CY 121 Final Exam Study Guide Fall 2024 PDF

CY 121 FINAL EXAM TOPICS – FALL 2024 FINAL EXAM WILL BE 40 MULTIPLE-CHOICE QUESTIONS (MCQs) AND 4 MULTI-PART FREE-RESPONSE QUESTIONS (FRQs)ON THE TOPICS LISTED BELOW. Possible FRQ topics are in bold below. 50% of the final exam will cover topics from Lessons 1-26 (Exams 1-3); the remain...

CY 121 FINAL EXAM TOPICS – FALL 2024 FINAL EXAM WILL BE 40 MULTIPLE-CHOICE QUESTIONS (MCQs) AND 4 MULTI-PART FREE-RESPONSE QUESTIONS (FRQs)ON THE TOPICS LISTED BELOW. Possible FRQ topics are in bold below. 50% of the final exam will cover topics from Lessons 1-26 (Exams 1-3); the remaining 50% of the final exam will cover topics from Lessons 27-37 as well as the Beer-Lambert Law (Beer’s Law) and Solution Composition. A Periodic Table and a reference sheet of General Solubility Rules for Ionic Compounds and select math equations will be provided. In addition to the periodic table and formula reference sheet, you may also prepare and use both sides of a 3” x 5” index card with HANDWRITTEN notes only. LESSONS 1-26: 50% OF EXAM Scientific inquiry o Variables (independent, dependent, constants, replicates) o Control group Classification of matter o Pure substances ▪ Elements vs. compounds ▪ Atoms, ions, and subatomic particles – protons, neutrons, electrons, quarks ▪ Molecules o Mixtures ▪ Homogeneous vs. heterogeneous mixtures ▪ Mixtures vs. Compounds o Classification of properties and change ▪ States of matter 4 states – solid, liquid, gas, plasma 6 changes of state – melting, freezing, vaporization, condensation, sublimation, deposition ▪ Physical properties and change ▪ Intensive and extensive properties ▪ Chemical properties and change, including common signs of chemical change – color change, gas formation, etc. Qualitative and quantitative data Units: SI/metric system o 6 base units – kg, m, s, K, mol, Pa o Prefix names and symbols: micro- (), milli- (m), centi- (c), deci- (d), kilo- (k) Accuracy and precision o Accuracy and precision o Random, systematic, and gross errors o Average (mean) calculation o Making accurate and precise measurements on instruments – Ex. Graduated cylinder, buret, pipet, etc. o Significant figures – counting, rounding, and calculating ▪ Multiplication and division ▪ Addition and subtraction Density o D = m/V (equation provided on formula sheet) o Calculating density, volume, or mass given 2 of these 3 variables o Water displacement (Vwater = Vsolid) Temperature o Absolute zero o Celsius ↔ Kelvin conversions: TK = T⁰C + 273.15 (equation provided on formula sheet) Dimensional analysis o Conversion factors  o SI-SI conversions ▪ Memorize 5 SI/metric prefix meanings ( m, c, d, k); Ex. 1 m = 100 cm, 1 L = 1000 mL, etc. o English-English conversions (Ex. 1 mi = 5280 ft, 1 ft = 12 in, etc.) o SI-English conversions (equalities will be provided in problems; Ex. 1 in = 2.54 cm, 1 kg = 2.205 lb, etc.) o Square (area) and cubed (volume) units: square and cube all numbers and units in conversion factor Introduction to the Periodic Table (copy of periodic table WITHOUT element names provided on formula sheet) o Element tiles ▪ Atomic number and number of protons ▪ (Average) atomic mass ▪ Memorize “Top 50” element symbols and names o Metals, nonmetals, and metalloids ▪ Locations on the periodic table – “staircase” ▪ Basic properties – appearance, conductivity, luster, etc. o Groups ▪ International (1-18) and USA (I-VIII/A-B Roman Numeral) systems ▪ Main-group, transition, and inner transition ▪ Alkali metals, alkaline-earth metals, halogens, and noble gases o Periods (1-7) Basics of atomic structure