Organic chemistry - chp. 1 (The basics)

Isotopess

• Same atomic number (same electrons and protons) different atomic mass (different amount of neutrons

Electronegativity and periodic trend

• atoms ability to attract electrons

• Electronegativity generally increases as you go to the top right of the periodic table

Octet rule violations

• Elements past the third period can accommodate for more than 8 valence electrons

• Elements with less than 8 electrons

Formal charge calculation

Formal charge = number of valence electrons - number of bonds - unbonded electrons

constitutional (structural) isomers

molecules with same molecular formula but different bonding patterns

• Leads to different physical (bp, mp, density) and chemical properties (reactivity)

different models

• ball and stick model

• Electron dot formula

• Dashed formula

• condensed formula

• Bond-line formula (Most common in orgo)

Condensed formula

Partially condensed - All hydrogens following a carbon are written after them

Fully condensed - all elements attached to a carbon are written directly after the C

ex. C3H8O —→ CH3CH(OH)CH3

groupings of atoms attached to C are written in parenthesis and multiple of same groupings are denoted by subscript

Bond line structures

Be familiar with rules and how to draw them

cycloalkanes

non linear structures

• can be written as different connected shapes

Bond line structure of alkenes

C’s are not fully saturated (do not have max number of H’s bonded) due to double bond

Bond line structures of alkynes

Not all C’s are fully saturated by H’s due to presence of triple bond

Resonance structures

alternative Lewis structures for a single compound that differ in electron positioning

• unshared electrons and those in double and triple bonds are capable of being moved around a single atom

Wave functions

equation denoted by Greek letter psi used to calculate energies and probability of electron location

probability of finding electron within an orbital

Electrons have a 90-95%

Energies within bonding and repulsion of electrons

• Highest energy when nuclear repulsion occurs

• Lowest energy when bonds form (Orbitals overlap)

Constructive interference

two waves combine to form a wave of larger amplitude (in phase) creating a molecular orbital

(Orbitals in = orbitals out)

destructive interference

Two waves combine to cancel each other out (out of phase) creating a node where there is zero electron density between two orbitals (anti bonding)

Order of orbital formation

Electrons fill in from low energy to high energy

Electrons fill and form bonds at molecular orbitals first at low energy before anti bonding at muc higher energy

Electron configuration review

Electrons fill lowest energy orbitals first with opposite spins

Electrons fill degenerate orbitals (orbitals of same energy, p, d, etc) singly before pairing up