Significant Figures – start at the left and proceed to the right 1.If the number does not have a decimal point count until there are no more non zero numbers 2.If the number has a decimal point start counting at the first non-zero number and continue counting until you run out of decimal places Vocabulary 1.Observation 2.Hypothesis 3.Experiment 4.Theory 5.Law 6.Chemistry 7.Matter 8.Energy 9.Chemical Properties 10.Physical Properties 11.Extensive Properties 12.Intensive Properties 13.Scientific (natural) law 14.Anion 15.Cation 16.Molecular Geometry 17.Accuracy 18.Precision 19.Exact numbers 20.Stoichiometry 21.Ionic Compound 22.Covalent Compound 23.Inorganic Compound 24.Organic Compound 25.Law of Conservation of Mass 26.Law of Conservation of Energy 27.Law of Definite Proportions 28.Law of Multiple Proportions 29.compounds 30.Molecules 31.Isotopes 32.Atomic number 33.Atomic Mass 34.chemical formula 35.empirical formula 36.molecular formula 37.structural formula 38.bond line formula 39.ball and stick model 40.space filling model 41.mole 42.Electronic Geometry 43.percent weight 44.percent error 45.percent composition 46.percent yield 47.%RSD 48.limiting reactant 49.VSEPR 50.Valence Bond Theory 51.Stoichiometric Coefficient 52.Electron Affinity 53.Electronegativity 54.Covalent Bond 55.Ionic Bond 56.Dipole 57.London Dispersion Forces 58.Resonance 59.Hybrid orbital 60.area of high electron density 61.Aufbrau Principle 62.Hund’s Rule 63.Pauli Exclusion Principle 64.Valence Electrons 65.Atomic Radii 66.Ionic Radii 67.Ionization energy 68.Oxidation Number 69.Formal Charge 70.Polarizability 71.Dipole Moment 72.Polar Covalent Bond 73.Coordinate Covalent Bond 74.Molecular Geometry

MonovalentDivalentTrivalent HydroniumH3O+H3O+ MagnesiumMg 2+ AluminiumAl 3+ (or hydrogen)H+H+ CalciumCa 2+ Antimony IIISb 3+ LithiumLi + StrontiumSr 2+ Bismuth IIIBi 3+ SodiumNa + BerylliumBe 2+ PotassiumK+K+ Manganese IIMn 2+ RubidiumRb + BariumBa 2+ CesiumCs + ZincZn 2+ FranciumFr + CadmiumCd 2+ SilverAg + Nickel IINi 2+ AmmoniumNH 4 + Palladium IIPd 2+ ThaliumTl + Platinum IIPt 2+ Copper ICu + Copper IICu 2+ Mercury IIHg 2+ Mercury IHg 2 2+ Iron IIFe 2+ Iron IIIFe 3+ Cobalt IICo 2+ Cobalt IIICo 3+ Chromium IICr 2+ Chromium IIICr 3+ Lead IIPb 2+ Tin IISn 2+ Table of Common Ions Common Positive Ions (Cations)

MonovalentDivalentTrivalent HydrideH-H- OxideO 2- NitrideN 3- FluorideFl - PeroxideO 2 2- ChlorideCl - SulfideS 2- BromideBr - SelenideSe 2- IodideI-I- OxalateC 2 O 4 2- HydroxideOH - ChromateCrO 4 2- PermanganteMnO 4 - DichromateCr 2 O 7 2- CyanideCN - TungstateWO 4 2- ThiocynateSCN - MolybdateMoO 4 2- AcetateC2H3O2-C2H3O2- tetrathionateS 4 O 6 2- NitrateNO 3 - ThiosulfateS 2 O 3 2- BisulfiteHSO 3 - SulfiteSO 3 2- BisulfateHSO 4 - SulfateSO 4 2- BicarbonateHCO 3 - CarbonateCO 3 2- Dihydrogen phosphateH 2 PO 4 - Hydrogen phosphateHPO 4 2- PhosphatePO 4 3- NitriteNO 2 - AmideNH 2 - HypochloriteClO - ChloriteClO 2 - ChlorateClO 3 - PerchlorateClO 4 - Table of Common Ions Common Negative Ions (Anions)

