General Chemistry    Organic Chemistry     Physical Chemistry




General Chemistry





Chemical Formulas and Composition Stoichiometry

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


Types of Chemical Reactions

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


The Structure of Atoms

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


Chemical Periodicity

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)


Chemical Bonding

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


Molecular Structure and Covalent Bonding

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 Orbitals

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


Acids, Bases, and Salts

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


Liquids and Solids

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


Acids and Bases

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 Principle

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

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 Chemistry

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






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)


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


DG0 = -RTln(K)

van't Hoff equation: ln(KT2/KT1) = DH0/R (1/T1 - 1/T2) ; estimate equilibrium constant at another


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)



Constants and Numbers


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




Polyatomic Ions

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


Ternary acids names: perXic (perXate), Xic (Xate), Xous (Xite), hypoXous (hypoXite)

Strengths (inc.) : NH3, H2O, NH4, HCN, CH3COOH, HF, HNO3, HCl, HBr, HI, HclO4


Quantum Numbers

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


Orbital Orderings

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p68s2   diagonals


Bond Orders

H2 1, He2 0, B2 1, N2 3, O2 2


Solubility Rules

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


Three Laws of Thermodynamics

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


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





First Twenty

H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca



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


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


Group IIIA

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)


Noble Gases

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


Group VIA

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


Group VA

Nitrogen (N) - 78% of atmosphere, nitrogen cycle (nitrogen-fixation)

Phosphorus (P) - present in all living things; used in fertilizers


Group IVA

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




Organic Chemistry



Organic Compounds


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


                                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


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),


                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


                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



Types of Reactions


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

                Nucleophile has unshared electron pair

                Leaving groups (nucleofuges) such as halides become stable weak base, triflate ion

CF3SO3- > I > Br > Cl > F

                SN2 reaction rate proportional to concentrations of both reactants, nucleophile approaches leaving

group carbon from backside causing R/S inversion of configuration, goes via transition state

                SN1 three-steps with first slowest so rate depends only on alkyl halide concentration, product is

racemic since carbocation intermediate is achiral

                Structure of substrate can slow rate if it blocks access (steric hindrance); methyl > primary >

secondary >> tertiary; stability of carbocation determines SN1 rate; vinyl and phenyl halides


                Hammond-Leffler postulate: structure of transition state resembles the stable species that is nearest

it in free energy

                Polar protic solvents may slow reaction by solvating nucleophile so SN2 reactions may be faster in

polar aprotic solvents (such as DMSO dimethyl sulfoxide, DMF N,N-dimethylformamaide, DMA


Elimination reactions

                E2 reaction rate proportional to concentrations of both reactants; SN2 is favored with primary

halides but E2 with secondary halides; high temperature and strong sterically hindered base

favor elimination

                E1 reaction forms carbocation like SN1 and makes some substitution product also






Hydrogenation - make alkanes from alkenes with H2 and Ni or Pt

Reduction of alkyl halides - make alkanes from alkyl halides with H+, Zn, and acid

Corey-Posner Whitesides-House Synthesis - make alkanes from organic halides via lithium dialkylcuprate

(R2CuLi) in diethyl ether

Enantioselective reactions produce more of one enantiomer than another

Hydrolysis: ester + water --> carboxylic acid + alcohol


Alkyl Halides

Functional group transformations - make almost anything from alkyl halide

Dehydrohalogenation - make alkenes from alkyl halides with ethanol and sodium alkoxide, ab or 1,2

eliminations; Zaitsev's Rule: forms most stable, most highly substituted alkene; Hofmann's Rule: exception

to Zaitsev's Rule, forming less substituted alkene; anti periplanar transition conformation preferred to syn




Catalytic hydrogenation - make alkanes from alkenes with H2 and fine metal (Ni, Pd or Pt)

Syn Hydrogenation - make Z-alkenes from alkynes with H2 /Ni2B or Lindlar's catalyst

Anti Hydrogenation - make E-alkenes from alkynes with Li, etheamine and ammonium chloride, vinylic

radical intermediate with one radical carbon across double bond

Dehydration of Alcohols - make alkenes from alcohols with strong acid, 30 alcohols fastest because they

make more stable E1 carbocation

Debromination of Vicinal Dibromides - make alkenes from alkyl halides with bromides vic (on adjacent

carbons, not gem, on same carbon) with Zn and formaldehyde or NaI and acetone

Addition of Hydrogen Halides to Alkenes - make alkyl halide from alkene with HX; Markovnikov's Rule:

hydrogen atom adds to carbon of double bond that already has greater number of hydrogen atoms, yielding

more stable carbocation intermediate; regioselective reaction, producing mostly one of two possible

constitutional isomers; exception if HBR is added with peroxides; forms radicals

Addition of Sulfuric Acid - make alcohols from alkenes with cold sulfuric acid via alkyl hydrogen sulfate

intermediate, then heating

Acid-Catalyzed Hydration - make alcohols from alkene with water and acid

Addition of Bromine and Chlorine to alkenes - make vic alkyl halide from alkane and dimolecular Br or Cl

with sunlight and CCl4; red brown Br color goes away as test for alkenes; stereospecific to anti addition

to cycloalkene

Halohydrin Formation - make halohydrin (-CXCOH-) from alkene with dimolecular Br or Cl and water

Oxidations of Alkenes - make glycols (1,2 diols) from alkenes with KMnO4 and OH- or OsO4 (osmium

tetroxide) via syn hydroxylation and osmate intermediate

Oxidative Cleavage of Alkenes - make carboxylic acid from alkene with hot permanganate

Ozonolysis of Alkenes - make aldehydes and/or ketones from alkenes with ozone and Zn / water

Addition of Bromine and Chlorine to Alkynes - make trans-dihaloalkenes or tetrahaloalkanes with one or

two equivalents of dimolecular halogen

Addition of Hydrogen Halides to Alkynes - make gem-dihalide from alkyne with 2 HX; anti-Markovnikov

intermediate if peroxides used


Radical Reactions

Radical Reactions - alkane + halogen --> halo, dihalo, trihalo, and/or tetrahaloalkane + HX with light;

initiation, propagation, and terminating steps; bromine less reactive but more selective

