Organic Chemistry Vocabulary Chapter (1, 2, 3, 4, 5, 6, 7, 8 ,9)

Organic chemistry

Organic chemistry

the study of carbon compounds

Organic compounds

are made from mostly C and H atoms may contain other nonmetals (O, S, N, P, or halogens)

Carbon atoms are the basis for life because they:

atom

a dense nucleus that holds protons and neutrons electrons orbit in a large, empty space around the nucleus

Atomic number

the number of protons present in EVERY atom of that element

Quantum Mechanics

The behavior of an electron is described by a wave equation

orbital

The solution to the wave equation is called a wave function

Principal Quantum Numbers

Assigned energy levels for electrons (aka shells). Lower energy levels are closer to the nucleus. Energy levels increase in energy as the value of n increases

Subtypes of Orbitals

s, p, d, f

s orbitals

spherical, nucleus at center.

p orbitals

There are three different p orbitals, depending on their orientation within the atom. dumbbell-shaped, with the nucleus at middle.

d Orbitals

Five d different d orbitals 4 orbitals are Four-leaf clover shaped: lobe orientations vary. 1 orbital has two lobes with a donut shaped region called a torus in the x-y plane.

Pauli Exclusion Principle

if two electrons share the same orbital, they must have opposite spins

Aufbau Principle:

electrons will always go into the lowest-energy orbitals available (maximum of 2 e- per orbital).

Hund’s Rule

Orbitals of the same type have the same energy (e.g., all 2p are equal) The lowest-energy configuration maximizes the number of unpaired electrons.

Valence electrons

electrons are in the outer most energy shell (furthest from the nucleus) most unstable involved in chemical bonding usually electrons is the s & p orbitals of the highest energy shell (n)

Group Number

gives the number of valence electrons for the representative elements Exception: helium

Electron-dot symbol or Lewis symbols

represent the valence electrons as dots placed on the top, bottom, and sides of a chemical symbol

Lewis Theory of Chemical Bonding

Stable Configurations: 8 electrons in the outermost valence shell (most Noble Gases). The octet rule generally applies to all main-group elements except: Hydrogen (wants a duet, 2e-, in outermost shell) Lithium Beryllium Boron

Chemical bonds

form when atoms lose, gain, or share valence electrons to fulfill the octet rule making bonds releases energy, breaking bonds absorbs energy

Ionic bond

atoms gain/lose valence electrons (ionic compounds, salt crystals)

Covalent bond

nonmetal atoms share electrons to form molecules

Lone pair

valence electrons not used in bonding

sigma (σ) bond

head-on overlap between two orbitals and electrons are located between the nuclei of the bonding atoms

Bond length

ideal distance between nuclei that leads to maximum stability

Hybridization

combination of orbitals to form new ones

Tetrahedral geometry

C-H bond strength = 439 kJ/mole Bond Angle 109.5* Bond Length 109 pm

Double bonds

one sigma and one pi bond, 4 electrons shared

Triple bonds

one sigma and two pi bonds, 6 electrons shared

Pi bond

p orbitals have a sideways overlap above and below the internuclear axis A covalent bond in which electron density is greatest around—not along—the bonding axis

Double C-C Bonds

Intermediate bond length, medium bond strength Trigonal planar geometry

Triple Carbon bond

Shortest bond length, strongest bond Linear geometry

molecular formula

which give the total number of element

Bond-Line Formula

Structural Formulas shows each bond

Condensed Formula

Structural Formulas shows each C atom and its attached H atoms as a group

Skeletal formulas

show the carbon skeleton bonds are represented by straight lines C atoms are represented at each corner or vertex H atoms attached to C are omitted, but other heteroatoms are shown

Electronegativity

attraction for shared electrons in a bond

Pauling Electronegativity Scale

Higher the number = more Electronegative

Ionic bond

complete electron transfer between metal and nonmetal ions large eN difference (greater than 2.0)

Nonpolar covalent bond

equal or almost equal sharing of electrons with small eN difference (less then 0.5)

Polar covalent bond

unequal sharing of electrons between nonmetals with moderate eN difference (0.5-2)

Inductive effect

the shifting of e- in a sigma bond in response to the eN of nearby atoms

dipole moment

Happens in Polar Molecules vector summation of individual bond polarities and lone-pair contributions can result in uneven charge distribution throughout the molecule

High Symmetry Carbon Molecules

no lone pairs on central atom identical atoms bonded to central molecule possible: nonpolar molecules that have polar

Symmetrical molecules

individual bond dipoles pointing equally in opposite directions. local dipoles cancel each other out. nonpolar molecules

Noncovalent interactions

(intermolecular forces) Electrostatic interactions BETWEEN molecules due to unevenly distributed electrons​ DIFFERENT from of intramolecular forces which are WITHIN a molecule​ relatively weak compared to covalent bonds​

Dipole-dipole interactions

occur between polar covalent molecules The partial negative end is attracted to the partial positive end

