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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
Note 
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