FINAL REVIEW Chapter 2 Atoms Molecules and Ions Scientists Dalton s Postulates 1 Each element is composed of extremely small particles called atoms 2 All atoms of a given element are identical wrong isotopes 3 Atoms of an element are not changed into a different type of element by chemical reactions atoms are neither created nor destroyed in chemical reactions 4 Compounds are formed when atoms of more than one element combine a given compound always has the same relative number and kind of atoms J J Thomson Discovered the electron Plum pudding model of atom Overall neutral Plums were negative electrons Pudding positive else No nucleus no organization R A Milikan discovered mass of electron 9 10 x 10 28 g with oil drop Ernest Rutherford Gold foil experiment found a nucleus that was positively charged Electrons are randomly dispersed through a positive medium Niels Bohr expands on Rutherford by saying electrons travel in stationary orbits defined by their angular momentum James Chadwick proved existence of neutrons Dmitri Mendeleev First modern periodic table Organized by increasing atomic number Amadeo Avogadro contributions in the theory of molarity molecular weight and Avogadro s Law V1 n1 V2 n2 Math Stuff Average atomic mass given isotopes atomic mass 1 1 atomic mass 2 2 atomic mass n n Periodic Table Alkali Metals 1A Alkaline Earth Metals 2A Halogens 7A Noble Gases 8A Transition Metals the low ones Lanthanides top Actinides bottom Naming 1 Metal nonmetal where metal has a fixed predictable charge ex LiCl lithium chloride note the name doesn t reflect the quantity of either atom present 2 Metal nonmetal where metal is a transition metal Use the above method plus Roman numerals to indicate the charge on the metal ex FeCl2 iron II chloride note the name doesn t reflect the quantity of either atom present 3 Nonmetal nonmetal use prefixes know 1 10 given in class in front of each element but you can drop the mono if in front of the first element ex SO3 sulfur trioxide note the name DOES reflect the quantity of both atoms present Prefixes 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca 4 Oxyanions oxyacids know them be able to apply the naming to oxyacids that contain an element other than Cl ate ending on the anion becomes an ic ending on the acid Ex NO3 nitrate HNO3 nitric acid ite ending on the anion becomes an ous ending on the acid Ex NO2 nitrite HNO2 nitrous acid Name the ionic compound then use the prefix hydrate to indicate the quantity of water molecules occluded in the crystal structure Ex MgSO4 7H2O is magnesium sulfate heptahydrate 5 Hydrates 6 Exceptions HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid HF hydrofluoric acid HCN hydrocyanic acid H2S hydrogen sulfide H2Se hydrogen selenide H2Te hydrogen telluride B2H6 diborane PH3 phosphine SiH4 silane H2O water NH3 ammonia 7 Anions to know Look at chart from notes Chapter 3 Calculations with Chemical Formulas and Chemical Equations be able to determine empirical formulas molecular formulas given s or amounts be able to calculate limiting reagents theoretical yield yield be able to do problems that involve yields less than 100 be able to do combustion analysis problems Chapter 4 Chemical Reactions Electrolytes Electrolyte a substance that dissolves in water giving a solution that conducts electricity Strong electrolytes HCl HBr HI HNO3 H2SO4 HClO4 HClO3 Weak electrolytes molecular compounds all other acids Nonelectrolyte a substance that dissolves in water giving a solution that does NOT conduct electricity Reaction Types combustion hydrocarbon C H maybe O O2 CO2 H2O metathesis AB CD AD CB or AB C A CB driven by formation of gas formation of precipitate reactivity series with metals combination A B C 2 or more little things combine decomposition C A B something breaks down into smaller bits neutralization HX YOH YX H2O acid base salt water Redox Reactions be able to write net ionic equations be able to do dilution problems using molarity calculate molarity be able to apply it in a stoichiometry problem Activity series problem Chapter 5 Gases Gas Laws Boyle s Law P1V1 P2V2 Charles Law V1 T1 V2T2 Avogadro s Law V1 n1 V2 n2 Ideal Gas Law PV nRT R 08206 L atm Kmol R 8 314 J mol Can use d PM RT where M molar mass P atms V L T K n moles Conversions K C 273 5 9 F 32 273 1 atm 760 mm Hg 760 torr 1 01325 x 105 Pa STP 1 atm 273 K Partial Pressure Ptotal P1 P2 Pn P1 n1RT V etc Ptotal n1 n2 nn RT V Kinetic Molecular Theory RMS RMS sqrt 3RT M R 8 314 J K T K M kg mole 1 The size of a particle is negligibly small 2 The avg kinetic energy of a particle is proportional to temp in K 3 The collisions between particles are completely elastic the speed of gas molecules increases as T goes up and decreases as MW goes up Effusion when a gas leaks through a hole Diffusion a gas spreads out through another gas to occupy the space uniformly Graham s Law r1 r2 u1 u2 rates of effusion proportional to rms r1 r2 sqrt M2 M1 if T constant van der Waal s equation Gas behave non ideally at low temperature and high pressure because they get close and stick This creates nonelastic conditions which goes against kinetic molecular theory 1st term corrects for IMF 2nd for particle volume Gas collected over water problem RMS problem Partial pressure problem Chapter 6 Thermodynamics First Law of Thermodynamics energy is conserved if H is positive the reaction in endothermic if H is negative the reaction is exothermic be able to calculate H by these methods using the enthalpies given Hrxn n Hproducts n Hreactants using Hess s Law manipulate the reactions puzzle method using bond enthalpies Hrxn bond enthalpies of bonds broken bond enthalpies of bonds formed be able to use specific heat to calculate the energy needed to heat cool substances be able to do bomb calorimeter and coffee cup calorimeter problems q energy mS T Bomb calorimeter problem Coffee cup calorimeter problem Practice bond enthalpies Hess s law enthalpies given Energy needed to heat cool substances problems Chapter 7 Quantum Theory and the Electronic Structure of Atoms Max Planck discovered smallest unit is quantum Planck s constant h 6 63 x 10 34 J s E hv began quantum theory proposed that radiant energy is emitted or absorbed in discrete quantities only Neils Bohr electron is only allowed to occupy certain orbits of specific energies circular Schrodedinger current model of atom light has dual nature both as particles and waves Schroedinger equation gives