l3 Molecular Shapes, Forces & Non-Covalent Interactions

• The lecture covers molecular shape, molecular forces, and non-covalent interactions
• It aims to describe simple molecule structures and use MO theory to explain properties of biomolecular structures like peptide bonds and aromatic rings
• It aims to integrate terminology for molecules/chemical bonding and catalogue atom interactions in biomolecules

Recap of Last Lecture

• Electrons are in orbitals and distinguished by quantum numbers n, l, m₁, and ms
• Electrons fill orbitals from low to high energy, following the n+l rule
• Fully filled shells are the core and don't contribute to bonding
• Partially filled shells/subshells with higher energy contain valence electrons relevant for bonding
• Molecular orbitals form by combining atomic orbitals (AOs) in- and out-of-phase
• These combinations can be bonding, non-bonding, or antibonding
• MO theory is used to study bond polarization and conjugated systems like aromatic rings

Lecture Synopsis Topics

• Hybridization of atomic orbitals
• VSEPR theory
• Chemical bonding determines molecular geometries
• Molecular geometry descriptions
• Atom interactions (non-covalent)
• Biomolecule interactions summary

Hybrid Orbitals

• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals
• Atomic orbitals transform into hydrogen-like orbitals; orbitals in the same shell possess similar energy levels
• Combining orbitals may result in an overall lower energy state than original atomic orbitals
• Hybrid orbitals better describe the orientation of bonding interactions.
• Hybrid orbitals derive from mixing atomic orbitals close in energy on the same atom.
• Hybrid orbitals inherit characteristics of the atomic orbitals they are based on, in proportion to their mixture
• They are directional

sp³ Hybridization

• sp³ hybridization involves mixing all p orbitals with the s orbital
• p and s orbitals have similar energies and can mix to lower electron energy
• Mixing three p and one s orbital yields four sp³ hybrid orbitals arranged tetrahedrally
• Methane (CH4) and alkyl carbons serve as examples
• All hybrid orbitals have the same energy
• Low energy orbitals do not mix

sp² Hybridization

• sp² hybridization involves mixing two p orbitals with the s orbital, forming three sp² hybrid orbitals
• The central atom is in a plane, arranged in a triangle
• The remaining p orbital is perpendicular
• C=C double bonds are examples
• Carbonyl (C=O) groups and conjugated systems like aromatic rings also demonstrate this
• The p orbital is slightly higher in energy

sp Hybridization

• sp hybridization mixes only one p orbital with the s orbital, resulting in two sp hybrid orbitals
• Hybrid atoms align along a line in opposite directions with two p orbitals perpendicular
• A carbon-carbon triple bond/Cyanide (CN) exemplify this hybridization

VSEPR Theory

• VSEPR (Valence Shell Electron Pair Repulsion) theory explains molecular arrangement
• Atoms with hybridized valence shell orbitals optimize the distance between repulsing electron pairs (σ bonds/lone pairs)
• Atom valency determines correct hybridization
• Determining geometry requires identifying bound atoms and lone electron pairs
• Determine the number of valence electrons, identify single/double/triple bonds, and identify charges.
• Single bonds distribute first valence electrons (X), with remaining electrons forming lone pairs (E)
• This gives the structure of a central atom A as AXmEn
• Methane (CH4) includes carbon with 4 valence electrons forming single bonds with each H atom
• It has 8 electrons in total, with four bonds and no lone pairs, giving it a shape of AX4
• Nitrogen, in NH3, possesses 5 valence electrons and three bonds to H
• 8 electrons create three bonds, leaving one lone pair and molecule is categorized as AX3E

Relevant Structures for AXmEn

• Linear (AX2): 180° bond angle
• Trigonal Planar (AX3): 120° bond angle
• Tetrahedral (AX4): 109.5° bond angle
• Bent (AX2E): <120° bond angle
• Trigonal Pyramidal (AX3E): <109.5° bond angle
• Bent (AX2E2): <109.5° bond angle

