Honors Chemistry Final Exam Study Guide Fall 2024 PDF

HONORS CHEMISTRY FINAL EXAM STUDY GUIDE – Fall 2024 1. Determine formula mass and number of atoms & molecules (use Avogadro’s number) 2. Determine average atomic mass using isotopic mass and percent abundances. 3. Determine (mass) percent composition from chemical formulas....

HONORS CHEMISTRY FINAL EXAM STUDY GUIDE – Fall 2024 1. Determine formula mass and number of atoms & molecules (use Avogadro’s number) 2. Determine average atomic mass using isotopic mass and percent abundances. 3. Determine (mass) percent composition from chemical formulas. portion (Hint: think of fractions, % = total mass x 100% = ___% ) Example: what is theoretical percent composition of CH4? Get mass of C and mass of all H’s from periodic table. Then, get total mass of CH4 12 4 %C = 16 x 100% = 75% C %H = 16 x 100% = 25% H Mass percent composition: Example below is for glucose, C6H12O6. 4. Five things you need to determine molecular formula: 1. Empirical formula 2. Mass from Empirical formula 3. Molar mass from Molecular formula 4. Ratio of Molar mass/Empirical mass 5. Molecular formula: Multiply the ratio found in Step 4 to the empirical formula (subscripts) 5. Determine the empirical formula that is 88.8% copper and 11.2% oxygen 6. Determine the molecular formula of caffeine (molar mass 194.19 g/mol), given that a sample contains 49.47 g carbon, 28.85 g nitrogen, 16.48 g oxygen, and 5.20 g hydrogen. 7. A compound with an empirical formula of C2H8N and a molar mass of 46.0 g/mol. What is the molecular formula of this compound? mass 8. Know how to use the density equation, d = volume. What is the unit for density? 9. Know how to use dimensional analysis method to convert between units, for example, to convert nm to m or m to nm, etc. This method also applies to converting feet to yard, meter to cm, vice-versa and other forms of units, as well as converting between moles and grams and between moles and number of stuff using Avogadro’s number. Converting mole to grams 44.01 g CO2 0.688 mol CO2 ( ) = 30.3 g CO2 mol CO2 Converting grams to mole 1 mol CO2 16.0 g CO2 ( ) = 0.364 mol CO2 44.01 g CO2 10. Solid lithium hydroxide has been used in space vehicles to remove exhaled carbon dioxide from the living environment. The products are solid lithium carbonate and liquid water. A student is trying to determine the mass of gaseous carbon dioxide that is needed to react with 1.00 x 103 g of lithium hydroxide absorb. Which step(s) in the student’s work below is (are) incorrect? 2LiOH(s) + CO2(g) → Li2CO3 (s) + H2O(l) 1 mol LiOH 2 mol LiOH 1 mol CO2 1.00 x 103 g LiOH (23.95 g LiOH )( 1 mol CO )(44.01 g CO ) 2 2 Step 1 Step 2 Step 3 Step 4 11. Limiting Reactant Calculate the mass of magnesium oxide that could be produced if 2.40 g of Mg reacts with 10.0 g of O2. Mg(s) + O2(g) → MgO(s) Determine limiting reactant to get amount of product produced (theoretical yield). 1 mol Mg 2 mol MgO 40.31 g MgO 2.40 g Mg ( )( )( ) = 3.98 g MgO 24.31 g Mg 2 mol Mg 1 mol MgO 1 mol O2 2 mol MgO 40.31 g MgO 10.0 g O2 ( )( )( ) = 25.2 g MgO 32.0 g O2 1 mol g O2 1 mol MgO Based on the limiting reactant, the possible amount of MgO produced is 3.98 g. (Limiting reactant is the reactant that produces the smaller product yield [moles or grams]) 12. Given the following reaction: 2 Fe + 3 Cl2 → 2FeCl3. You wish to determine the limiting reactant for the formation of FeCl3 from 2.30 g of Fe and 4.00 g of Cl2. a) How many moles of FeCl3 can be produced from the given mass of Fe? b) How many moles of FeCl3 can be produced from the given mass of Cl2? c) Based on (a) and (b) above, which is the limiting reactant, Fe or Cl2? ________ d) What is percent yield if the actual yield is 5.22 g of FeCl3? (Hint: calculate theoretical yield from limiting reactant) actual yield % yield = theoretical yield x 100% 13. Be able to name compounds and write chemical formulas of compounds. - Group 1, 2 and 3 ionic compounds: NaCl (sodium chloride), CaBr2 (calcium bromide) KNO3 (potassium nitrate), MgNO3)2 (magnesium nitrate), Al(NO3)3 (aluminum nitrate)  - Transition metal ionic compounds (notice the use of Roman Numerals): CuCl – copper(I) chloride CuCl2 – copper(II) chloride Fe(OH)2 – iron(II) hydroxide Fe(OH)3 – iron(III) hydroxide Fe2(CO3)3 – iron(III) carbonate Ni3(PO4)2 – nickel(II) phosphate - Covalent compounds (nonmetal and nonmetal) – use of Greek prefixes CO – carbon monoxide CO2 – carbon dioxide N2O4 – dinitrogen tetroxide PCl5 – phosphorus pentachloride P4O10 – tetraphosphorus decoxide IF7 – iodine heptafluoride - Binary Acids HF – hydrofluoric acid HCl – hydrochloric acid HBr – hydrobromic acids H2S – hydrosulfuric acid - Oxyacids H2SO4 – sulfuric acid H2SO3 – sulfurous acid H2CO3 – carbonic acid HClO4 – perchloric acid HNO3 – nitric acid CH₃COOH – acetic acid - Ionic Compound hydrates Pb(ClO4)2 3H2O – lead(II) perchlorate trihydrate Ba(OH)2 8H2O – barium hydroxide octahydrate CuCl2 2H2O – copper(II) chloride dihydrate Name the following compounds. NaF CaCl2 K2O Fe(NO3)3 I2F7 SO3 SnBr2∙2H2O Ni(NO3)2 Write the chemical formulas for the following compounds. Lithium bromide magnesium hydroxide Calcium sulfide ammonium carbonate Sulfur hexafluoride cobalt(III) sulfate dinitrogen pentoxide lead(IV) phosphate 14. Know isotope symbols. What is an isotope? For example: 12 C, 13 C, and 14 C or 12C, 13C, 14C 6 6 6 Determine the mass number, number of protons, neutrons and electrons for the following isotopes. i. isotope Mass # # neutrons # protons #electrons 12 C 13 C 58 Ni 61 Ni 64 Ni Be able to identify atomic number, mass number and number of protons, electrons and neutrons. 31 Example: P : mass number = 31, atomic number = 15, # protons = 15, 15 # electrons = 15, and # neutrons = 31 – 15 = 16 15. Distinguish and give examples of chemical and physical properties. Physical Properties (composition does not change): Boiling Point, Melting Point, Density, Size, ductile, malleable, conductivity, luster (shine), color Chemical Properties (leads to change in composition of compounds): Reactivity, flammable, sensitivity to air (oxygen), corrosive, acidic or basic 16. Know elements, pure substances, molecules, compounds, homogeneous mixtures (solutions) and heterogeneous mixtures. Example: element (anything in periodic table, Na, C, S, O2 etc) Homogeneous mixture (solution) – NaCl dissolved in water, sugar and water Solid solution - brass, stainless steel (Fe, Ni and Cr) Heterogeneous mixture – sand water, salad dressing Pure substance – 24-karat gold, glucose (C6H12O6) 17. How are ions formed? Which part of the periodic table tends to form positive ions and which part form negative ions? Give examples of symbols of ions. Remember the ions to memorize. 18. Identify the parts of the periodic table, for example, alkali metals, transition metals, gases, etc. - Group number, valence electrons, which elements loses electrons and which gain electrons 19. What are strong and weak electrolytes? What are strong acids and strong bases? Give examples. 20. Know how to write and balance chemical equations – molecular equation, complete ionic, and net ionic equations. Write chemical equations from word equations. Know how to complete word equations. - know aqueous ionic compounds dissociate which means the ionic compounds separate into ions (positive and negative) when dissolved in water. Give examples. A. Show the dissociation of the following aqueous solutions NaOH(aq) → + Na2CO3(aq) → + Na3PO4(aq) → + _ CuSO4(aq) → + ___Fe(NO3)3(aq) → + B. Determine the chemical formulas of the ions when they combine. Charges must add to zero. (Hint: Use Criss-Cross Method) Na+ + OH- → Na+ + CO32- → Na+ + PO43- → Cu2+ + SO42- → Fe3+ + NO3- → Zn2+ + NO3- → C. Balancing Chemical Equations Balance the chemical equations: 1 H3PO4 + 3 KOH →___ K3PO4 + ___ H2O 2 NaOH + 1 H2CO3 → ___Na2CO3 + ___ H2O 16 Al + 3 S8 → ___ Al2S3 1 C3H8 + 5 O2 → ___ CO2 + ___ H2O FeCl3 + ____MgO → ___Fe2O3 + MgCl2 Li + _H3PO4 → H2 + ___Li3PO4 H2(g) + ___V2O5(s) → ___V2O3(s) + ___H2O(ℓ) Balance and write complete ionic and net ionic equation: NaOH(aq) + Fe(NO3)3(aq) → ___NaNO3(aq) + ___Fe(OH)3(s) Complete Ionic Equation: Net Ionic Equation: 21. Be able to identify different types of reactions, for example, acid-base, oxidation-reduction, precipitation, combustion, gas forming etc. Balance the following equations and state what reaction type is taking place: Combustion, double- replacement, single-replacement, neutralization (acid-base), decomposition, or redox. 