General
Chemistry Organic
Chemistry Physical
Chemistry
Allotropes
- different forms of element in same physical state, eg O2 and O3
Law
of Definite Proportions - compound always has ratio of elements same by mass
Law
of Multiple Proportions - ratio of masses of elements in compound is small
whole number ratio
Stoichiometry
- quantitative relationships, composition or reaction
Chemical
Equations and Reaction Stoichiometry
Law
of Conservation of Matter - matter is not created or destroyed, only
rearranged
Limiting
reactant - the reactant that is used up completely in the reaction
Solution
- solute dissolved in solvent
Titration
- titrant reactant slowly added to solution of another reactant and measure
amount for complete
reaction; plot
curve of added volume vs. pH; at equivalence point equal amounts of acid and
base
reacted, should coincide with end point, when
indicator color changes; use buret
Periodic
Law - properties of elements are periodic functions of atomic number
Metals
- high conductivity (inc. with inc. temp.), high thermal conductivity, solid
except mercury (Ce and
Ga melt), malleable, gray except Ag and Au, few
electrons in outer shell, metallic character inc.
down and left on PT
Electrolytes
- substances whose aqueous solutions conduct electricity well, incl. strong
acids, strong
soluble bases, most soluble salts
Precipitates
- settle out of solution
Oxidation
number - number of electrons gained or lost by atom in binary compound
Oxidation
- loss of electrons
Reduction
- gain of electrons
Oxoacids
- ternary acids
Photoelectric
effect - electromagnetic radiation causes electron emission from metal surface
Heisenberg
Uncertainty Principle - can't know both momentum and position of small
particle
Aufbau
Principle - electrons added into orbitals in way giving lowest total energy
Pauli
Exclusion Principle - no two electrons in atom have same 4 quantum numbers
Hund's
Rule - electrons mus toccupy all orbitals of a sublevel before pairing
Paramagnetic
- unpaired electrons weakly attracted into magnetic fields
Diamagnetic
- all electrons paired and are very weakly repelled by magnetic fields
Ferromagnetic
- Fe, Co, and Ni permanently magnetized as spins align with field
Screening
causes effective nuclear charge to be less than actual nuclear charge
Combustion
reaction - oxygen combines rapidly, very exothermic, hydrocarbon+oxygen yields
carbon
dioxide water and heat
Roasting
- extracting free metals by heating an ore in air (oxygen)
Ionic
compounds - high melting pt., soluble in polar solvents, insoluble in
nonpolars, molten and aqueous
solutions conduct electricity; large
electronegativity difference between atoms
Lewis
dot formulas - show valence electrons
Octet
Rule - most compounds achieve noble gas configurations
Resonance
- two or more Lewis structures describe bonding
Formal
charge - charge on atom in a molecule or polyatomic ion
Polar
covalent bond - electrons shared unequally; creates dipole
Sigma
bond - head on overlap; all single bonds are sigma
Pi
bond - side on overlap; may include unhybridized p orbital
Molecular
orbital - an orbital resulting from overlap and mixing of atomic orbitals on
different atoms;
belongs to molecule as whole
Antibonding
orbital - molecular orbital higher in energy than any of atomic orbitals from
which it is
derived; lends stability when populated; marked with
asterick
Nonbonding
orbital - orbital derived only from an atomic orbital of one atom; lends no
stability
Delocalization
- formation of set of molecular orbits that extend over more than two atoms
Nodal
plane - region of zero probability of finding electrons
Protonic
acids - acids with acidic hydrogen atoms
Arrhenius
theory - acid produces H+ in aqueous solution; base produces OH-
in solution
Bronsted-Lowry
theory - acid is proton donor; base is proton acceptor
Lewis
theory - acid accepts a share in electron pair, base donates a share in
electron pair
Conjugate
acid-base pairs - differ by proton; weak acid yields strong conjugate base and
vice versa
Amphoterism
- ability to react as either acid or base
Coordinate
covalent bond - both electrons furnished by one atom
Standardization
- process to determine concentration by measuring volume required to react
with known
amount of primary standard
Equivalent
weight of an acid - mass needed to furnish 6.022*1023 hydrogen ions
Half-reaction
- either reduction or oxidation part of redox reaction
Fluids
- liquids and gases; flow freely
Vapor
