CP Chemistry

Semester 2 Final Review

Information about the final:

  • Your final will be multiple choice. Bring a pencil, eraser, and SCIENTIFIC CALCULATOR. Graphing calculators will NOT be allowed.
  • You will get a formula sheet with equations and constants, periodic table, polyatomic ion sheet, standard enthalpy of formation chart, bond energies chart, and standard reduction potential table.
  • See the board for 7th period review
  • Answer key to final review will be posted on my website.

Resources for studying:

  • Notes
  • HW assignments
  • Previous test review guides
  • Quizzes
  • Textbook
  • Semester 2 Final Review

Unit 5: Stoichiometry

  • Stoichiometry
  • Mole ratio
  • Limiting reactants
  • Calculations to find limiting reactant
  • Calculations to find amount product formed
  • Calculations to find amount excess reactant left over
  • Percent composition
  • Empirical formula
  • Molecular formula
  • Hydrates and formulas of hydrates

Practice Problems

  1. Consider the following unbalanced equation: __NaBr + __Ca(OH)2 __CaBr2 + __NaOH
  2. If you have 8g of NaBr and 10g of Ca(OH)2, what is the limiting reactant?
  3. What is the reactant in excess?
  4. How many grams NaOH can you make?
  5. How many grams of the excess reactant do you have left?
  6. Calculate the percent composition of deoxyribose C5H10O4
  7. A compound is 65.5% C, 5.5%H, and 29.0% O. It has a molar mass of 110g/mol. What is the empirical formula? What is the molecular formula?
  8. A hydrate is found to be 45.53% water and 54.57% cobalt (II) chloride. Find the formula of the hydrate. What is the name of the hydrate?
  9. A 1.04g sample of a hydrate of nickel (II) sulfate is heated. The mass after heating is 0.61g. Find the formula of the hydrate. What is the name of the hydrate?
  10. A chemist analyzing a compound finds that it contains .8007g C, .9333g N, .2016g H, and 2.133g of O. Determine the empirical formula.
  11. Consider the following unbalanced equation: ___Fe(s) + ___O2(g)  ___Fe2O3(s)
  12. If you have 50g of iron, how many grams of Iron (III) oxide is produced?
  13. How many moles of oxygen are needed to completely react with 10 g of iron?
  14. A chemist analyzing a compound finds that it contains .8007g C, .9333g N, .2016g H, and 2.133g of O. Determine the empirical formula.
  15. Consider the combustion of ethane: ___C2H6 + ____O2 ___CO2 + ___H2O. If 4.5 g of ethane were combusted completely, how many grams of carbon dioxide would be produced? How many moles of water would be produced?
  16. Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) What is the mole ratio of zinc to hydrochloric acid (HCl). How many moles of hydrochloric acid is needed to completely react with 2 moles of zinc?

Unit 6: Gases and gas laws

  • Properties of gases
  • Kinetic Molecular Theory
  • Kinetic energy
  • Temperature
  • Maxwell-Boltzman Distribution
  • Diffusion
  • Barometer
  • Monometer
  • Gas Variables
  • Temperature (Celsius to Kelvin conversion)
  • Amount
  • Volume
  • Molar volume
  • Pressure (Conversion factors will be given to you)
  • Gas laws
  • Boyle’s
  • Charles’s
  • Gay-Lussac’s
  • Combined gas law
  • Ideal gas law
  • Gas constants (will be given to you, just be able to pick correct one out)
  • Gas laws and stoichiometry
  • Dalton’s law of partial pressures
  • Collecting a gas over water

Practice Problems

  1. How does a mercury barometer work? What happens to atmospheric pressure as you climb up a mountain?
  2. Convert 450 kPa to mmHg, atm, torr, and psi.
  3. What is the conversion from ºC to K?
  4. What is absolute zero? Describe a gas at absolute zero.
  5. What are the units for P, V, n, and T when R = 0.0821?
  6. What are the assumptions of the Kinetic Molecular Theory?
  7. Draw and label a manometer that shows a confined gas pressure of 800mmHg when the atmospheric pressure is 760mmHg.
  8. Use the diagram below to answer these questions:
  9. A and B are fluorine and chlorine gas at 27ºC. Which is which? Explain.
  10. A and B are gasses: one at 200 K and the other at 350 K. Which is which? Explain.
  1. What is the relationship between the temperature of a gas and the kinetic energy? What happens to the velocity of a particle as the temperature increases?
  2. Why do hotter gasses exert more pressure? Give 2 reasons.
  3. Argon and Krypton are in a closed container at the same temperature. Which one has a greater speed? Why?
  4. What is Charles’s law? What is one example of Charles’s law in real life?
  5. What is Boyle’s law? Which constants are held constant?
  6. What is Gay-Lussac’s law? Which constants are held constant?
  7. What is the ideal gas law and combined gas law?
  8. What is STP?
  9. What pressure in atm is required to compress 12.5 L of gas at 760 mmHg to a volume of 3.8L?
  10. A gas has a volume of 2.4 Liters at 60ºC. When the gas is heated to 78ºC, what will the new volume be?
  11. A gas is in a container with a volume of 9.8 L. When the temperature is 45ºC, the pressure inside is 1.5 atm. What is the pressure when the temperature is cooled to 24ºC?
  12. What is the volume of 10 moles of gas at STP? At 1.5 atm and 25ºC?
  13. A gas at 780mmHg has a volume of 500mL. What will the volume be when the pressure is decreased to 500mmHg?
  14. A gas at 350 K has a pressure of 2.0 atm. What will the temperature be when the pressure is increased to 3.5 atm?
  15. A balloon at sea level has a pressure of 1.0 atm, temperature of 28ºC, and a volume of 1.3 L. The balloon ascends and the pressure decreases to 0.71 atmand the temperature decrease to 24ºC. What is the new volume?
  16. Water can be decomposed to hydrogen and oxygen gas. If 4.0L of oxygen gas is made at 1.1 atm and 30ºC, how many grams of hydrogen gas were produced?
  17. A container has a volume of 12.0 L. The gas inside has a temperature of 300 K and exerts a pressure of 0.8 atm. The container is compressed to a volume of 10.1 L and the pressure increase to 1.3 atm. What is the new temperature in K? In ºC?
  18. Three gases, A, B, and C, are in a container. The total pressure of the gasses is 2.0atm. If gas A exerts a pressure of 1.0atm and gas B exerts a pressure of 0.2 atm, how much pressure does gas C exert? Which type of gas is the majority in the container?
  19. Consider this equation: NaHCO3(s) + C2H4O2(aq)  NaC2H3O2(aq) + H2O(l) + CO2 (g). 100ml of carbon dioxide gas was collected over water at 19ºC. At 19ºC, the partial pressure of water 16.5 mmHg. The total pressure is 750mmHg.
  20. What is the partial pressure of CO2?
  21. How many moles of CO2 were produced?
  22. How many grams of NaHCO3was used in the reaction?
  23. Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)
  24. If 3.0 L of hydrogen gas were produced at STP, how many grams of zinc were used in the reaction?
  25. If 3 moles of HCl were used in the reaction, how many liters of hydrogen gas was produced at STP? At 1.5 atm and 20ºC?