o Subatomic particles – protons, neutrons, electrons, quarks o Dimensions (relative volumes and masses) – nucleus vs. electron cloud o Forces – strong force, electromagnetic force, weak force, gravity o Scanning Tunneling Microscope (STM) Counting subatomic particles o Neutral atoms and ions: protons and electrons ▪ Ex. Neutral atom of sodium is Na has 11 protons and 11 electrons ▪ Ex. Sodium ion is Na+ and has 11 protons and 10 electrons o Isotopes: protons, neutrons, and electrons ▪ Atomic number ▪ Mass number ▪ (Average) atomic mass (carbon-12 standard for definition of 1 amu) ▪ Longhand notation (Ex. magnesium-24) ▪ Abbreviated notation (Ex. Mg-24) ▪ Atomic symbols (Ex. atom: 24 24 24 2+ 12Mg or just Mg, Ex. ion: 12Mg or 24 2+ just Mg ) Atomic mass o Mass spectrometer o Element mass spectrum ▪ Sketch graph ▪ Identify isotopes and symbols from graph o Conceptual understanding of (average) atomic mass when only 1 element isotope or multiple element isotopes o Calculate average atomic mass of an unknown element o Calculate percent abundances of a known element with 2 isotopes Covalent and ionic bonding o Elements ▪ CNN: Covalent = Nonmetal (or metalloid) + Nonmetal ▪ MIN: Ionic = Metal + Nonmetal  o Properties of ionic vs. covalent substances: differences in melting points, conductivity, solubility in water, etc. o Representative particles: molecules (covalent), formula units (ionic), and atoms o Chemical formulas: counting numbers of atoms or ions Chemical nomenclature o Ionic substances (metal + nonmetal, unless it contains NH4+) ▪ Binary ionic (and hydrogen compounds not dissolved in water) Cation (+) o Metal (or H) o Roman numerals: ▪ 3 transition metal exceptions that DO NOT need one: Ag, Cd, Zn (1+, 2+, 2+) ▪ 2 main-group metals that DO need one (2+ or 4+): Sn, Pb Anion (-) o Monatomic: Nonmetal stem + -ide suffix Traditional name exceptions: H2O (water) ▪ Ternary ionic Same rules as binary ionic except use polyatomic ion names and formulas o Memorize “Top 20” Polyatomic Ions o Non-“Top 20” to know: H2O2 (hydrogen peroxide) o Covalent substances ▪ Element nomenclature Diatomic molecules (Br2-I2-N2-Cl2-H2-O2-F2 “twins”) Ozone (O3) o Binary covalent ▪ Prefixes No “Mono” rule ▪ Second element + -ide suffix –o/a ending before “oxide” ▪ Traditional nitrogen exceptions: NH3 (ammonia), N2O (nitrous oxide) ▪ Alkane hydrocarbons: CnH2n + 2 Methane, ethane, propane, butane, pentane, etc. o Acids (HX, where X = monatomic or polyatomic anion) ▪ Common properties of acids and bases ▪ Classifications of acids Monoprotic Polyprotic (diprotic, triprotic) Oxyacids o From polyatomic ions (-ate/-ite) ▪ Nomenclature Rules o Anion ends in –ide → hydro-stem-ic acid o Anion ends in –ate → stem-ic acid o Anions ends in –ite → stem-ous acid Use “sulfur” and “phosphor” as stems Chemical composition of compounds o Definition of the mole o Formula mass (amu) and molar mass (MM, g/mol)  o Stoichiometric conversions for substances: ▪ Mol ↔ Particles (atoms, molecules, formula units): convert using Avogadro’s number (N A provided on exam formula/equation sheet) ▪ Mol ↔ Grams: convert using MM ▪ Apply kilo- and milli- metric prefixes as needed o Mass percent composition of compounds ▪ Mass percent of an element in a compound given the compound name or formula ▪ Mass of element in a given mass of compound o Empirical formulas ▪ Determine x and y in AxBy From masses of elements in compound From % of elements (assume 100 g) o Molecular formulas ▪ Determine n in (AxBy)n given MM of the compound o Solution composition and molarity (M = mol/L) ▪ Dilute vs. concentrated ▪ Molarity of compound and ions ▪ Moles ↔ Liters: convert using M ▪ Apply kilo- and milli- metric prefixes as needed Chemical reactions o Evidence of a chemical reaction ▪ Color change, energy change (heat/light/temperature), precipitation, gas/odor evolution o Chemical equation symbols ▪ (s), (l), (g), (aq) ▪ +, →, , (∆ = heat, Pt = catalyst) ▪ Reactant(s), product(s) Remember Br2-I2-N2-Cl2-H2-O2-F2 “twins” Remember periodic trend in ion charges Remember polyatomic ion formulas o Writing balanced chemical equations ▪ Write chemical formulas from names in “reaction statement” if necessary → skeleton equation ▪ Balance skeleton equation by adding coefficients Reduce coefficients to smallest whole numbers! Keep polyatomic ions together and treat H2O as HOH as needed Balance elements by themselves last, like O2 in CH/O combustion reactions Can use fractional coefficients temporarily to balance, but multiply the entire equation by 2 at the end to get to smallest whole numbers ▪ Particle diagrams o Types of reactions and predicting products ▪ Synthesis/Combination Ex. 2Na + Cl2 → NaCl ▪ Combustion (Reactant 1 + O2 ) Reactant 1 = element is also synthesis/combination o Ex. 2Be + O2 2BeO Reactant 1 = organics (CH/O) → CO2 + H2O o Ex. CH4 + 2O2 CO2 + 2H2O  ▪ Decomposition Binary compounds o Ex. 2AlCl3 → 2Al + 3Cl2 Metal carbonates → met oxides + carbon dioxide o Ex. CaCO3 → CaO + CO2 ▪ Single replacement Metal (and H) replacement o Ex. Zn + 2HCl → ZnCl2 + H2 Nonmetal replacement o Ex. Cl2 + 2NaBr → Br2 + 2NaCl ▪ Oxidation-reduction (redox) Oxidation states/numbers o Rules for assigning (memorize!) Elements oxidized (OIL = Oxidation Is Loss of electrons) vs. reduced (RIG = Reduction Is Gain of electrons) ▪ Double replacement Precipitation o Identify precipitate (ppt) ▪ General Solubility Rules for Ionic Compounds in Water provided on exam formula/equation sheet ▪ Ex. NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq) (AgCl = ppt) Acid-base o Acid + base → salt + water ▪ 7 strong acids ▪ Ex. HCl(aq) + KOH(aq) → KCl(aq) + H2O(l)) o Net ionic equations ▪ Dissociation of soluble ionic compounds (aq) and strong acids ▪ Identify spectator ions (cancel in equation) ▪ Ex.  General Rules for Solubility of Ionic Compounds in Water – PROVIDED ON EXAM FORMULA/EQUATION SHEET COMMON ION(S) SOLUBILITY IN WATER NOTABLE EXCEPTIONS Group 1 ions, NH4+ high none NO3−, C2H3O2− high none Cl−, Br−, I− high Pb2+, Hg22+, Ag+ SO42− high Pb2+, Hg22+, Ag+, Ca2+, Sr2+, Ba2+ OH−, S2− low Group 1 ions, NH4+, Ca2+, Sr2+, Ba2+ CrO42−, CO32−, PO43− low Group 1 ions, NH4+ Reaction stoichiometry o Particle diagrams o Dimensional analysis ▪ Mole ratio (coefficients from balanced equation) ▪ Particles to moles: Avogadro’s number (6.022 x 1023) ▪ Grams to moles: MM (molar mass) ▪ Liters (aq) to moles: M (molarity) o Limiting and excess reactants o Reaction yields: theoretical yield, actual yield, and percent yield (equation on exam formula sheet) o Titrations Thermochemistry o First Law of Thermodynamics and system vs. surroundings ▪ Exothermic vs. endothermic processes (including graphs) o Thermochemical equations and stoichiometry ▪ Dimensional analysis Thermochemical ratio: H/mol reactant and product Quantity (from moles) of reactant or product Heat produced (-) or consumed (+) o Specific heat capacity and calorimetry ▪ cp = specific heat capacity at constant pressure ▪ q = mcpT (equation on exam formula sheet) qsurr = -qsys = mcpT ▪ H = qsys/mol o Hess’s Law ▪ Combination of multiple reaction equations to obtain H for single (target) reaction equation ▪ Standard states and