Heat Energy Internal energy Kinetic Energy Potential Energy Endothermic Exothermic Thermodynamics

1. 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

2. 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 sugars fats polymers solution solvent solute concentration molarity ppm ppb wt% vol% molecular equations ionic equations net ionic equationsspectator 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

3. Significant Figures – start at the left and proceed to the right • If the number does not have a decimal point count until there are no more non zero numbers • 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 • Observation • Hypothesis • Experiment • Theory • Law • Chemistry • Matter • Energy • Chemical Properties • Physical Properties • Extensive Properties • Intensive Properties • Scientific (natural) law • Anion • Cation • Molecular Geometry • Law of Conservation of Mass • Law of Conservation of Energy • Exact numbers • Accuracy • Precision • compounds • molecules • chemical formula • empirical formula • molecular formula • structural formula • bond line formula • ball and stick model • space filling model • mole • Electronic Geometry • percent weight • percent error • percent composition • percent yield • %RSD • limiting reactant • Stoichiometry • Stoichiometric Coefficient • Electron Affinity • Electronegativity • Covalent Bond • Ionic Bond • Dipole • London Dispersion Forces • Resonance • Hybrid orbital • area of high electron density

4. Table of Common Ions Common Positive Ions (Cations)

5. Table of Common Ions Common Negative Ions (Anions)

8. Standard P  1.00000 atm or 101.3 kPa • Standard T  273.15 K or 0.00oC • K = 273 + oC • 1 mm Hg = 1 torr 760 torr = 1 atm • The standard molar volume is 22.4 L at STP • PV = nRT • R = 0.08206 L atm mol-1 K-1 • Ptotal = PA + PB + PC + ..... • At low temperatures and high pressures real gases do not behave ideally. • The reasons for the deviations from ideality are: • The molecules are very close to one another, thus their volume is important. • The molecular interactions also become important.

9. 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.

10. Specific heat capacity (J/(g∙K) = heat lost or gained by system (Joules) mass(grams) DT (Kelvins) ∆H = Hfinal - Hinitial • The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions. • ∆Horxn =  ∆Hfo(prod) -  ∆Hfo(react) q cP = m(Tf –Ti)

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

13. organic nomenclature and structure inorganic nomenclature and structure

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

15. Quantum Numbers n and l define the energy of the electron • 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 ms which characterizes the single electron. The spin quantum number only has two possible values. ms = +½ or -½ one spin up ↑ and one spin down↓ 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.

16. Ionic Polar Covalent Covalent Determine Inductive effect Count the number of electrons the element should have Determine how equally electrons are shared (DEN) >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 • draw the Lewis dot structure • 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. • 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 • list all the elements follow with an equal sign • follow with the number of atoms of that type in the • molecule • follow with a multiplication sign • If the element is O follow with a -2 • If the element is H follow with a +1 • any other element enter a ? • follow with an = sign, do the math • draw a total line, then enter the charge on the molecule • Do the algebra backwards to solve for ?

17. VSEPR Theory • Lone pair to lone pair is the strongest repulsion. • Lone pair to bonding pair is intermediate repulsion. • Bonding 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 109.5o. • Electronic geometryis 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 geometrydetermined by the arrangement of atoms around the central atom(s). Summary of Electronic & Molecular Geometries