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 Intramolecular forces Intermolecular forces Hydrogen Bonding Polarization Polarizability Vapor Pressure Equilibrium Heat of Vaporization Phase Diagram Solid Liquid Gas Triple Point Critical Point Super Critical Fluid

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 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

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.Law of Conservation of Mass 18.Law of Conservation of Energy 19.Exact numbers 20.Accuracy 21.Precision 22.compounds 23.molecules 24.chemical formula 25.empirical formula 26.molecular formula 27.structural formula 28.bond line formula 29.ball and stick model 30.space filling model 31.mole 32.Electronic Geometry 33.percent weight 34.percent error 35.percent composition 36.percent yield 37.%RSD 38.limiting reactant 39.Stoichiometry 40.Stoichiometric Coefficient 41.Electron Affinity 42.Electronegativity 43.Covalent Bond 44.Ionic Bond 45.Dipole 46.London Dispersion Forces 47.Resonance 48.Hybrid orbital 49.area of high electron density

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)

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

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. 1Cond. 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.

∆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 = VariableSystem 1System 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)

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

organic nomenclature and structure inorganic nomenclature and structure

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).