allotropes Formation chemical equations ▪ Standard enthalpies of formation, Hf° o H° calculated from Hf° values (table provided) ▪ H° = (moles x H°f, products) – (moles x H°f, reactants) (equation provided on formula sheet) LESSONS 26-37 + BEER’S LAW AND SOLUTION COMPOSITION: 50% OF EXAM Quantum Atomic Theory o Waves, Light, and Energy ▪ Parts and properties: wavelength, amplitude, frequency, and velocity ▪ Electromagnetic (EM) radiation and EM spectrum  ▪ The quantum ▪ The photon ▪ Relationships (direct/inverse) between: (equations provided on formula sheet) Wavelength and frequency (c = n) Frequency and energy (E = h) Wavelength and energy (E = hc/) o Absorption, emission, and transmission of light by matter ▪ Bohr model of atom ▪ Line emission spectra of the elements o Quantum model of the atom ▪ Quantum numbers: n, l, ml, ms Principal energy levels/shells: n = 1, 2, 3, 4, … ∞ o Size and energy of orbitals Sublevels/subshells (s, p, d, f): l = 0, 1, 2, 3, … n-1 o Types of orbitals (s, p, d, f) Orbitals o Orientation of orbitals: ml = -l, …, 0, …+l o # orbitals in each sublevel/subshell: s = 1, p = 3, d = 5, f = 7 Electrons: o Spin: ms = -½ or +½ (↑ or ↓) o Max # electrons in each sublevel: s = 2, p = 6, d = 10, f = 14 o Max # electrons in each orbital = 2 Ground-state electron configurations o The 3 rules/principles: aufbau, Pauli, and Hund o 3 notations ▪ Orbital diagrams (Ex. ) ▪ Complete notation (Ex. S: 1s 2s22p63s23p4) 2 ▪ Condensed (noble gas) notation (Ex. S: [Ne]3s23p4) o Valence electrons (highest occupied n) (Ex. S has 6 valence electrons) o Unpaired electrons and paramagnetism vs. diamagnetism o Exceptions (Cr and Cu = 4s1 … elements) o Reading from the Periodic Table: ▪ Periods = n = 1-7 ▪ Groups = s, p, d, and f blocks o Ions ▪ Cations (Ex. Na+ = 1s22s22p6, Ga3+ = [Ar]3d10) ▪ Anions (Ex. F- = 1s22s22p6) Periodic trends o Number of valence electrons (and core electrons) o Ionic charges o Atomic and ionic radii o First and successive (2nd, 3rd, etc.) ionization energies o Electron affinity  o Explanations in terms of atomic structure and/or Coulomb’s law (equation provided on formula sheet) ▪ Period trends: effective nuclear charge (Zeff) determines Coulombic force between protons and electrons ▪ Group trends: size/energy of orbitals/# electron shells dictates the distance (r) and Coulombic force between protons and electrons Lewis symbols for elements Electronegativity o Explanations in terms of Coulomb’s law (Zeff, r) o Pauling values for elements Difference in electronegativity (EN) and bond character o Bond polarity and dipole moments (+, - or ) o Ionic, polar, and nonpolar covalent bonds Ionic compounds and properties o Electron transfer o Lattice energy (Coulomb’s law): ionic charges vs. size o Ex. Melting point, hardness Lewis structures (molecules and polyatomic ions) o Symbols: lines (bonding electron pairs), dots (nonbonding electrons), brackets and charge (ions only) o Octet rule o Exceptions to the octet rule ▪ Electron-deficient (Ex. BeF2) ▪ Odd-electron (Ex. NO) ▪ Electron-rich (Ex. SF6) o Resonance o Bond order Covalent bond dissociation energy o Potential energy diagram/well o Relative bond lengths, strengths, and energies based on bond order and atomic size Beer’s law o max determination  o A = bc (A = absorbance,  = molar absorptivity, b = path length (1.00 cm), c = concentration; equation provided on formula sheet) o Calibration line (A vs. c) o Determination of unknown concentration (c) Solution preparation o From solids o From solutions: M1V1 = M2V2 (M = molarity, V = volume; equation provided on formula sheet)