- gas formed by evaporation or sublimation
Pressure
- force per unit area; measured by barometer (1 torr = 1 mm Hg), manometer
U-shaped tube
Dumas
method - used to find molecular weights of volatile liquids using boiling
water bath
Kinetic-molecular
theory - by Rudolf Clausius; collisions are elastic, molecules travel in
straight line with
constant velocity until collide; gases consist of
discrete molecules
Effusion
- escape of gas through tiny hole
Diffusion
- movement of gas into a space or mixing with another gas
London
forces - weak attractive forces in molecules; vary as 1/d7; only
intermolecular forces among
symmetric nonpolars
Dipole-dipole
interactions - attraction of opposite partial charges; vary as 1/d4
Hydrogen
bonding - H to F, O, or N; like dipoles
Viscosity
- resistance to flow of a liquid; can measure with Ostwald viscometer
Surface
tension - inward force overcome to expand surface are of liquid
Meniscus
- surface of liquid
Cohesive
forces - hold liquid together; adhesive forces hold liquid to another surface
Evaporation
- opposite of condensation; molar heat of vaporization and heat of
condensation
Vapor
pressure - partial pressure of vapor molecules above liquid surface; easily
vaporized are volatile
Boiling
point - vapor pressure = external pressure
Melting
- opposite of freezing; molar heat of fusion and heat of solidification
Sublimation
- opposite of deposition
Phase
diagrams - temperature vs. pressure; triple point all 3 states at equil. (4.6
torr, 0.01 C for water);
can't liquefy gas above critical point
Amorphous
solids - no well-defined structure (like rubber, some plastics)
Crystals
- unit cells repeat and can be replaced with lattice point; 7 systems incl.
Cubic, tetragonal,
orthorhombic, monoclinic, triclinic, hexagonal,
rhombohedral
Isomorphous
- substances that crystallize in same type of lattice
Polymorphous
- substance that crystallizes in multiple forms
Coordination
number - number of neighbors in solid packing
Metallic
bonding - band theory describes continuous bands of closely spaced molecular
orbitals
Conduction
band - a band electrons must move into to allow conduction; insulators have
band gap;
semiconductors have filled bands that are slightly
below empty bands
Solvation
- process of solvent molecules surrounding solute ions or molecules; called
hydration if water
Miscibility
- ability of a liquid to dissolve in another; add acid to water
Saturated
- solid and dissolved ions in equilibrium
Supersaturated
- high solute prepared at high temperature then cooled
Colligative
properties - physical properties depending on number not kind of solute
particles
Fractional
distillation - separate liquid
mixture by boiling points
Boiling
point diagram - mole fraction vs. temperature; bowed curves for vapor and
liquid; intercepts show
boiling points
Colloids
- dispersed phase (solutes) suspended in dispersing medium (solvent); solid in
solid solid sol,
liquid in solid solid emulsion, gas in solid solid
foam, solid in liquid sols and gels, liquid in liquid emulsion,
gas in liquid foam, solid in gas solid aerosol,
liquid in gas liquid aerosol
Tyndall
effect - scattering of light by collodial particles
Micelles
- cluster of molecules with hydrophobic tails in center and hydrophilic heads
outward
Surfactant
- has ability to suspend and wash away oil and grease
Hard
water - contains Fe3+, Ca2+, and/or Mg2+ ions
Emulsifiers
- coat particles of dispersed phase to prevent coagulation into separate phase
Synthetic
detergents - soap-like emulsifiers with sulfonate or sulfate instead of
carboxylate
Eutrophication
- overgrowth of vegetation because of high phosphorous concentration
State
function - value depends only on current state not how it got there
Calorimetry
- measuring heat transfer between system and surroundings using calorimeter;
coffee-cup and
bomb caliometers (constant volume)
Enthalpy
- heat content
Standard
molar enthalpy of formation - enthalpy change for reaction in which one mole
is formed from its
elements at their standard states
Bond
energy - energy needed to break one mole of bonds
Transition
state theory - activation energy to form transition state
Mechanism
- step by step reactions; rate determined by slowest, rate-determining step
Heterogeneous
catalysts - speed up reaction but are in different phase than reactants, such
as powdered
noble metals and metal oxides in catalytic converters
Enzymes
- biological catalysts; bind substrates
Chemical
equilibrium - two opposing reactions occur simultaneously at same rate;
