ORGANIC CHEMISTRY - CHAPTER 3: REACTIONS OF ALKANES

IMPORTANT CONCEPTS

1. The ΔH^o of bond homolysis is defined as the bond-dissociation energy, DH^o. Bond homolysis gives radicals or free atoms.

2. The C–H bond strengths in the alkanes decrease in the order

3. Catalysts speed up the establishment of an equilibrium between starting materials and products.

4. Alkanes react with halogens (except iodine) by a radical chain mechanism to give haloalkanes. The mechanism consists of initiation to create a halogen atom, two propagation steps, and various termination steps.

5. In the first propagation step, the slower of the two, a hydrogen atom is abstracted from the alkane chain, a reaction resulting in an alkyl radical and HX. Hence, reactivity increases from I₂ to F₂. Selectivity decreases along the same series, as well as with increasing temperature.

6. The Hammond postulate states that fast, exothermic reactions are typically characterized by early transition states, which are similar in structure to the starting materials. On the other hand, slow, endothermic processes usually have late (productlike) transition states.

7. The ΔH^o for a reaction may be calculated from the DH^o values of the bonds affected in the process as follows:

               ΔH^o = ∑DH^o_{bonds broken} - ∑DH^o_{bonds formed}

8. The ΔH^o for a radical halogenation process equals the sum of the ΔH^o values for the propagation steps.

9. The relative reactivities of the various types of alkane C – H bonds in halogenations can be estimated by factoring out statistical contributions. They are roughly constant under identical conditions and follow the order

The reactivity differences between these types of CH bonds are greatest for bromination, making it the most selective radical halogenation process. Chlorination is much less selective, and fluorination shows very little selectivity.

10. The ΔH^o of the combustion of an alkane is called the heat of combustion, ΔH^o_comb. The heats of combustion of isomeric compounds provide an experimental measure of their relative stabilities.

From "Organic Chemistry" Textbook of VOLLHARDT & SCHORE

ORGANIC CHEMISTRY - CHAPTER 2: STRUCTURE AND REACTIVITY

IMPORTANT CONCEPTS

1. Chemical reactions can be described as equilibria controlled by thermodynamic and kinetic parameters. The change in the Gibbs free energy, ΔG^o, is related to the equilibrium constant by ΔG^o = - RT ln K = - 1.36 log K (at 25^oC). The free energy has contributions from changes in enthalpy, ΔH^O, and entropy, ΔS^o: ΔG^o = ΔH^o – TΔS^o. Changes in enthalpy are due mainly to differences between the strengths of the bonds made and those of the bonds broken. A reaction is exothermic when the former is larger than the latter. It is endothermic when there is a net loss in combined bond strengths. Changes in entropy are controlled by the relative degree of energy dispersal in starting materials compared with that in products. The greater the increase in energy dispersal, the larger a positive ΔS^o.

2. The rate of a chemical reaction depends mainly on the concentrations of starting material(s), the activation energy, and temperature. These correlations are expressed in the Arrhenius equation: rate constant k = Ae^-Ea/RT.

3. If the rate depends on the concentration of only one starting material, the reaction is said to be of first order. If the rate depends on the concentrations of two reagents, the reaction is of second order.

4. Brønsted acids are proton donors; bases are proton acceptors. Acid strength is measured by the acidity constant K_a; pK_a = - log K_a. Acids and their deprotonated forms have a conjugate relation. Lewis acids and bases are electron pair acceptors and donors, respectively.

5. Electron-deficient atoms attack electron rich atoms and are called electrophiles. Conversely, electron-rich atoms attack electron-poor atoms and are called nucleophiles. When a nucleophile, which may be either negatively charged or neutral, attacks an electrophile, it donates a lone electron pair to form a new bond with the electrophile.

6. An organic molecule may be viewed as being composed of a carbon skeleton with attached functional groups.

7. Hydrocarbons are made up of carbon and hydrogen only. Hydrocarbons possessing only single bonds are also called alkanes. They do not contain functional groups. An alkane may exist as a single continuous chain or it may be branched or cyclic. The empirical formula for the straight-chain and branched alkanes is C_nH_2n+2.

8. Molecules that differ only in the number of methylene groups, CH₂, in the chain are called homologs and are said to belong to a homologous series.

9. An sp³carbon attached directly to only one other carbon is labeled primary. A secondary carbon is attached to two and a tertiary to three other carbon atoms. The hydrogen atoms bound to such carbon atoms are likewise designated primary, secondary, or tertiary.

10. The IUPAC rules for naming saturated hydrocarbons are (a) find the longest continuous chain in the molecule and name it; (b) name all groups attached to the longest chain as alkyl substituents; (c) number the carbon atoms of the longest chain; (d) write the name of the alkane, citing all substituents as prefixes arranged in alphabetical order and preceded by numbers designating their positions.