Electron Pair Repulsion and Bond Angles

• Lone pair repulsion is stronger than valence pair repulsion
• This leads to alteration of bond angles

Carbonyl Group (sp² Geometry)

• Both carbon/oxygen atoms are sp² hybridized
• Carbon possesses 4 valence electrons, and oxygen has 6
• There's a double bond between C and O, with one σ-bond/one π-bond, where each atom donates two electrons
• Carbon has two electrons left for two sp² orbitals= forms two more σ-bonds
• Oxygen has four electrons left= forms two lone pairs
• Orbital energies are mismatched, resulting in a polarized bond

Bond Rotations

• Single bonds enable free rotation between molecule parts
• Steric clashes i.e. bulky groups hinder rotation
• Cyclic compounds limit rotation within rings
• Conjugated π-systems reduce rotational freedom
• Steric hindrance leads to staggered or eclipsed methyl groups, causing steric clashes

Ethane Configuration

• Ethane prefers the staggered arrangement for non-steric reasons
• Repulsion between electrons in the bonds between C and H has been proposed but isn't strong enough
• The reason for this behavior is hyperconjugation

Double/Triple Bonds

• Single bonds have no nodal plane correlating with the bond, allowing rotation that doesn't affect the bonding orbital
• Pi(π)-bonds possess a nodal plane
• Rotating the atoms will break the bond
• Atoms cannot rotate around double/triple bonds

Peptide Bonds and Delocalisation

• N is normally sp³ hybridized, but sp² hybridization is possible through resonance
• Sp² hybridization allows interaction with π-systems which lowers overall energy
• This interaction can be accounted for by resonance structures
• As a consequence, N, C, and O end up being sp² centres

Peptide Bond Geometry

• Atoms attached to an sp² center exists in the same plane
• For a peptide bond, the carbonyl (C=O) and amine (N-H) groups are in a single plane
• Atoms connected to the two Cα atoms are also in the plane
• Visualizing these planes reveals how the protein backbone consists of these planes

Ramachandran Plot

• Due to restricted rotations in peptide bonds, protein conformations depend on dihedral angles (phi φ and psi ψ)
• These angles indicate the relative position of "peptide planes"
• The rotation range is restricted by interactions between said atoms
• Steric hindrance limits these dihedral angles together with favorable hydrogen bond production
• Bond angles are displayed by a Ramachandran plot

Molecular Geometry Communication

• The Protein Data Bank stores atoms as XYZ coordinates to define the precise location of each atom
• However, this does not give the connectivity/origin of structure
• Alternatively, bond lengths, angles, and dihedrals describe the structure
• This reveals the chemistry behind molecular structure

Bond Length

• Bond Single bonds are longer than double bonds, and double bonds are generally longer than triple bonds
• This is due to additional π interactions
• Covalent radius of atoms influences bond lengths

Bond Angles

• Bond angles result from repulsion between bonds and lone pairs
• Bonds are roughly equally spaced around the central atom for symmetric arrangements
• Lone pairs have a stronger repulsion than regular bonds, so the angles get compressed

Ring Flips

• Sugar rings have restricted motion
• Steric repulsion defines ring conformations, such as cyclohexane
• Axial substituents prefer the equatorial config. to minimize steric repulsion
• A chair config. is the lowest energy structure but A twisted boat(B) is an alternative structure
• These restricted conformations apply to 5 membered rings but are more constrained due to small size

Sugar/Ring Flips

• Ring conformation is dictated by: substituents, temperature, puckering
• Preferential pucker conformations of ribose (C3'-endo) and deoxyribose (C2'-endo) lead to distinct structures

Interactions Between Atoms

• Bonding interactions involve: chemical bonds, bond angles, and dihedrals
• Whereas non-bonding interactions involve: hydrogen bonds, dipole interactions, induced dipoles, dispersion forces, π-π interactions, and hydrophobicity

