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

Chapter 22 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. =>

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 23

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 24

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

Chapter 26 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

Chapter 27 π Bonding Pi (π) bond : formed by constructive side-to-side overlap of parallel p orbitals.

Chapter 28 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 29 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. =>

σ bonds vs π bonds Chapter 210 σ bondπ bond 1) overlaphead-to-headside-to-side 2) e- densitycylindrically symmetric about bond axis maximum in regions about & below internuclear axis 3) # of bondsonly one can exist between 2 atoms 1 or 2 bonds between atoms 4) free rotationyesno

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 90 o angles because 2p x, 2p y, and 2p z are orthogonal.  Actual angles are ~ 109 o, 120 o or 180 o

Chapter 212 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.

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 213

Chapter 214 sp Hybrid Orbitals (add an s and p orbital together) 2 atomic orbitals = 2 hybrid orbitals Linear electron pair geometry 180° bond angle

Chapter 215 sp Hybrid Orbitals (add an s and p orbital together) Example: BeH 2  Bond angle is 180 o

Chapter 216 sp 2 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 217 sp 2 Hybrid Orbitals (add an s and 2 p orbitals together) Example: BeH 3  Bond angle is 120 o

Chapter 218 sp 3 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 (CH 4 )

Summary of s-p type orbitals Hybrid S character P character Bond AngleGeometry Bond Formed sp1/2 180 o Linearσ sp 2 1/32/3120 o Trigonal planar σ sp 3 1/43/4109 o Tetrahedralσ

Chapter 220 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! NH 2 NH 2 CH 3 -C  C-CHO

Chapter 221 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 222 Structural Isomers

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

Chapter 224 Rotation around Bonds σ bonds freely rotate. π bonds cannot rotate unless the bond is broken.  only one conformation  leads to 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 sp 2 carbon. Cis - same side Trans - across Chapter 225 Stereoisomers No cis-trans isomers possible

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),  x  (electron charge) x d(angstroms)

Bond Dipole Moments(μ)

Chapter 228 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 229 More examples:

Chapter 230 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 231 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 232 Dipole-Dipole =>

Chapter 233 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 234 Dispersions =>

Chapter 235 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 236 H Bonds =>

Chapter 237 Hydrogen Bonding & BP Higher bp’s for NH 3 & H 2 O which can H- bond H 2 O bp is so high because… Compoundbp o C CH NH H2OH2O+100

Chapter 238 Boiling Points and Intermolecular Forces =>

Chapter 239 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 240 Ionic Solute with Polar Solvent Hydration releases energy. Entropy increases. =>

Chapter 241 Ionic Solute with Nonpolar Solvent Attraction for solvent much weaker than ion-ion interactions

Chapter 242 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 243 Nonpolar Solute with Polar Solvent Dissolving solute would require breaking up organized solvent structure. (crisco in water)

SoluteSolventSoluble? Polar Yes PolarNon-polarNo Non-polar Yes Non-polarPolarNo

Chapter 245 Classes of Compounds Classification based on functional group Three broad classes  Hydrocarbons  Compounds containing oxygen  Compounds containing nitrogen =>

Chapter 246 Hydrocarbons Compounds containing only C and H. Includes alkane, cycloalkane, alkene, cycloalkene, alkyne, and aromatic hydrocarbons Typically hydrophobic

Chapter 247 Alkanes Hydrocarbons that contain only single bonds, sp 3 carbons (CH 3 -CH 2 -CH 3 ) Can be cyclic : then called cycloalkane Alkanes are very unreactive Alkane portion of molecule is called an alkyl group

Chapter 248 Alkenes Hydrocarbons that contain C=C double bonds, sp 2 carbons. More reactive than alkanes Show geometric isomerism Cycloalkene: double bond in ring

Chapter 249 Alkynes Hydrocarbons with C—C triple bond, sp carbons More reactive than alkanes & alkenes

Chapter 250 Aromatic Hydrocarbons Hydrocarbons with alternating C—C single bond and C=C double bonds. Have unusual stability compared to alkenes and cycloalkenes

Chapter 251 Compounds Containing Oxygen Alcohol: R-OH Ether: R-O-R ’ Aldehyde: RCHO Ketone: RCOR '

Chapter 252 Alcohol Organic compounds that contian the hydroxyl group (-OH) Common solvents Have polar & non-polar parts

Chapter 253 Ethers Organic compounds that containing 2 alkyl groups bonded to a bridging O atom More polar than hydrocarbons because of the O atom.

Chapter 254 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 255 Carboxylic Acids and Their Derivatives Organic compounds containing the carboxyl group Example: CH 3 COOH = These are weak acids, often used in buffers

Chapter 256 Carboxylic Acids and Their Derivatives Carboxylic Acid: RCOOH Acid Chloride: RCOCl Ester: RCOOR ' Amide: RCONH 2 =>

Chapter 257 Carboxylic Acid Derivatives Have the formula RCOL All react with H2O to give RCOOH LCompound —OHacid —XX = halide acid/halide —ORester — NH 2 amide

Chapter 258 Nitrogen Compounds Amines: RNH 2, RNHR ', or R 3 N Amides: RCONH 2, RCONHR, RCONR 2 Nitrile: RCN

59 Amines Alkylated derivatives of ammonia (NH 3 ) General formulas RNH 2, RNHR ', or R 3 N Polar because of C—N and N—H bond. Example: ethylamine = CH 3 CH 2 -NH 2 1,5-pentanediamine

Chapter 260 Amides General formulas : RCONH 2, RCONHR, RCONR 2 The amide C-N linkage in proteins is known as the pepide bond

Chapter 261 Wave Function Wave Equation: H  Solution gives  the wave function   = 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 - + _ + - =>

 for a 1s orbital Chapter 262

 for a 2p orbital Chapter 263 P x, P y and P z are energetically equivalent