dynamic equilibrium
LeChatelier's
Principle - system responds to stress at equilibrium in a way that reduces
stress and reaches
new state of equilibrium
Haber
process - N2 + 3H2 <-> 2NH3
Common
ion effect - behavior of solution in which same ion is produced by two
different compounds
Buffers
- minimize changes in pH because basic component can react with H3O+
ions and acidic
component can react with OH- ions
Polyprotic
acids - furnish two or more hydronium ions per mole
Solvolysis
- reaction of substance with the solvent in which it is dissolved; hyrolysis
if water
Solubility
product constant Ksp - equilibrium constant for reactions involving
slightly soluble compounds
Solubility
Product Principle - like equilibrium expression, but can take solids to be one
Fractional
precipitation - remove some ions from solutions while leaving others in
Molar
solubility - number of moles of solute that dissolve to produce liter of
saturated solution
Electrolytic
cells - external electricity causes nonspontaneous reactions by electrolysis
Voltaic
cells (galvanic cells) - spontaneous chemical reactions produce electricity
Electrodes
- surfaces upon which oxidation (anode) or reduction (cathode) half reaction
occurs
Downs
Cell - electrolysis of molten sodium chloride
Faraday's
Law of Electrolysis - amount that oxidizes or reduces at each electrode is
directly prop. to
amount of electricity that passes through cell
Faraday
- amount of electricity that reduces one equivalent weight at cathode and
reduces at anode
Electroplating
- using using electrolysis to plate metal onto surface
Salt
bridge - circuit between two solutions in a voltaic cell
Standard
cell - all species are in thermodynamic standard states (1 M , 1 atm)
Standard
Hydrogen Electrode (SHE) - reference electrode relative to which electric
potentials are measured
as reduction at 25 C; if Eo > 0
reduction occurs more readily than 2H+ to H2
Corrosion
- redox process by which metals are oxidized by oxygen in presence of
moisture; prevent by
plating or galvanizing (coating steel with zinc)
Primary
voltaic cells - cannot be recharged; includes Georges Leclanche's dry cell
(ZN(NH4)3) and
alkaline dry cells
Secondary
voltaic cells - reversible; can be recharged, such as lead storage battery in
cars (PbSO4), nickel-
cadmium (nicad) cells, and hydrogen-oxygen fuel cells
Native
ores - uncombined free state of less active metals, like Cu, Ag, Au
Ores
- contain minerals mixed with gangue (sand, rock, etc)
Metal
separation includes flotation, roasting (heating with oxygen), reaction with
coke (carbon) or CO, and
electolysis of molten salt
Hall-Heroult
process - cell for electolyzing Al
Iron
- blast furnace with CO converts to limestone flux, which reacts with silica
gangue to form slag of
calcium silicate; iron from blast furnace contains
carbon (pig iron); remelted and cooled to cast iron; add
other metals like Mn, Cr, Ni, W, Mo, V to make steel
Coordination
compounds - compouns with bonds in which both shared electrons are donated by
same atom
Ligand
- a Lewis base in a coordination compound
Polydentate
- ligands with multiple donor atoms
Chelate
- a ligand that utilizes two or more donor atoms in bonding to metals
Nuclear
fission - splitting of heavy nucleus into lighter nuclei
Nuclear
fusion - combination of light nuclei to produce heavier nucleus
Mass
deficiency - difference between sum of masses of electrons/proton/neutrons and
actual mass
Scintillation
counter - detects radiation using fluorescence
Cloud
chamber - detects radiation using water vapor; developed by Wilson
Gas
Ionization chamber - such as Geiger-Muller counter
Disintegration
series - sequence of atoms during decay
Radiocarbon
dating - C14, K-Ar, U-Pb methods
Radioactive
tracers - Na24 blood, Th201 and Tc99 heart, I131
thyroid liver and brain, Pl238 pacemakers
Cyclotrons
- devise for accelerating charged particles along spiral path
Linear
accelerators - device used for accelerating charged particles along straight
line path
Uranium-235
decay - to Uranium-236 to Sm/Zn, La/Br, Ba/Kr, Cs/Rb, Xe/Sr
Fission
reactors - use U3O8 fuel rods enriched in uranium-235,
water and graphite moderators (and He and
heavy water), B/Li control rods, cooling systems,
concrete shielding
Thermonuclear
bombs (fusion bombs, hydrogen bombs) - activation energy of fusion obtained by
fission
Plasma
- state of matter at high temperatures at which all molecules are dissociated
and most ionized
D=m/V
Sp.
