LTHS Chemistry Honors Chemistry 2023-2024 Second Semester Exam Review PDF

Name: ___________________________________ Date: ____________ Period: _______ LTHS: Chemistry Second Semester Exam Review Honors Chemistry 2023-2024 Unit 7 – The Mole 7.1 I can apply the mole concept to estimate quantities. 7.2 I can calculate the...

Name: ___________________________________ Date: ____________ Period: _______ LTHS: Chemistry Second Semester Exam Review Honors Chemistry 2023-2024 Unit 7 – The Mole 7.1 I can apply the mole concept to estimate quantities. 7.2 I can calculate the molar mass of a compound given the name or the formula. 7.3 I can convert between particles, mass and volume. 7.4 I can calculate the percent composition of a compound based on the formula and/or mass of individual elements. 7.5 I can determine the empirical and/or molecular formula of a compound based on composition data. 1. Which of the following has the greatest number of atoms? a. 26.98 g of Al b. 1 mole of S c..5 moles of O2 d. 18.02 g H2O e. All are the same 2. Calculate the molar mass of manganese (IV) oxalate. 3. How many moles are in 53.0 g of CH4? 4. What is the mass of 6.9x1025 molecules of NO? 5. Which of the following iron compounds contain the greatest percentage of iron: pyrite (FeS2), hematite (Fe2O3), or siderite (FeCO3)? 6. Determine the percent composition of Ca(NO3)2 7. A fluoride of aluminum is 67.87% F. What is the formula of this compound? 8. Give the empirical formula of the following based on the molecular formula given: a. C2H5OH b. C6H12O6 9. A compound is 40.00% C, 6.713% H and 53.28% O by mass. a. What is the empirical formula for the compound? b. If the molecular mass of the compound is 60.052 g, what is its molecular formula? 10. A compound with the empirical formula CH2 is known to have a molar mass of 56.12 g/mole. What is the molecular formula of the compound? Unit 8 – Balancing Equations and Stoichiometry 8.1 I can identify parts of a chemical reaction and can identify types of reactions. 8.2 I can balance a chemical equation. 8.3 I can write a chemical equation from a word problem. 8.4 I can use a balanced equation to determine the ratio between chemicals in a reaction. 8.5 I can convert from grams/particles/liters of a substance to grams/particles/liters of another. 8.6 I can determine the limiting reactant. 8.7 I can determine the amount of excess reactant left over and the theoretical yield of products. 8.8 I can determine the percent yield. 11. Identify the type of reaction for each reaction shown below. a. ____ AgNO3 + ____ NaCl → ____ AgCl + ____ NaNO3 __________ b. ____ MnO2 + ____ Al → ____ Mn + ____ Al2O3 __________ c. ____ Cr + ____ S8 → ____ Cr2S3 __________ 12. Determine if the following statements are true or false. If it is false, correct it. a. Subscripts can be added/deleted to balance an equation. b. Coefficients can be added/changed in order to balance an equation. c. The number of moles of reactants will equal the number of moles of products. d. The number of atoms of each element must be the same on both sides. 13. Balance each of the following equations using whole numbers: a. ____ H2O → ____ H2 + ____ O2 b. ____ C6H12O6 + ____ O2 → ____ CO2 + ____ H2O c. ____ Al + ____ H2SO4 → ____ Al2(SO4)3 + ____ H2 14. Write a complete balanced equation from the following: Solid zinc is placed in aqueous hydrochloric acid which produces aqueous zinc chloride and hydrogen gas. The following questions deal with the following balanced equation for the burning of propane gas: C3H8(g) + 5 O2(g) → 4 H2O(g) + 3 CO2(g) 15. How many moles of CO2 are produced when 17 moles of O2 react? 16. How many moles of C3H8 must react to produce 1500.0 g of water? 17. If 38.5 grams of oxygen react, how many molecules of carbon dioxide are produced? 