mass of molecule Molar Mass given or calculated from periodic table Mass of element, or reactant or product Number of atoms, or molecules of reactant or product Avogadro's Number Number of molecules Molar Ratio moles of element, or other reactant or product moles of molecule Avogadro's Number Calculate from molecular formula or balanced equation Molar Mass given or calculated from periodic table Given or determined from balanced stoichiometric equation Vol solution density Concentration solution molarity, ppm, molality, normality, etc. These concepts lead to solving problems determining limiting reactant and percent yield.

The principal quantum number has the symbol ~ n which defines the energy of the shell n = 1, 2, 3, 4, “shells” The angular momentum quantum number has the symbol ~ which defines the subshells. = 0, 1, 2, 3, 4, 5, (n-1) = s, p, d, f, g, h, (n-1) The symbol for the magnetic quantum number is m which defines the orbital. m = -, (- + 1), (- +2),.....0, , ( -2), ( -1), The last quantum number is the spin quantum number which has the symbol m s which characterizes the single electron. The spin quantum number only has two possible values. m s = +½ or -½ one spin up ↑ and one spin down↓ Quantum Numbers n and define the energy of the electron The Nucleus: Build by adding the required number of protons (the atomic number) and neutrons (the mass of the atom) Pauli’s Exclusion Principle states that paired electrons in an orbital will have opposite spins. Electrons: Hund’s Rule states that each orbital will be filled singly before pairing begins. The singly filled orbitals will have a parallel spin. Fill the electrons in starting with the lowest energy level adhering to Hund’s and Pauli’s rules.

IonicPolar Covalent Covalent Determine Inductive effect Count the number of electrons the element should have Determine how equally electrons are shared (  EN) >1.7 consider it ionic Oxidation number Formal charge Never Have a Full Octet Always Have a Full Octet Sometimes Have a Full Octet Sometimes Exceed a Full Octet To calculate a formal charge 1.draw the Lewis dot structure 2.draw circles around each atom and the electrons associated with it. Remember that formal charges are associated with covalent bonds and that all electrons are shared equally. 3.compare to the group number for that atom. If the number is larger the formal charge is negative, smaller the formal charge is positive. To calculate an oxidation number 1.list all the elements follow with an equal sign 2.follow with the number of atoms of that type in the molecule 1.follow with a multiplication sign 2.If the element is O follow with a -2 3.If the element is H follow with a +1 4.any other element enter a ? 5.follow with an = sign, do the math 6.draw a total line, then enter the charge on the molecule 7.Do the algebra backwards to solve for ?

Summary of Electronic & Molecular Geometries Regions of High Electron Density Electronic GeometryHybridization 2Linearsp 3Trigonal planarsp 2 4Tetrahedralsp 3 5Trigonal bipyramidalsp 3 d 6Octahedralsp 3 d 2 VSEPR Theory Lone pair to lone pair is the strongest repulsion. 2Lone pair to bonding pair is intermediate repulsion. 3Bonding pair to bonding pair is weakest repulsion. Mnemonic for repulsion strengths lp/lp > lp/bp > bp/bp Lone pair to lone pair repulsion is why bond angles in water are less than o. Electronic geometry Electronic geometry is determined by the locations of regions of high electron density around the central atom(s). Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used. Molecular geometry Molecular geometry determined by the arrangement of atoms around the central atom(s).