Chain growth polymers - monomers combine using peroxides to form radical intermediates

Autoxidation - organic compound reacts with oxygen to form hydroperoxide

Ozone depletion - radical reactions with CF2Cl2 convert O3 to O2 via radicals


Alcohols and Ethers

Oxymercuration / Demercuration - make alcohols from alkenes with mercuric acetate and THF

Hydroboration - make organoborane from alkene and boron hydride with THF via borane intermediate;


Alcohol Reactions - make protonated alochols from alcohols with strong acid or protonated ethers from


Alcohol Reactions - make mesylates (methanesulfonates) and tosylates (toluenesulfonates) from alcohols

and sulfonyl chlorides

Alcohols into Alkyl Halides - make alkyl halides from alcohols with hydrogen halides, phosphorus

tribromide, or thionyl chloride

Synthesis of Ethers - make ethers from alcohols with H+

Williamson Synthesis of Ethers - make ethers from sodium alkoxide with alkyl halide, alkyl sulfonate, or

alkyl sulfate; can make sodium alkoxides from phenols

Silylation - make trimethylsilyl ethers from alcohols with chlorotrimethylsilane, protecting OH

Ether Reactions - make oxonium salts from ethers and hydrogen halide

Epoxidation - make epoxide from alkene and peroxy acid

Alcohols by Reduction - make alcohols from carboxylic acids and lithium aluminum hydride or esters and

high pressure or aldehydes/ketones and sodium borohydride

Oxidation of Alcohols - make aldehydes from alcohols with potassium dichromate and sulfuric acid or

pyridinium chlorochromate (PCC)  and dichloromethane

Oxidation of Alcohols - make carboxylic acids from alcohols with potassium permanganate and hydroxide

Oxidation of Secondary Alcohols - make ketones from secondary alcohols with sodium dichromate and


Grignard Reagents - make Grignard reagents from organic halide and Mg with ether

Grignards with Oxiranes - make primary alcohols from oxiranes with Grignard

Grignards with Carbonyls - make primary/secondary alcohols from aldehydes with Grignard or tertiary

alcohols from ketones with Grignard

Organolithium reagents and sodium alkynides work like Grignard


Conjugated Unsaturated Systems

Shell Process - make allyl chloride from propene and chlorine via radicals

Bromination of Allyl - make allyl bromide from propene with N-Bromosuccinimide (NBS)

Electrophilic Attack on Conjugated Dienes - HX adds to one of two double bonds, or H to one and X to the

other and double bond in between (1,2 and 1,4 additions)

Diels-Alder Reaction - make an adduct from conjugated diene and double bonded dienophile; forms two

sigma bonds at expense of two pi bonds; syn addition and cis and endo


Electrophilic Aromatic Substitution Reactions of Aromatic Compounds

Bromine with Benzene - reacts with Lewis acid catalyst by substitution not addition

Electrophilic Aromatic Substitution Reactions - arene + electrophile --> arene-electrophile + H+via

nonaromatic carbocation called arenium ion

Halogenation - make halobenzene from benzene and Br2 or Cl2 with Lewis acid FeBr2 or FeCl2

Nitration - make nitrobenzene from benzene and hot nitric acid with sulfuric acid

Sulfonation - make benzene-sulfonic acid from benzene and fuming sulfuric acid (extra SO3)

Friedel-Crafts Alkylation - make alkylbenzenes from alkyl halides and benzene with AlCl3, or from alkene

and acid, or from alcohol and acid

Friedel-Crafts Acylation - make acylated benzene from benzene and acetyl halogen or carboxylic

anhydride with AlCl3; poor yield if strong withdrawing groups are present on ring; major product form more

stable carbocation; often get polyalkylations

Clemmensen Reduction - make alkyl benzene from ketone from Friedel-Crafts Acylation with

amalgamated zinc and HCl reflux

Birch Reduction - reduce benzene to 1,4-Cyclohexadiene with alkali metal, ammonia, and alcohol


Aldehydes and Ketones

Aldehyde Synthesis - make aldehydes from primary alcohols with PCC and dichloromethane

Aldehyde Synthesis - make aldehydes from acyl chlorides with lithium tri-tert-butoxy-aluminum hydride

Aldehyde Synthesis - make aldehydes from esters or nitriles with DIBAL-H and hexane

Ketone Synthesis and Tautomerization - make ketones from alkynes with sulfuric acid and mercuric ions,

resulting in keto-enol forms which tautomerize; Markovnikov; enol form more common with beta-dicarbonyl

compounds; ketone can lose its optical activity by converting to achiral enol form

Ketone Synthesis - make ketones from acyl chlorides with lithium dialkylcuprate or from nitriles with

Grignard or organolithium

Nucleophilic Substitutions of Carbonyl - RCOH + Nu-H --> RCNuHOH; aldehydes more reactive

Hydrates - make hydrates (gem diols, RCH(OH)2) from aldehyde and water with acid or base

Hemiacetals - make hemiacetals (RCH(OR')(OH)) from aldehyde or ketone and alcohol

Acetals - make acetals (RCH(OR')2) from adehyde or ketone and alcohol with gaseous HCl; may serve as

a protecting group for subsequent reactions, as it can be removed with acid and water

Thioacetals - make thioacetals (HCR(SR')2) from aldehyde or ketone and thiols with acid

Desulfurization - make hydrocarbons from thioacetals and hydrogen with Raney nickel

Imines - make imines (C=N-R) from aldehyde or ketone and primary amine

Wolff-Kishner Reaction - reduce C=O to CH2 in aldehyde or ketone with hydrazine (H2NNH2) and base,

via hydrazone (C=NNH2)

Semicarbazide - reduce C=O to CH2 in aldehyde or ketone with semicarbazide (H2NNHCONH2) via


Cyanohydrins - make cyanohydrins (RHC(OH)(CN)) from aldehydes or ketones with hydrogen cyanide (HCN)

