Chemistry Final Exam Review 2006-2007

Honors Chemistry Final Exam Review 2012-2013 2nd Semester

The exam will cover vocabulary, concepts and calculations from the entire semester. Below are the units for the 2nd half of the semester.

Your exam date is June 4th.

Unit 7: CHEMICAL REACTIONS Define these terms: (9 points) Single replacement reaction Double replacement reaction Combination reaction Diatomic molecule Decomposition reaction Catalyst Precipitate Yield sign Reactants Activity series Products Law of conservation of mass Coefficient Subscript Chemical equation Combustion reaction Balanced chemical equation Aqueous

(22 points) 1. What are the 5 types of chemical reactions? 2. Know how to identify the different types of chemical reactions. Examples: a) FeCl3 + NaOH → Fe(OH)3 + NaCl b) Al + O2 → Al2O3 c) C2H2 + O2 → CO2 + H2O d) Na + H2O → NaOH + H2 e) KClO3 → KCl + O2 3. Know how to balance equations. Example: a) __ FeCl3 + __ NaOH → __ Fe(OH)3 + __ NaCl b) __ Al + __ O2 → __ Al2O3 c) __ C2H2 + __ O2 → __ CO2 + __ H2O d) __ Na + __ H2O → __ NaOH + __ H2 e) __ KClO3 → __ KCl + __ O2

4. Know how to predict products as in the problems below. Example: a) Al + N2 → b) H2O → c) Ca + H2O → d) Cl2 + NaBr → e) FeS + HCl → 5. What is an activity series chart? What type of reaction do you use it for? a) Using the activity chart, why can sodium replace hydrogen? 6. What are 5 indicators/observations of a chemical reaction? 7. List the chemical formulas for the 7 diatomic molecules. 8. Know how to translate chemical equations and balance them appropriately. Example: a) ammonium chloride reacts with calcium hydroxide to form calcium chloride and nitrogen trihydride (ammonia) and water b) sodium oxide and water yield sodium hydroxide

Unit 8: STOICHIOMETRY Know these terms: (.5) Mole ratio Actual yield Theoretical Yield Percent Yield

(6 points) 1. What do the coefficients mean in a chemical equation? 2. Know how to calculate the mole ratio between reactants and products in a chemical formula. a) What is the mole ratio for calcium and oxygen in 2Ca + O2 → 2CaO 3. Know how to solve mole to mole, mole to mass, mass to mole, mass to mass problems. a) How many moles of lithium hydroxide are required to react with 20. mol of carbon dioxide? CO2 + 2LiOH → Li2CO3 + H2O b) What mass, in grams, of glucose is produces when 3.00 mol of water react with carbon dioxide? 6CO2 + 6H2O → C6H12O6 + 6O2 c) How many moles of NO are formed when 824 g of ammonia reacts with an excess of oxygen? (balance the equation first) NH3 + O2 → NO + H2O d) How many grams of SnF2 are produced from the reaction of 30.0 g HF with Sn? Sn + 2HF → SnF2 + H2 4. Know how to solve percent yield problems. a) When 36.8 g of C6H6 react with an excess of Cl2. The actual yield of C6H5Cl is 38.8 g. What is the percent yield of C6H5Cl? C6H6 + Cl2 → C6H5Cl + HCl

Unit 9: GASES Know these terms (5.5) Ideal gas law directly proportional inversely proportional molar volume STP Ideal gas constant Partial pressure Dalton’s law kinetic molecular theory elastic collision law of combining volumes