3 The following questions deal with the following balanced equation for the production of ammonia: N2 + 3H2 → 2NH3 18. If 14.0 g N2 is mixed with 9.0 g H2, a. What is the limiting reactant? I C E b. How much excess reactant is left over? c. What is the theoretical yield? d. What is the percent yield if 16.1 g NH3 actually forms? Unit 9 – Reactions and Solutions Extension: I can differentiate between saturated, unsaturated, supersaturated solutions based on descriptions and/or models. 9.1 I can explain the concept of solubility (macroscopic and particle level) for ionic/molecular compounds. Extension: I can determine if substances would mix together/dissolve based on molecular polarity. 9.2 I can use solubility rules to determine if an ionic substance is soluble/insoluble and write a solvation equation to represent the dissociation of a soluble ionic compound. 9.3 I can write and interpret molecular equations, complete ionic equations, and net ionic equations for a reaction. 9.4 I can solve concentration calculations using molarity and/or dilution equations. 19. When an ionic substance dissolves in water, the ions _____________ and behave _______________ of one another. 20. What is the explanation behind why the lightbulb is lit up when placed in a saltwater solution? 4 21. Determine which diagram best represents the ionic compound Na2CO3 dissolved in water. 22. Classify each of the following substances as soluble ionic, insoluble ionic, or molecular. NaNO3 __________________ AgCl __________________ 23. Using the Solubility Rules for Ionic Compounds, determine which of the following substances will be soluble in water: a. Aluminum chloride b. BaSO4 c. Iron (III) hydroxide d. Na2CO3 24. Write the solvation equation for the following: a. NaCl (s) → b. BaSO4 (s) → c. (NH4)3PO4 (s) → 5 25. For each of the following reactions, complete and balance the molecular equation indicating clearly which product is the precipitate. If no reaction occurs, state “No Reaction.” If there is a reaction write the complete ionic equation and net ionic equation. a. Lead (II) nitrate + sodium chloride → b. NaNO3 (aq) + KCl(aq) → 26. What is the formula of the precipitate that forms when FeCl3(aq) reacts with KOH(aq)? Write a molecular equation (ME), Complete Ionic Equation (CIE), and Net Ionic Equation (NIE) for this reaction. ME: ___________________________________________________________________ CIE: ___________________________________________________________________ NIE: ___________________________________________________________________ 27. If 0.60 mole of NaCl is dissolved in enough water to make 0.50 L of solution, calculate the molarity? 28. How many grams of NaOH are needed to make 10.0 L of a 0.550 M solution? 29. What volume of 2.5 M HCl must be used to prepare 250mL of 1.0 M hydrochloric acid? 6 Unit 10- Equilibrium 10.1 I can identify factors that would affect the rate of a reaction. 10.2 I can explain how chemical equilibrium is achieved and how it relates to reversible reactions. 10.3 I can write an equilibrium constant expression, and determine the value of K using the equilibrium constant expression for a given reaction. 10.4 I can analyze the extent of a reaction based on the magnitude of K. 10.5 I can use the equilibrium expression to determine equilibrium concentrations based on given equilibrium concentrations. 10.6 I can determine the value of Q using the reaction quotient expression for a system. I can use the value of Q to determine if the system is at equilibrium or how it will proceed to get to equilibrium. 10.7 I can use ICE tables to determine equilibrium concentrations and/or K values. 10.8 I can use LeChatelier’s principle to describe how changes in energy, concentration, pressure, and temperature affect a reaction at equilibrium. 30. What is collision theory? What are the requirements for a successful collision? 