Wittig Reaction - make alkenes and triphenylphosphine oxide from aldehyde or kentone and phosphorus

ylide (phosphorane, (C6H5)3P-CRR') via betaine intermediate

Reformatsky Reaction - make beta-hydroxy esters from aldehydes or ketones and alpha-bromo ester with

zinc and benzene

Baeyer-Villiger Oxidation - make carboxylic esters from ketones and peroxy acid

Haloform Reaction - make multiple halogen substitutions on alpha-carbon of methyl ketone with X2 and


Aldol Additions - make aldol (aldehyde alcohol) from aldehyde with dilute NaOH; can have crossed aldol

reactions if start with two different carbonyl compounds

Aldol Condensation - make enal (unsaturated aldehyde) from aldol via dehydration; can also condensate to

make unsaturated amines or nitriles from nitroalkenes or nitriles

Claisen-Schmidt Reaction - crossed aldol reactions using one ketone; make unsaturated ketones from

aldehyde and ketone with base; used with geranial and acetone to make Vitamin A

Cyclizations via Aldol Condensations - make 5 or 6 membered rings from dialdehyde, diketone, or keto

aldehyde with base

Additions to alpha-beta Unsaturated Aldehydes and Ketones - simple addition and conjugate addition (with

keto and enol forms) reactions both occur

Michael Additions - conjugate addition of enolate ions to alpha-beta unsaturated carbonyl compounds with


Robinson Annulation - uses Michael addition and simple aldol condensation to build one ring onto another


Carboxylic Acids, Esters, and Amides

Grignards - make carboxylic acids from Grignard reagent and carbon dioxide with acid

Nucleophilic Substitutions of Acyl Carbon - replace leaving group on acyl carbon with nucleophile

Acyl Chloride Synthesis - make acid chlorides from carboxylic acids and thionyl chloride (SOCl2) or

phosphorus pentachloride (PCl5)

Synthesis of Carboxylic Acid Anhydrides - make carboxylic acid anhydrides (R-COOCO-R') from

carboxylic acid and acyl chloride with pyridine

Esterification - make esters by condensation of carboxylic acids and alcohols with acid or from acyl

chlorides and alcohols or from carboxylic acid anhydrides and alcohols

Saponification - make alcohol and carboxylate salt from hydrolysis of ester by base

Amide Synthesis - make amides from acid chlorides or acid anhydrides or esters and amines or ammonia,

or from carboxylic acids and ammonia with dicyclohexylcarbodiimide (DCC)

Amide Hydrolysis - make carboxylic acid and ammonia from amide and acid or base

Amide Dehydration - make nitriles from amides with phosphorus pentoxide P4O10

Nitrile Hydrolysis - make carboxylic acids from nitriles with acid or base

Hell-Volhard-Zelinski Reaction - make alpha-halo carboxylic acids from aliphatic carboxylic acids and Br2

or Cl2 with phosphorus

Decarboxylation of Carboxylic Acids - remove carboxyl group from beta-keto carboxylic acids by heating,

or from carboxyl radicals

Dicarbonyl Compounds

Claisen Condensation - make beta-keto esters from esters and sodium ethoxide, via aldol addition, enolate

anion attack, and acid-base reaction; can be crossed with two esters if one ester has no alpha hydrogens

Dieckmann Condensation - make 5 or 6 membered rings by an intramolecular Claisen condensation

Acetoacetic Ester Synthesis - make substituted acetones from acetoacetic ester

Malonic Ester Synthesis - make substituted acetic acid from malonic ester

Knoevenagel Condensation - active hydrogen compounds condense with aldehydes and ketones, like aldol

condensations, with weak base

Mannich Reaction - make Mannich bases from enols and formaldehyde and primary or secondary amine

Stork Enamine Reaction - make enamines from aldehydes or ketones with secondary amines, and then

acylate or alkylate the enamines or use in Michael Additons



Nucleophilic Substitution Reaction of Amines - make amines from alkyl halide and ammonia, optionally

via azide (N3-) ion intermediate

Preparing Aromatic Amines - make aromatic amine from arene with nitric acid and sulfuric acid and then

reduction with H2 or iron and HCl

Reductive Amination - make amines from aldehyde or ketone and ammonia or an amine

Hofmann Rearrangement - make amines from amides and X2 with NaOH

Curtius Rearrangement - make amine from acyl chloride with NaN3 via acyl azide and isocyanante


Diazotization Reaction - make unstable aliphatic diazonium salts from primary aliphatic amines and

nitrous acid (HONO, made from HCl and NaNO2 in situ)

Sandmeyer Reaction - make Cl, Br, or CN substituted arenes from arenediazonium salts and CuCl, CuBr,

or CuCN

Diazonium Replacements - make I, F, or OH substituted arenes from arendiazonium salts and KI, HBF4, or


Deamination by Diazotization - replace diazonium group with hydrogen using hypophosphorous acid

Diazo Coupling Reactions - make azo compounds (Ar-N=-N-Ar) from arenediazonium ions and reactive

aromatic compounds

Synthesis of Sulfonamides - make sulfonamides (R-NH-SO2Ar) from primary or secondary amine and

sulfonyl chloride

Hofmann Elimination - make alkene, water, and tertiary amine from quaternary ammonium hydroxide in

E2 elimination

Cope Elimination - eliminate dialkylhydroxylamine from tertiary amine oxides



Dow Process - make phenol from chlorobenzene and NaOH by heating at high pressure and using HCl

Alkali Fusion - make phenol from sodium benzenesulfonate and NaOH at 350 C

Cumene Hydroperoxide - make phenol and acetone from benzene and propene via cumene by Friedel-

Crafts alkylation

Kolbe Reaction - make salicylic acid from phenol and carbon dioxide; salicylic acid and acetic anhydride

form aspirin (acetylsalicylic acid)

Claisen Rearrangement - make o-Allylphenol from allyl phenyl ether by heating

SNAr Mechanism - replace halogen on arene by hydroxyl if strong electron withdrawing groups (such as

NO2) are ortho or para to halogen; via caranion with delocalized electrons, called Meisenheimer complex





Bromine in carbon tetrachloride - red/brown disappears if carbon-carbon double or triple bonds