Isomers structural isomers constitutional isomers stereo isomers racemic mixture entantiomers geometric isomers positional isomers chiral molecules chiral centers optical isomers cis mer trans fac hydration isomers ionization isomers coordination isomers linkage isomers hydrocarbons unsaturated hydrocarbons saturated hydrocarbons alkanes alkenes alkynes aromatic compounds alkyls phenyls phenols alcohols esters ethers carbonyl groups aldehydes ketones carboxylic acids acyl chlorides organic halides amines amides resonance Arrhenius acids/bases Brönsted/Lowery acids/bases Lewis acids/bases Electrolytes Non electrolytes sugarsfats polymerssolution solventsolute concentrationmolarity ppmppb wt%vol% molecular equations ionic equations net ionic equations spectator ion metathesis reaction combination reaction decomposition reaction displacement reaction redox reaction addition polymerization condensation polymerization ligand donor atom unidentate polydentate chelate coordination number coordination sphere titration titrant primary standard secondary standard end point equivalence point pH oxidation numbers

Naming Saturated Hydrocarbons 1.Choose the longest continuous chain of carbon atoms which gives the basic name or stem. 2Number each carbon atom in the basic chain, starting at the end that gives the lowest number to the first group attached to the main chain (substituent). 3For each substituent on the chain, we indicate the position in the chain (by an Arabic numeric prefix) and the kind of substituent (by its name).  The position of a substituent on the chain is indicated by the lowest number possible. The number precedes the name of the substituent. 4When there are two or more substituents of a given kind, use prefixes to indicate the number of substituents.  di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, hepta = 7, octa = 8, and so on. 5The combined substituent numbers and names serve as a prefix for the basic hydrocarbon name. 6Separate numbers from numbers by commas and numbers from words by hyphens.  Words are "run together".

Alcohols and Phenols The stem for the parent hydrocarbon plus an -ol suffix is the systematic name for an alcohol. A numeric prefix indicates the position of the -OH group in alcohols with three or more C atoms. Common names are the name of the appropriate alkyl group plus alcohol. Ethers Common names are used for most ethers. Aldehydes and Ketones Common names for aldehydes are derived from the name of the acid with the same number of C atoms. IUPAC names are derived from the parent hydrocarbon name by replacing -e with -al. The IUPAC name for a ketone is the characteristic stem for the parent hydrocarbon plus the suffix -one. A numeric prefix indicates the position of the carbonyl group in a chain or on a ring. Amines Amines are derivatives of ammonia in which one or more H atoms have been replaced by organic groups (aliphatic or aromatic or a mixture of both). There are three classes of amines. Carboxylic Acids IUPAC names for a carboxylic acid are derived from the name of the parent hydrocarbon. –The final -e is dropped from the name of the parent hydrocarbon –The suffix -oic is added followed by the word acid. Many organic acids are called by their common (trivial) names which are derived from Greek or Latin.

Priority Functional groupFormulaPrefixSuffix 1 Cations e.g. Ammonium –NH 4 + -onio- ammonio- -onium -ammonium 2Carboxylic acids–COOHcarboxy--oic acid* 3 Carboxylic acid derivatives Esters Acyl chlorides Amides –COOR –COCl –CONH 2 R-oxycarbonyl- chloroformyl- carbamoyl- -oyl chloride* -amide* 4 Nitrites Isocyanides –CN –NC cyano- isocyano- -nitrile* isocyanide 5 Aldehydes Thioaldehydes –CHO –CHS formyl- thioformyl- -al* -thial* 6 Ketones Thioketones >CO >CS oxo- thiono- -one -thione 7 Alcohols Thiols –OH –SH hydroxy- sulfanyl- -ol -thiol 8Amines–NH 2 amino--amine 9 Ethers Thioethers –O– –S– -oxy- -thio- When compounds contain more than one functional group, the order of precedence determines which groups are named with prefix or suffix forms. The highest precedence group takes the suffix, with all others taking the prefix form. However, double and triple bonds only take suffix form (-en and -yn) and are used with other suffixes.

