General Chemistry Reference Sheet

This reference sheet addresses some of the more peculiarpieces of information that need to be memorized in a gen-eral chemistry course. It also contains a simple set of es-sential formulas in chemistry with cautions, explanations,and general tips.This sheet is meant to be as concise as possible, and manyinformation in the textbook is left out in favor of cautionsand tips. This sheet is, therefore, best used as a supple-ment to, not a replacement of, the textbook.

SI Fundamental UnitsMass Kilogram (kg)Length Meter (m)Time Second (s)Temperature Kelvin (K)Amount of substance Mole (mol)Electric current Ampere (A)Luminous intensity Candela (cd)

Atomic Experiments and ModelsJ. J. Thomson Discovered e−; Cathode ray

Plum pudding modelR. A. Millikan Measured charge of e−; Oil dropH. Becquerel/M. Curie Discovered radioactivityE. Rutherford Discovered α, β, and γ rays

Discovered nucleus; Gold foil experimentJ. Chadwick Discovered neutronsN. Bohr Bohr model (electron orbits)Quantum mechanicists Quantum model

Polyatomic IonsNH4

+ ammonium OH− hydroxideCN− cyanide C2O4

2− oxalateO2

2− peroxide CNO− cyanateHSO4

− hydrogen sul-fate

C2H3O2− acetate

SCN− thiocyanate NO3− nitrate

SO32− sulfite ClO4

− perchlorateCO3

2− carbonate ClO3− chlorate

PO43− phosphate ClO− hypochlorite

S2O32− thiosulfate HPO4

2− hydrogenphosphate

CrO42− chromate H3O+ hydronium

Cr2O72− dichromate PO3

3− phosphiteMnO4

− permanganate Hg22− mercury (I)

N3− azide C2

2− carbideC4H4O6

2− tartrate S22− disulfide

O2− superoxide AsO3

3− arsenitePO2

3− hypophosphite AsO43− arsenate

SiO32− silicate P2O7

4− pyrophosphate

Ionic Solubility ChartSoluble ExceptionsNO3

− NoneCH3COO− NoneCl− Ag+, Hg2

2+, Pb2+

Br− Ag+, Hg22+, Pb2+

I− Ag+, Hg22+, Pb2+

SO42− Sr2+, Ba2+, Hg2

2+, Pb2+

Insoluble ExceptionsS2− NH4

+, alkali metal cations, Ca2+,Sr2+, Ba2+

CO32− NH4

+, alkali metal cationsPO4

3− NH4+, alkali metal cations

OH− NH4+, alkali metal cations, Ca2+,

Sr2+, Ba2+

Strong Acids and BasesStrong acids and bases dissociate in water completely.

Strong Acids Strong BasesHCl HClO4 Alkali metal hydroxidesHBr HNO3 Ca(OH)2HI H2SO4 Sr(OH)2HClO3 Ba(OH)2

Activity SeriesMetals below H+cannot react with acids to form H2. Moreactive metals are better reducing agents.From most active to least active:Li+, K+, Ba2+, Ca2+, Na+, Mg2+, Al3+, Mn2+, Zn2+,Cr3+, Fe2+, Co2+, Ni2+, Sn2+, Pb2+, H+, Cu2+, Ag+,Hg2+, Pt2+, Au3+

Flame ColorsCalcium Brick redCopper (I) BlueCopper (II) Green or blue-greenPotassium LilacLithium Dark redSodium Bright yellowStrontium RedBarium Light greenIron (III) GoldCesium Blue–VioletIndium BlueLead BlueRubidium Red–Violet

Phase ChangesFrom solid From liquid From gas

To solid - freezing depositionTo liquid melting - condensationTo gas sublimation vaporization -

Solution ColorsCopper (II) BlueNickel GreenPermanganate PurpleChromate YellowDichromate OrangeIron (II) Light blueIron (III) Rusty yellow

Thermodynamic LawsFirst Law: Energy cannot be created nor destroyed. Itcan only be transferred in the form of either heat or work.Second Law: Any spontaneous reaction increases the en-tropy of the universe.Third Law: An ideal solid crystal at 0 K has an entropyof 0.