(19 points) 1. Know the 5 assumptions of the kinetic molecular theory. 2. Know the difference between an ideal gas and a real gas. 3. Explain a gas based on the following properties: density, compressibility, diffusion, effusion, fluidity, shape , IMF, particle arrangement, and volume expansion. 4. Define pressure. What are some common pressure units? 5. Know how to convert pressure units: a) convert .200 atm to mmHg b) convert 345.8 kPa to atm c) convert 760 mmHg to kPa 6. What is standard temperature and standard pressure? 7. Know how to solve problems using Boyle’s law, Charles law, Gay-Lussac, Combined, Ideal, Density and Molar mass using the Ideal gas law and Dalton’s law of partial pressure. a) A gas occupies a volume of 200. ml at 100. mmHg. What volume will the gas occupy at 300. mmHg? b) Air has a total pressure of 20.6 atm and contains carbon monoxide, oxygen, and nitrogen. If air is made up of 0.6 atm of carbon monoxide, 12.6 atm of oxygen, what would be the partial pressure of nitrogen? c) If a sample of gas occupies 15.9 L at 34 C, what will its volume be at 27 C if the pressure does not change? d) The volume of a sample of oxygen gas is 300.0 ml when the pressure is 1.00 atm and the temperature is 27.0 C. At what temperature would the volume change to 1.00 L and the pressure change to 0.500 atm? e) A sample of gas at 25.0 C has a volume of 11.0 L and exerts a pressure of 660.0 mmHg. How many moles of gas are in the sample? f) A sample of gas in a closed container at a temperature of 100. C and a pressure of 3.0 atm is heated to 300. C. What pressure does the gas exert at the higher temperature? 8. Use the law of combining volumes, Avogadro’s law, and molar volume to solve these problems. a) 3O2 → 2O3 Both gases are measured at the same temperature and pressure. How many liters of O2 are required to make 24 L of O3 ? b) How many liters of O3 are formed from 12 mol of O2 at STP? 11. Know these answers: a) As the temperature of a gas decreases, the volume of a gas will ______. b) As the temperature of a gas decreases, the pressure of the gas will ______. c) As the volume of the gas decreases, the pressure of the gas will ______.

Unit 10: THERMOCHEMISTRY/SOLIDS & LIQUIDS Know these terms(9 points) Vaporization Condensation Evaporation Melting point Freezing point Sublimation Triple point Melting Freezing Deposition Phase Diagram Boiling Endothermic reaction Exothermic reaction Heating &Cooling Curve Specific heat capacity temperature heat

(13 points) 1. State the 6 phase changes of state and which ones work in opposition to each other. i.e. sublimation and deposition 2. Explain how a solid melts into a liquid using kinetic energy in your explanation. 3. What 2 temperatures measure the same amount during a phase change of a liquid pure solvent to a solid? 4. Know how to read phase diagrams. Sketch a quick diagram locating the triple point, critical point, the melting point /freezing point line and the boiling point/condensation point line. Also label the 3 sections as solid , liquid, and gas. 5. Know how to read a heating and cooling curve. What do the plateaus tells you? What do the slopes tell you? Where is the KE of the substance constant? 6. Sketch an endothermic reaction graph, labeling the reactants, products, activation energy, activated complex, and the heat of reaction. 7. What is the sign of an endothermic reaction and exothermic reaction? 8. Using the specific heat values for water and iron, which one would have the largest temperature change if they have the same mass? 9. Know how to calculate the heat released or absorbed during a physical change. a. Calculate the heat absorbed when 15.0 g of ice melts to liquid.

See reference sheet for Hfus b. Calculate the heat released when 75.4 g of vapor condenses into

liquid. See reference sheet for Hvap 10. Know how to calculate the heat released or absorbed in a chemical reaction? a) What is the specific heat of a metal that releases 2500 J of energy. The metal has a mass of 25 g and had a temperature change of 5C. b) How much heat is released when iron is dropped in a beaker of water. The mass of the metal was 43 g and the initial temperature of the metal was 78 C. The water temperature changed from 25 C to 32 C. The specific heat of the metal is .45J/gC. c) What is the amount of heat absorbed by water if 23.4 g of water is heated from 34C to 78 C. See reference sheet for specific heat of water.

Unit 11 : KINETICS & EQUILIBRIUM Know these terms (3 pts) Rate Activation energy Collision theory Activated complex Transition state Catalyst

(18 points) 1. Explain the three criteria of the collision theory.

