OChem Unit 1 Chad's Prep

Organic Chemistry

the study of carbon containing compounds

Organic compounds

carbon containing compounds

Inorganic compounds

generally defined as carbon-lacking compounds

Valency

the number of bonds a compound makes

Tetravalent

valency of 4

Trivalent

valency of 3

Divalent

Valency of 2

Monovalent

Valency of 1

Octet rule

atoms aim to have 8 ve- in their shell

Lone pair

a pair of unshared/non-bonding e-

Covalent bond

sharing of e-

Formal charge

associated with any atom that doesn’t xhibit the appropriate number of ve-

FC Equation

(Number of ve-) - (dots and lines)

+FC

Missing e-

-FC

Excess e-

Sigma (σ) bonds

Present in all bonds, end to end overlap along internuclear axis

Internuclear axis

imaginary line connecting the nuclei of two bonded atoms

Valence bond theory

a bond is simply the sharing of e- density bw 2 atoms as a result of constructive interference of their atomic orbitals

e- density

how likely an electron is to be found at a specific point in space

Pi (π) bonds

Can only occur bw p orbitals, occrus above and below the internuclear axis, and is any additional bond

sp Hybridization

2 e- groups and bond angles of 180

sp² Hybridization

3 e- groups and bond angles of 120

sp³

4 e- groups and bond angles of 109.5

ve-

e- involved in chemical bonding

Wave (Ψ) function

takes into account the wave-like behavior of an e- that is in the magnetic field of a p+

Ψ²

represents the probability of finding an e- in a particular location

Orbital

a region of space that can be occupied by an e-

Atomic orbital (AO)

a region of space filled with respect to the nucleus of a single atom

Molecular orbital theory

uses mathematics as a tool to explore the consequences of AO overlap

Bonding MO (π, σ)

the result of constructive interference of the original two AOs

Antibonding MO (π*, σ*)

the result of deconstructive interference (explains why it’s higher in E)

Nodes

locations where Ψ = 0, no possibility of finding an e-

2py orbital

2px orbital

2pz orbital

1s orbtial

2s orbital

Aufbau Principle

lowest E obrital filled first

Pauli Exclusion Principle

each orbital can acomodate a max of 2 e- hat have opposite spin

Hund’s Rule

1 e- is placed in each degenerate orbital first, before e- are paired up

Degenerate orbitals

orbitals with the same E level Ex: 2s and 2p)

Constructive interference

overlap with orbitals of the same phase, produced a wave with larger amplitude

Destructive interference

overlap with differencing pases, cancelling each other out, and always produced a node

Diatomic

molecule composed of 2 atoms

Heterodiatomic

molecule composed of 2 different atoms

Homodiatomic

molecule composed of 2 of the same element

LUMO

“Lowest unoccupied Mo”: MO with the highest E in the molecular orbital diagram that’s empty

HOMO

“Highest occupied MO”: MO with the highest E in the MO diagram with e-

Bond order (BO)

the difference between the number of bonding and antibonding electrons divided by two.

Fronteir MOs

LUMO and HOMO, usually the ones involved in chemical RXNs

Needs to donate e- → donate highest E ones in the LUMO

Needs to receive e- → receive in LUMO (most stable E spot)

Non-polar covalent

Difference of EN 0-0.5(small difference in EN)

Polar Covalent

Difference of EN 0.5-1.7

Ionic

Difference of EN 1.7< (larger difference in EN)

Dipole moment

measure of polarity

Molecular dipole moment

the vector sum of all bond dipoles in a molecule, indicating its overall polarity

δ Partial charge

representing partial charges in a polar bond

Dipoles that cancel

Linear in opposite directions, and symmectrical MGs (trigonal planar, tetrahedral, and octahedral)

C-C and C-H

Non-polar bonds

Dispersive forces

Temporary dipole moment, random movement of additive, random e- movement, and the more e- → the more polarizable the molecule

Dipole-dipole forces

Permanent dipole, polar

Hydrogen-bonding

H- O, N, F

Ion-dipole forces

IMF bw an ion and the partial charge of a molecule

Constitutional isomers

compounds that have the same molecular formula but difffere in the way the atoms are connected (usually differ in surface area)

Branching

Gives you a more compact structure

smaller SA → lower dispersive forces

Less branched

Higher BP, usually lower MP

More branched

Lower BP, usually higher MP

Solubility

“like dissolves like” in terms of polarity.

Ex: A molecule with the least amount of non-polar bonds →greatest solubility with a nonpolar molecule

A molecule with the greatest amount of polar bonds → greatest polarity and would have the greatest solubility in a polar molecule