Organic Chemistry: Structure and Bonding Overview

Atomic Number

Atomic Number

Number of protons in the nucleus and electrons surrounding

Mass Number

Number of protons plus neutrons in the nucleus

Cation

Positively charged ion with fewer electrons than protons

Anion

Negatively charged ion with more electrons than protons

Isotopes

Atoms of the same element with different numbers of neutrons

Periodic Table

Arrangement of elements by increasing atomic number

Orbitals

Regions of space with high electron density

S Orbital

Orbital with a sphere shape and lower energy

P Orbital

Orbital with a dumbbell shape and higher energy

Valence Electrons

Outermost electrons, more loosely held

Bonding

Joining of atoms to attain stable electron configurations

Octet Rule

Elements tend to gain, lose, or share electrons to have 8 valence electrons

Compounds

Combining two or more elements, can have ionic or covalent bonds

Molecule

Structure with only covalent bonds

Lewis Structures

Diagrams showing bonding between atoms using dots

Resonance

Delocalization of electrons in molecules

Determining Molecular Shape

Process of identifying the 3D structure of a molecule

Degrees of Unsaturation

Number of rings and/or pi bonds in a molecule

Nucleus

Central part of an atom containing protons and neutrons

Electron Cloud

Area around the nucleus with negatively charged electrons

Ions

Charged atoms due to loss or gain of electrons

Isotopes of Hydrogen

Variants of hydrogen with different numbers of neutrons

Period

Horizontal row in the periodic table

Group

Vertical column in the periodic table

Orbital Types

s, p, d, and f orbitals with specific shapes and energies

Covalent bonds

Results from the sharing of electrons between two nuclei

Ionic bond

Occurs when elements of the far left side combine with elements in the far right side of the periodic table, forming ions held together by strong electrostatic interactions

Covalent bonding

Occurs with elements like carbon in the middle of the periodic table, forming a two-electron bond and molecules

Hydrogen

Forms one covalent bond using its one valence electron

Lewis structures

Electron dot representations for molecules, drawn following specific rules

Formal charge

Charge assigned to individual atoms in a Lewis structure based on the number of bonds and lone pairs

Isomers

Different molecules with the same molecular formula but different connectivity of atoms

Resonance structures

Multiple Lewis structures representing a molecule due to electron delocalization, with a resonance hybrid showing characteristics of all structures

Resonance theory

Explains that resonance structures are not real, not in equilibrium with each other, and not isomers

Resonance hybrid

Composite of all possible resonance structures, with a major contributor being the 'better' structure

Bond length

The average distance between the centers of two bonded nuclei; decreases across a row of the periodic table as the size of the atom decreases

Bond angle

Determines the shape around any atom bonded to two other atoms; increases down a column of the periodic table as the size of an atom increases

VSEPR

Valence shell electron pair repulsion theory based on the repulsion between electron pairs, arranging groups around an atom as far apart as possible

Linear

Geometry with two groups around an atom, having a bond angle of 180 degrees

Trigonal planar

Geometry with three groups around an atom, having bond angles of 120 degrees

Tetrahedral

Geometry with four groups around an atom, having bond angles of approximately 109.5 degrees

Condensed structures

Representation used for compounds with a chain of atoms bonded together, omitting two-electron bond lines and lone pairs

Skeletal structures

Representation used for organic compounds with rings and chains of atoms, assuming tetravalent carbons and drawing in heteroatoms directly bonded

Heteroatom

Atom other than carbon or hydrogen, directly bonded to the carbon atom in skeletal structures

Charged atoms

A charge on a carbon atom replaces one hydrogen atom, determining the number of lone pairs

Negatively charged carbon atoms

Carbon atoms with one lone pair; positively charged carbon atoms have none

Sigma bond

Cylindrically symmetrical bond formed by two hydrogen atoms sharing electrons

Ground state

The lowest electron arrangement for an atom

Hybridization

Combining atomic orbitals to form an equal number of hybrid orbitals with the same shape and energy

Sp3 hybridization

Hybridization of an atom surrounded by four groups

Sp3 hybrid orbitals

Formed by one 2s orbital and three 2p orbitals

Sp2 hybrid orbitals

Formed by one 2s orbital and two 2p orbitals

Sp hybrid orbitals

Formed by one 2s orbital and one 2p orbital

Electronegativity

Measure of an atom's attraction for electrons in a bond

Polar bond

Bond with an electronegativity difference between two atoms ≥ 0.5 unit

Bronsted-Lowry acid

Proton donor in a reaction

Bronsted-Lowry base

Proton acceptor in a reaction

Acid strength

Tendency of an acid to donate a proton

Acidity constant

Measure of acidity in an equilibrium reaction

Element effects

Factors affecting acidity based on electronegativity

Inductive effects

Pull of electron density through sigma bonds due to electronegativity differences

Resonance effects

Acidity factor based on resonance stabilization

Hybridization effects

Stability of conjugate base based on percent s-character of hybrid orbital

Lewis acid

Electron pair acceptor in a Lewis acid-base reaction

Lewis base

Electron pair donor in a Lewis acid-base reaction

Nucleophile

Electron-rich species in a Lewis acid-base reaction

Electrophile

Electron-loving species in a Lewis acid-base reaction

Ethane

A compound where each carbon atom is tetrahedral, sp3 hybridized, and all bonds are sigma σ.

Bond Length and Bond Strength

Shorter and stronger bonds result from an increased number of electrons between nuclei, with triple bonds being shorter and stronger than double bonds, which are shorter and stronger than single bonds.

Bronsted-Lowry Acids and Bases

Acids donate protons, while bases accept protons in proton transfer reactions, with acids forming conjugate bases and bases forming conjugate acids.

Lewis Acids and Bases

Lewis acids accept electron pairs, while Lewis bases donate electron pairs, with Lewis acid-base reactions involving the formation of new covalent bonds.

Functional group

An atom or group of atoms with characteristic properties that is the reactive part of a molecule, often containing heteroatoms, lone pairs, and creating electron-deficient sites on carbons.

Hydrocarbon

Compounds composed of only carbon and hydrogen elements, classified as aliphatic or aromatic, with alkanes, alkenes, and alkynes as subgroups of aliphatic hydrocarbons.