31. How would each of the following affect the rate of a reaction? a. Increasing the temperature ____________________________________ b. decreasing Surface Area ____________________________________ c. decreasing the Concentration of Reactants ________________________ d. Adding a Catalyst __________________________________________ 32. Indicate whether the statements are True or False. CHANGE any false ones to make them true! a. At equilibrium the amount of products must equal the amount of reactants. b. During equilibrium, the concentration of the products does not change. c. A collision between reactants is all that is needed to cause a reaction. d. At the beginning of a reaction the forward reaction is faster than the reverse reaction, but then the forward reaction slows down as the reverse reaction speeds up. 33. If K > 1, __________ are favored. If K< 1, ______________ are favored. 7 34. Write equilibrium expression for the following reactions: a. 4NH3 (g) + 7O2 (g) ↔ 4NO2 (g) + 6H2O (g) b. NH4NO3 (s) ↔ N2O (g) + 2H2O (g) 35. For the reaction, N2 (g) + O2 (g) ↔ 2NO (g) at a certain temperature the equilibrium concentrations are found to be [N2] = 0.041 M, [O2] = 0.0078 M, and [NO] = 4.7 x 10-4 M. Calculate K for the reaction. 36. What factors can change the value of K? 37. How is Q used to determine if a reaction is at equilibrium? If K < Q the reaction will shift ________. If K > Q, the reaction will shift _______. 38. Iodine molecules react reversibly with iodide ions to produce triiodide ions. If a solution with the concentrations of I2 and I− both equal to 0.001M before reaction gives an equilibrium concentration of I2 of 0.000661M, what is the equilibrium concentration of I3- for the reaction? I2(aq) + I-(aq) ↔ I3-(aq) Initial: Change: Equilibrium: 39. For the synthesis of ammonia: the K is 1.2 at 375°C. Starting with [H2]=0.76M , [N2]=0.60M, and [NH3]=0.48M. Calculate Q. Which direction will the reaction shift to reach equilibrium? Calculate the equilibrium concentrations of each gas. N2(g) + 3H2(g) ↔ 2NH3(g) Initial: Change: Equilibrium: 8  40. For the following reaction, which of the following changes would shift the reaction to the right? heat + H2(g) + I2(g) ↔ 2HI(g) a. Increase the temperature ___________ b. Add I2 ___________ c. Increase the volume ___________ d. Add HI ___________ e. Increase the pressure ___________ f. Remove H2 ___________ g. Decrease the volume ___________ 41. For the reaction, heat + N2O4 (g) ↔ 2NO2 (g) use Le Châtelier’s Principle to predict shifts in the equilibrium position for each of the following changes: Equilibrium Shift [NO2] K value Addition of N2O4 Addition of NO2 Increase volume/ decrease pressure Increase in temperature Unit 11 – Acids and Bases 11.1 I can use both the Arrhenius and Bronsted-Lowry definitions to classify acids and bases. 11.2 I can identify and predict acid/base conjugate pairs. 11.3 I can identify the common physical and chemical properties of acids and bases. 11.4 I can calculate the concentration of the H and OH of a solution. (Kw = [H ] [OH ] & + - + - [H ] = 10 ) + -pH 11.5 I can describe the pH scale and calculate the pH of an acid or base solution. (pH = - log [H ]) + Extension: strengths of acids/bases, weak acids/bases (Ka/Kb and pH calculations) 11.6 I can explain acid-base reactions. 11.7 I can describe and execute a titration and perform calculations based on data. 42. What are the definitions for acids & bases: Acid: Base: 43. Write the equation for HCO3- acting as a base when combined with H2O. 44. Write the equation for HCO3- acting as an acid when combined with H2O. 9 45. Complete the following table listing the properties of acids and bases: Acid Base 46. In each of the following chemical equations, identify the conjugate acid-base pairs a. HNO3 (aq) + H2O (l) ↔ H3O+(aq) + NO3-(aq) b. CH3NH2 (aq) + HCl (aq) ↔ CH3NH3+(aq) + Cl-(aq) 47. What is meant by the term amphiprotic? 