Thermodynamic FormulasStandard thermodynamic conditions 298 K; 1 atm; 1 MKinetic energy K = mv2/2Electrostatic potential energy UE = (kCQ1Q2)/dInternal energy ∆E = q + wEnthalpy H = E + PVSpecific heat s = q/(m ·∆T )Entropy in reversible reaction ∆Ssystem = (∆H)/T

∆Ssurrounding = −(∆H)/TMicrostate-entropy relationship S = k lnWGibbs free energy G = H − TSGibbs free energy change ∆G = ∆H − T∆S

∆G = ∆G◦ +RT lnQHess’s Law ∆Htotal = Σ∆Hi

ConstantsBoltzmann’s constant kB = 1.381× 10−23 m2kg · s−2K−1

Coulomb’s constant kC = 1/(4πε0) = 8.988× 109 J ·m/C2

Avogadro’s number NA = 6.022× 1023 mol−1

Faraday’s constant F = 9.649× 104 C/molPlanck’s constant h = 6.626× 10−34 J·sIdeal gas constants R = 0.0821 (L · atm)/(mol ·K)

R = 8.314 J/(mol ·K)Vacuum permittivity ε0 = 1/(µ0c

2) = 8.854× 10−12 F/mVacuum permeability µ0 = 1.257× 10−6 N·A−2

Atomic mass 1 amu = 1.661× 10−24 gElectron charge e = 1.602× 10−19 CElectronvolt 1 eV = 1.602× 10−19 JAtmospheric pressure 1 atm = 1.013× 105 PaAbsolute zero 0 K = -273.15 ◦CSpeed of light in vacuum c = 2.998× 108 m/s

Quantum Mechanical FormulasEnergy of a quantum E = hνWavelength-frequency relationship c = ν · λProbability distribution PV =

∫∫∫V

|ψ(x, y, z)|2dxdydz

Laws of Quantum MechanicsHeisenberg’s Uncer-tainty Principle

∆x ·∆p ≥ h/2

Corollary: It is impossible to de-termine both the position andthe momentum for a sufficientlysmall particle like an e−.

Pauli Exclusion Prin-ciple

No two e−in an atom can sharethe same set of four quantumnumbers.Corollary: A suborbital can holda maximum of 2e−.

Hund’s Rule Energy is the lowest when thenumber of e− with the same spinis maximized.Corollary: e− will first half-fillall the empty suborbitals, thengo back and fill the half subor-bitals.

Quantum Numbers of e−

Principal (n) The energy shell of the e−, e.g. 4 in4d1.

Azimuthal (l) The suborbital shape, with s=0,p=1, d=2, f=3, e.g. 2 in 4d1.

Magnetic (ml) The suborbital, ranging from −l tol, e.g. −2 in 4d1.

Spin (ms) The spin of e−. Two e−in the samesuborbital has either −1/2 or 1/2.

Molecular GeometryHybridization Nonbonding Geometry

electrons

sp 0 linear

sp2 0 trigonal planar1 bent

sp3 0 tetrahedral1 trigonal pyramidal2 bent

sp3d 0 trigonal bipyramidal1 seesaw2 T-shaped3 linear

sp3d2 0 octahedral1 square pyramidal2 square planar

Atomic PropertiesAtomic size 1

2the distance between two ad-

jacent nucleii.Ionic size Cations are smaller than their

parent atoms. Anions arelarger than their parent atoms.

N-th ionization energy The energy required to re-move the n-th electron from aground state gaseous atom.

Electron affinity The energy released by addingan electron to a gaseous atom.

Metallic character The qualities of a metal.Metals are shiny and heat andelectricity-conducting; theyhave malleble solid form, formbasic ionic oxides, and tendto form cations in an aqueoussolution.

