Molecular Shape, Intermolecular Forces, & Intro to Organic Chemistry

Lecture 1: Molecular Shape and Intermolecular Forces

Introduction

• The lecture is divided into two parts:
  • Molecular shape (continued from the previous lecture).
  • Intermolecular forces (applications of bonding).
• Organic chemistry will be discussed, focusing on carbon, hydrogen, oxygen, nitrogen, and halides.

Molecular Shape and VSEPR

• Review of the summary table from the last lecture on molecular shapes using VSEPR.
• Lewis structures show bonds and electron pairs but not molecular shape.
• VSEPR rationalizes molecular shapes based on electron group repulsion (bonds and lone pairs).

Electronegativity and Dipoles

• Covalent bonds between atoms with different electronegativities lead to unequal electron sharing.
• Electronegativity Trends: Increases across and up the periodic table.
• Fluorine and oxygen are the most electronegative elements.
• Unequal sharing creates a permanent dipole, with more electron density on the more electronegative element. The bond is polarized.
• This concept mainly applies to covalent compounds, not cations/anions (full charge overwhelms local bond effects).
• Examples of Dipoles:
  • HF: Dipole moment points towards F (more electronegative).
  • H2O: Bent molecule, O pulls electron density from H, creating a positive side and a negative side.
  • CHCl3 (Chloroform): Dipole moment towards the chlorines, opposite the hydrogen.

Molecular Symmetry and Polarity

• Molecules with polar bonds can be nonpolar overall if the dipoles cancel out due to symmetry.
• Example: CO2 is nonpolar because the dipoles from C=O bonds are equal and opposite.
• Supercritical CO2 can be used as a solvent to extract fats.
• $CCl_4$ (Carbon Tetrachloride): Nonpolar due to symmetrical tetrahedral shape, even though C-Cl bonds are polar.

Examples and VSEPR Application

• BF3: Trigonal planar, nonpolar because dipoles cancel out, memorize that Boron doesn't need 8 electrons around the central atom.
• ClF3: T-shaped, polar molecule. Fluorine is more electronegative than the lone pairs, so the dipole moment points towards the fluorine.

Advanced VSEPR Thinking

• Example: $BrF_3$: Determining the correct arrangement of atoms and lone pairs.
• Three possible arrangements of three fluorines and two lone pairs around bromine in $BrF_3$: all fluorines equatorial, two equatorial positions as lone pairs, and one axial and one equatorial position as lone pairs
• Lone pair-lone pair interactions are more repulsive than lone pair-bonding pair interactions.
• The favored geometry is the one that minimizes unfavorable repulsive interactions.
• The correct structure of $ClF_3$ has lone pairs in equatorial positions to minimize 90-degree interactions, resulting in a T-shape and a permanent dipole.

More Examples: $SiF<em>4$ and $SF</em>4$

• $SiF<em>4$: Tetrahedral, nonpolar (similar to $CCl</em>4$). Very volatile gas, byproduct of silicon etching.
• $SF_4$: Sulfur has a lone pair, resulting in a trigonal bipyramidal arrangement. The lone pair is in the equatorial position to minimize interactions, creating a seesaw shape and a polar molecule.

Application: Molecular Polarity and Drug Action

• Molecular shape and polarity determine how molecules interact with each other.
• Examples:
  • Morphine:
  • Heroin: Acetylated morphine, less polar, crosses the blood-brain barrier faster.
  • Methadone: Different structure, but similar shape and electrostatics, acts as an opioid receptor agonist, used for treating opioid addiction.
  • Naloxone: Binds strongly to opioid receptors, displacing heroin without eliciting a response, used to treat overdoses.

Intermolecular Forces: States of Matter

• Intermolecular forces determine whether a substance is a solid, liquid, or gas at room temperature.
• Example: Halides (Group 7): Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
• Boiling points, melting points, and crystal packing depend on the strength of intermolecular interactions.
• Stronger intermolecular forces lead to harder or tougher materials.

Boiling Point and Kinetic Energy

• Increasing temperature increases kinetic energy, causing molecules to wiggle and jiggle.
• Melting point: Temperature at which a substance transitions from solid to liquid.
• Boiling point: Temperature at which a substance transitions from liquid to gas (vapor).
• 101 kilopascals (kPa) or $1.013 </li> <li>10^5$ Pascals is standard pressure (approximately one atmosphere).