Gr. = D/Dwater
Sp.
Heat = (heat in J)/((mass in g)*(temp. change in C))
Molarity
= moles/Liter
V1M1
= V2M2
v
= fl
E
= hv
Rydberg
equation: 1/l
= R(1/n12-1/n22) relating H
spectrum wavelengths
De
Broglie equation: l
= h/(m*f) showing small particles can display wave properties
Schrodinger's
equation: in terms of electron wave function y,
solutions are possible energy states for
electron in atom; Dirac incorporated relativity
Number
of atomic orbits = (energy level n)2
Formal
charge = (group number) - (number of bonds) - (number of unshared electrons)
Dipole
moment = (distance)*(magnitude of charge)
Bond order = (bonding electrons - antibonding electrons)/2
Normality
= (number of equivalent weights of solute)/(L of solution)
Boyle's
Law : P1V1 = P2V2 ; volume
inversely prop. to pressure
Charles'
Law: V1/T1 = V2/T2 ; volume
directly prop. to temperature
Combined
gas law : P1V1/T1 = P2V2/T2
Avogadro's
Law: V1/n1 = V2/n2 ; volume
directly prop. to number of moles of gas
Ideal
Gas Law: PV = nRT
Dalton's
Law of Partial Pressures: Ptotal = PA + PB +
PC + ... ; partial
pressure of each gas is its mole
fraction times total pressure of mixture
Average
molecular kinetic energy is directly prop. to absolute temperature
Van
der Waals equation: (P+n2a/V2)(V-nb) = nRT ; extends
ideal gas law to real gases using two empiricals
Coulomb's
Law: F=kq1q2/d2
Clausius-Clapeyron
equation: relates temperature to vapor pressure and molar heat of vaporization
Bragg
equation: nl
= 2*d*sin(q), relates reflections for X-rays to
wavelength and distance
Henry's
Law: Pgas = kCgas ; pressure of gas above solution is
prop. to concentration of gas in solution
Molality
= (number of moles of solute)/(number of kilograms of solvent)
Raoult's
Law: Psolvent = Xsolvent/P0solvent
; vapor pressure of solvent is directly prop. to mole fraction of
solute
Boiling
point elevation: DTb
= Kbm ; boiling point directly prop. to molality of solute
Freezing
point depression: DTf
= Kfm; freezing point depression directly prop. to molality of
solute
Osmotic
pressure p
= MRT
KE
= mv2/2
Hess'
Law: DHrxn0 = DHa + DHb + ... ; enthalpy change is same
as series of steps as if one reaction
DHrxn0
= S(bond energies of reactants) - S(bond energies of products)
DH
= DE + PDV
DE
= q + w = q - PDV; difference in internal energy = heat and
work
Gibbs
free energy: DG
= DH - TDS
Rate-law
expression: xA + yB -> C + D
rate = k[A]x[B]y
Arrhenius
equation: k = Ae-Ea/RT ; relates rate constant to activation
energy, temperature, and collision freq.