Cold potassium permanganate - purple turns brown if carbon-carbon double or triple bonds

Silver ion - forms precipitate with alkyl halides

Silver nitrate in ammonia - forms precipitate with alkynes

Chromic oxide - turns from orange to green with primary or secondary alcohols

Tollens' Test (Silver Mirror Test) - metallic silver precipitates from silver nitrate and aqueous ammonia if

aldehydes or alpha-hydroxy ketones are present

Iodoform Test - iodine in sodium hydroxide precipitates bright yellow iodoform (CHI3) if COCH3 or

CH(OH)CH3 groups are present

Hinsberg Test - excess KOH and then acid used to demonstrate whether amine is primary, secondary, or


Benedict's Test - alkaline Fehling solution with cupric citrate complex ion gives red Cu2O precipitates if

aldoses are present; acetal carbohydrates give negative result





Visible and UV spectroscopy - plot wavelength vs. absorbance; multiple bonds absorb radiation;

nonconjugated molecules have maxima below 200 nm and can't be measured; conjugated systems of at

least 8 are in visible range; electron iexcited from bonding pi highest occupied molecular orbital to

antibonding pi lowest unoccupied molecular orbital; C=O absorb in UV region

Infrared spectroscopy - plot wavenumber (1/wavelength) vs. transmittance (1/absorbance); atomic masses

and bond stiffness give different functional groups absorbances at characteristic frequencies; dipole

moment changes as vibration absorbs IR energy; bond strength: sp > sp2 > sp3; stretching and out-of-plane


Nuclear Magnetic Resonance spectroscopy - older sweep (CW) and newer Fourier Transform (FT)

methods; area of peaks indicates number of hydrogens of that type; signal splitting (doublet, etc.) is one

more than the number of non-equivalent hydrogen neighbors within three sigma bonds (from spin-spin

coupling); induced field of electrons shileds proton from external field; tetramethylsilane (TMS) reference

compound; chemical shifts based on shielding; homotopic and enantiotopic hydrogens have same chemical

shift; separation is called coupling constant J; C13 NMR has no signal splitting since only one of 100 is C13,

different carbons have different chemical shifts; off-resonance decoupling splits carbon signal based on

number of attached hydrogens; simplified DEPT (distorionless enhanced polarization transfer) method







Aldoses - contain aldehyde

                                                Aldohexoses (six carbons) - Glucose (RLR), Galactose (RLL)

                                                Aldopentoses (five carbons) - Ribose (RR)

                                                Aldopetroses (three carbons) - Erythrose (R), Threose (L)

                                Ketoses - contain ketone

                                                Ketohexoses (six carbons) - Fructose

                                                Ketopentoses (five carbons) - Ribulose


                                Sucrose - glucose and fructose

                                Maltose - glucose and glucose (alpha glycosidic linkage)

                                Cellobiose - glucose and glucose (beta glycosidic linkage)

                                Lactose - glucose and galactose


                                Starch - alpha 1,4 links; amylose (unbranched) and amylopectin (branched)

                                Glycogen - alpha 1,4 links

                                Cellulose - beta 1,4 links

Other Information

                D(+) and L(-) enantiomers

Haworth formulas of cyclic hemiacetyl forms

Mutarotation between alpha and beta forms

Glycosides are carbohydrate acetals

Lobry de Bruyn-Alberda van Ekenstein transformations dissolve monosaccharides in base

                Monosaccharides react with phenylhydrazine to form osazones

                Kiliani-Fischer Synthesis lengthens the chain of an aldose

                Ruff Degradation shortens the chain of an aldose

                Emil Fischer proved the configuration of D-(+)-Glucose





Greek "lipos" means fat

Triacylglycerols (fatty acids) consist of three-hydroxyl alcohol glycerol hydrolysed with carboxylic acids,

forming esters

Unsaturated fats have at least one double bond; polysaturates have multiple (unconjugated) double bonds

Saponification of triacylglycerols produces glycerol and salts of carboxylic acids

Terpene lipids have repeating isoprene (-CH2CH(CH3)CH2CH2-) units; terpenoids have oxygen

Carotenes are tetraterpenes

Natural rubber is a 1,4 addition polymer of isoprene

Steroids are lipids with fused ring system (3 six-carbon rings and one five-carbon ring), including

cholesterol and hormones Prostaglandins are C20 carboxylic acids with five-membered ring and at least one

double bond


Amino Acids and Proteins


Twenty of 22 amino acids are used in protein synthesis; hydroxyproline and cystine are made after the

chain is intact

Amino acids are of form H2NCHRCO2H, with side chain R = hydrogen in  glycine, methyl in alanine,

CH(CH3)2 in valine, CH3-Ph in phenylalanine, CH2OH in serine, CH2SH in cysteine, CH2CH2SCH3 in


Strecker Synthesis of amino acids from ammonia, an aldehyde, and hydrogen cyanide

Amino acids link by peptide bonds, forming (di,tri,oligo,poly)peptide proteins; linear polymers with N-

terminal and C-terminal residues

Sanger method and Edman degradation determine the N-terminal amino acid residue

Protein Structure

                Primary - sequence of amino acids

                Secondary - local conformation of polypeptide backbone; alpha-helix or beta-pleated sheet

                Tertiary - further foldings, exposing polar groups to aqueous environment, caused by things such

as disulfide bonds

                Quaternary - combination of multiple polypeptides into a protein



Nucleic Acids


Nucleosides consist of five-carbon monosaccharide (D-ribose or 2-deoxy-D-ribose) and heterocyclic base

(purine [adenine or guanine] or pyrimidine [cytosine or thymine or uracil])

Nucleotides consist of nucleoside and phosphate ion

In DNA, phosphate esters link 3' OH of one ribose with 5' OH of another; A-T and C-G pairings

Protein synthesis takes place via transcription (messenger RNA made from DNA) and translation (mRNA

codon matches with transfer RNA-amino acid anticodons at ribosomes, forming polypeptides)





Carbon ions - carbocation (posititve), carbanion (negative); carbocations are trigonal planar and are more

stable the more carbons are connected (R3C+ > HR2C+)

