Organic Chemistry McMurry 7th

Organic chemistry
the study of carbon compounds
atomic number (Z)
the number of protons in the nucleus of an atom
mass number (A)
the total of protons plus neutrons
isotopes
atoms of the same element that have different mass numbers
atomic weight
the average mass number of the atoms of an element
orbital (Ψ)
a wave function, which describes the volume of space around a nucleus in which an electron is most likely found
node
a surface of zero electron density within an orbital
ground-state electron configuration
the most stable, lowest energy electron configuration of a molecule or atom
electron configuration rules
1) the lowest energy orbitals fill first, according to the order 1s->2s->2p->3s->3p->4s->3d, a statement called the aufbau principle; 2) electrons act as if they were spinning around an axis, in much the same way that the earth spins. This spin can have two orientations, denoted as up and down arrows; only two electrons can occupy an orbital, and they must be of opposite spin, a statement called the Pauli exclusion principle; 3) if two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full,a statement called Hund’s rule
s orbital

spherical shape; one per shell

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p orbital

dumbell shaped; three mutually perpendicular per shell, denoted as px, py, pz

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d orbital

4 of 5 are cloverleaf shaped and the fifth is an elongated dumbell shape with a donut around the middle; five per shell

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electron shells
the layers in which the orbitals of an atom are organized into; they become progressively larger and higher in energy the further from the atom they are
valence shell
the outermost shell of an atom
covalent bond
a shared-electron bond between atoms
molecule
the neutral collection of atoms held together by covalent bonds
electron-dot structures
aka Lewis structures; a way of representing the valence electrons of an atom as dots
line-bond structures
aka Kekule structures; a two-electron covalent bond is indicated as a line drawn between atoms
lone-pair electrons/nonbonding electrons
valence electrons that are not used for bonding
valence bond theory

when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom, a covalent bond forms between the two atoms

ex., H-H bond

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sigma (s) bonds
a covalent bond formed by head-on overlap of atomic orbitals
bond strength
an alternative name for bond dissociation energy (the energy required to break a bond)
bond length
the equilibrium distance between the nuclei of two atoms that are bonded to each other
sp3 hybrid orbital
a hybrid orbital derived by combination of an s atomic orbital with three p atomic orbitals; the four orbitals that result are directed toward the corners
;of a regular tetrahedron at angles of 109; to each other; they form stronger bonds than regular s or p orbitals
bond angle

the angle formed between two adjecent bonds; e.g., the tetrahedral orientation of methane (CH4) where each H is 109.5; from the other

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tetrahedral sp3 orbitals

contain 109.5; angles, e.g., methane

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trigonal sp2 orbitals

120; angles between orbitals

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sp2 hybrid orbitals
a hybrid orbital derived by combination of an s atomic orbital with 2 p atomic orbitals; the three orbitals that result lie in a plane at angles of 120; to each other, and the other p orbital remains unchanged; shown as double bonds
pi (p) bond
the covalent bond formed by sideways overlap of atomic orbitals; e.g., carbon-carbon double bonds
sp hybrid orbital
a hybrid orbital derived from the combination of an s and a p atomic orbital; the two orbitals that result are oriented at an angle of 180; to each other and the remaining two p orbitals are left unchanged; shown as triple bonds
linear sp bond

180; angles; e.g., acetylene

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molecular orbital (MO) theory
a description of covalent bond formation as resulting from a mathematical combination of atomic orbitals (wave functions) to form molecular orbitals
bonding MO
a molecular orbital that is lower in energy than the atomic orbitals from which it is formed
antibonding MO
a molecular orbital that is higher in energy than the atomic orbitals from which it is formed
condensed structure
a shorthand way of writing structures in which carbon-hydrogen and carbon-carbon bonds are understood rather than shown explicitly; e.g., propane = CH3CH2CH3
skeletal structures

a shorthand way of writing structures in which carbon atoms are assumed to be at each intersection of two lines (bonds) and at the end of each line; e.g., phenol, C6H6O

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rules for drawing skeletal structures
1) carbon atoms aren’t shown but are assumed to be at each intersection and at the end of lines; 2) hydrogen bonded to carbon is not shown – carbon always has a valence of 4, thus the hydrogens are known to be there; 3) atoms other than H and C are shown
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