Periodic PropertiesProperty Left to Right Top to BottomAtomic size Decreasing IncreasingIonization energy Increasing DecreasingElectron affinity Large if adding

to a previouslyempty orbital

No apparentchange

Metallic character Decreasing Increasing

Types of Crystalline SolidsType IMFs PropertiesMolecular Van der Waals forces,

dipole-dipole inter-actions, hydrogenbonds

Soft, low meltingpoint, poor conduc-tion

Covalent-network

Covalent bonds Very hard, highmelting point, poorconduction

Ionic Electrostatic interac-tions

Hard, high meltingpoint, poor conduc-tion

Metallic Metallic bonds Soft to very hard,low to very highmelting point, ex-cellent conduction

Boiling Points of Molecular CompoundsThe relative boiling points of molecular compounds can bedetermined by their IMFs. The stronger the IMFs are,the higher the boiling point. (Note that linear compoundslike straight-chain hydrocarbons have higher van der Waalsforces than non-linear compounds because their moleculeshave a greater area of contact.)

IMFs in Molecular SolidsLondon dispersionforce (van der Waalsforces, induced dipole-dipole interactions)

Interactions between dipolespartially charged through themovement of shared electron.Presents in all compounds.Weakest of the three.

Dipole-dipole interac-tions

Interactions between dipolespartially charged through theelectronegativity difference oftwo bonding atoms.

Hydrogen bonds A special kind of dipole-dipoleinteractions present in com-pounds that have hydrogenand either oxygen or nitrogen.Strongest of the three.

Acid-Base TheoriesArrhenius Brønsted-

LowryLewis

Acids [H+] >[OH−] Protondonors

Electronacceptors

Bases [OH−] >[H+] Protonacceptors

Electrondonors

Acids have H+ Hydrogenatom

Electronacceptingatom

Bases have OH− Unsharedelectron pair

Unsharedelectronpair

Acid + Base→

Salt + H2O Conjugateacid +Conjugatebase

Kinetic Molecular Theory (KMT)There is a very large number of particles;Particles are in constant random motion and collide con-stantly with the wall;Collisions of particles with the wall are perfectly elastic;Particles exert no force upon each other.

Properties of SolutionsSolvation The uniform dispersion of a solute in a solvent.Hydration Solvation in water.Crystallization The reverse reaction of solvation.Saturated A solution in equilibrium.Unsaturated A solution with less solute than saturation.Supersaturated A solution with more solute than

saturation. (Will undergo crystallizationif a crystal seed is present.)

Miscible Two liquids that dissolve in any proportion.Henry’s Law S ∝ P (S: solubility)

2

Colligative PropertiesPhysical properties of a solution that depends on the con-centration of solutes. More solutes will lead to:

1. Lower vapor pressure: PA = XAP◦A (Raoult’s

Law)

2. Higher boiling point: Tb = T ◦b + kbm (Molality)

3. Lower freezing point: Tf = T ◦f − kfm (Molality)

4. Higher osmotic pressure: π = RT ·M (Molarity)

Colligative properties also depend on the van ’t Hoff factor(i = number of particles after reaction / number of parti-cles before reaction). The greater i is, the more colligativeproperties it exerts on the solution.

Reaction RateThe reaction rate r = d[X]/dt can be determined from thereaction by the rate law

r = k[A]a[B]b...

Where a, b, etc. are reaction orders for the reactants. Re-action orders can only be determined experimentally, be-cause reactions will in theory go through several steps, theslowest of which is the rate-determining step. Reaction or-der is determined by the number of atoms participating inthe rate-determining step. The sum of these orders is theoverall order.Concentration function [X] can be determined as

[X]t =

∫ t

0

rdτ + [X]0

Therefore, for first-order reactions,

ln[X]t = −kt+ ln[X]0

Graphically, t is proportional to ln[X]t with the slope −k.And for second-order reactions,

1

[X]t= kt+

1

[X]0

Graphically, t is proportional to 1/[X]t with the slope k.

