Structure and Properties Help
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.)
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