Boiling Point as a Measure of Intermolecular Forces

• Boiling point indicates the strength of intermolecular forces.
• The higher the boiling point, the stronger the intermolecular forces.
• The concept is most relevant for covalently bonded molecules; ionic solids require much higher temperatures to boil.

Types of Intermolecular Forces

• Dispersion forces (London forces): Weak attractive forces between all molecules.
• Dipole-dipole forces: Stronger than dispersion forces, occur in polar molecules.
• Hydrogen bonding: A special, strong type of dipole-dipole interaction, essential for life.

Nature of Intermolecular Forces

• Electrostatic interactions (positive attracts negative) are the fundamental basis for attractive forces between molecules.

Dispersion Forces (London Forces)

• Present in all molecules because electrons are in everything.
• Molecules are dynamic, with constant rotation, vibration, and movement of electron clouds.
• Instantaneous asymmetry in electron density leads to temporary dipoles (delta positive and delta negative), even in nonpolar molecules like hydrogen.
• The more electrons in the system, the stronger the dispersion forces.
• Boiling points increase with the number of electrons (e.g., dihalides: fluorine to iodine). This is the reason that larger particle sizes, such as nano-particles as in gold nano-particles, becomes gold dust and not golden gas. It just has too many atoms to not be affected heavily by dispersion forces.

Surface Area and Dispersion Forces

• Available surface area is important for dispersion forces.
• Example: Pentane and its isomer (same formula, different structure) have different boiling points.
  Pentane has a more extended structure with greater surface area, leading to stronger dispersion forces.

Real-World Examples and Applications

• Carbon chains: Methane (gas) to waxes and lubricants (long chains) to plastics like polyethylene (very long chains).
• Long linear polyethylene: Extremely long chains create strong dispersion forces, resulting in bulletproof plastics.

Dipole-Dipole Forces

• Stronger than dispersion forces due to permanent dipoles in polar molecules.
• The larger and more polarized the bond, the stronger the interaction.
• Example: HCl (delta positive and delta negative ends).
• Dipole-dipole interaction strength: 5 to 25 kJ/mol, less than covalent bonds (e.g., C-H bond ~154 kJ/mol).

Comparison of Polar and Nonpolar Molecules

• Methylpropane (nonpolar) vs. Acetone (polar ketone): Similar molecular weight but acetone has a much higher boiling point due to dipole-dipole forces.
• Changing functional groups: Changing the electronegativity and bonding arrangements in molecules results in different boiling points and properties (e.g., propane to dimethyl ether to methyl chloride to acetaldehyde to acetonitrile).

Hydrogen Bonding

• A special case of dipole-dipole interaction, crucial for life.
• Occurs when hydrogen is attached to an electronegative atom (O, N, F).
• The electronegative atom pulls electron density from hydrogen, exposing the hydrogen nucleus (proton).
• Results in an anomalously high dipole-dipole interaction.
• For hydrogen bonding to occur, Hydrogen must be attached directly to O, N, or F.
• Water, hydrogen fluoride, can do hydrogen bonding.
• $H_2S$ has dipole-dipole but not hydrogen bonding.

Drawing Hydrogen Bonds

• Dotted lines indicate hydrogen bonds, not full lines.
• Lone pairs and hydrogen should align directly for optimal interaction.
• Examples: Water, ammonia.

Importance of Hydrogen Bonding

• Allows water to be liquid: Hydrogen sulfide ($H_2S$) (sulfur below oxygen) is a gas; water's high boiling point (100°C) is due to four hydrogen bonds per molecule.
• Enables chemistry: Polar liquid solvent for salts and polar organic compounds, essential for biology.
• Water density: Ice is less dense than liquid water due to the hydrogen bonding arrangement, causing ice to float and allowing aquatic life to survive.

Intramolecular Hydrogen Bonding

• Hydrogen bonds can occur between different polar parts of the same molecule.
• Example: Ortho-hydroxybenzoic acid: Intramolecular hydrogen bond controls the shape of the molecule.