2. On the pathway below, label the activated complex, activation energy with catalyst, and activation energy without catalyst 3 What are the five factors that affect the rate of a reaction? 4. Which of the five factors change collision frequency? 5 Which factor changes collision frequency and the energy of the collisions? 6. How does rate change if you increase the concentration of the reactants? 7. How does rate change if you increase the surface area? 8. How does rate change if you decrease the temperature? 9. How does rate change if you add a catalyst? 10. Write the equilibrium expression for the following reaction. a) H2(g) + Cl2(g) 2HCl(g) + heat 11. In the process of chemical equilibrium, what stays constant at equilibrium? 12. In the process of equilibrium, are the rates equal to each other? 13 Using the reaction above, answer the following questions regarding Le Chatelier’s principle. a) Which direction does the reaction shift if temperature increases? b) Which direction does the reaction shift if hydrogen gas is increased? c) Which direction does the reaction shift if HCl is removed? d) Which direction does the reaction shift if the volume is decreased? e) Which direction doe the reaction shift if temperature is decreased? 14. If K = .00045, what side of the reaction will be favored?

Unit 12: SOLUTIONS Know these terms:(8 pts) Solution solute soluble insoluble miscible immiscible electrolytes (strong and weak) non-electrolytes solubility solvent supersaturated solution aqueous solution unsaturated solution saturated solution Henry’s law molarity

(13 points) Know the following: 1. Explain the like dissolves like rule and give an example following the rule. 2. Name 3 factors that increase the rate of dissolution of a substance. 3. Describe solution equilibrium. 4. Name substances that are considered electrolytes and non-electrolytes. 5. What is the effect of temperature and pressure on gas solubility? 6. What is the effect of temperature on the solubility for most ionic solids? 7. Know how to calculate molarity a) What is the molarity of 4.5 moles of Ba(OH)2 in 10.0 L? b) A solution has a molarity of 2.8 M and a volume of 250 ml. How many moles of solute are in the solution? 8. Know how to read a solubility graphs. a. Using the solubility graph from the notes, how much of NaCl can be dissolved at 45C b. Using the solubility graph from the notes, 50 g of KClO3 is dissolved in 100 g of water at 45C. Is the solution saturated or unsaturated? 9. Know how to solve dilution problems. a) How many ml of a 2.0 M NaBr solution are needed to make 200 ml of a 0.50 M solution? 10. Which types of substances produce electrolytes? 11. Which type of substances produce non electrolytes?

UNIT 13: ACIDS & BASES Know these terms: (6 pts) Arrhenius acid Arrhenius base Bronsted-Lowry acid Bronsted-Lowry base pH conjugate acid Conjugate base hydroxide ion hydronium ion neutralization titration equivalence point

(23 points) 1. List some common properties of an acid. 2. List some common properties of a base. 3. Define self-ionization of water. 4. Know how to predict the products and balance neutralization (double replacement) reactions. a) H2CO3 + Fe(OH)3 → 5. Know how to calculate the pH from hydrogen and hydroxide ion concentrations a) What is the pH of a [OH-] = 1 x 10-5 M? b) What is the pH of a [H+] = 1 x 10-5 M? c) What is the pOH of a [H+] = 1 x 10-1 M? d) What is the pOH of a [OH-] = 1 x 10-12 M? - 6. What is the hydrogen ion concentration of 0.001 M HNO3? What is the [OH ]? 7. What is the hydrogen ion concentration of [OH-] = 3.0 x 10-2 M? What is the pH? 8. What is the pH of a solution if the [H+] = 3.4 x 10-5 M? What is the hydroxide concentration? -2 9. Determine the pH of a 2.0 x 10 M Sr(OH)2? 10. The pH of a solution is measured and determined to be 7.52? What is the hydrogen ion concentration? Is the solution acidic or basic? 11. Know how to look at an equation and predict Bronsted-Lowery acids and bases and conjugate acids and conjugate bases. a) + + NH4 + H2O → NH3 + H3O What is the base? What is the conjugate base? What is the acid? What is the conjugate acid? 12. What are the products of neutralization? 13. Know how to name acids and bases a) HF b) H2SO4 c) NaOH d) HNO2 e) Fe(OH)2

14. In a titration, how much of .15 M NaOH is needed to neutralize 20 ml of .500M HCl solution? HCl + NaOH  H2O + NaCl 15. In a titration, what is the molarity of HNO3 if 25 ml of it neutralized 15 ml of . 60M Ca(OH)2 2 HNO3 + Ca(OH)2  2 H2O + Ca(NO3)2 16. What is the difference between end point and equivalence point?