Reaction Half-timeThe half-time of a reaction is the amount of time neededto consume half of the reactants. It is denoted t1/2.For first order reactions, t1/2 ≈ 0.693/k. For second orderreactions, t1/2 = 1/(k[X]0).

Activation EnergyCollision model Reactions occur as a result of collisions

between molecules.Activation energy (Ea) The minimum energy required

for a reaction to occur.Arrhenius equation ln k = lnA− Ea/RT

(This means that ln k ∝ 1/T )Activation energy is lowered when a catalyst is present. In-organic catalysts usually provide a site on which reactantscan adsorb; organic catalysts, or enzymes, bind specific tosubstrate molecules (“lock-and-key”).

Spectrophotometry of ConcentrationBeer’s law, A = εlc, relates concentration and light absorp-tion.Absorbance (A) − log10(I/I0) (liquids)

− ln(I/I0) (gases)Absorption coefficient (ε) Depends on the solution.Length of path (l) The length of the path travelled

by light.Concentration (c) The concentration of the solution.In spectrophotometry, the length of path is fixed. There-fore, when using the same solution, A ∝ c.

Gas LawsSTP 273 K; 1 atmBoyle’s Law P ∝ 1/VCharles’s Law P ∝ TAvogadro’s Law P ∝ nIdeal Gas Equation PV = nRTLaw of Partial Pressure Pn = XnPt

Effusion Rate u =√

(3RT )/MGraham’s Law u1/u2 =

√M2/M1

Density Formula d = (PM)/(RT )Deviation from Ideal Behavior (PV )/(RT )Van der Waals Equation (

P + (n2a)/V 2)

(V − nb) = nRTClausius-Clapeyron Equation lnP = −∆Hvap/(RT ) + C

Caution: When using the effusion rate formula, the Rvalue must be in joules (8.314), and the M value must beconverted to kg/mol.

Equilibrium FormulasIon-product constant of water

Kw = [H+][OH−] = 1.0× 10−14 (278 K)Henderson-Hasselbalch equation

pH = pKa + log([base]/[acid])Van ’t Hoff equation d(lnK)/dT = (∆H◦)/(RT 2)

lnK = −∆H◦/(RT ) + ∆S◦/R

ConcentrationNotation DefinitionMolarity (M) (Moles solute)/(Liters solution)Molality (m) (Moles solute)/(Kilogram solvent)Mole fraction (X) (Moles solute)/(Moles solution)Mass percentage (Mass of solute)/(Mass of solution)Volume percentage (Volume of solute)/(Volume of solu-

tion)Caution: It is an extremely common mistake to confusemolarity with molality. Check your R’s and L’s!

Le Chatelier’s PrincipleAn equilibrium reaction will spontaneously balance an out-side effect added to it. For example,Change in amount of reactants or products: Thereaction will consume more of the substance in excess tobalance the change;Change in volume or pressure: The reaction will formmore gas if volume increases or if pressure decreases, andwill form less gas if volume decreases or if pressure in-creases;Change in temperature: Endothermic reactions willshift left for lower temperatures and shift right for highertemperatures. Exothermic reactions will shift right forlower temperatures and shift left for higher temperatures.

Equilibrium ConstantFor reactions in a solution, the equilibrium constant of areaction

aA+ bB ⇀↽ sS + tT

Is defined as

Kc =[S]s[T ]t

[A]a[B]b

When the reaction is in equilibrium.If the reaction is a equilibrium between a solid and its ionsin solution, then Kc is its solubility product constant, Ksp.If the reaction is a dissociative reaction of a weak acid, thenKc is its acid dissociation constant, Ka. Polyprotic (hav-ing more than one H) acids have multiple Ka, but usuallyKa1 determines the pH.For a conjugate acid-base pair, Ka ·Kb = Kw.If the reactants are gases, then the equilibrium constant isdefined as

Kp =P sSP

tT

P aAP

bB

When the reaction is in equilibrium.The Formulas above, when applied to non-equilibrium sit-uations, gives Q. The reaction forms products if Q < K,reactants if Q > K, and nothing if Q = K (already inequilibrium).