Chemical
equilibrium: aA + bB -> cC + dD
Keq = ([C]c[D]d)/ ([A]a[B]b)
; reaction quotient Q is same
form for a specific time; can also use partial
pressures rather than concentrations
KP
= KC(RT)Dn
DG0
= -RTln(K)
van't
Hoff equation: ln(KT2/KT1) = DH0/R (1/T1 - 1/T2)
; estimate equilibrium constant at another
temperature
Kw
= [H3O+][OH-] = 10-14
pH
= -log([H3O+])
pKa
= -log(Ka) ; large Ka
-> small pKa -> strong acid
Henderson-Hasselbalch
equation: pH = pKa + log([conj. base]/[acid])
Nernst
equation: E = E0 - (2.303*R*T)/(n*F)*log(Q)
= E0 -
(0.0592*T)/n*log([Red]y/[Ox]x)
; calculates
electrode potentials for concentrations and partial
pressures other than standard values
nFE0
= 2.303*R*T*log(K)
DG
= -nFEcell
nuclear
binding energy = (mass deficiency)*(speed of light)2
Half-life
decay: t1/2 = ln(2)/k
Reaction
orders:
Zero
rate=k
[A] = [A]0 - akt
t1/2 = [A]0/(2*a*k)
First
rate=k[A]
ln([A]0/[A]) = akt
t1/2 = ln(2)/ak
Second rate=k[A]2
1/[A] - 1/[A]0 = akt
t1/2 = 1/(ak[A]0)
Mole
= 6.022*1023 particles
Electron
= 1.75882*108 C/g, 9.109*10-28 g
1
g = 6.022*1023 amu
Planck's
constant h = 6.6262*10-34 Js
Rydberg's
constant 1.097*107 m-1
Standard
molar volume of ideal gas at STP: 22.414 liters per mole
Universal
gas constant R = 0.08206 (L*atm)/(mol*K)
Heat
of vaporization of water = 2.26 kJ/g
Specific
heat of water = 4.18 J/(g*C)
Heat
of fusion of water = 334 J/g
1
faraday = 96485 Coulombs
Joule
= kg*m2/s2
Plus
one: Na, K, NH4 ammonium, Ag, Cu+ cuprous
Plus
two: Fe2+ ferrous, Cu2+ cupric, Zn, Mg, Ca, Hg mercuric,
Hg2 mercurous
Plus
three: Fe3+ ferric, Al
Minus
one: CH3COO acetate, F, Cl, Br, OH, NO2 nitrite NO3
nitrate, CN cyanide, ClO hypochlorite, ClO2
chlorite, ClO3 chlorate, ClO4
perchlorate
Minus
two: SO3 sulfite, SO4 sulfate, CO3 carbonate,
CrO4 chromate, Cr2O7 dichromate
Minus
three: PO4 phosphate, AsO4 arsenate
Acids:
HNO3 nitric, HclO4 perchloric, HClO3 chloric,
H2SO4 sulfuric, H3PO4 phosphoric,
H3PO2
hypophosphorous
Ternary
acids names: perXic (perXate), Xic (Xate), Xous (Xite), hypoXous (hypoXite)
Strengths
(inc.) : NH3, H2O, NH4, HCN, CH3COOH,
HF, HNO3, HCl, HBr, HI, HclO4
primary
n (main energy level, 1,2,3...), subsidary or azimuthal l (shape of region,
0..n-1 = s,p,d,f,etc), magnetic ml
(spatial
orientation -l..l orbitals), spin ms (1/2 or -1/2)
Oxidation
numbers
+1/-1:
H
+1:
Li, Na, K
+2:
Be, Mg, Ca, Cu, Zn
+3:
B, Al, Ga, Se
+4:
C, Si, Ge, Ti
+5/-3:
N
+5:
P
+6/-2:
S, Se
-2:
O
-1:
F, Cl, Br
None:
He, Ne, Ar
Oxides:
O2- oxides, O22- peroxides, O2-
superoxides
Methyl
red: <4 red, >7 yellow; Bromthymol blue: <6 yellow, >8 blue;
Phenolphthalein: <8 colorless, >10 red
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p68s2
diagonals
H2
1, He2 0, B2 1, N2 3, O2 2
Soluble
- common inorganic and low molecular weight organic acids, compounds of Group
IA metals, nitrates,
acetates,
chlorates, perchlorates; Insoluble - most hydroxides, carbonates, phosphates,
arsenates, sulfides
First
- total energy in universe is constant
Second
- in spontaneous reactions universe tends towards state of greater disorder
(greater entropy)
Third
- entropy of pure, perfect crystalline substance is zero at 0 K
Periodic
trends
Inc.
up and right: ionization energy, electron affinity negativeness (easily
becomes anion), electronegativity
(Fr least, F most, none for nobles)
Inc.
