Published on March 30, 2014
Chem 105 Final Review DR. SHIRTS-WINTER 2013 JOANNA WILLIAMS
Disclaimer All the problems on here come from your textbook! If you need more, I suggest working through micro-exams and practice sheets/parallel example problems, seeing your Learning Community Mentor to run through previous exams, and looking in the book for more problems to work through over specific things your struggling with.
Measurements Accurate: numbers close to the actual value Precise: numbers close to each other Significant Figures: All non-zeros are sig. figs. Zeros between two non-zeros are sig. figs. Zeros left of first non-zero are NOT sig. figs. If #>or=1, all zeros right of decimal are sig. figs. If #<1, all zeros at end of # and between non-zeros are sig. figs. Trailing zeros may or may not be sig. figs. (That’s why we use scientific notation)
Nomenclature Metal + Nonmetal = Ionic compound Charges designate formula, name the elements and add –ide to the end Nonmetal + Nonmetal = Covalent molecule Use prefixes and add –ide to the end Polyatomic Ions Organic Functional Groups Organic Prefixes Acids If you “–ate” too much you feel “–ic”ky “-ite”s like Nephites and Lamanites are people like “-ous” Hypo-ous, ous, ic, per-ic increasing O Hydro-ic
Dimensional Analysis Use the Mole to Mole Ratio from stoichiometric coefficients in balanced chemical equation Find limiting reactant
Practice Problems 3.51) Determine the empirical and molecular formulas of each of the following substances: Styrene, a compound used to make Styrofoam cups and insulation, contains 92.3% C and 7.7% H by mass and has a molar mass of 104 g/mol Caffeine, a stimulant found in coffee, contains 49.5% C, 5.15% H, 28.9% N, and 16.5% O by mass and has a molar mass of 195 g/mol Monosodium glutamate (MSG), a flavor enhancer in certain foods, contains 35.51% C, 4.77% H, 37.85% O, 8,29% N, and 13.60% Na, and has a molar mass of 169 g/mol
Practice Problems 3.69) A piece of aluminum foil 1.00 cm square and 0.550 mm thick is allowed to react with bromine to form aluminum bromide. How many moles of aluminum were used? (density of aluminum=2.699 g/mL) How many grams of aluminum bromide form, assuming the aluminum reacts completely?
Practice Problems 3.76) Aluminum hydroxide reacts with sulfuric acid as follows: 2Al(OH)3(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 6H2O(l) Which is the limiting reactant when 0.500 mol Al(OH)3 and 0.500 mol H2SO4 are allowed to react? How many moles of Al2(SO4)3 can form under these conditions? How many moles of the excess reactant remain after the completion of the reaction?
Reactions Know the solubility rules and the exceptions for precipitation reactions Oxidation-Reduction (RedOx) 123FHO7654 LEO goes GER / OIL RIG Oxidizing agents/Reducing agents
Reactions Acid/Base/Neutralization Titrations and Dilutions: M1V1=M2V2 (molarity= mol/L)(volume=L) Net Ionic Equations Strong Acids (ionize completely) H2SO4 , HNO3, HCl, HBr, HI, HClO4 Strong Bases (dissociate completely) Group 1-OH, Group 2-OH from Ca down
pH pH=-log[H+] pOH=-log[OH-] pH+pOH=14 [H+]+[OH-]=1x1014 pH<7 acidic pH=7 neutral ph>7 basic
Practice Problems 4.83) Some sulfuric acid is spilled on a lab bench. You can neutralize the acid by sprinkling sodium bicarbonate on it and then mopping up the resultant solution. The sodium bicarbonate reacts with sulfuric acid as follows: 2NaHCO3(s) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) + 2 CO2(g) Sodium bicarbonate is added until the fizzing due to the formation of CO2(g) stops. If 27 mL of 6.0 M H2SO4 was spilled, what is the minimum mass of NaHCO3 that must be added to the spill to neutralize the acid?
Practice Problems 4.40) Write the balanced molecular and net ionic equations for each of the following neutralization reactions: Aqueous acetic acid in neutralized by aqueous barium hydroxide Solid chromium(III) hydroxide reacts with nitrous acid Aqueous nitric acid and aqueous ammonia react
Practice Problems 4.51) Which element is oxidized and which is reduced in the following reactions? N2(g) + 3H2(g) 2NH3(g) 3Fe(NO3)2(aq) + 2Al(s) 3Fe(s) + 2 Al(NO3)3(aq) Cl2(aq) + 2NaI(aq) I2(aq) + 2NaCl(aq) PbS(s) + 4H2O2(aq) PbSO4(s) + 4H2O(l)
Thermochemistry Ek = ½(mv2) 1 Cal = 1000 cal 1 cal = 4.184 J (specific heat of water) ∆E=q+w H=E+PV Bond enthalpies: reactants – products aka bonds broken – bonds formed ∆Hf o: products – reactants (diatomics in natural state = 0) q=mCs∆t Hess’s Law
Practice Problems 5.43) Consider the following reaction: 2Mg(s) + O2(g) 2MgO(s) ∆H = -1204 kJ Is this reaction exothermic or endothermic? Calculate the amount of heat transferred when 3.55g of Mg(s) reacts at constant pressure How many grams of MgO are produced during an enthalpy change of -234 kJ? How many kilojoules of heat are absorbed when 40.3g of MgO(s) is decomposed into Mg(s) and O2(g) at constant pressure?