Acidity increases down group (HF weakest and HI strongest, ethyne more acidic than ethane, carboxylic

acids more acidic than alcohols because of resonance and inductive effects

Protic solvent - has hydrogen atom attached to strongly electronegative element (eg O or N)

Protonated alcohol or alkyloxonium ion - conjugate base of alcohol

Catalytic cracking - alkane gas oil mixture heated at 500C and break apart, rearranging into smaller, highly

branched alkanes

Thermal cracking - like catalytic cracking but no catalysts so chains unbranched

Catenanes - chains of large rings of cycloalkanes

Pheromones - odorous chemicals used by animals for communication

Chiral - not identical with its mirror image ("hand"), achiral are superposable on mirror image

Stereocenter - atom bearing groups such than any interchange produces a steroisomer

Solvolysis - nucleophile is molecule of the solvent; hydrolysis if water

Dielectric constant - measure of solvent's ability to insulate opposite charges and be polar

Index of Hydrogen Defiiciency - number of rings and pi bonds

Radicals - have unpaired electron; 30 > 20 > 10 stability

Ziegler-Natta catalysts - used to make polyehylene

Vitamin A - alpha-tocopherol

Azeotropes - mixture with boiling point different from either pure component

Conjugated unsaturated systems - molcules with delocalized pi bonds; p orbital on atom adjacent to double

bond; multiple bonds may be cumulated (allene; one carbon participates in both), conjugated (adjacent

carbon), or isolated

Resonance energy - difference between amount of heat actually released and that predicted based on


Aliphatic - hydrocarbon or a derivative of a hydrocarbon

Urea - H2NCONH2

Carbamates (urethanes) - RO-CO-NHR'

Polyamides - chains of amides, including nylon and proteins

Polyesters - chains of esters, including poly(ethylene terephthalate), called Dacron, Terylene, or Mylar

Polyurethanes - formed from alchol and isocyanate

Barbiturates - formed from diethyl malonate and urea with sodium ethoxide

Sulfanilamide - used in chemotherapy

Naphthols and phenanthrols - hydroxyl group attached to polycyclic benzenoid ring

Benzyne - benzene with one triple bond

Zwitterions - dipolar ions (both positive and negative ions in same molecule, such as amino acids)



Physical Chemistry and Thermodynamics



The Boltzmann constant is R/NA , or 1.38*10-23 J/K.

A Joule is kg*m2/s2.

The de Broglie relation states that the product of linear momentum and wavelength is a constant, Planck's

constant, 6.6*10-34 Js

The Boltzmann distribution gives the ration of the numbers of particles in states with given energies, Ni/Nj

= e-(Ei-Ej)/kT

The Maxwell distribution gives the proportion of molecules that have a specific speed at a particular

temperature, based on Boltzmann distribution.



Zeroth Law of Thermodynamics - If A is in thermal equilibrium with B, and B is in thermal equilibrium

with C, then C is also in thermal equilibrium with A

Ideal Gas Law: PV = nRT

PV = 1/3nMc2 ; c = root mean square speed of molecules, M = molar mass mNA

Maxwell distribution of speeds

Van der Waals equation, using compression factor and virial coefficients, adjusts ideal gas law to real gases

Maxwell construction replaces unrealistic van der Waals loops.

Principle of corresponding states: real gases at same volume and temperature exert same reduced pressure.


First Law

The First Law states that the total energy in the universe is constant.

Diathermic boundary permits heat transfer; adiabatic boundary permits energy transfer but not heat


Heat capacity at constant volume is partial of internal energy with respect to temperature; at constant

temperature is partial of internal energy with respect to volume.

Enthalpy is sum of internal energy and pressure times volume.

Hess' Law: DHrxn0 = DHa + DHb + ... ; standard enthalpy is the sum of the standard enthalpies of the

individual reactions into which a reaction may be divided.

Kirchoff's Law is used to estimate standard enthalpies of formation from molar heat capacities and reaction

enthalpies at some other temperature.

James Joule tried to measure internal pressure by observing the change in temperature of a gas expanding

in a vacuum.

Joule-Thompson coefficient is partial of temperature with respect to pressure at constant enthalpy.

The Joule-Thompson effect is the cooling by adiabatic expansion, such that the temperature difference is

proportional to the pressure difference.


Second Law

The Second Law states that he entropy (disorder) of an isolated system increases in the course of a

spontaneous change.

Carnot cycle consists of isothermal reversible expansion, reversible adiabatic expansion, isothermal

reversible compression, and adiabatic reversible compression.  Efficiency of a Carnot engine is 1-Tc/Th

Clausius inequality states that change in entropy is greater than or equal to the heat supplied to the system

during the process divided by temperature.

Trouton's rule states that the standard molar enthalpy of vaporization is about the same for a wide range of

liquids (85 J/(K*mol).

The Debye extrapolation fits measurements of heat capacity at lower temperatures.

The Nernst heat theorem states that the entropy change accompanying any physical or chemical

transformation approaches zero as temperature approaches zero.

The Third Law states that the entropy of pure, perfect crystalline substance is zero at 0 K.

The Helmholtz energy A is internal energy minus temperature times entropy; it equals the maximum work

accompanying a process.

The Gibbs energy G is the enthalpy minus temperature times entropy; it equals the maximum non-

expansion work.

The standard Gibbs energy of formation is the standard reaction Gibbs energy for the formation of a

compound from its elements in their reference states.

dU = TdS-pdV

The Maxwell relations are derived from the fact that enthalpy, Gibbs energy, and Helmholtz energies are

all state functions.  For example, the partial of temperature with respect to volume at constant entropy equals

negative the partial of pressure with respect to entropy at constant volume.

The Gibbs-Helmholtz equation shows that if the enthalpy of the system is known, then the temperature

dependence of Gibbs energy over temperature is also known; the partial of the Gibbs energy over the

temperature with respect to temperature at constant pressure equals negative the Helmholtz energy divided

by the temperature squared.

Chemical potential is the partial of the Gibbs energy with respect to the number of molecules at constant

temperature and pressure.