Hydrogen bonding

is the strongest dipole-dipole interaction Large differences in electronegativity between H and O, N

Hydrogen bond

a weak electrostatic attraction between an electronegative atom (such as oxygen or nitrogen) and a hydrogen atom covalently linked to a second electronegative atom

Dispersion forces

are the weakest intermolecular force occur in both polar and nonpolar molecules Attractive force caused by temporary dipoles that develop when molecules bump into each other

Formal charges

electron bookkeeping keeps track of electrons on a molecule do not imply the presence of actual charges

Formal Charge Equation

Formal Charge = Number of Valence Electrons - Number of Bonding Electrons - Number of Nonbonding Electrons (aka lone pairs)

Resonance forms

structures that differ by the placement of their pi or nonbonding electrons not different chemical species, but different depictions of one chemical species

Resonance hybrid

the actual structure of a molecule with resonance forms. composite – single unchanging structure that averages the structures of all resonance forms

Resonance

explains how some electrons are distributed over more than two atoms (delocalized)

Conjugated double bonds

repeating pattern of a double bond followed by a single bond

Alkanes

hydrocarbons that contain only C—C and C—H bonds continuous chain of C atoms have names that end in ane

Alkyl halides

contain an alkyl group attached to a group 7A element (halogens

Aromatic compound

contains a ring of 6 C atoms, each bonded to 1 H atom with 3 alternating double bonds

Alcohols contain

contain – OH (hydroxyl) groups can form hydrogen bonds

Phenols

alcohols that contain a hydroxyl group directly attached to a benzene ring

Thiols

contain an —SH group

Alkenes

hydrocarbons that contain double bonds

Alkynes

hydrocarbons that contain triple bonds

Sulfides

contain a C–S–C group

Ether

(C-O-C) contain an —O— between two C groups that are alkyl or aromatic

Nitriles

contain a carbon atom triple bonded to a nitrogen

carbonyl group

consists of a C=O polar double bond

aldehyde

carbonyl group is attached to one C group and one H atom

ketone

carbonyl group is attached to two C groups

Carboxylic acids

contain a hydroxyl group —OH attached to the C in a carbonyl group

Esters

contain an O bonded to a carbonyl and an alkyl group

Amines

contain N attached to one or more alkyl or aromatic groups

Amides

contain a carbonyl directly attached to a N group

Acid chlorides

contain a carbonyl attached to a chlorine

Anhydrides

contain an O sandwiched between two carbonyls

Bronsted Acid

proton (H+) donor

Bronsted Base

proton (H+) acceptor must have a lone pair of electrons to bond to the proton.

Conjugate acid–base pairs

related by the loss and gain of H+

conjugate base

forms after the acid loses a proton

conjugate acid

forms after the base has gained a proton

Inorganic Bronsted-Lowry Acids

eN halogen/oxygen attached to a H+

Organic Bronsted-Lowry Acids

carboxylic acid functional group alcohol functional groups amine functional group

Strong Bronsted Acids

greater ability to donate their acidic proton completely ionize in water have a weak conjugate base

Strong Bronsted bases

greater ability to accept a proton completely protonated in water have a weak conjugate acid

Strong acids form

a stable conjugate base acid is more likely to lose its proton.

Weak acids form

an unstable conjugate base the acid is less likely to lose its proton.

5 Different Stability Factors to Consider

CARIO Charge, Atom, Resonance, Inductive Effects, Orbitals

CAIRO Charge

acidity increases with increasing positive charge

CAIRO Atom

periodic trends Acidity increases across a PERIOD (from left to right) as eN increases Acidity increases down a GROUP as atom size increases

CAIRO Resonance

delocalization of a negative charge on a conjugate base increases acidity

CAIRO Inductive Effects

electron withdrawing groups (EWG) increase acidity of nearby atoms Acidity increases with increasing electronegativity

CAIRO Orbitals

increasing s character of atom attached to acidic proton increases acidity

Arrhenius

Type of Acid/Base reactions must occur in water (aq) acids increase [H+] bases increase [OH-]

Bronsted-Lowry

Type of Acid/Base acid: H+ donor base: H+ acceptor

Lewis

Type of Acid/Base acid: electron pair acceptor base: electron pair donor

Lewis Acid

electron pair acceptor Must contain a low energy, vacant orbital to accept electrons. Examples: H+, metal cations (Li+, Mg2+), group 3A elements that have formed 3 bonds (BF3, AlCl3), transition metal complexes (TiCl4, FeCl3, ZnCl2).

Lewis Base

electron pair donor Examples: atoms with electron lone pairs (phosphates, sulfates, halide anions, hydroxides, amines, amides, alcohols, ethers, carboxylic acids, ester, amides, aldehydes, nitriles, thiols)

Naming Alkyl groups

partial structure when an H is removed from an alkane named by removing –ane ending and replacing with -yl

Naming Halogen substituents

get are treated like alkyl substituents, but their substituent name ends with an -o. bromo, chloro, iodo, fluoro