Organic structure and bonding

Book: Soderberg Chapter: 2

Molecular orbital theory Valence bond theory – explaining bonding geometry Compounds that contain alternating double and single bonds – no use of valence theory ideas of molecular orbital (MO) theory – to think about chemical bonding in a new way

Example: simplest possible covalent bond H2 Using valence orbital theory - two 1s orbitals from each atom overlap, allowing the two electrons to be shared and thus forming a covalent bond In molecular orbital theory - two atomic 1s orbitals don’t just overlap, they combine to form two completely new orbitals They describe the location of an electron pair around two or more nuclei

when orbitals combine, the number of orbitals before the combination takes place must equal the number of new orbitals that result – principle of quantum mechanical theory Orbitals don’t just disappear! two atomic 1s orbitals combine in the formation of H2, the result is two molecular orbitals called sigma (σ) orbitals

Orbitals first σ orbital is lower in energy than either of the two isolated atomic 1s orbitals - bonding molecular orbital The second, σ* orbital is higher in energy than the two atomic 1s orbitals - antibonding molecular orbital In MO theory * sign always indicates an antibonding orbital

aufbau (building up) principle - place the two electrons in the H2 molecule in the lowest energy orbital, which is the (bonding) σ orbital egg-shaped, encompassing the two nuclei, and with the highest likelihood of electrons being in the area between the two nuclei. The high-energy, antibonding σ* orbital can be visualized as a pair of droplets, with areas of higher electron density near each nucleus and a ‘node’, (area of zero electron density) midway between the two nuclei

Electron behavior as wave behavior two separate waves combine – constructive interference (where the two amplitudes reinforce one another) Destructive interference - where the two amplitudes cancel one another out Bonding molecular orbitals – consequence of constructive interference between two atomic orbitals ->attractive interaction and an increase in electron density between the nuclei Antibonding MO-consequence of destructive interference which -> in a repulsive interaction and a ‘canceling out’ of electron density between the nuclei

Pi-bonds Covalent chemical bonds Two lobes of an orbital on one atom overlap two lobes of an orbital on another atom Overlap occurs laterally Each of atomic orbitals has zero electron density at shared nodal plane passing through the two bonded nuclei Pi- referrs to p orbitals; orbital symmetry is the same as of the p orbitals when seen down the bond axis

Conjugated pi-bonds Two or more pi bonds must be separated by only one single bond there cannot be an intervening sp3-hybridized carbon, because this would break up the overlapping system of parallel 2pz orbitals

Conjugated systems can involve oxygens and nitrogens Also carbon Example: metabolism of fat molecules, some of the key reactions involve alkenes that are conjugated to carbonyl groups

organic molecules that contain extended systems of conjugated π bonds often have distinctive colors β-carotene (orange color of carrots) has an extended system of 11 conjugated π bonds

Aromaticity MO theory helpful in in explaining the unique properties of a class of compounds called aromatics Benzene - simplest example of an aromatic compound the π bonds in benzene are significantly less reactive than isolated or conjugated π bonds in most alkenes Heat of hydrogenation - illustrate this unique stability

Exothermic process the alkane is lower in energy than the alkene, so hydrogenating the double bond results in the release of energy in the form of heat heat of hydrogenation of benzene is only 49.8 kcal/mol – quite less than expected Something about the structure of benzene makes these π bonds especially stable - ‘aromaticity’

What exactly is this ‘aromatic’ property that makes the π bonds in benzene so much less reactive than those in alkenes? benzene is a cyclic molecule in which all of the ring atoms are sp2-hybridized This allows the π electrons to be delocalized in molecular orbitals that extend all the way around the ring, above and below the plane of the ring The ring must be planar - otherwise the 2pz orbitals couldn’t overlap properly

Do all cyclic molecules with alternating single and double bonds have this same aromatic stability? NO! Is not flat Pi-bonds are more reactive that those of benzene

conditions that must be satisfied for a molecule to be considered aromatic: 1) It must have a cyclic structure. 2) The ring must be planar. 3) Each atom in the ring must be sp2-hybridized, so that π electrons can be delocalized around the ring. 4) The number of π electrons in the ring must be such that, in the ground state of the molecule, all bonding MOs are completely filled, and all nonbonding and antibonding MOs are completely empty.

Non-covalent interactions Between the molecules or between different functional groups within a single molecule Dipoles Ion-ion, dipole-dipole and ion-dipole interactions van der Waals forces Hydrogen bonds

Dipoles Electronegativity refers to “ the power of an atom in a molecule to attract electrons to itself” In organic chemistry - covalent bonds between two atoms with very different negativities, and in these cases the sharing of electrons is not equal more electronegative atom pulls the two electrons closer to itself

Result of unequal sharing is bond dipole or polar covalent bond has both negative and positive ends, or poles electron density is lower (the positive pole) and higher (the negative pole) degree of polarity depends on the difference in electronegativity between the two atoms in the bond.

