Structure and Properties of Organic Molecules Highland Hall Biochem Block Handout #2 Structure and Properties of Organic Molecules Adapted from Chapter 2 of Organic Chemistry by L. G. Wade, Jr.

Molecular Orbitals (MO’s) Definition: interaction between orbitals on different atoms Generated by Linear Combination of Atomic Orbitals (LCAO) Conservation of orbitals # of MO’s created = # of AO’s combined Interactions can be constructive or destructive. => Chapter 2

Bonding MO (between two 1S orbitals) Bonding MO : formed by constructive combination of AO’s with high e- density between the two nuclei. Results in MO with lower energy. Chapter 2

Antibonding MO (between two 1S orbitals) Antibonding MO: MO formed by destructive combination of AO’s with low e- density between the two nuclei Results in MOs higher in energy than the AO’s. Chapter 2

MO Energy Level Diagram Two atomic orbitals will always overlap so as to give 1 bonding MO and 1 antibonding MO. Example: H2 => Chapter 2

The σ bond Sigma (σ) bond : single bond formed by constructive overlap of orbitals in which the e- density is centered about a line connecting the nuclei (i.e. cylindrically symmetrical) Constructive orbital overlap → σ bonding MO Destructive orbital overlap → σ* anti-bonding MO 1S 1S Chapter 2

π Bonding Pi (π) bond : formed by constructive side-to-side overlap of parallel p orbitals. Chapter 2 7

π Bonding Not cylindrically symmetrical Parallel overlap required for π bonds π bonds cannot rotate with respect to each other π bonds formed in multiple bonds π bonds form after a σ bond. Chapter 2

Multiple Bonds A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. => Chapter 2

σ bonds vs π bonds σ bond π bond 1) overlap head-to-head side-to-side 2) e- density cylindrically symmetric about bond axis maximum in regions about & below internuclear axis 3) # of bonds only one can exist between 2 atoms 1 or 2 bonds between atoms 4) free rotation yes no Chapter 2

Hybrid Orbitals Hybrid orbitals : formed by mixing orbitals on the same atom. Mixtures of s and p orbitals Why do we need hybrids? Bond angles cannot be explained with simple s and p orbitals, if so we would predict carbon to have 90o angles because 2px, 2py, and 2pz are orthogonal. Actual angles are ~ 109o, 120o or 180o

VSEPR Valence Shell Electron Pair Repulsion (VSEPR) theory : e- pairs will repel each so as to attain a maximal distance of separation. The observed bond angles are not possible using only s & p orbitals thus we use hybrids. Hybridized orbitals are lower in energy because electron pairs are farther apart. Hybridization is LCAO within one atom, just prior to bonding. Chapter 2

Hybridization Rules 1) Both σ bonding e- and lone pairs occupy hybrid orbitals: # hybrid orbitals = (# σ bonds) + (# lone pairs) 2) The observed hybridization will be that which gives maximum separation of σ bonds and lone pairs 3) For multiple bonds, one bond is a σ bond using hybrid orbitals, while additional bonds are π bonds formed between p orbitals. Chapter 2

sp Hybrid Orbitals (add an s and p orbital together) 2 atomic orbitals = 2 hybrid orbitals Linear electron pair geometry 180° bond angle Chapter 2

sp Hybrid Orbitals (add an s and p orbital together) Example: BeH2 Bond angle is 180o Chapter 2 15

sp2 Hybrid Orbitals (add an s and 2 p orbitals together) 3 atomic orbitals = 3 hybrid orbitals Trigonal planar e- pair geometry 120° bond angle Chapter 2

sp2 Hybrid Orbitals (add an s and 2 p orbitals together) Example: BeH3 Bond angle is 120o Chapter 2 17

sp3 Hybrid Orbitals (add an s and 3 p orbitals together) 4 atomic orbitals = 4 hybrid orbitals Tetrahedral e- pair geometry 109.5° bond angle Example : methane (CH4) Chapter 2

Summary of s-p type orbitals Hybrid S character P Bond Angle Geometry Bond Formed sp 1/2 180o Linear σ sp2 1/3 2/3 120o Trigonal planar sp3 1/4 3/4 109o Tetrahedral

Sample Problems Predict the hybridization, geometry, and bond angle for each atom in the following molecules: Caution! You must start with a good Lewis structure! NH2NH2 CH3-CC-CHO Chapter 2

Isomers Isomers : molecules which have the same molecular formula, but differ in the arrangement of their atoms Constitutional (or structural) : isomers which differ in their bonding sequence. Stereoisomers : molecules that differ only in the arrangement of the atoms in space. Chapter 2

Structural Isomers Chapter 2

Conformations Conformations : structures differing only in rotations about a single bond. Example: ethane Cylindrical symmetry of the σ bond allows rotation about a single bond “eclipsed” “staggered” side view side view

Rotation around Bonds σ bonds freely rotate. π bonds cannot rotate unless the bond is broken. only one conformation leads to stereoisomers Chapter 2

Stereoisomers Stereoisomer : distinct compounds differing only in the spatial arrangement of atoms attached to the C=C double bond. Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp2 carbon. Cis - same side Trans - across No cis-trans isomers possible Chapter 2 25

Dipole Moment (μ) Polar bond : bond in which e- are shared unequally. Non-polar : bond in which e- are shared equally. Dipole moment (μ) : measure of polarity. are due to differences in electronegativity. depend on the amount of charge and distance of separation. In debyes (D),  = 4.8 x  (electron charge) x d(angstroms)