down and left: atomic radii
Radioactive
decay
beta emission (electron ejected from nucleus as
neutron is converted to proton),
positron emission or electron K-capture (positron
ejected from nucleus as proton is converted to
neutron),
alpha emission (helium nucleus with 2 protons and 4
neutrons is ejected)
Valence
Shell Electron Pair Repulsion Theory (VSEPR)
Bonds+electron pairs = 2 (linear, sp, 180), 3 (trigonal planar, sp2,
120), 4 (tetrahedral, sp3, 109.5),
5
(trigonal bipyramidal, sp3d or dsp3, 90,120,180), 6
(octahedral, sp3d2 or d2sp3,
90,180)
hybrid - mixing of orbitals
Elements
in the Earth
O
49.5%, Si 25.7, Al 7.5, Fe 4.7, Ca 3.4, Na 2.6
Most
Commercially Used Acids
sulfuric,
lime (CaO and Ca(OH)2), ammonia, NaOH, phosphoric, nitric
H,
He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca
f-Transitions
Lanthanides
- 57, 58 cerium to 71 lutetium
Actinides
- 89, 90 thorium to 103 lawrencium
Alkali
Metals (IA)
Sodium
(Na) - yellow glowing highway lamps, needed for life, soda lye (NaOH), baking
soda (NaHCO3),
table salt (NaCl)
Lithium
(Li) - highest heat capacity, Li-Al aircrafts, dry cells, mental drugs,
nuclear reactor heat transfer
Potassium
(K) - needed for life, saltpepper KNO3 fertilizer
Others:
rubidium, cesium, francium
Alkaline
Earth Metals (IIA)
Calcium
(Ca) - reducing agent, remove impurites, cheap base slaked lime Ca(OH)2,
mortar, plaster of Paris
2CaSO4*H2O
Magnesium
(Mg) - burns white in air; photo flashs, fireworks, anti-oxidation coating,
plentiful in oceans
Beryllium
(Be) - X-ray window tubes
Strontium
(Sr) - red glow; fireworks and flares
Barium
(Ba) - spark plugs
Boron
(B)
Aluminum
(Al) - most abundant in earth's crust and third overall; buildings, electrical
transmission lines,
reducing agent including thermite reaction with Fe2O3
in welding steel
Gallium
(Ga) - melts in the hand; largest liquid state; transistors and high-temp.
thermometers
Indium
(In) - soft bluish; electronics
Thallium
(Tl)
Helium
(He) - hot-air balloons, He/O2 deep-sea breathing, cryogenics
Neon
(Ne) - neon signs
Argon
(Ar) - inert atmosphere for welding, incandescent light bulbs
Krypton
(Kr) - airport lights
Xenon
(Xe) - short-exposure photographs
Radon
(Ra) - radiotherapy of cancer
Halogens
("salt formers") (VIIA)
Fluorine
(F) - pale yellow gas; prepared in Monel metal cell
Chlorine
(Cl) - "green", yellow-green gas; made from NaCl; chlorinates
hydrocarbons (chain eactions with
radicals and termination steps), household bleaches,
swimming pools
Bromine
(Br) - "stench", dark-red liquid; eyeglasses, film, sedatives
Iodine
(I) - "purple", violet-black crystalline; from dried seaweed; in
growht-regulating hormone thyroxine
Oxygen
(O) - breathing, oxidizing, many other uses
Sulfur
(S) - mined by Frasch "hot water" process, "brimstone",
yellow, stable rhombic and monoclinic
forms; contact process used to make 40 million tons
of sulfuric acid annually
Selenium
(Se) - red glass coloring, copy machines, solar cells
Tellurium
(Te) - added to metals to increase electrical resistance
Nitrogen
(N) - 78% of atmosphere, nitrogen cycle (nitrogen-fixation)
Phosphorus
(P) - present in all living things; used in fertilizers
Carbon
(C) - part of all organic compounds
Silicon
(Si) - Al-Si alloys for aircraft, silicon dioxide occurs as quartz and flint;
glass and computer chips
Hydrocarbons
- compounds of only carbon and hydrogen
alkanes - no multiple bonds between carbons
(saturated), 1.54 A bond, 109.50, originally called
"paraffins"
(little affinity)
CnH2n+2, methane, ethane, propane, butane,