Fugacity is an effective pressure.


Physical Transformations

Phase diagrams plot temperature vs. pressure.

At the triple point all 3 states are at equilibrium (4.6 torr, 0.01 C for water).

You cannot liquefy gas above critical point.

The Clapeyron equation is the slope of the phase boundary.

The Clausius-Clapeyron equation describes the variation of vapor pressure with temperature.

Ehrenfest classification groups phase transitions into first-order, second-order, and lambda-transition.

Surface tension is the constant relating work to change in surface area of a liquid.

The Laplace equation states that the pressure on the concave side of an interface is alwasys greater than on

the convex side.


Simple mixtures

The Gibbs-Dunhem equation states that the chemical potential of one component of a mixture cannot

change independently of the chemical potentials of the other components.

Raoult's Law states that the ratio of the partial vapour pressure of each component to its vapour pressure as

a pure liquid is about equal to the mole fraction of the component in the mixture.

Henry's law states that the vapor pressure of a solute is proportional to its mole fraction but the constant of

proportionality is not the pure substance's vapor pressure.

In ideal-dilute solutions, solvent obeys Raoult's Law and solute Henry's Law.

Cryoscopy measures molar mass from freezing point depression.

The van't Hoff equation states that the osmotic pressure equals the molar concentration of the solute times

R times the temperature.


Phase Diagrams

Gibbs' Phase Rule states that the variance equals the number of components minus the number of phases

plus two in a phase diagram.

Vertical isopleth and horizontal tie lines on phase diagram, used in lever rule.

Number of theoretical plates on temperature-composition diagram determines efficiency of fractional


In an azeotrope, evaporation occurs without a change in composition ("boiling without changing").

At eutectic composition, a liquid mixture freezes at a single temperature.


Chemical Equilibrium

Gibbs reaction energy is slope of Gibbs energy vs. extent of reaction.

G = -RTln(K)

LeChatelier's Principle states that a system responds to stress at equilibrium in a way that reduces stress

and reaches new state of equilibrium.

Henderson-Hasselbalch equation: pH = pKa + log([conj. base]/[acid])


Equilibrium Electrochemistry

The Born equation identifies the Gibbs energy of solvation with the electrical work of transferring an ion

from a vacuum into the solvent treated as a continuous dielectric of relative permittivity.

The Debye-Huckel limiting law calculates the activity coefficient.  The long range of Coulombic

interactions dominates contributions to nonideality in ionic solutions.

Faraday's constant equals electron charge times Avogadro's number 96.485 kC/mol; it is the amount of

electricity that reduces one equivalent weight at cathode and reduces at anode

The Nernst equation calculates electrode potentials for concentrations and partial pressures other than

standard values;  E = E0 - (2.303*R*T)/(n*F)*log(Q)  =  E0 - (0.0592*T)/n*log([Red]y/[Ox]x)  ;

In electrolytic cells external electricity causes nonspontaneous reactions by electrolysis.

In voltaic cells (galvanic cells) spontaneous chemical reactions produce electricity.

Electrodes are surfaces upon which oxidation (anode) or reduction (cathode) half reaction occurs.

Faraday's Law of Electrolysis states that the amount that oxidizes or reduces at each electrode is directly

proportional to amount of electricity that passes through cell.

In a standard cell, all species are in thermodynamic standard states (1 M , 1 atm).

The Standard Hydrogen Electrode (SHE) is a 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

pH = (E+E(cal))/(-59.16 mV)


Quantum Theory

Wien Displacement Law

The Stefan-Boltzmann law states that the energy density of the electromagnetic field varies with fourth

power of temperature.

The Rayleigh-Jeans law uses the equipartition principle to calculate the average energy of each oscillator;

predicts infinite energy density at short wavelengths.

The Planck Distribution quantizes energy.

The Einstein formula relates heat capacity to frequency.

In the photoelectric effect, electrons are ejected from metals when exposed to UV radiation.

The deBroglie relation states that particles with high linear momentum have short wavelength.

The Davisson-Germer experiment shows particles have wave-like properties.

Schrodinger's equation finds wavefunction of particle.

The Born interprestion squares the wavefunction to get probability distribution.

Heisenberg's uncertainty principle states it is impossible to known precisely both the momentum and

position of a particle.

Gaussian function is of the form e-x*x


Atomic Structure and Atomic Spectra

Rydberg combined the Balmer series (visible), Lyman series (UV), and Paschen series (infrared) about

wavenumbers transitions.

The Ritz combination principle states that the wavenumber of any spectral line is the difference between

the two terms.

The Bohr frequency condition states that an atom's energy change is carried away by a photon of frequency


The Bohr radius is 52.9177 pm.

Quantum numbers: 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)

Orbital Orderings: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p68s2   diagonals

The Aufbau Principle states that electrons added into orbitals in way giving lowest total energy.

The Pauli Exclusion Principle states that no two electrons in atom have same 4 quantum numbers.

Hund's Rule states that electrons mus toccupy all orbitals of a sublevel before pairing.

The Hartree-Fock self-consistent field procedure finds numerical solutions to electron-electron interaction

terms of Schrodinger's equation.

Ionization energy is the energy needed to remove an electron.

Signal splitting multiplicity (doublet, etc.) is one more than the number of non-equivalent hydrogen

neighbors within three sigma bonds (from spin-spin coupling); induced field of electrons shields protons

from external field.

The Clebsch-Gordon series defines the total orbital angular momentum quantum number L.

The Zeeman effect is the modification of an atomic spectrum by the application of a strong magnetic field.


Molecular Structure

The Born-Oppenheimer approximation assumes nuclei is stationary and electrons move around.

Valence Shell Electron Pair Repulsion Theory (VSEPR) describes formation of hybrid orbitals.

If 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)

In a polar covalent bond, electrons are shared unequally; creates dipole.

In a sigma bond there is head on overlap; all single bonds are sigma.

In a pi bond, there is side on overlap; may include unhybridized p orbital.

