Organic Chemistry Bonding Review

Topics: Atom, Chemical bond, Electron configuration Pages: 15 (3435 words) Published: December 6, 2011

- organic chemistry is the study of life at the molecular level; to many it is the key to understanding life “The language of chemistry- an international language, a language without dialects, a language for all of time, and a language that explains where we came from, what we are, and where the physical world will allow us to go” (Nobelist Arthur Kornberg, a biochemist, 2000)

- but its study has been challenging for students for centuries “Organic chemistry nowadays almost drives me mad. To me it appears like a primeval tropical forest full of the most remarkable things, a dreadful endless jungle into which one does not dare to enter for there seems to be no way out” (Freidrich Wohler, 1835 - he was the first to synthesize an organic molecule, urea, from inorganic materials in 1828)

- but YOU can enjoy it if you make the effort to understand it. It is NOT “rocket science”


1. COVALENT BONDING & SHAPES OF MOLECULES [text 1.1-1.7] ( © Chemistry Dept, University of Western Ontario, 2011) - largely a review of essential material from year-1 chem A. Electronic Structure of Atoms - The bonding behaviour of atoms depends entirely on electron configuration, as revealed by an atom's position in the Periodic Table. - The "Organic Chemist's" periodic table.

Key to numbers: - upper left = atomic number = number of electrons - lower left = number of outer shell electrons (= group number) - upper right = valence = number of unpaired electrons in valence shell - lower right = electronegativity value [a full periodic table with atomic numbers (only) is provided on tests]


- intro chem showed us how quantum numbers fix the identity of electrons in atoms, ie, 1. Principal Quantum Number, symbol n n may have any positive integral value > 0, e.g. 1, 2, 3, 4... As n increases, the energy of the electron and its distance from the nucleus increase. Within each principal level there are sublevels denoted by 2. Orbital Quantum Number, symbol R R may have integral values from 0 to (n ! 1); thus if n=2, R can only have values 0 and 1. If n=1, only R=0 is allowed. The general shape of the electron cloud is determined by R. We have ‘names’ for electrons with different values of R :

R = 0, R = 2,

‘s electrons’ ‘d electrons’

R = 1, R = 3,

‘p electrons’ ‘f electrons’

3. Magnetic Quantum Number, symbol m R m R may have integral values from !R through zero to + R. thus if R = 1, a p electron, m R can be -1, 0, or +1. The magnetic quantum number sets the number of suborbitals for each value of R. 4. Spin Quantum Number, symbol m s Regardless of the values of the other quantum numbers, m s may be only +½ or -½ . when two electrons come together, their spins may be either the same (both + or both !) or different (one + and one !). We


refer to these cases as, respectively, parallel spins and opposite spins, and they are usually represented by arrows pointing up or down

€ € (parallel, not allowed)

€  (opposite, allowed).

Ground State Electron Configuration - the lowest energy state for electrons in an atom can be determined by placing electrons in pairs into atomic orbitals, filling orbitals in succession from lowest energy to highest, ie,

But, the 4s and the 3d are very close in energy

- this approach works for the first several elements, such as 17Cl, whose configuration is . It gives correct electron configurations for all the atoms normally encountered in organic chem (see B&P, Table 1.3) - it is easy therefore to write out the valence electron structure, called the Lewis structure, for the first 18 atoms of the periodic table. Their valence shell electron configuration (Brown, Table 1.4) can be shown as


- NB: for elements in groups IIA, IIIA and IVA, the two paired electrons become unpaired before forming bonds so the valence of these elements becomes equal to the total number of valence electrons, eg C=4, Mg=2,...
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