Most molecules contain both polar and nonpolar covalent bonds Depending on the location of polar bonds and bonding geometry – molecules can have an overall dipole moment Water results from the combined dipoles of its two oxygen-hydrogen bonds

Bond Dipole Moments Dipole moments are due to differences in electronegativity. They depend on the amount of charge and distance of separation. They are measured in debyes (D). Chapter 2

Bond Dipole Moments for Some Common Covalent Bonds Chapter 2

Molecular Dipole Moments The molecular dipole moment is the vector sum of the bond dipole moments. Depend on bond polarity and bond angles. Lone pairs of electrons contribute to the dipole moment. Chapter 2

Ion-ion, dipole-dipole and ion-dipole interactions common types of ion-ion interaction in biological chemistry – between Mg ion and a carboxylate or phosphate group 2-phosphoglycerate - intermediate in the breakdown of glucose Polar molecules – align themselves in such a way to allow their respective positive and negative poles to interact with each other dipole-dipole interaction

Dipole–Dipole

charged species (an ion) interacts favorably with a polar molecule or functional group - ion-dipole interaction Example in biological chemistry: between a metal cation, most often Mg+2 or Zn+2, and the partially negative oxygen of a carbonyl

Van der Waals forces Nonpolar molecules - weak but still significant attractive intermolecular forces also called London dispersion forces, hydrophobic interactions, or nonpolar interactions result from the constantly shifting electron density in any molecule Even a nonpolar molecule will, at any given moment, have a weak, short-lived dipole will induce a neighboring nonpolar molecule to develop a corresponding transient dipole of its own Result: weak dipole-dipole force is formed

relatively weak by nature constantly forming and dissipating among closely packed nonpolar molecules when added up the cumulative effect can become significant

Dispersions Chapter 2

Hydrogen bonds extremely important form of noncovalent interaction result from the interaction between a hydrogen that is bound to a electronegative heteroatom (N, O or F) and a lone pair on a second nitrogen or oxygen When H loses electron density in a polar bond it essentially becomes an approximation of a ‘naked’ proton capable of forming a strong interaction with a lone pair on a neighboring oxygen or nitrogen atom

depicted with dotted lines in chemical structures Group that provides a proton to a hydrogen bond is said to be acting as a hydrogen bond donor A group that provides an oxygen or nitrogen lone pair is said to be acting as a hydrogen bond acceptor water molecule, for example, can be both a hydrogen bond donor and acceptor Also many common organic functional groups

strongest type of noncovalent interaction enormous implications in biology Genetic inheritance - hydrogen bonding arrangement between DNA bases

Effect of Branching on Boiling Point The long-chain isomer (n-pentane) has the greatest surface area and the highest boiling point. As the amount of chain branching increases, the molecule becomes more spherical and its surface area decreases. The most highly branched isomer (neopentane) has the smallest surface area and the lowest boiling point. Chapter 2

Boiling Points and Intermolecular Forces H 3 O dimethyl ether, b.p. = -25°C C H 3 2 O ethanol, b.p. = 78°C Hydrogen bonding increases the b. p. of the molecule. C H 3 2 O ethanol, b.p. = 78°C ethyl amine, b.p. 17°C C H 3 2 N O—H is more polar than N—H, so alcohols have stronger hydrogen bonding. Chapter 2

Solved Problem 4 Rank the following compounds in order of increasing boiling points. Explain the reasons for your chosen order. Copyright © 2006 Pearson Prentice Hall, Inc. Chapter 2

Solution To predict relative boiling points, we should look for differences in (1) hydrogen bonding, (2) molecular weight and surface area, and (3) dipole moments. Except for neopentane, these compounds have similar molecular weights. Neopentane is the lightest, and it is a compact spherical structure that minimizes van der Waals attractions. Neopentane is the lowest-boiling compound. Neither n-hexane nor 2,3-dimethylbutane is hydrogen bonded, so they will be next higher in boiling points. Because 2,3-dimethylbutane is more highly branched (and has a smaller surface area) than n-hexane, 2,3 dimethylbutane will have a lower boiling point than n-hexane. The two remaining compounds are both hydrogen-bonded, and pentan-1-ol has more area for van der Waals forces. Therefore, pentan-1-ol should be the highest-boiling compound. We predict the following order: neopentane < 2, 3-dimethylbutane < n-hexane < 2-methylbutan-2-ol < pentan-1-ol The actual boiling points are given here to show that our prediction is correct. Copyright © 2006 Pearson Prentice Hall, Inc. 10°C 58°C 69°C 102°C 138°C 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

Polar Solute in a Polar Solvent Dissolves Hydration releases energy; entropy increases. Chapter 2

Polar Solute in Nonpolar Solvent The solvent cannot break apart the intermolecular interaction of the solute, so the solid will not dissolve in the solvent. Chapter 2

Nonpolar Solute with Nonpolar Solvent The weak intermolecular attractions of a nonpolar substance are overcome by the weak attractions for a nonpolar solvent. The nonpolar substance dissolves. Chapter 2

Nonpolar Solute with Polar Solvent If a nonpolar molecule were to dissolve in water, it would break up the hydrogen bonds between the water molecules. Therefore, nonpolar substances do not dissolve in water. Chapter 2

Summary MO theory Valence theory Interactions Dipole moment Solubility