48. Rank the following solutions from the most acidic to most basic? a. [H+] = 0.00015 M b. pOH = 6.7 c. [OH-] = 4.39 x 10-5 M d. pH = 4.2 49. Calculate [H+] in each of the following solutions and indicate whether the solution is acidic, basic or neutral: a. [OH-] = 3.99 x 10-5M b. pH = 13.21 50. Calculate the pH in each of the following solutions and indicate whether the solution is acidic, basic or neutral: a. [H+] = 0.00100 M b. [OH-] = 9.18 x 10-11M 10 51. Write the neutralization reaction for the reaction of H2SO4 and NaOH. 52. What is an indicator and what is it used for? 53. How many mL of 0.50M barium hydroxide (Ba(OH)2) are required to fully titrate a 100.0 mL solution of 0.15M sulfuric acid (H2SO4)? 54. In a titration of sulfuric acid with sodium hydroxide, 32.20mL of 0.250M NaOH is required to neutralize 26.60mL of H2SO4. Calculate the molarity of the sulfuric acid. Unit 12 - Thermochemistry 12.1 I can distinguish between exothermic and endothermic reactions in terms of enthalpy. 12.2 I can draw and interpret energy diagrams (emphasis on ΔH and activation energy) 12.3 I can explain how specific heat capacity affects thermal properties of different substances. 12.4 I can perform calculations using the specific heat equation. 12.5 I can describe how a calorimeter determines energy of a system and perform calorimetry calculations. 12.6 I can determine the ΔH of a reaction using calorimetry. 12.7 I can use the ΔH of a reaction in order to calculate the heat involved for any substance used in a chemical reaction (thermostoich). 12.8 I can apply Hess’s Law to determine the ΔH of an overall reaction. 55. Label the following on the diagram a. Exothermic or endothermic b. Energy of reactants c. Energy of products d. ∆H 56. Determine the ∆H of the reaction above. What does the ∆H tell you? 11 57. If a reaction is exothermic, then: a. is the ΔH negative or positive? ____________________ b. should energy be included on the right side or the left side of the equation? __________ c. energy is absorbed or produced by the reaction? ___________________ d. the temperature of the surroundings _________________________ 58. What is specific heat capacity? What does the size of the specific heat value indicate? 59. When the sun is shining on equal amounts of sand and water, they are absorbing the same amount of energy. Explain why water has a smaller temperature change than sand. 60. How many joules of energy are required to raise the temperature of 125 g of water from 25.0ºC to 99ºC? 61. A sample of metal has a mass of 65.5 grams. If 1.90 kJ of energy flow into the metal, the temperature changes by 55.45ºC, what is the specific heat and identity of the metal? 62. An unknown mass of water at 98.3°C is mixed with 57.5 g of ethanol at 34.0°C. The final temperature of the mixture is 45.8°C. The specific heat of water and ethanol is 4.184 J/g°C and 2.44 J/g°C respectively. What is the mass of water? 12 63. A 1.50 g sample of KCl is added to 35.0 g H2O in a styrofoam cup and stirred until dissolved. The temperature of the solution drops from 24.8∘C to 22.4 ∘C. Assume that the specific heat and density of the resulting solution are equal to those of water and assume that no heat is lost to the calorimeter itself, nor to the surroundings. KCl(s)+H2O(l)⟶KCl(aq) ΔH=? a. Is the reaction endothermic or exothermic? b. What is the heat of solution(∆H) of KCl expressed in kilojoules per mole of KCl? 64. Use the following balanced equation for the burning of propane gas: C3H8 (g) + 5 O2 (g) → 4 H2O (g) + 3 CO2 (g) ΔH = - 2043 kJ/mol a. If 350 grams of propane burn, how much heat energy is produced? b. In order the produce 10,000.0 kJ of energy, how many grams of oxygen must react? 65. Calculate the ∆H for the reaction 2N2(g) + 6H2O(g) → 3O2(g) + 4NH3(g) Given the following data NH3(g) → 1/2N2(g) + 3/2H2(g) ∆H = 46 kJ 2H2(g) + O2(g) → 2H2O(g) ∆H = -484 kJ 13