3

Equilibrium Constant (cont.)For all equilibrium reactions, there are more reactants thanproducts if K < 1, more products than reactants if K > 1,and the same amount of reactants and products if K = 1.

Acid Character of Hydrogen AtomsHydrogen atoms are acidic when they are weakly bonded,and when the molecule/atom they are bonded to formsstable anions.In organic compounds, the hydrogen atoms in carboxylgroups (COOH) are usually the most acidic.

Indicators for Acid-Base TitrationIndicator Small pH Color change Large pHMethyl violet Yellow 0.0–1.6 VioletBromophenolblue

Yellow 3.0–4.6 Blue

Methyl orange Red 3.1–4.4 YellowMethyl red Red 4.4-6.2 YellowLitmus Red 5.0–8.0 YellowBromothymolblue

Yellow 6.0–7.6 Blue

Phenolphthalein Colorless 8.3–10.0 Pink

Oxidation and ReductionOxidation ReductionLoss of electrons Gains electronsOxidation num. increases Oxidation num. decreasesOccurs at anode Occurs at cathode

Mnemonic devices:

• “OIL RIG” (Oxidation Is Loss, Reduction Is Gain)

• “What an ox loses, a red cat gains” (An =anode; ox = oxidation; red = reduction; cat = cath-ode)

Electrochemical FormulasElectromotive force E = −(∆G)/(nF )Nernst equation E = E◦ − (RT/nF ) lnQ

E = E◦ − (0.0592/n) logQStandard cell potential E◦ = E◦

redcathode− E◦redanode

Energy of a charged particle E = qVFaraday’s Law of Electrolysis m = (Q/F )(M/z)

Types of Magnetic MaterialsParamagnetic: Can be magnetized to attract externalmagnetic fields, but cannot retain magnetism. Param-agnetic materials have a magnetic permeability of morethan µ0. They usually have free electrons, especially dand f electrons. Their magnetization follows Curie’s Law(M = C · B/T ). Examples of paramagnets are tungstenand cesium.

Types of Magnetic Materials (cont.)Diamagnetic: Can be magnetized to repulse externalmagnetic fields, but cannot retain magnetism. Diamag-netic materials have a magnetic permeability of less thanµ0. Examples of diamagnets are bismuth and antimony.Ferromagnetic: Can be magnetized and retain mag-netism. Ferromagnetism depends both on the chemicalcomposition and the structure of the material (iron is aferromagnet, while stainless steel is not). Examples of fer-romagnets include cobalt and iron.

Nuclear ChemistryAlpha particles (α) Helium nuclei (42He)Beta particles (β−) Electrons (0−1e)Positrons (β+) Antielectrons (01e)Gamma radiation (γ) High energy radiation (00γ)Units of radioactivity SI: Becquerel (Bq): 1 nucleus/s(Disintegration per second) Curie (Ci): 3.7 ×1010 nuclei/sUnits of absorbed radiation SI: Gray (Gy): 1 J/kg(Energy per kilogram tissue) Rad: 0.01 Gy

MetallurgyMetallurgy is the extraction of minerals from ores.Pyrometallurgy: The use of heat to convert ores to met-als. (Example: Production of iron)Hydrometallurgy: The use of chemical processes in a so-lution to separate a metal from its ore. (Example: Bayerprocess for producing aluminum)Electrometallurgy: The use of electrochemical processesto separate a metal. (Example: Hall process for producingaluminum)

HydrocarbonsName Common Formula HybridizationAlkane CnH2n+2 sp3

Cycloalkane CnH2n sp3

Alkene CnH2n sp2

Alkyne CnH2n−2 spAromatic CnH2n−6 sp2

In a hydrocarbon with n carbons, the number of hydrogensis 2n+ 2, minus 2 for each π bond or carbon ring.

StereoisomerismStereoisomerism occurs at bonds such as C=C, where bothends have two different substituents, because the rotationof these substituents are restricted.Cis-trans isomerism: If both ends have a hydrogenatom substituent, then the compound exhibits cis-transisomerism. The cis-isomer has both hydrogen atoms on thesame side, and the trans-isomer has the hydrogen atomson different sides.

