Structure and Properties Help (page 2)
Organic chemistry is the study of carbon (C) compounds, molecules which have covalent bonds. Carbon atoms can bond to each other to form open-chain compounds, or cyclic (ring) compounds. Both types can also have branches of C atoms. Saturated compounds have C's bonded to each other by single bonds, C—C; unsaturated compounds have C's joined by multiple bonds. Examples with double bonds and triple bonds are shown below. Cyclic compounds having at least one atom in the ring other than C (a heteroatom) are called heterocyclics. The heteroatoms are usually oxygen (O), nitrogen (N), or sulfur (S).
Most carbon-containing molecules are three-dimensional. In methane, the bonds of C make equal angles of 109.5° with each other, and each of the four H's is at a vertex of a regular tetrahedron whose center is occupied by the C atom. Other shapes do occur: ethene, for example, is planar, and ethyne (acetylene) is linear.
Organic compounds show a widespread occurrence of isomers, which are compounds having the same molecular formula but different structural formulas. Isomers have different chemical and physical properties. This phenomenon of isomerism is exemplified by isobutane and n-butane. The number of isomers increases as the number of atoms in the molecule increases.
Hydrocarbons contain only C and hydrogen (H). H's in hydrocarbons can be replaced by other atoms or groups of atoms. These replacements, called functional groups, are the reactive sites in molecules. Double and triple bonds are considered to be functional groups. Some common functional groups are given in the Functional Group Table. The "R" group is a generic group, and is not part of the functional group of interest. Compounds with the same functional group form a homologous series having similar chemical properties and often exhibiting a regular gradation in physical properties with increasing molecular weight.
The formal charge on a covalently bonded atom equals the number of valence electrons of the unbonded atom minus the number of electrons assigned to the atom in its bonded state. The assigned number is one half the number of shared electrons plus the total number of unshared electrons. The sum of all formal charges in a molecule equals the charge on the species. In this outline formal charges and actual ionic charges are both indicated by the signs + and –. The structures shown below are called Lewis dot structures (or simply dot structures). Each dot represents an electron in the outer shell of the atom. These drawings can be highly useful in determining if an atom bears a formal charge.
An atomic orbital (AO) is a region of space about the nucleus in which there is a high probability of finding an electron. For organic molecules, the atomic orbitals of most interest are the s orbital and the p orbitals.
The s orbital is a sphere around the nucleus, as shown below. A p orbital has 2 lobes touching on opposite sides of the nucleus. The three p orbitals are labeled px, py, and pz because they are oriented along the x-, y-, and z-axes, respectively. In a p orbital there is no chance of finding an electron at the nucleus—the nucleus is called a node.
Three principles are used to distribute electrons in orbitals.
- "Aufbau" or building-up principle. Orbitals are filled in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, etc.
- Pauli exclusion principle. No more than two electrons can occupy an orbital and then only if they have opposite spins.
- Hund's rule. When filling orbitals of equal energy, place one electron in each orbital (using parallel spins) before pairing electrons. (Substances with unpaired electrons are paramagnetic—they are attracted to a magnetic field.)
Hybridization and Bonding
A carbon atom must provide four equal-energy orbitals in order to form four equivalent bonds, as in methane, CH4. It is assumed that the four equivalent orbitals are formed by blending the 2s and the three 2p AO's to give four new hybrid orbitals, called sp3 orbitals. The larger lobe, the "head," having most of the electron density, overlaps with an orbital of its bonding mate to form the bond. The smaller lobe, the "tail," is often omitted when depicting hybrid orbitals. However, at times the "tail" plays an important role in an organic reaction.
The s and p orbitals of carbon can hybridize in ways other than sp3, as shown below. Repulsion between pairs of electrons causes these hybrid orbitals to have the maximum bond angles. The sp2 and sp hybrid orbitals induce geometries about the C's as shown below.
Head-to-head overlap of AO's gives a sigma (s) bond. The bond angles (angles between s-bonds) at sp3 carbons are 109.5°, leading to a tetrahedral geometry, The bond angles at sp2 carbons are 120°, leading to a trigonal planar geometry, and the bond angles at sp carbons are 180°, leading to a linear geometry. The imaginary line joining the nuclei of the bonding atoms is the bond axis, whose length is the bond length.
Two parallel p orbitals overlap side-by-side to form a pi (π) bond. The bond axis lies in a nodal plane (plane of zero electronic density). Single bonds are σ bonds. A double bond is one σ and one π bond. A triple bond is one σ and two π bonds.
Electronegativity and Polarity
The electronegativity of an atom is its tendency to attract bonding electrons toward itself. The higher the electronegativity, the more strongly the atom attracts and holds electrons. A nonpolar covalent bond exists between atoms having a very small or zero difference in electronegativity. A few relative electronegativities are
F (4.0) > O (3.5) > Cl, N (3.0) > Br (2.8) > S, C, I (2.5) > H (2.1)
A bond formed by atoms of dissimilar electronegativities is called polar due to partial charge separation. The more electronegative element of a covalent bond is relatively negative in charge, while the less electronegative element is relatively positive. The symbols δ+ and δ– represent partial charges in polar bonds. These partial charges should not be confused with ionic charges. Polar bonds are indicated by ; the arrow points toward the more electronegative atom.
The vector sum of all individual bond moments gives the net dipole moment of the molecule. H2O has polar bonds. Since the molecule has a bent shape, the dipoles of the bonds do not cancel and the molecule has a net dipole moment.
Resonance and Delocalized π Electrons
Resonance theory describes species for which a single structure does not adequately describe the species' properties. As an example, consider the cation on the next page (called the allyl cation):
A comparison of the calculated and observed bond lengths shows that the 2 C–C bonds are the same length. Neither resonance structure alone can explain this similarity in bond length. When resonance structures form a resonance hybrid, we obtain a structure consistent with the observed bond length. The resonance hybrid has some double-bond character between the central carbon and both outside carbons. This state of affairs is described by the non-Lewis structure in which dotted lines stand for the partial bonds in which there are delocalized π electrons in an extended π bond created from overlap of p orbitals on each atom. The symbol ↔ denotes resonance, not equilibrium.
The hybrid is more stable than any single resonance structure. The more nearly equal in energy the contributing structures, the greater the resonance energy. When contributing structures have dissimilar energies, the hybrid looks most like the lowest-energy structure. Contributing structures (a) differ only in positions of electrons (atomic nuclei must have the same positions); (b) must have the same number of paired electrons; and (c) must not place more than 8 electrons on any second period atom. Relative energies of contributing structures are assessed by the following rules.
- Structures with the greatest number of covalent bonds are most stable. However, for second-period elements (C, O, N) the octet rule must be observed.
- With few exceptions, structures with the least charge separation are most stable.
- If all structures have formal charge, the most stable (lowest energy) one has – on the more electronegative atom and + on the more electropositive atom.
- Structures with like formal charges on adjacent atoms have very high energies.
- Resonance structures with electron-deficient, positively charged atoms have very high energy, and are usually ignored.
Practice problems for these concepts can be found at: Structure and Properties Practice Problems
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