Carbon Atom HybridizationC usesC formsExample sp 3 tetrahedral4 sp 3 hybrids 4  bonds CH 4 sp 2 trigonal planar3 sp 2 hybrids & 1p orbital 3  bonds 1  bond C2H4C2H4 sp linear2 sp hybrids & 2 p orbitals 2  bonds 2  bonds C2H2C2H2

Nomenclature Rules for Naming Complex Species 1.Cations (+ ions) are named before anions (- ions). 2.Coordinated ligands are named in alphabetical order. –Prefixes that specify the number of each kind of ligand (di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, etc.) are not used in alphabetizing –Prefixes that are part of the name of the ligand, such as in diethylamine, are used to alphabetize the ligands. 3.For complicated ligands, especially those that have a prefix such as di or tri as part of the ligand name, these prefixes are used to specify the number of those ligands that are attached to the central atom. –bis = 2 tris = 3 tetrakis = 4 pentakis = 5 hexakis = 6 4.The names of most anionic ligands end in the suffix -o. –Examples of ligands ending in –o are: Cl - chloro S 2- sulfido O 2- oxo 5.The names of most neutral ligands are unchanged when used in naming the complex. –There are several important exceptions to this rule including: NH 3 ammine H 2 Oaqua 6.The oxidation number of a metal that exhibits variable oxidation states is designated by a Roman numeral in parentheses following the name of the complex ion or molecule. 7.If a complex is an anion, the suffix "ate" ends the name.  No suffix is used in the case of a neutral or cationic complex.  Usually, the English stem is used for a metal, but if this would make the name awkward, the Latin stem is substituted. ferrate instead of ironate plumbate instead of leadate

Ion/MoleculeNameName as a Ligand NH 3 ammoniaammine COcarbon monoxidecarbonyl Cl - chlorideChloro CN - cyanidecyano F-F- fluoridefluoro OH - hydroxidehydroxo NOnitrogen monoxidenitrosyl NO 2 - nitritenitro PH 3 phosphine

SystemAcid (HCl)Base (NaOH) Arrhenius Brönsted- Lowry Lewis

∆H = H final - H initial The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C 5 H 12 reacts with 8 mol of O 2 to produce 5 mol of CO 2, 6 mol of H 2 O, and releasing 3523 kJ is referred to as one mole of reactions. ∆H o rxn =  ∆H f o (prod) -  ∆H f o (react) Specific heat capacity (J/(g∙K)= heat lost or gained by system (Joules) mass(grams)  T (Kelvins) m  T f –T i ) q c P = Variabl e System 1 System 2 CpCp TfTf TiTi m q

heat transfer out (exothermic), -q heat transfer in (endothermic), +q SYSTEM ∆E = q + w w transfer in (+w) w transfer out (-w)

Heat Energy Internal energy Kinetic Energy Potential Energy Endothermic Exothermic Thermodynamics Thermal Equilibrium System Surroundings Law of Conservation of Energy Heat Capacity Specific Heat Capacity First Law of Thermodynamics Melting Freezing Deposition Sublimation Evaporation Condensation State Function Standard state temperature Standard state pressure Standard states matter Enthalpy Hess’s Law Thermochemical Equation Enthalpy of Formation

Standard P  atm or kPa Standard T  K or 0.00 o C K = o C 1 mm Hg = 1 torr 760 torr = 1 atm The standard molar volume is 22.4 L at STP PV = nRT R = L atm mol -1 K -1 P total = P A + P B + P C At low temperatures and high pressures real gases do not behave ideally. The reasons for the deviations from ideality are: 1.The molecules are very close to one another, thus their volume is important. 2.The molecular interactions also become important. VariableCond. 1 Cond. 2 P (atm) V (L) N (moles) R (L atm mol -1 K -1 ) T (K)

The Kinetic-Molecular Theory The basic assumptions of kinetic-molecular theory are: Postulate 1 –Gases consist of discrete molecules that are relatively far apart. –Gases have few intermolecular attractions. –The volume of individual molecules is very small compared to the gas’s volume. Proof - Gases are easily compressible. Postulate 2 –Gas molecules are in constant, random, straight line motion with varying velocities. Proof - Brownian motion displays molecular motion. Postulate 3 –Gas molecules have elastic collisions with themselves and the container. –Total energy is conserved during a collision. Proof - A sealed, confined gas exhibits no pressure drop over time. Postulate 4 –The kinetic energy of the molecules is proportional to the absolute temperature. –The average kinetic energies of molecules of different gases are equal at a given temperature. Proof - Brownian motion increases as temperature increases.