.

  • Bonding: Involves a σ-bond and a π-bond; the double bond features contributions from hybrid orbitals and lone pairs.

Bond Rotations

Single Bonds and Rotations

  • Example: Butane.

    • Free rotation is popular around central bonds, though steric clashes can limit this.

    • Bond rotation can result in staggered or eclipsed configurations, impacting energy profiles due to sterics.

Ethane Configuration

  • Preference for staggered configurations driven by hyperconjugation rather than conventional steric considerations.

Double and Triple Bonds

  • π Bonding: Restricts rotational freedom due to presence of a nodal plane.

  • Atoms cannot rotate around double and triple bonds without breaking those bonds.

Peptide Bonds and Delocalization

  • Peptide bonds typically characterized by sp2 hybridization of nitrogen, allowing interaction with π-systems, leading to resonance stabilization.

Geometry of Peptide Bonds

  • The planar arrangement in peptide bonds includes carbonyl (C=O) and amine (N-H) along with Cα atoms, creating a visually planar backbone for proteins.

Ramachandran Plot

  • Dihedral Angles (${1}$ and ${2}$): Restricted by sterics and optimal interactions like hydrogen bonding.

  • Core Concept: Defines configurations of protein backbones through dihedral angles reflecting adjacent peptide planes.

Molecular Geometry Discussion

  • Key considerations in molecular geometry discussions:

    • Bond lengths: distances between atoms in a bond.

    • Bond angles: angles between bonds.

    • Dihedral angles: spatial configurations involving multiple atoms.

  • Connectivity Representation: Atoms are also described using bond lengths, angles, and dihedrals, capturing the underlying chemistry in biomolecules.

Bond Lengths

  • General Range: Bond lengths span around 1 Å (10^(-10) m).

  • Length Influencing Factors:

    • Single bonds > double bonds > triple bonds in length.

    • The effective charge also impacts the covalent radii.

  • Examples of Specific Bond Lengths:

    • C-C single bond: 154 pm

    • C-C double bond: 134 pm

    • C-C triple bond: 121 pm

    • C-H bond: 110 pm

Bond Angles

  • Bond angles arise from electron pair repulsion:

    • Symmetrical arrangements around a central atom without lone pairs yield perfect angles (e.g., tetrahedron: 109.5°).

    • Presence of lone pairs compresses angles (e.g., H-N-H angle in NH3: 106.8°, H-O-H in H2O: 104.5°).

Sugar and Ring Flips (I)

  • Ring conformations hinder motion, as seen in cyclohexane:

    • Chair form (A): lowest energy, alternating hydrogens minimize steric clashes.

    • Twisted boat (B): alternative configuration.

    • Substituents prefer equatorial positioning to reduce steric repulsion.

Sugar and Ring Flips (II)

  • Various sugars adopt cyclic structures, exhibiting distinct conformations:

    • Examples: α-D-glucopyranose, β-D-ribofuranose, 2-deoxyribofuranose, β-D-fructofuranose.

    • Specific puckering patterns impact structures based on their ring configuration.

Non-Covalent Interactions (Part III)

Molecule Interactions Overview

  • Bonding Interactions:

    • Chemical bonds, bond angles, and dihedrals arise from covalent bonds, influenced by electron pair repulsion.

  • Non-Bonded Interactions:

    • Hydrogen bonds

    • Dipole interactions

    • Induced dipoles

    • London dispersion forces

    • π-π interactions

    • Hydrophobic interactions

Van der Waals Forces

  • Types of Van der Waals forces include:

    • Induction (Debye) forces and dispersion (London) forces: Attractive;

    • Permanent multipoles: Depend on molecular orientation.

  • Characteristics:

    • Weaker than ionic and covalent bonds.

    • Additive and non-directional.

    • Short-ranged and dependent on distance.

Dipoles

  • Molecular Dipoles: Result from the combination of individual dipole moments in molecules (e.g., water).

  • Bond Dipoles: Polarized bonds giving rise to partial charges.

    • Evaluated by product of bond length and charge magnitude.

Permanent Dipole Interactions

  • Permanent dipoles lead to interactions based on fixed molecular geometries, and these interactions can dictate solubility and reactivity in polar environments.

Induced Dipoles

  • Induced dipoles occur in neutral atoms/molecules influenced by polar entities, termed as Debye forces, leading to temporary charge separation.

London Dispersion Forces

  • Nature: Electron motion results in transient asymmetries in charge distribution, creating instantaneous dipoles.

  • These fluctuations lead to interactions between neighboring atoms.

π Interactions

  • π Systems can engage in attractive interactions; the foundation of these interactions may originate from electrostatic or orbital overlaps.

π-Cation Interactions

  • Aromatic rings can facilitate interactions with cations via regions of positive and negative charge, leading to interactions important in biochemical environments.

Hydrogen Bonding

  • Governed by interactions between hydrogen (H) atoms covalently bonded to electronegative atoms (N, O) and other electronegative atoms.

  • Directionality promotes some covalent-like character, positioning hydrogen bonds as crucial for structural integrity in biomolecules.

Special Cases of Hydrogen Bonds

  • C-H···O Interactions: Weakly polar C-H bonds may engage in H-bonding under specific conditions.

  • Salt Bridges: Complex interactions combining hydrogen bonding and electrostatics, often found in protein structures, are contingent on pH levels due to side chain charges.

Hydrophobicity

  • Hydrophobic substances exhibit minimal solubility in aqueous media and tend to coalesce into lower energy structures to minimize unfavorable surface interactions.

  • Hydrophobicity arises from the entropic effects of solvation, influencing molecular conformations significantly.

Strengths of Interactions Summary

  • Comparison of interaction types:

    • Covalent bond C-H: 95 to 115 kcal/mol

    • Covalent bonds C-C (single to triple): 90 to 240 kcal/mol

    • Hydrogen bonds in biomolecules: 3.0 to 7.0 kcal/mol

    • Dipole-dipole: 0.5 to 2.0 kcal/mol

    • London forces: <<1 kcal/mol

Summary of Lecture

  1. Hybrid orbitals from atomic orbitals enhance descriptions of bond geometries around central atoms.

  2. VSEPR theory elucidates geometries based on electron pair repulsion minimizing interactions.

  3. Freedom of rotation differs with bond type; single bonds vs. double/triple bonds affect molecular flexibility.

  4. Cyclic structures limit available conformations, particularly with substituents induced.

  5. Non-covalent forces expand interactions beyond covalent bonds, influencing molecular behavior.

  6. The interaction of diverse forces creates complex biological behaviors, including hydrogen bonding and hydrophobicity.