CHEM 12A Organic Chemistry

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Lewis base

An electron pair donor

Lewis acid

An electron pair acceptor

Lewis adduct

The product formed when a Lewis base donates an electron pair to a Lewis acid to make a new bond

Electrophile

A Lewis acid (an electron pair acceptor)

Nucleophile

A Lewis base (an electron pair donor)

Leveling effect

When a very strong acid or base is neutralized by the solvent before it reaches its target

Counterion

An unreactive ion added alongside an ionic species so the overall mixture stays neutral

Chemical equilibrium

The state where forward and reverse reaction rates are equal so there is no net change in concentrations

Equilibrium constant (Keq)

The ratio of product to reactant concentrations at equilibrium

Acid dissociation constant (Ka)

a number that measures how strongly an acid breaks apart (dissociates) in water

pKa

The negative log of Ka; lower values mean stronger acids

Brønsted-Lowry acid-base theory

The idea that an acid-base reaction is a transfer of a proton (H+)

Brønsted-Lowry acid

A species that donates a proton

Brønsted-Lowry base

A species that accepts a proton

Conjugate acid

The species formed when a base accepts a proton

Conjugate base

The species formed when an acid donates a proton

Reaction mechanism

The step-by-step description of how a reaction occurs

Localized electrons

Electrons confined to one or two atoms that cannot take part in resonance or delocalization

Delocalized electrons

Electrons spread over a larger area than two atoms

Vitalism

The disproven theory that substances from living organisms carry a "vital force" and are fundamentally different from other substances

Partially condensed structure

A structure that leaves out the explicit C-H bonds

Condensed structure

A structure that omits bonds entirely with groups written in sequence and branches in parentheses

Bond line structure (skeletal structure)

A drawing where carbons are vertices and hydrogens on carbon are not shown

Wedged and dashed bonds

Bonds that show atoms coming out of the plane (wedge) or going into the plane (dash)

Kinetics

The study of reaction rates

Rate equation

An equation showing how the rate of a reaction depends on reactant concentrations

Order

The power a concentration is raised to in the rate law (for one reactant or overall)

Rate constant

The proportionality constant k in the rate equation that depends on factors like activation energy and temperature

Activation energy (Ea)

The height of the energy barrier that must be crossed for a reaction to occur

Reaction free energy diagram

A plot of free energy versus reaction coordinate showing kinetic and thermodynamic information

Transition state

The highest energy point along the reaction coordinate

Exergonic

A reaction with negative ΔG where products are lower in free energy than reactants (spontaneous)

Endergonic

A reaction with positive ΔG where products are higher in free energy than reactants (nonspontaneous)

Enthalpy diagram

A diagram showing whether a reaction is exothermic or endothermic

Exothermic

A reaction that releases heat (negative ΔH)

Endothermic

A reaction that absorbs heat (positive ΔH)

Thermodynamics

The study of whether a reaction will happen in principle

Equilibrium

The state where there is no net change in reactant and product concentrations

State functions

Properties that depend only on the state of a system (like G or H or S)

Gibbs free energy

The quantity (ΔG) that determines whether a reaction at constant T and P is spontaneous

Spontaneous

A reaction that is thermodynamically favorable (negative ΔG)

Entropy (S)

A measure of how many ways molecules can occupy available states in a statistical distribution

Enthalpy (H)

A measure of heat content related to bond strengths that makes some states more likely than others

Molecular dipole moment

The vector sum of the individual bond dipoles in a molecule

Intermolecular forces

Attractive forces between molecules that determine properties like melting and boiling points

Dipole-dipole interactions

Attractions between the partially charged ends of polar molecules

Hydrogen bonds

Strong attractions that occur when hydrogen is attached to an electronegative element (N or O or F)

London dispersion forces

holds together non polar compounds

strength accumulates with size/surface area

more branched < more linear
in bond strength

VSEPR theory

Valence shell electron pair repulsion theory which predicts molecular geometry because electron pairs repel and spread apart

Steric number

The number of lone pairs plus sigma bonds on an atom

Tetrahedral

The geometry of four groups around an atom (sp3 hybridization)

Trigonal planar

The geometry of three groups around an atom (sp2 hybridization)

Valence bond (VB) theory

A model where a chemical bond is the constructive overlap of exactly two atomic orbitals

Sigma (σ) orbital

A bonding orbital with cylindrical symmetry about the bond axis

Pi (π) orbital

A bonding orbital with a nodal plane along the bond axis

Molecular orbital (MO) theory

A model where atomic orbitals from any number of atoms combine into molecular orbitals

Linear combination of atomic orbitals (LCAO)

Adding atomic orbitals together so they interfere constructively or destructively

Bonding MO

A lower energy molecular orbital formed from more constructive overlap

Antibonding MO

A higher energy molecular orbital formed from more destructive overlap

Highest occupied molecular orbital (HOMO)

The molecular orbital with the highest energy and most reactive electrons

Lowest unoccupied molecular orbital (LUMO)

The most stable empty orbital able to receive additional electrons

Molecular formula

A formula showing which atoms make up a single molecule of a compound

Covalent bonds

Bonds involving the sharing of electrons between atoms

Isomers

Different molecules with the same molecular formula

Lewis structure

A drawing that shows all valence electrons as dots and bonds as lines

Valence electrons

The outer shell electrons of an atom (equal to the group number for a neutral atom)

Octet

A full set of eight valence electrons that atoms form bonds to obtain

Formal charge

The charge from comparing an atom's valence electrons (bonds split in half) to the expected number for that element

Electronegativity

An atom's ability to pull electron density toward itself in a bond

Induction

The pulling of electrons away from the less electronegative atom in a bond

Polar bond

A bond where electrons are unequally shared because of an electronegativity difference

Bond dipole moment

A measure of bond polarity equal to the charge magnitude times the distance between the separated charges

Partial charges

The δ+ and δ- charges on the atoms of a polar bond

Electrostatic potential map

A visual map of the charge distribution in a molecule

Alcohol

Ether

Ketone

Aldehyde

Thiol

Sulfide

Carboxylic acid

Alkyl Halide

Amine

Ester

Amide

Equilibrium will favor the formation of the acid with the…

higher pKa (the weaker acid)

linear

trigonal planar

tetrahedral

trigonal bipyramidal

sigma bond

pi bond

Negative Charges are more satisfied/stable when…

in a lower sp orbital than the other because it is closer to the nucleus

Alkane

A compound containing only C & H atoms

Aklene

A compound that contains a C-C double bond

Alkyne

A compound that contains a triple bond

Phenyl group

a six membered ring w/ alternating pi bonds and signals an aromatic compound

Cylic

Acylic

Resonance

the delocalization (spreading out) of pi electrons and lone pairs across multiple atoms in a molecule