pentane, hexane, heptane, octane, nonane,
decane,
eicosane (20), triacontane (30), hectane (100);
branching: iso - one carbon off main chain, tert -
two carbons off main chain,
neopentane
C(CH3)4;
alkyl groups - alkane attached to another group; iso
- connecting carbon in middle of
side
chain, sec - 2o connecting carbon, tert - 3o connecting
carbon
cycloalkanes - rings, CnH2n,
substituted at axial and equatorial positions (switch in ring
flip)
so can be cis/trans
bicycloalkanes - two fused or bridged rings, decalin C10H18
alkenes - at least one double bond between carbons, three sp2
hybrid orbitals, rotation breaks pi
bond,
1.34 A bond, 1200
CnH2n,
ethene (ethylene), propene, butene, pentene, hexene, heptene, octene, nonene
vinyl
group CH2=CH- , allyl group CH2=CHCH2- ,
5,5-Dimethyl-2-hexene CH3CH=CHCH2C(CH3)3
, alkadiene has two double bonds,
alkatriene
has three double bonds
alkynes - at least one triple bond between carbons, 1.2 A bond, 1800
CnH2n-2,
ethyne (acetylene), propyne, butyne, pentyne, hexyne, heptyne, octyne, nonyne
5-methyl-1-hexyne CH3CH(CH3)CH2CH2C=-CH
, alkadiyne has two triple bonds,
alkatriyne
has three triple bonds
arenes (aromatic) - unsaturated cyclic hydrocarbons
annulene - monocyclic compounds with alternating single and double
bonds
Huckel's
Rule - planar monocyclic rings with 4n+2 delocalized electrons are aromatic
antiaromatic - greater pi-electron energy than open chain; nonaromatic
same; aromatic
less
benzene - C6H6
Kekule structure of alternating single/double C bonds
phenyl group - benzene ring attached to another group
benzyl - benzene-CH2- attached to another group
benzenoid polycyclic aromatics including naphthalene C10H8
nonbenzenoid aromatic compounds including azulene C10H8
fullerenes
- Kroto, Curl , and Smalley found C60 buckminsterfullerene, 20
hexagons and
12 pentagons, each sp2, can make salt with
K
heterocyclic aromatic compounds including pyridine C5H5N,
pyrrole C4H5N, furan
C4H4O,
thiophene C4H4S
benzene
derivatives
fluorine - fluorobenzene
methyl - toluene
hydroxyl - phenol
amine - aniline
hydrogen sulfate - benzenesulfonic acid
carboxyl - benzoic acid
CH=CH2 (phenylethene) - styrene
COCH3 (ester) - acetophenone
OCH3 (ether) - anisole
two methyls - xylene (ortho, meta, para)
two hydroxyls - benzenediol (hydroquinone if 1,4)
methyl and hydroxyl - cresol
carbonyl - benzaldehyde
carbonyl and meta OCH3 and para hydroxyl - vanillin
CONH2 - benzamide
C=-N - benzenecarbonitrile
EAS
benzene activating ortho-para directors (eg OH, O) and deactivating meta
directors
(eg NO2, have partial or full positive
charge), halo groups are deactivating ortho-para
directors
Functional
groups
alkyl halides - halogen (F, Cl, Br, I) replaces hydrogen on an alkane;
primary, secondary, or
tertiary
depending on number of carbons connected to the carbon bound to the halide
chloroethane
CH3CH2Cl, vinyl halide C=C-X, phenyl or aryl halide
phenyl-X
alcohols - hydroxyl group (OH) attached to sp3 carbon, R-OH;
primary, secondary, or
tertiary
depending on number of carbons connected to the carbon bound to the halide
methanol CH2OH, ethanol CH3CH2OH,
4-Methyl-1-hexanol CH3CH2C(CH3)CH2CH2CH2COH,
1,2-Ethanediol or Ethylene glycol HO-CH2CH2-OH,
3-penten-2-ol CH3CH=CHCHOHCH3 ,
2-Methyl-4-pentyn-2-ol
CH3C(OH)(CH3)CH2=-CH,
not as strong of acids as phenols
ethers - oxygen between carbons, R-O-R'
ethyl methyl ether CH3OCH2CH3 ,
2-Methoxypentane CH3CH(OCH3)CH2CH2CH3
,
epoxides (oxiranes) are 3-member cyclic ethers, crown ethers are cyclic
polymers of
ethylene