Stereoisomerism (cont.)E/Z isomerism: If the two ends of the bond do not have acommon hydrogen atom, then the compound exhibits E/Zisomerism. The Z isomer has the “larger” substituents(defined by the CIP Rules) of both ends on the same side,while the E isomer has the larger substituents on differentsides.

Cahn-Ingold-Prelog RulesThe CIP Rules are used to compare two substituent groupsin the E/Z and R/S groups of naming isomers.

1. Direct comparison: If the atoms that are di-rectly connected to the stereocenter are different,then the atom with a higher atomic number receiveshigher priority.

2. Tiebreaker I: If there is a tie, then a list of atomstwo bonds away from the stereocenter is compiledfor each of the two substituent groups. The atomswith the greatest atomic number from each list arethen compared. If they tie, then the second greatestatoms from each list are compared. This process isrepeated until the tie is broken.

3. Tiebreaker II: If there is still a tie after consideratoms two bonds away from the center, then atomsthree bonds away are considered in the same way inTiebreaker I. This process is repeated until the tie isbroken.

4. Isotopes: If two groups differ only in isotopes (andare otherwise identical), then mass number is usedinstead of atomic number in the process.

5. Double and triple bonds: If there is a dou-ble bond in the substituent group, then the dou-ble bond is treated as a bond with “ghost atoms”(e.g. R-A=B-R’ is treated as R-(A-B)-(B-A)-R’).Triple bonds, similarly, have two ghost atoms foreach atom.

6. Cycles: To handle a molecule containing one ormore cycles, one must first expand it into a tree(called a hierarchical digraph by the authors) bytraversing bonds in all possible paths starting atthe stereocenter. When the traversal encounters anatom through which the current path has alreadypassed, a ghost atom is generated in order to keepthe tree finite.

4

Criteria of AromaticityIf a hydrocarbon

1. Is cyclic, i.e. possesses a carbon ring;

2. Is planar, i.e. all carbons on the ring are on the sameplane;

3. Has an uninterrupted cloud of π electrons;

4. The number of pairs of π electrons in the cloud isan odd number, i.e. the number of π electrons in thecloud is 4n+ 2;

then the hydrocarbon is aromatic. Aromatic compoundsare highly stable (cannot undergo addition reactions), butcan undergo substitution reactions.

Functional GroupsFunctional Group Name Suffix/PrefixR-OH (hydroxyl) alcohol -olR-O-R’ (ether) ether etherR-X (halo) haloalkane halo-R-NH2 (amino) amine -amineR-COH (aldehyde) aldehyde -alR-COX (haloformyl) acyl halide -oyl halideR-CO-R’ (carbonyl) ketone -oneR-COOH (carboxyl) carboxylic acid -oic acidR-COO(carboxylate) carboxylate -oateR-COO-R’ (ester) ester -oateR-CONH2 (amide) amide -amideR-CNH-R’ (ketimine) ketimine imino-R-CHNH (aldimine) aldimine imino-R-CONCO-R’(imide)

imide imido-

R-N3 (azide) azide azido-R-N2-R’ (azo) azo azo-R-OCN (cyanate) cyanate cyanato-R-NCO (isocyanate) isocyanate isocyanato-R-CN (nitrile) nitrile cyano-R-NC (isonitrile) isonitrile isocyano-R-NO (nitroso) nitroso nitroso-R-NO2 (nitro) nitro nitro-R-ONO (nitrosooxy) nitrite nitrosooxy-R-ONO2 (nitrate) nitrate nitroxy-R-SH (sulfhydryl) thiol -thiolR-SCN (thiocyanate) thiocyanate thiocyanato-R-NCS (isothio-cyanate)

isothiocyanate isothiocyanato-

R-CSH (carbonoth-ioyl)

thial -thial

R-PH3 (phosphino) phosphine -phosphaneR-C6H5 (phenyl, Ph) benzene der. phenyl-

Functional Groups (cont.)Functional Group Name Suffix/PrefixR-CH2C6H5 (benzyl,Bn)

toluene der. benzyl-

R-C5H4N (pyridyl) pyridine der. pyridin-x-ylNote: In actual compounds, change all instances of “halo”above to halogen names (fluoro, chloro, bromo, iodo).