1) 1 C5H12 (g) + 8 O2 (g) →____ CO2 (g) + ____ H2O (g) + heat Reaction type: 2) 2 K3PO4 (aq) + 3 BaCl2 (aq) →____ Ba3(PO4)2 (s) + ____ KCl (aq) Reaction type: 3) 2 Al(OH)3 (aq) + 3 H2SO4 (aq) → ____ Al2(SO4)3 (aq) + ____ H2O (ℓ) Reaction type: 4) 2 HCl + 1 K2CO3 →____ CO2 + ____ H2O + ____ KCl Reaction type: 22. Know acid-base, strong and weak electrolyte, strong and weak acid. Know Bronsted-Lowry acid-base and conjugate acid-base. In the following acid-base reactions, identify the acid, base, conjugate acid, and conjugate base. a. HOCN(aq) + H2O(l) ⇌ OCN- (aq) + H3O+(aq) b. HClO4(aq) + H2O(l) ⇌ H3O+(aq) + ClO4–(aq) 23. Be able to determine oxidation numbers, identify oxidation and reduction processes given a chemical equation, write half-reactions, and write a balanced overall oxidation-reduction equation. Be able to determine oxidation numbers of elements, especially elements within chemical formulas. For example, what is oxidation number for S in SO2 and in SO3. Determine the Oxidation Number of each of the elements that is in bold. a) NH3 _____ b) H2SO4 _____ c) ZnSO3 _____ d) Al(OH)3 _____ e) Na _____ f) Cl2 _____ g) AgNO3 _____ h) ClO4- ____ i) SO2 _____ j) K2Cr2O4 ____ k) K2Cr2O7 _____ i) S ____ Fe2+(aq) + Co(s) → Co2+(aq) + Fe(s) (i) Oxidation: Co(s) → Co2+ (aq) + 2e- (ii) Reduction: Fe2+ (aq)+ 2e- → Fe(s) Cr2O7 2— + ClO2— → Cr3+ + ClO4— Substance oxidized Oxidizing agent Substance reduced Reducing agent LIGHT and Electronic Structure c = 3.00 x108 m/s, h = 6.626 x10-34 J∙ s, 1 nm = 1.00 x 10-9 m 24. Know the properties of a wave and identify parts of a wave. Know the electromagnetic spectrum, high/low energy, high/low frequency, short/long wavelength, type of light UV, Visible, Infrared, etc). Energy level diagrams. A. True or False. Gamma rays have long wavelengths. B. True or False. Visible light has shorter wavelengths than UV light. C. True or False. Infrared light has higher energy than radio waves. D. True or False. Microwave has higher frequency than x-rays. E. What is the speed of light equation? What is λ? What is ν? 25. Know how to use the speed of light equation, c =  , and the energy equation, 𝒉𝒄 E = h =  A. The wavelength of a diagnostic x-ray is only 0.01 nm. What frequency does the doctor’s machine operate with? B. A bright line spectrum contains a line with a wavelength of 518 nm. Determine ____. i. the wavelength, in meters ii. the frequency iii. the energy C. A hypothetical wave has 6.6 x 10-19 J of energy. What is its hypothetical, approximate frequency? 26. What is the Bohr Model of the atom? Identify and draw energy level diagrams for absorption and emission. Label parts of diagram. 27. Be able to draw the different types of orbitals. Remember to always label the x-, y-, and z- coordinate axes. 28. What are valence electrons and core electrons? Can you identify valence and core electrons from electron configurations? How would you determine valence electrons from the periodic table. 29. Be able to explain wave-particle duality of light and the photoelectric effect. 30. Be able to write electron configurations (example: Carbon has atomic number of 6, thus it has 6 electrons. Electron configuration: C (6 e-) - 1s22s22p2 K (19 e-) - 1s22s22p63s23p64s1 And draw orbital diagrams, for example: Carbon Potassium - apply Aufbau principle (order of filling of orbitals from lowest energy to highest energy) - Pauli Exclusion principle - Hund’s Rule - Draw energy level diagram for orbitals 31. Know and explain the periodic patterns/trends with regard to size of atoms, ionization energies, electron affinity, electronegativity and metallic/non-metallic character of elements. 1. Graph the following information in a Line graph. Label and number the x- and y-axes appropriately. Experimental Techniques 2. What is the measured volume of the liquid in the 50-ml graduated cylinder? a. 50.7 mL b. 56.0 mL c. 44.0 mL d. 43.0 mL 3. A student measured the length of an object to be 41.38 cm. Which digit is uncertain? a. “8” b. “3” c. “1” d. “4” 4. Identify the target that shows high accuracy and high precision ________. Based on the data provided below, which students has the most accurate data? a. Student A b. Student B c. Student C 4. Convert the following I. To scientific notation. II. To decimal notation. a. 6431 ___________________ a) 6.56 x 10-4 5. Solve the following. Answers must have the correct significant figures. a) 9.20 – 4.00 = __________ (b) (0.50)(15.8) = _____________ 6. Identify the different physical processes (e.g., melting, freezing, vaporization, condensation, deposition, sublimation) 8. Identify and describe the different ways to separate matter. a. b. c. d. evaporation chromatography filtration decanting e. distillation