Organic Chemistry

A Review of General Chemistry

1.2 Structural Theory of Matter

Organic Compounds (derived from living organisms)

Inorganic Compounds (derived from nonliving sources)

Compounds can have the same chemical formula, yet they differ in their interactions

Tetravalent - four bonds

Trivalent - three bonds

Divalent - two bonds

Monovalent - one bonds

Stronger intermolecular forces means higher boiling point

How to draw Constitutional Isomers

  • Determine Valency: Atoms with highest Valency are connected first

    • Carbon : Tetravalent

    • Hydrogen: Monovalent

    • Oxygen: Divalent

  • Connect atoms with highest valency, place monovalent atoms at the periphery (at the end)

1.3 Electrons, Bonds, and Lewis Structures

Covalent bond - two atoms sharing a pair of electrons

  • the force of repulsion between the negatively charged electrons

  • the force of repulsion between the positively charged electrons

  • the force of attraction between two oppositely charged electrons

Electrons are capable of moving in a way that minimize repulsive forces and maximize attractive forces

Hydrogen atoms have an intermolecular distance of 0.74 angstroms

Atomic Structure:

  • Every atom is composed of protons and neutrons.

  • For a neutral atom, the number of protons is equal to the number of elections

  • First shell - 2 electrons, Second shell - 8 electrons

  • Outermost shell - Valence electrons

Drawing Lewis Structures

Octet Rule: Second-Row elements will form the necessary number of bonds in order to gain 7

4 Rules

  • Draw all individual atoms

  • Connect atoms that form more than one bond

  • Connect the hydrogens

  • Pair unpaired electrons so atom achieves an octet

1.4 Identifying Formal Charges

Formal Charge - Correct number of valence electrons

Extra electrons (negative formal charge)

Too few electrons (positive formal charge)

If you are having trouble paying attention during a long lecture, your levels of acetylcholine (a neurotransmitter) may be to blame.

1.5 Polar Covalent Bonds

Electronegativity - measure of an atom to attract electrons

If electronegativity is below 0.5, electrons are considered to be equally shared

Examples: C-C and C-H

Covalent

If electronegativity is between 0.5 and 1.7, electrons are considered to be polar.

Examples: O-H and Li-C

If the difference in electronegativity is higher than 1.7, electrons are considered to be ionic pair

Examples: Na-Cl

C-BR and C-I are considered polar covalent

Resulting Partial Charges from Induction

  • Identify all covalent bonds

  • Determine direction of each dipole

  • Indicate the direction of partial charges

Bond Strength increases with more bonds btw

charges are on the element and not on the bond.

the donating atom has a positive charges, and the atom that takes the charge has a negatively charges

Electrostatic Potential Maps

Reading Bond-Line Structures

Bond-line structures - Each corner or endpoint

Triple bonds have linear geometry

Hydrogen atoms are not shown because it is assumed that each carbon atom will have enough

1.7 Atomic Orbitals

Electrons were discovered and identified as the source of bonding and Lewis structures were the generally accepted method of mapping electron interactions.

1924: Louis de Broglie - suggested that electrons had the same properties as waves.

Wave Equation - Energy of an atom is “"quantized”, meaning that all electrons are in orbitals (s,p,d,f)

Closest to the nucleus - less energy - less nodes

Farthest from the nucleus - most energy - most nodes

3 rules

  • Aufau Principle: Lowest Energy Orbital is filled first

  • Pauli Exclusion Principle: Each orbital can accomodate a maximum of 2 orbitals with opposite spin

  • Hund’s Rule: One electron is placed BEFORE electrons are paired

Positive charges have empty orbitals, negative charges have positive orbitals

Nodes are the intersections between positively and negatively spun orbital

Valence bond theory - sharing electron density

Sigma bonds - Single bonds

Pi bonds - Not sigma bonds

Molecular Orbital theory - implies that orbitals are combined to create new orbitals

An atomic orbital is associated with a singular atom, while a molecular atom is associated with a singular molecule.

Molecule = multiple atoms

Bonding MO: Constructive interference - lower energy

Antibonding MO: destructive interference - higher energy

Smallest Polar molecule - best soluability