Chapter 1: Structure and Bonding
Atoms share electron pairs to form covalent bonds
An atom consists of a small, dense nucleus of protons and neutrons surrounded by electrons that occupy orbitals of increasing energy. The number and arrangement of the valence (outermost) electrons determine how an atom bonds. The octet rule states that neutral second-row atoms common in organic molecules — carbon, nitrogen, oxygen, and their neighbors — become especially stable when surrounded by eight valence electrons. Hydrogen is an exception to this count: as a first-row atom with only a 1s orbital available, it becomes stable with two valence electrons rather than eight. Second-row…
sp3 hybridization and tetrahedral geometry
Carbon has four valence electrons and forms four bonds, but its ground-state 2s and 2p orbitals are not equivalent. To account for four identical bonds, one 2s and three 2p orbitals are combined into four equivalent sp3 hybrid orbitals. The four sp3 orbitals point toward the corners of a regular tetrahedron, separated by the maximum possible angle of about 109.5 degrees. In methane, each sp3 orbital on carbon overlaps end-on with the 1s orbital of a hydrogen atom to form a σ bond, a bond whose electron density lies along the internuclear axis. The same sp3 description applies to any carbon…
sp2 hybridization and the π bond
When a carbon atom forms a double bond, it has three regions of electron density — three σ-bonding connections — so one 2s and two 2p orbitals combine into three sp2 hybrid orbitals. These lie in a common plane, separated by about 120 degrees, and form the σ framework. One 2p orbital remains unhybridized and stands perpendicular to that plane. Side-on overlap of this p orbital with the corresponding p orbital on the adjacent carbon produces a π bond, whose electron density lies above and below the σ framework. A carbon–carbon double bond therefore consists of one σ bond and one π bond.…
sp hybridization and the triple bond
A carbon–carbon triple bond contains two π bonds. Each π bond forms by side-on overlap of a pair of parallel 2p orbitals; the two pairs lie in perpendicular planes around the carbon–carbon axis. Building those two π bonds uses two of each carbon’s three 2p orbitals. That leaves one 2s and one 2p orbital on each carbon. They combine to form two sp hybrid orbitals, which point in opposite directions, 180 degrees apart. The two sp orbitals provide the two σ-bonding directions: one forms the carbon–carbon σ bond, and the other forms the bond to the atom on the far side (hydrogen in acetylene). A…
Molecular shape and bond polarity
The three-dimensional shape of a molecule follows from the arrangement of electron pairs around each atom. Bonding and nonbonding pairs orient themselves to be as far apart as possible: four pairs give a tetrahedral arrangement, and lone pairs push the bonded atoms into pyramidal (ammonia) or bent (water) shapes. Bonds are not always symmetric. When two bonded atoms differ in electronegativity — the tendency to attract shared electrons — the bonding pair is displaced toward the more electronegative atom. The bond is then polar, carrying a partial negative charge (δ−) on one atom and a partial…
Condensed and line-angle structures; functional groups
Because full Lewis structures are cumbersome for larger molecules, organic chemists use condensed and line-angle structures. A condensed structure lists the atoms in sequence and groups each carbon with its hydrogens (CH₃CH₂OH). A line-angle, or skeletal, structure goes further: carbon atoms are represented by the ends and vertices of a zig-zag of lines, hydrogen atoms on carbon are not drawn, and only heteroatoms (and the hydrogens bonded to them) appear explicitly. Reading these abbreviations fluently is essential, because the reactive behavior of a molecule is concentrated in its…
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