A molecular orbital is an orbital resulting from overlap and mixing of atomic orbitals on different atoms;

belongs to molecule as whole

An antibonding orbital is a molecular orbital higher in energy than any of atomic orbitals from which it is

derived; lends stability when populated; marked with asterick

Nonbonding orbitals are orbitals derived only from an atomic orbital of one atom; lends no stability

Delocalization is the formation of set of molecular orbits that extend over more than two atoms

Nodal planes are the regions of zero probability of finding electrons

Electronegativity is a measure of the power of an atom to attract electrons to itself when it is part of a


The variation principle states that if an arbitrary wavefunction is used to calculate energy, the value is never

less than the true energy.

Walsh diagrams show the variation of orbital energy with molecular geometry.

The Huckel approximations (overlap integrals and resonance integrals between non-neighbors are zero and remaining

resonance integrals are equal) allow calculation of molecular orbital energy levels.

Bonding HUMO (highest occupied molecular orbital) and antibonding LUMO (lowest unoccupied

molecular orbital).

Conjugated systems are stabilized by delocalization energy.

Semiconductor conductivity increases with temperature; conductor decreases with temperature.


Molecular Symmetry

Group theory is the discussion of symmetry.

Point groups, space groups, identity, n-fold rotation, reflection; Schoenflies and Hermann-Mauguin

(crystal) notations for symmetry elements.

Character tables characterize the symmetry types possible in a point group.



Emission spectroscopy measures change in molecule from high to low energy state, emitting excess energy

as photon.

Absorption spectroscopy monitors net absorption of nearly monochromatic incident radiation.

Michelson interferometer in Fourier transform machine analyses the frequencies.

Elements of spectroscopy include radiation source, dispersing element, FT technique, detectors, and the


Raman spectroscopy examines frequencies present in radiation scattered by molecules, low-frequency

Stokes and high frequency anti-Stokes; resonance and coherent anti-Stokes variations of Raman.

Beer-Lambert law states that absorbance equals extinction coefficient times concentration times column


The Stark effect states that the energy of a state depends on the square of the permanent electric dipole

moment; Stark modulation.

Molecules are modeled as rigid rotors (spherical, symmetrical, linear, or asymmetric), bodies that do not

distort under the stress of rotation.

The Morse potential energy curve reproduces the general shape of a molecular potential energy curve.

Birge-Sponer plots may be used to determine dissociation energy.

Tumbling is the random changing of orientation of a molecule.

The exclusion rule states that no modes of a symmetrical molecule can be both infrared and Raman active.

The Franck-Condon Principle states that electronic transitions take place much faster than nuclei can

respond since nuclei are so much more massive.

The Laporte selection rule states that the only allowed transitions are those that are accompanied by a

change of parity.

In fluorescence, radiation emission ceases immediately but in phosphorescence it may continue after the

exciting radiation is extinguished.

Jablonski diagrams show the relative positions of the electronic energy levels of a molecule.

Laser light is coherent (in step).

The active medium of a solid-state laser is a single crystal or a glass (including Maiman's first ruby laser

and neodymium lasers).

Gas lasers include He-Ne, Ar ion, Kr ion, CO2.

Other lasers include chemical, exciplex (combo of two atoms surviving in an excited state), dye, light-

emitting diodes, and semiconductors.  

Photoelectron spectroscopy finds orbital energies by measuring ionization energies when electrons are

ejected from different orbitals.

Koopman's theorem states that ionization energy equals orbital energy of ejected electron.

Shielding of nuclei decreases chemical shift in nuclear magnetic resonance spectroscopy.

Observed shielding constant is sum of local (from Lamb formula), neighbor (number of non-equivalent

adjacent hydrogens), and solvent contributions.

The splitting of resonances into individual lines is the fine structure of the spectrum.

The Karplus equation defines the coupling constant.

Fermi contact interactions occur when s-orbital electrons come very close to the nucleus.

Magnetically equivalent nuclei are chemically equivalent and have identical spin-spin interactions.

The nuclear Overhauser effect uses spin relaxation to enhance resonance line intensities.

Electron spin resonance studies molecules with unpaired electrons by observing the magnetic fields at

which they come into resonance with monochromatic radiation, measuring their hyperfine structure.


Statistical Thermodynamics

The Boltzmann distribution can be written in terms of the molecular partition function, which can be used

to calculate the proportion of molecules in different states.

The Boltzmann formula states that entropy is directly proportional to the log of the weight of the most

probable configuration of the system.

The canonical ensemble is an imaginary collection of replications of a system in thermal contact with a

constant temperature.

The Sackur-Tetrode equation gives the entropy of a monatomic gas.

The internal energy with  and entropy of a system may be calculated from its canonical partition function.


Diffraction Techniques

Diffraction is the interference caused by an object in the path of waves.

Crystals consist of repeating unit cells; 7 systems include cubic, tetragonal, orthorhombic, monoclinic,

triclinic, hexagonal, and rhombohedral.

The Miller indices are the reciprocals of intersection distances in a lattice.

Bragg's law states that a bright reflection should occur when two times the distance times the sine of the

glancing angle is an integral multiple of the wavelength.

Von Laue and Debye - Scherrer used X-rays with crystals.

The Wierl equation calculates the angular variation of the total intensity by summing the contributions from

all pairs.


Electric and Magnetic Properties of Molecules

The polarization of a sample is the electric dipole moment density, and a dielectric is a polarizable,

nonconducting medium.

Differences in atomic radii of overlapping atoms causes homopolar contribution to the dipole moment.

Induced dipole moments are proportional to the field strength.

The relative permittivity is also called the dielectric constant and is the square of the refractive index,

which is the ratio of the speed of light in a vacuum to the speed in the medium.

The Clausius-Mossotti equation assumes no permanent dipole moment, simplifying the Debye equation for

molar polarization.

The Keesom interaction is the interaction of two rotating molecules.

The London formula approximates the interaction energy in an induced dipole - induced dipole dispersion

(London) interaction.

The Lennard-Jones formula is a specific case of the Mie formula, which estimates the potential energy from

repulsions and attractions.

The magnetic flux density is related to the applied field strength and the magnetization.

