Chapter 2 Structure and Properties of Organic Molecules

Organic Chemistry, 6th EditionL. G. Wade, Jr. Chapter 2Structure and Propertiesof Organic Molecules Jo Blackburn Richland College, Dallas, TX Dallas County Community College District ã 2006,Prentice Hall

=> Wave Properties of Electrons • Standing wave vibrates in fixed location. • Wave function, , mathematical description of size, shape, orientation. • Amplitude may be positive or negative. • Node: amplitude is zero. Chapter 2

Wave Interactions • Linear combination of atomic orbitals • on different atoms produce molecular orbitals • on the same atom give hybrid orbitals. • Conservation of orbitals. • Waves that are in phase add together.Amplitude increases. • Waves that are out of phase cancel out. => Chapter 2

=> Bonding Region • Electrons are close to both nuclei. Chapter 2

Sigma Bonding • Electron density lies between the nuclei. • A bond may be formed by s-s, p-p, s-p, or hybridized orbital overlaps. • The bonding MO is lower in energy than the original atomic orbitals. • The antibonding MO is higher in energy than the atomic orbitals. => Chapter 2

Bonding Molecular Orbital Two hydrogens, 1s constructive overlap => Chapter 2

Anti-Bonding Molecular Orbital Two hydrogens, destructive overlap. => Chapter 2

H2: s-s overlap => Chapter 2

Cl2: p-p overlap Constructive overlap along the same axis forms a sigma bond. => Chapter 2

=> HCl: s-p overlap Question: What is the predicted shape for the bonding MO and the antibonding MO of the HCl molecule? Chapter 2

=> Pi Bonding • Pi bonds form after sigma bonds. • Sideways overlap of parallel p orbitals. 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

Molecular Shapes • Bond angles cannot be explained with simple s and p orbitals. Use VSEPR theory. • Hybridized orbitals are lower in energy because electron pairs are farther apart. • Hybridization is LCAO within one atom, just prior to bonding. => Chapter 2

=> sp Hybrid Orbitals • 2 VSEPR pairs • Linear electron pair geometry • 180° bond angle Chapter 2

=> sp2 Hybrid Orbitals • 3 VSEPR pairs • Trigonal planar e- pair geometry • 120° bond angle Chapter 2

=> sp3 Hybrid Orbitals • 4 VSEPR pairs • Tetrahedral e- pair geometry • 109.5° bond angle Chapter 2

=> 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

=> Rotation around Bonds • Single bonds freely rotate. • Double bonds cannot rotate unless the bond is broken. Chapter 2

Isomerism • Same molecular formula, but different arrangement of atoms: isomers. • Constitutional (or structural) isomers differ in their bonding sequence. • Stereoisomers differ only in the arrangement of the atoms in space. => Chapter 2

=> Structural Isomers Chapter 2

Trans - across Cis - same side No cis-trans isomers possible => Stereoisomers Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp2 carbon. Chapter 2

are due to differences in electronegativity. depend on the amount of charge and distance of separation. In debyes,  = 4.8 x (electron charge) x d(angstroms) => Bond Dipole Moments Chapter 2

=> Molecular Dipole Moments • Depend on bond polarity and bond angles. • Vector sum of the bond dipole moments. Chapter 2

Effect of Lone Pairs Lone pairs of electrons contribute to the dipole moment. => Chapter 2

Intermolecular Forces • Strength of attractions between molecules influence m.p., b.p., and solubility, esp. 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 Dispersions • Between nonpolar molecules • Temporary dipole-dipole interactions • Larger atoms are more polarizable. • Branching lowers b.p. because of decreased surface 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 stronger hydrogen bonding. => Chapter 2

H Bonds => Chapter 2

ethanol, b.p. = 78° C ethyl amine, b.p. = 17 ° C Boiling Points and Intermolecular Forces Chapter 2

Solubility • Like dissolves like. • 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 withNonpolar Solvent => Chapter 2

Nonpolar Solute withNonpolar Solvent => Chapter 2

Nonpolar Solute with Polar Solvent => Chapter 2

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

Hydrocarbons • Alkane: single bonds, sp3 carbons • Cycloalkane: carbons form a ring • Alkene: double bond, sp2 carbons • Cycloalkene: double bond in ring • Alkyne: triple bond, sp carbons • Aromatic: contains a benzene ring => Chapter 2

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

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

=> Compounds Containing Nitrogen • Amines: RNH2, RNHR', or R3N • Amides: RCONH2, RCONHR, RCONR2 • Nitrile: RCN Chapter 2

End of Chapter 2 Chapter 2

Chapter 2 Structure and Properties of Organic Molecules