Bond Dipole Moments(μ)

Molecular Dipole Moments Depends on bond polarity and bond angles. Vector sum of the bond dipole moments. Lone pairs of electrons contribute to the dipole moment. μ = 1.9 D μ = 2.3 D μ = 0 D Chapter 2

More examples: Chapter 2 29

Intermolecular Forces Strength of attractions between molecules influence m.p., b.p., and solubility; especially for solids and liquids. Classification depends on structure. Dipole-dipole interactions London dispersions Hydrogen bonding => Chapter 2

Dipole-Dipole Forces Between polar molecules Positive end of one molecule aligns with negative end of another molecule. Lower energy than repulsions, so net force is attractive. Larger dipoles cause higher boiling points and higher heats of vaporization. Chapter 2

Dipole-Dipole => Chapter 2

London Dispersion Induced between non-polar molecules Temporary dipole-dipole interactions Forces are larger for molecules with large surface area. large molecules are more polarizable Branching lowers b.p. because of decreased surface area contact between molecules. Chapter 2

Dispersions => Chapter 2

Hydrogen Bonding Strong dipole-dipole attraction Organic molecule must have N-H or O-H. The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule. O-H more polar than N-H, so there is stronger hydrogen bonding. Chapter 2 35

H Bonds => Chapter 2

Hydrogen Bonding & BP Higher bp’s for NH3 & H2O which can H-bond H2O bp is so high because… Compound bpoC CH4 -162 NH3 -33 H2O +100 Chapter 2 37

Boiling Points and Intermolecular Forces => Chapter 2

Solubility The solubility rule: “Like dissolves like” Ionic & polar solutes dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents. Molecules with similar intermolecular forces will mix freely. Chapter 2

Ionic Solute with Polar Solvent Hydration releases energy. Entropy increases. => Chapter 2

Ionic Solute with Nonpolar Solvent Attraction for solvent much weaker than ion-ion interactions Chapter 2

Nonpolar Solute with Nonpolar Solvent Small attractive forces for solvent are still strong enough to over come weak ion-ion intermolecular forces in solid. Chapter 2

Nonpolar Solute with Polar Solvent Dissolving solute would require breaking up organized solvent structure. (crisco in water) Chapter 2

Solute Solvent Soluble? Polar Yes Non-polar No

Classes of Compounds Classification based on functional group Three broad classes Hydrocarbons Compounds containing oxygen Compounds containing nitrogen => Chapter 2

Hydrocarbons Compounds containing only C and H. Includes alkane, cycloalkane, alkene, cycloalkene, alkyne, and aromatic hydrocarbons Typically hydrophobic Chapter 2

Alkanes Hydrocarbons that contain only single bonds, sp3 carbons (CH3-CH2-CH3) Can be cyclic : then called cycloalkane Alkanes are very unreactive Alkane portion of molecule is called an alkyl group Chapter 2 47

Alkenes Hydrocarbons that contain C=C double bonds, sp2 carbons. More reactive than alkanes Show geometric isomerism Cycloalkene: double bond in ring Chapter 2 48

Alkynes Hydrocarbons with C—C triple bond, sp carbons More reactive than alkanes & alkenes Chapter 2 49

Aromatic Hydrocarbons Hydrocarbons with alternating C—C single bond and C=C double bonds. Have unusual stability compared to alkenes and cycloalkenes Chapter 2 50

Compounds Containing Oxygen Alcohol: R-OH Ether: R-O-R’ Aldehyde: RCHO Ketone: RCOR' Chapter 2

Alcohol Organic compounds that contian the hydroxyl group (-OH) Common solvents Have polar & non-polar parts Chapter 2 52

Ethers Organic compounds that containing 2 alkyl groups bonded to a bridging O atom More polar than hydrocarbons because of the O atom. Chapter 2 53

Aldehydes + Ketones Compounds containing the C=O carbonyl group Somewhat polar because of high μ of C=O bond. Aldehyde : has 1 alkyl group and one H atom on each side of the carbonyl (RCOH) Ketone : has 2 alkyl groups (RCOR’) R = alkyl groups Chapter 2 54

Carboxylic Acids and Their Derivatives Organic compounds containing the carboxyl group Example: CH3COOH = These are weak acids, often used in buffers Chapter 2

Carboxylic Acids and Their Derivatives Carboxylic Acid: RCOOH Acid Chloride: RCOCl Ester: RCOOR' Amide: RCONH2 => Chapter 2 56

Carboxylic Acid Derivatives Have the formula RCOL All react with H2O to give RCOOH L Compound —OH acid —X X = halide acid/halide —OR ester — NH2 amide Chapter 2 57

Nitrogen Compounds Amines: RNH2, RNHR', or R3N Amides: RCONH2, RCONHR, RCONR2 Nitrile: RCN Chapter 2 58

Amines Alkylated derivatives of ammonia (NH3) General formulas RNH2, RNHR', or R3N Polar because of C—N and N—H bond. Example: ethylamine = CH3CH2-NH2 1,5-pentanediamine

Amides General formulas : RCONH2, RCONHR, RCONR2 The amide C-N linkage in proteins is known as the pepide bond Chapter 2 60

Wave Function Wave Equation: H  = E  Solution gives  , the wave function 2 = amplitude of  at a given point in space electron density at a given point in space is a mathematical description of size, shape, orientation, can be + or - + _ - => Chapter 2

 for a 1s orbital Chapter 2

 for a 2p orbital Px, Py and Pz are energetically equivalent Chapter 2