If the molar magnetic susceptibility is positive, the material is paramagnetic; if negative, diamagnetic.

The Curie law defines the molar magnetic susceptibility, which is measured by a Gouy balance or a super-

conducting quantum interference device (SQUID).

At Curie temperature, spins may align making ferromagnetic transition, and at Neel temperature, spins may

alternate to antiferromagnetic phase.


Macromolecules and Colloids

Monodisperse molecules have a single, definite molar mass.

Solutions are virtually ideal at Flory theta temperature.

Donnan equilibrium is the equilibrium distribution of ions in two compartments in contact through a

semipermeable membrane, one with a polyelectrolyte (strings of acids or bases)

Sedimentation is the fall of heavy particles due to gravity.

The Stokes-Einstein relation states that the frictional cofficient is directly proportional to temperature and

indirectly to the diffusion coefficient.

In electrophoresis, charged macromolecules move due to an electric field, as in through a cross-linked

polyacrylamide gel (gel electrophoresis).

Viscosity is often measured by Ostwald or rotating drum viscometers.

Rayleigh scattering is the scattering by particles with diameters much smaller than the wavelength of the

incident radiation.

The Corey-Pauling rules describe the secondary structure of proteins, which may form alpha helices, beta

pleated sheets, or random coils.

Ramachandran plots contours of the potential energy of an entire molecule.

Colloids are purified by (electro)dialysis.

Micelles form above the Krafft temperature and may form orderly lyotropic mesomorphs.

Coagulation is the blending together of distinct particles into large particles.

The Schulze-Hardy rule states that hydrophobic colloids are flocculated (aggregation of particles)

efficiently by ions of high charge numdber and opposite charge type.

One molecule thick monolayers transferred to solid supports are Langmuir-Blodgett films.


Molecules in Motion

Graham's law of effusion states that the rate of effusion is inversely proportional to the square root of the

molar mass.

Fick's first law of diffusion states that the flux of matter is proportional to the concentration gradient.

Newtonian flow is a series of layers moving past one another.

Diffusion equals one-third the product of wavelength and mean speed for a perfect gas, and viscosity equals

one-third the product of the molar mass, wavelength, mean speed, and molar concentration.

Monte Carlo methods are used to measure the change in total potential energy of particles moved over

small but random distances in a box.

The conductance (in siemens) of a sample equals its conductivity constant times its cross-sectional area

divided by its length.

Kohlrausch's law states that at low concentrations the molar conductivities of strong electrolytes vary

linearly with the square root of the concentration.

The Grotthuss mechanism describes the motion of a proton involving the rearrangement of bonds in a

group of water molecules.

The transport number is the fraction of total current carried by the ions of a given type.

The Debye-Huckel-Onsager theory quantifies electrophoretic effects.

The Green-Kubo relation expresses a transport property in term sof teh fluctuations in microscopic

properties of a system.

The Einstein relation links the molar conductivity of an electrolyte to the diffusion coefficients of its ions.

The Nernst-Einstein equation determines ionic diffusion coefficients from conductivity measurements;

Walden's rule.

The Einstein-Smoluchowski equation states that the diffusion constant equals the square of the step length

of a one-dimensional random walk divided by twice the time.


The Rates of Chemical Reactions

Reaction progress may be monitored using real-time, quenching, flow, stopped-flow, and flash photolysis


The reaction order is the sum of the orders for each component.

Reaction orders (from integrated rate laws):

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)

Half-life decay: t1/2 = ln(2)/k

The Arrhenius equation relates rate constant to activation energy, temperature, and collision frequency; 

k = Ae-Ea/RT

The steady-state approximation assumes that after an initial induction period and the major part of the

reaction, the rates of change of concentration of all reaction intermediates are negligibly small.

In the Michaelis-Menten mechanism of enzyme action, the rate depends on enzyme concentration even

though it undergoes no net change;  E + S = ES --> P + E

A Lineweaver-Burk plot of the reciprocal of the reaction rate against the reciprocal of the substrate

concentration in a Michaelis-Menten mechanism gives the rate of the second step and the

Michaelis constant.

The Lindemann-Hinshelwood mechanism describes unimolecular reactions.

The Rice-Herzfeld mechanism describes chain reactions.

The Lotka-Volterra mechanism describes oscillating reactions at steady-state.

Oscillating reactions, such as the brusselator and the oregonator, must be far from equilibrium, have

autocatalytic steps, and be able to exist in two steady states (bistability).


Processes at Solid Surfaces

The Auger effect is the emission of a second electron after high-energy radiation has expelled another.

Adsorption is the attachment of particles to a surface.

In scanning tunnelling microscopy, a platinum-rhodium or tungsten needle is scanned across the surface of

a conducting solid.

In atomic force microscopy, a sharpened stylus attached to a beam is scanned across the surface.

In physisorption, there is a van der Waals interaction between the adsorbate and substrate.

In chemisorption, molecules stick to the surface by forming a chemical bond.

In accommodation, a molecule bouncing on the surface will lose its energy and eventually adsorb to it.

The BET (Brunauer-Emmett-Teller) isotherm deals with multilayer adsorption.

The Temkin isotherm supposes adsorption enthalpy changes linearly with pressure; the Freundlich isotherm


In the Eley-Rideal mechanism of a surface-catalysed reaction, a gas-phase molecule collides with another

molecule already adsorbed on the surface.


Dynamic Electrochemistry

Current density is the charge flux through a region.

The interface at an electrode can be modeled as an electrical double layer, including Helmholtz and Gouy-

Chapman, and combined Stern model.

The surface potential is the difference between the Volta (outer) and Galvani (inner) potentials.

The Butler-Volmer equation relates the current density to the Galvani potential difference.

Voltammetry measures the current as the potential of the electrode is changed to study the kinetics of

electrode processes, including linear-sweep, differential pulse, and cyclic methods.

Primary voltaic cells cannot be recharged; secondary voltaic cells (including fuel cells) may be.

Corrosion is a redox process by which metals are oxidized by oxygen in the presence of moisture.

Galvanizing is the coating of an iron object with zinc to prevent corrosion.