Hydrogen Bonding in DNA

• Nuclear bases in DNA hydrogen bond to each other, allowing the mechanism for complementarity of DNA strands.
• Essential for double-stranded DNA, stable DNA, heredity, evolution, complex life.
• Cooperative effect of multiple hydrogen bonds allows DNA to unwind and self-correct.

Summary of Intermolecular Forces

• Main forces: Hydrogen bonding, dipole-dipole, dispersion forces.
• Other forces: Combinations of these, such as dipole-induced dipole and ion-dipole interactions.
• Strength ranking: Hydrogen bonding is stronger than dipole-dipole and dispersion forces.
• Ion-dipole interactions are even stronger.
• Ion-induced dipole: occurs, example, charged species (fully negatively charged) interacting with the delta positive charge on another molecule.

Announcements

• Solutions to lecture questions and worksheets will be posted on Canvas.
• Introduction to organic chemistry lecture to follow.

Lecture 2: Introduction to Organic Chemistry

Introduction to Organic Chemistry

• Concerned with compounds containing C-H and C-C bonds.
• Includes biochemistry and synthetic organic chemistry.
• Aims to understand molecular structures and their relevance in biological processes and drug development. An example is a HIV-1 protease molecule which is a very large macromolecule.

What is Organic Chemistry

• The example given is the structure of Vitamin B12, which shows a complex array of lines, wedges, dashes, double bonds, and element names.
• Molecular orbital theory will be touched upon but not comprehensively covered.

Examples of organic molecules

• Examples: Vitamin C, chemical weapons (Novichok), aspirin.

Proteins and Macromolecules

• Proteins are complex molecules, often depicted in cartoons to show the amino acid backbone, helices, and sheets.
• Cryo-electron microscopy allows for detailed imaging of proteins at near-atomic resolution without crystallization.
• Cryo-electron images reveal proteins structures, with distinct electron density between the atoms, showing atomic resolution after the protein has been vitrified.

Unique Attributes of Carbon

• Carbon sits in the middle of the octet (Group 4), thus makes it want to create four bonds.
• Has intermediate electronegativity, capable of making covalent bonds.
• Forms strong materials to make materials and bits that form humans that do not dissolve in the rain.
• Forms strong bonds with itself that can create very large varying sizes of molecular bonds.

Variety of Molecular Structures

• Carbon forms various structural bits, ring, chain, multiple bonds and coordinate metals and heteroatoms.
• Has several geometries, thus enabling varying shapes.

Representing Organic Molecules

• Several ways to connect atoms when a formula of C2H6O is formulated.
• Several ways to depict molecules made with carbon arises, thus the question is: what's the best way one can draw it?
• Extended formula: Lewis structure (long and tedious to draw).
• Condensed formula: Grouping bits (CH3-CH2-OH) used for quick notes.
• Space-filling model: Shows 3D shape computationally.
• Line structure: Useful (quick to draw) but requires understanding of conventions.
• Implicit hydrogen: Hydrogens are assumed to be present if not drawn, carbon atoms are on the corners or at the end of a line.
• 3D Line structure: To also show the distribution of the atoms in it's 3D arrangement in space.

Stick Abbreviation

• Condensing long structures by condensing notations and drawing quicker lines based on atom valence.
• Corms 4 bonds, N forms 3 bonds, O forms 2 bonds, and Halides form a single bond with three lone pairs.
• Carbon skeleton is drawn without being written with all valences being assumed.
• Example of use: Citronella fragrance on a candle that keeps away mosquitoes.

Notes and conventions

• Carbon atoms are at the intersections of lines. Otherwise a heteroatoms are written.
• If we haven't put a heteroatom attached to a carbon there then it's attached to a hydrogen.
• However, by convention, hydrogens are labelled when attached to Heteroatoms.

Structurally Neat Structures

• The angles are also correct, with 120 degrees for SP2 and SP3 atoms. And 180 for SP atoms. (Explained later here what the SP is. )

Examples of Conversions from Extended Line Structure to Condensed Line Structure

• Draws C3-CH2-CH2-CH2-OH in a nice, lovely stick notation which contains all information and easy to draw.

Representing Molecules in 3D

• For lewis structure a solid line has every bond that is shown has 2 electrons only.
• If the bond exists in the plane of the board or the paper: solid line is used.
• If the bond is coming out of the board: wedge is used.
• If the bond is going backwards into the board: dash is used.