Practice Problems 5.56) When a 4.25g sample of solid ammonium nitrate dissolves in 60.0g of water in a coffee-cup calorimeter, the temperature drops from 22.0 C to 16.9 C. Calculate ∆H (in kJ/mol NH4NO3) for the solution process NH4NO3(s) NH4+(aq) + NO3-(aq) Assume that the specific heat of the solution is the same as that of pure water. Is this process endothermic or exothermic?
Practice Problems 5.65) From the enthalpies of reaction H2(g) + F2(g) 2HF(g) ∆H = -537 kJ C(s) + 2F2(g) CF4(g) ∆H = -680 kJ 2C(s) + 2H2(g) C2H4(g) ∆H = +52.3 kJ Calculate ∆H for the reaction of ethylene with F2: C2H4(g) + 6 F2(g) 2CF4(g) + 4HF(g)
Electrochemistry E=hv=hc/λ 1/λ=R(1/n1 2-1/n2 2) E=-Rhc(1/n2) λ=h/mv Bohr’s Model
Practice Problem 6.37) Calculate the energy of an electron in the hydrogen atom when n=2 and when n=6. Calculate the wavelength of the radiation released when an electron moves from n=6 to n=2. Is this line in the visible region of the electromagnetic spectrum? If so, what color is it?
Orbitals & Nodes s=spherical Radial nodes starting at 2s p=peanut 1 planar node and radial nodes starting at 3p d=dlover leaf? 2 planar nodes and radial nodes starting at 4d f
Quantum Numbers Pauli Exclusion Principle n=shell (1, 2, 3….) l=subshell (n-1 to 0) (0=s, 1=p, 2=d, 3=f…) ml=orientation (-l to l) ms=spin (-1/2 or +1/2)
Practice Problem 6.56) Which orbital goes with the following quantum numbers? Which are not allowed? 2, 1, -1 1, 0, 0 3, -3, 2 3, 2, -2 2, 0, -1 0, 0, 0 4, 2, 1 5, 3, 0
Electron Configuration Expanded Condensed using Noble Gas configuration Cu and Cr exceptions Why? Hund’s Rule
Practice Problem 6.61) For a given value of the principal quantum number, n, how do the energies of the s, p, d, and f subshells vary for Hydrogen? A many-electron atom?
Periodic Trends Electronegativity Size Ionization Energy Electron Affinity Effective Nuclear Charge Family Names Cation and Anion Size
Lewis Dot Structures Count total valence electrons Least electronegative atom in the middle Fill octet, create multiple bonds if too many electrons Resonance structures: none actually what the molecule looks like, it’s a hybrid of all of them Formal charges Bond strengths
Molecular Orbitals Bond Order MO Diagrams Paramagnetic vs. Diamagnetic VSPER Molecular shapes and Angles
Gases PV=nRT 22.4 L = 1 mole gas @ STP Ideal gas characteristics Small, high temp, low pressure Partial pressures Pa=XaPt Pt=P1+P2+P3… Effusion and Diffusion Rates Urms=√(3RT/M) r1/r2= √(M2/M1) J=kgm2/s2
Practice Problems 10.54) Calculate the density of sulfur hexfluoride gas at 707 torr and 21 C Calculate the molar mass of a vapor that has a density of 7.135 g/L at 12 C and 743 torr
Practice Problem 10.69) A piece of dry ice (solid carbon dioxide) with a mass of 5.50 g is placed in a 10.0 L vessel that already contains air at 705 torr and 24 C. After the carbon dioxide has totally vaporized, what is the partial pressure of carbon dioxide and the total pressure in the container at 24 C?
Intermolecular Forces London Dispersion Induced Dipole (Polarizability) Dipole-Dipole Hydrogen Bonding Ion-Dipole
Liquids Intermolecular force effect on Viscosity Surface Tension
Phase Changes Phase Diagrams Specific heats to heat the substance to melting or boiling point Heats of vaporization or fusion to melt or evaporate substance
Practice Problems 11.43) For many years drinking water has been cooled in hot climates by evaporating it from the surfaces of canvas bags or porous clay pots. How many grams of water can be cooled from 35 C to 20 C by the evaporation of 60 g of water? (The heat of vaporization of water in this temperature range is 2.4 kJ/g. The specific heat of water is 4.18 J/gK.)
Colligative Properties Depends on number of solute particles present, not identity Don’t forget ions dissociate! Vapor Pressure ↓ Pa=XaP° Boiling Point ↑ Freezing Point ↓ Osmotic Pressure ↑ π=iMRT
Solids Unit Cells Simple/Primitive Cubic 1 atom Face-Centered Cubic 4 atoms Body-Centered Cubic 2 atoms
Concentrations [ ]=M=moles solute/L solution=molarity m=moles solute/kg solvent=molality X=moles substance/total moles=mole fraction ppm=(mass substance/total mass)x106 mass %=(mass substance/total mass)x100
Equilibrium kc=[products]/[reactants] kp=Pproducts/Preactants aA + bB cC + dD kc=[C]c[D]d/[A]a[B]b Le Chatelier’s Principle Noble gases being pumped in to increase the pressure have no effect (don’t change individual partial pressures) Catalysts have no effect Solids and liquids have no effect
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