Amino AcidsHydrophobic amino acids:

Name Code Name CodeAlanine Ala Valine ValPhenylalanine Phe Methionine MetLeucine Leu Proline ProIsoleucine Ile Tryptophane Trp

Hydrophilic amino acids:Name Code Name CodeGlycine Gly Threonine ThrSerine Ser Cysteine CysTyrosine Tyr Asparagine AsnGlutamine Gln Arginine ArgLysine Lys Histidine HisAspartic acid Asp Glutamic acid Glu

Protein StructureProteins are large biochemical complexes that contain sev-eral polypeptide compounds (amino acid chains). Theyare organized into four levels of structure:Primary structure: The chain of amino acids that makeup the protein; this chain directly controls the other levelsof protein structure.Secondary structure: The patterns formed by segmentsof the polypeptide chain; can be either α-helices or β-pleated sheets.Tertiary structure: The folding of the polypeptide toproduce a certain shape.Quarternary structure: The geometrical bonding ofseveral polypeptides to form the protein.

ChiralityA molecule possessing a nonsuperimposable mirror imageis chiral.Two mirror images of a chiral molecule are enantiomers.A carbon that is bonded to 4 different groups is an asym-metric center. Chiral molecules have at least one asym-metric centers.Chiral molecules rotate polarized light. Two enantiomersrotate polarized light by the same degrees, one clockwiseand one counterclockwise. A mixture of two enantiomersin 1:1 does not rotate polarized light, and is racemic.

EnantiomerismSystem Name Based OnR/S Structure(+)/(−) Direction of rotation of polarized

lightD/L Enantiomer of glyceraldehyde the

molecule is derived fromR/S notation: Orient the enantiomer so that the small-est (by CIP Rules) substituent points backward (away fromthe viewer) and the largest substituent points upward. Ifthe larger substituent of the other two points toward theright, then the enantiomer is an R-enantiomer. If the largersubstituent points toward the left, then the enantiomer isan L-enantiomer.(+)/(−) notation: An enantiomer that rotates the planeof polarization clockwise is dextrorotary (+). An enan-tiomer that rotates the plane of polarization counterclock-wise is levorotary (−).D/L notation: An enantiomer that is derived from (+)-glyceraldehyde is the D-enantiomer. An enantiomer thatis derived from (−)-glyceraldehyde is the L-enantiomer.Note that nomenclature in a system cannot be determinedby that in another system.Caution: The (+)/(−) system is sometimes written as(d)/(l), which is easily confused with the D, L system. Asthese two systems sometimes conflict (a D-enantiomer canbe an (l)-enantiomer), the (+)/(−) notations are stronglypreferred.

Significant FiguresSignificant figures (“sig figs”) is the number of digits thatcarry precision in a number.Non-measured Numbers: Non-measured numbers,such as π, integer counts, definition of units, etc. alwayshave infinite sig figs. Other constants, such as NA, havelimited sig figs.Non-zero Digits: Nonzero digits are always significant,unless one or more of the other rules are violated.Zeros: Leading zeros are never significant; trailing zeros,however, are significant only if they are part of the mea-surement. Zeros between non-zero digits are always signif-icant.Reporting Numbers: Reported numbers are only signif-icant to the precision of the equipments with which theyare measured.Addition/Subtraction: When adding or subtractingtwo numbers, the result should have as many decimalplaces as the number with the smallest sig figs.Multiplication/Division: When multiplying or divid-ing, the result should have as many sig figs as the numberwith the smallest sig figs.

c©2009–2011 Zee Zuo. Licensed under CC-BY-SA 3.0, United States. Revision 3

5