Line and Solid Line Representation: Methane Example

• Two hydrogens are in the plane, one is coming forward, and one is going back.
• Allows for 3D representation.

Hybridization to Orbitals

• Electrons are isolated to the bonding network of that molecule, so called molecular orbitals.

SP3 Hybridization Example

• 4 orbitals in a bonding state: PX, PY, PZ, and S.

• 4 orbitals, mixing together, make SP3, the ‘3’ referring to the 3 P orbitals and the S referring to the single S orbital.

• 4, now resulting, SP3 orbitals form 4 covalent bonds.

• The Sigma is the greek letter S to signify that the bond looks a little bit like an S orbital, and thus marks the signal bond.

• In this methane example: 4 SP3 orbitals make 4 single bonds (4 sigma bonds).

• SP3 orbitals are symmetrically arranged around the carbon in a tetrahedral shape.

SP2 Hybridization Example

• If we need to have a double bond, or a four electron interaction, the 2S orbital mixes with two of the 2P orbitals. A P orbital is left free on the atom.
• Three SP2 orbitals and one a single P orbital are what remain.
• P orbitals do make double bonds and Sigma bonds form along plane atoms and the single P orbital sits above and below each carbon.
• Overlap of these carbons, creates the sigma bond and then a Pi bond (which forms right behind the sigma bond).
• This is why it's given the Greek letter Pi for P, as cutting it yields a P orbital.

SP Hybridization Example

• 1S orbital mixes with 1P orbital to create 2SP orbitals. This allows each carbon to allow 2P orbitals left for multiple bonding.
• Also forming the sigma bond between the two carbons and the SP orbital bond (holding in thios case, the hydrogens).
• With said arrangement, the 2 sets of P orbitals are at right angles of orthogonal (at 90 degrees). With SP, a linear arrangement exists.
• Triple bond molecules are triple bonds (thus a bond above / below / in front).

Hybridization Conclusion with Molecular Shape

• SP3: Tetrahedral bonding arrangement
• SP2: Trigonal Arrangement: one, two, three.
• Linear: SP Hybridization.
• Linear = SP / Trigonal = SP2 / Tetrahedral = SP3

Hybridization Continued

• Sigma: single bond can rotate
• Double or multiple: multiple or double bond, cannot rotate. Must put energy in to make a single bond and then reform into the double bond. Holds the electron density above and below.

Double Bonds Make a Different Side of the Molecule

• I can't rotate a molecule that has a double bond, different from those that can rotate.
• The molecule that I use is not the same as the molecule because it's not being able to rotate.
• Isomers do have the same formula, but differently arranged (E for Entgegen, also E zussarman, opposite sides, or Z).
• Transfats are bad, cis-fats are good.
• Real implication happens from changes in molecular geometry (trans or cis/ Z and E) in a complex organism like the human's.

More Doublebond Examples - EndoCannabinoid and an Enzyme Example

• ZZZZ in saturated fats looks like a questionmark with the double bonded parts preventing intermolecular stacking.
• Unsaturated fatty acids are oils at room temperature. It is unsaturated, as opposed to oil based lubricants.
• Active sites are chiral and operate on chiral atoms.
• Lock and Key is the model by which the chiral molecules fit in to active sites of enzymes, acting in an enantiomer or molecule that really distinguish between enantiomers.

Chirality: Key Conclusion

• Subtle is in more former isomerism coming from handedness.
• You can’t shake a hand coming from the wrong chirality
• This means it comes in the wrong shape.
• Handshaking example: Bodies full of chiral molecules need the proper 3D matching to be proper.
• The chiral molecule with 4 different things attached does not superimpose. It’s not able to be placed on top.

Thalidomide Example

• A German morning sickness drug whose enantiomer was very good (morning sickness) / other enantiomer (birth defects).
• Because enantiomers, were not sold separately, regulation changed to allow chiral drugs must be proven their enantiomers are safe, or must sell enantiomers separately. Thalidomide was sold in a mixture.
• Active sites are chiral, as seen with the Carraway and Spearmint Example (Smells Carraway vs Spearmint with the exact same chemical just a different alignment).
• Limonene smells different depending the type (one lemons one citrus).
• Enzymes active sites are chiral and operate on chiral enzymes.