Van Der Waals

• Van der Waals forces are distance-dependent interactions independent of chemical bonding
• These encompass electrostatic interactions (charge, dipoles, and hydrogen bonds) and the London dispersion force
• Induction (Debye) forces and dispersion (London) forces are attractive while multipoles are sign dependent
• They are anisotropic because they rely on the orientation of molecules
• They are weaker than covalent/ionic bonds and relatively short in range

Dipoles

• Multipoles are charge distributions that generate e-fields
• A monopole represents a single charge
• A dipole comprises two equal and opposite charges at a certain distance
• A quadrupole encompasses more complicated, charge arrangements
• Polarized bonds create dipoles with nonzero "partial charges"
• The dipole strength can be expressed as the charge times the bond length.
• Overall dipole moments stem from the compilation of individual dipole moments
• Symmetric molecules with polarized bonds are cancelled out (methane)

Dipole Interactions

• Permanent dipoles generate the molecular geometry
• Attraction between the partial negative and partial positive charges gives rise to dipole-dipole interactions
• Molecules with dipoles are known as polar

Induced Dipoles

• Inducing dipoles originates with the presence of other polar atoms/molecules in neutral species
• Moreover, permanent dipoles and charges may induce dipoles
• These interactions are referred to as Debye forces

London Dispersion Forces

• Electron motion introduces asymmetries that lead to instantaneous dipoles
• Electric fields may be felt by other atoms during fluctuations in electron densities
• But a time-averaged dipole of neutral atoms remains zero, even with instantaneous dipoles

Pi-Interactions

• Pi systems have the capacity to interact with other pi systems attractively
• The specific underlying mechanism driving these noncovalent forces has yet to be clarified
• Negative charged center may undergo cation interactions
• Multiple ring systems containing elements that are no C may undergo cyclic interactions

Hydrogen Bonding

• Positively charged hydrogen atoms in polar molecules form interactions with electronegative atoms (N/O)
• Hydrogen bonds are highly directional, so they carry covalent properties
• Hydrogen bonding is the key for biomolecular structure, especially base pairing and secondary structures

Special Hydrogen Bond Cases

• Small differences in electronegativity of C-H bonds means that they are weakly polarized.
• Weak hydrogen bonds can form with C-H as long as the hydrogen acceptor is very electronegative
• In protein-drug interactions and biomolecules, these cases apply
• Salt bridges happen when h-bonding and ionic associations form at the same time
• Proteins commonly have a specific arrangement of side chains on aspartic acid and glutamic acid, lysine ammonium ions, or arginine.
• Side chains are pH dependent, thus those interactions are also pH-dependent

Hydrophobicity

• Hydrophobic molecules will not mix with aqueous solvents while hydrophilic molecules dissolve readily
• This depends on a molecules capacity to form polar interaction because of charge and/or dipoles, or if it cannot form these interactions
• Hydrophobic substances collapse into small volume when placed inside aqueous media, thus reducing the unfavorable interaction surface
• This is caused by several interactions wherein mixing hydrophobic and hydrophilic molecules reduces interactions created, thus making it an unfavorable interaction

Strengths of Interactions

• Covalent Bond C-H: 95-115 kcal/mol
• Covalent Bonds C-C (single to triple): 90-240 kcal/mol
• Hydrogen Bonds (in biomolecules): 3.0 - 7.0 kcal/mol
• Dipole-Dipole: 0.5 - 2.0 kcal/mol
• London forces <<1 kcal/mol

Lecture Summary

• Hybrid orbitals offer accessible bond geometry descriptions around a given atom
• VSEPR theory explains geometry by minimizing electron repulsion
• Single bonds freely rotate, while double/single bonds don't, limiting conjugated system flexibility
• Cyclic systems are limited to what conformations they can adopt
• Noncovalent forces mediate interactions between molecules by leveraging electrostatics, Debye, and London forces
• Interactions combine to cause complex emergent behavior like π-interactions, hydrogen bonding, and hydrophobicity