11. Alkanes attract each other through weak London forces, polar molecules through stronger dipole – dipole interactions, and salts mainly through very strong ionic interactions.

12. Rotation about carbon – carbon single bonds is relatively easy and gives rise to conformations (conformers). Substituents on adjacent carbon atoms may be staggered or eclipsed. The eclipsed conformation is a transition state between staggered conformers. The energy required to reach the eclipsed state is called the activation energy for rotation. When both carbons bear alkyl or other groups, there may be additional conformers: Those in which the groups are in close proximity (60^o) are gauche; those in which the groups are directly opposite (180^o) each other are anti. Molecules tend to adopt conformations in which steric hindrance, as in gaucheconformations, is minimized.

From "Organic Chemistry" Textbook of VOLLHARDT & SCHORE

ORGANIC CHEMISTRY - CHAPTER 1: STRUCTURE AND BONDING IN ORGANIC MOLECULES

IMPORTANT CONCEPTS

1. Organic chemistry is the chemistry of carbon and its compounds.

2. Coulomb’s law relates the attractive force between particles of opposite electrical charge to the distance between them.

3. Ionic bonds result from coulombic attraction of oppositely charged ions. These ions are formed by the complete transfer of electrons from one atom to another, typically to achieve a noble-gas configuration.

4. Covalent bonds result from electron sharing between two atoms. Electrons are shared to allow the atoms to attain noble-gas configurations.

5. Bond length is the average distance between two covalently bonded atoms. Bond formation releases energy; bond breaking requires energy.

6. Polar bonds are formed between atoms of differing electronegativity (a measure of an atom’s ability to attract electrons).

7. The shape of molecules is strongly influenced by electron repulsion.

8. Lewis structures describe bonding by the use of valence electron dots. They are drawn so as to give hydrogen an electron duet and the other atoms electron octets (octet rule). Formal charge separation should be minimized but may be enforced by the octet rule.

9. When two or more Lewis structures differing only in the positions of the electrons are needed to describe a molecule, they are called resonance forms. None correctly describes the molecule, its true representation being an average (hybrid) of all its Lewis structures. If the resonance forms of a molecule are unequal, those which best satisfy the rules for writing Lewis structures and the electronegativity requirements of the atoms are more important.

10. The motion of electrons around the nucleus can be described by wave equations. The solutions to these equations are atomic orbitals,which roughly delineate regions in space in which there is a high probability of finding electrons.

11. An s orbital is spherical; a p orbital looks like two touching teardrops or a “spherical figure eight.” The mathematical sign of the orbital at any point can be positive, negative, or zero (node). With increasing energy, the number of nodes increases. Each orbital can be occupied by a maximum of two electrons of opposite spin (Pauli exclusion principle, Hund’s rule).

12. The process of adding electrons one by one to the atomic orbitals, starting with those of lowest energy, is called the Aufbau principle.

13. A molecular orbital is formed when two atomic orbitals overlap to generate a bond. Atomic orbitals of the same sign overlap to give a bonding molecular orbital of lower energy. Atomic orbitals of opposite sign give rise to an antibonding molecular orbitalof higher energy and containing a node. The number of molecular orbitals equals the number of atomic orbitals from which they derive.

14. Bonds made by overlap along the internuclear axis are called σ bonds; those made by overlap of p orbitals perpendicular to the internuclear axis are called π bonds.

15. The mixing of orbitals on the same atom results in new hybrid orbitals of different shape. One s and one p orbital mix to give two linear sp hybrids, used, for example, in the bonding of BeH₂. One s and two p orbitals result in three trigonal sp² hybrids, used, for example, in BH₃. One s and three p orbitals furnish four tetrahedral sp³ hybrids, used, for example, in CH₄. The orbitals that are not hybridized stay unchanged. Hybrid orbitals may overlap with each other. Overlapping sp³ hybrid orbitals on different carbon atoms form the carbon – carbon bonds in ethane and other organic molecules. Hybrid orbitals may also be occupied by lone electron pairs, as in NH₃.

16. The composition (i.e., ratios of types of atoms) of organic molecules is revealed by elemental analysis. The molecular formula gives the number of atoms of each kind.

17. Molecules that have the same molecular formula but different connectivity order of their atoms are called constitutional or structural isomers.They have different properties.

18. Condensed and bond-line formulas are abbreviated representations of molecules. Dashed-wedged line drawings illustrate molecular structures in three dimensions.

From "Organic Chemistry" Textbook of VOLLHARDT & SCHORE