glycol and can be phase transfer catalysts, tetrahydrofuran (THF)
amines - nitrogen attached to at least one carbon; primary R-NH2,
secondary two Rs and an H,
tertiary three Rs, ethylamine (ethanamine) CH3CH2NH2
, cyclic amines include pyrrole C4H4N,
pyridine C5H5N, pyrrolidine C4H8NH,
and purine C5H4N4 , more basic than amides,
biological
amines include nicotine, morphine, codeine, dopamine,
serotonin, adrenaline (epinephrine),
histamine
aldehydes - carbonyl group at end of chain, R-CO -H
formaldehyde CH2O, acetaldehyde (ethanal)
CH3CHO, benzaldeyde C6H5CHO,
5-Chloropentanal ClCH2CH2CH2COH
ketone - carbonyl group in middle of chain, R-CO -R
acetone
CH3COCH3, ethyl methyl ketone (butanone) CH3CH2COCH3,
4-Penten-2-one
CH3COCH2CH=CH2 ,
benzopehenone (diphenyl ketone) C6H5COC6H5
carboxylic acids - carboxyl group attached to carbon, R-CO -OH
formic acid HCOOH, acetic acid CH3COOH, benzoic acid C6H5COOH,
4-Hexenoic acid CH3CH=CHCH2CH2COOH,
dicarboxylic acids are called alkanedioic
acids
amides - nitrogen and oxygen bound to carbon, R-CO -NR'R"
acetamide (ethanamide) CH3CONH2,
N,N-Dimethylacetamide CH3CON(CH3)2 , cyclic
amides
are lactams
esters - two oxygens bound to carbon, RCOOR'
ethyl acetate (ethyl ethanoate) CH3COOCH2CH3
, tert-Butyl
propanoate
CH3CH2COOC(CH3)3 , malonate,
cyclic esters are lactones
nitriles - nitrogen triple-bonded to carbon, C=-N, ethanenitrile CH3C=-N
(Different
compounds that have the same molecular formula)
Constitutional
(structural) isomers - differ in connectivity; different physical properties
Stereoisomers
- differ only in arrangement of atoms in space, # of isomers < 2# of
stereocenters
Diastereomers - molecules are not mirror images of each other
Cis (same side) / trans (opposite sides, more stable) for disubstituted
alkenes
E
/ Z system to name by prioritizing groups (same as R/S)
Enaniomers - molecules are nonsuperposable mirror images of each other
R (rectus) / S (sinister) system to name by prioritizing groups
attached to stereocenter,
higher atomic number means higher priority; developed
by Cahn, Ingold, and Prelog
Optically active (rotate plane-polarized light)
Clockwise dextrorotatory or counterclockwise levorotatory
Specific rotation [a]
= a/(c*l)
observed/((g/mL) * (dm))
Equimolar mixture of two enantiomers is racemic
Meso compounds are achiral despite having tetrahedral atoms with 4
different attached
groups
because it has a plane of symmetry
Fischer projection formulas represent chiral molecules
Dimethylcyclohexanes: 1,4 diasteromers, 1,3 meso, 1,2 enantiomers
Resolution - separation of enantiomers of a racemic form
Allenes - chiral molecules with C=C=C instead of tetrahedron
(3D
aspects of molecular structure)
Conformational
analysis of alkanes - Newman projection and sawhorse formulas, torsional
strain:
anti < gauche < eclipsed
Ring
strain - measure by heat of combustion (greater heat means more potential
energy and less stable),
cyclohexane most stable and cyclopropane least stable
cycloalkanes, due to angle strain and torsional
strain; chair conformation of cyclohexane has no
angle or torsional strain; boat conformation has
torsional strain only; strain: chair < twisted
< boat; diaxial interactions cause steric strain, less if equatorial
than axial
Substitution,
addition, elimination, rearrangement
Heterolysis
produces ion, homolysis produces radicals
Electrophiles
seek extra electrons, nucleophiles seek proton or other positive center
Nucleophilic
substitution reactions
Nucleophile + Alkyl halide --> product + halide ion