Unit 7: Thermodynamics

  • Energy
  • Kinetic energy
  • Potential energy (chemical)
  • Law of conservation of energy
  • Heat
  • calorie, joule, nutritional calories
  • specific heat
  • calculating heat flow
  • thermochemistry
  • system
  • surrounding
  • exothermic
  • endothermic
  • bonds and energy
  • enthalpy
  • enthalpy of reaction
  • energy diagram
  • thermochemical equations
  • calorimetry
  • coffee cup calorimeter
  • finding the specific heat of a substance
  • finding enthalpy of reaction of a substance
  • Hess’s Law
  • Standard enthalpy of formation
  • Standard enthalpy of reaction
  • Using bond energies to calculate enthalpy of reaction

Practice Problems

  1. What is chemical potential energy? What is kinetic energy?
  2. What is specific heat? What are the units?
  3. Is bond forming exothermic or endothermic? What about bond breaking?
  4. How does the chemical potential energy of a system change during an exothermic reaction?
  5. In describing a chemical reaction, what is the system and what is the surrounding?
  6. If H is positive, is the chemical reaction exothermic or endothermic? How has the chemical potential energy of the system changed?
  7. Draw and label an energy diagram for an exothermic reaction. Label the axis, products, reactants, activation energy, the activated complex, and H.
  8. 1000 J of energy is added to 150ml of water. By how much does the temperature change? If the initial temperature of the water was 18ºC, what is the final temperature?
  9. Use the bond energy table to determine H for this reaction. Is it exothermic or endothermic?

2H2(g) + O2(g)  2H2O(g)

  1. A 50kg bar of pure gold absorbed 276000 J of heat. What is the final temperature of the gold if the initial temperature was 25ºC? Specific heat of gold is 0.129 J/(gºC)
  2. Calculate the amount of heat produced when 10.0g of methane burns in air according to the following unbalanced equation: ___CH4(g) + ___O2(g)  ____CO2(g) + ___H2O(l) H = -525kJ
  3. For the same equation in question 11, calculate H using standard enthalpies of formation.
  4. The temperature of an unknown metal with a mass of 30g increases from 25ºC to 28ºC when the metal absorbs 80 J of heat. What is the specific heat of the metal?
  5. A piece of metal with a mass of 65g is removed from boiling water and placed in a coffee cup calorimeter. The temperature of 75g of water increases from 24ºC to 28.1ºC. What is the specific heat of the metal?
  6. Refer to the diagram to the right
  • What is the ∆H of the forward reaction?
  • What is the activation energy of the forward reaction?
  • What is the ΔH of the reverse reaction is?
  • What is the enthalpy of the reactants of the forward reaction?
  • What is the enthalpy of the reactants of the reverse reaction?
  • What is the activation energy of the reverse reaction?
  1. A sandwich has 800 Calories. How many joules is that?
  2. Find ΔH for this reaction: CuCl2(s) + Cu(s)  2CuCl(s)

Given: Cu(s) + Cl2(g)  CuCl2(s) ΔH = -206kJ

2Cu(s) + Cl2(g)  2 CuCl(s) ΔH = -36kJ

  1. Find ΔºH for this reaction: 2NO(g) + O2(g)  2NO2(g) using standard enthalpies of formation.

Unit 8: Redox and Electrochemistry

  • Oxidation numbers
  • Oxidation
  • Reduction
  • Redox reactions
  • Half-reactions
  • Balancing reactions using the half-reaction method
  • Electrochemistry
  • Voltaic cell
  • Half-cell
  • Electrode
  • Anode
  • Cathode
  • Salt bridge
  • Cell notation
  • Reduction potential
  • Cell potential

Practice Problems

  1. What is a redox reaction, and what is transferred in a redox reaction?
  2. If the oxidation number of oxygen increases in a reaction, has it been reduced or oxidized?
  3. Sketch a voltaic cell and include the main components. Pick two electrodes of your choice, then use the standard reduction potential to determine the direction of electron flow.
  4. What does the salt bridge do?
  5. What is a cathode? Anode?
  6. What is the oxidation number of carbon in CO2? CH4? C6H12O6? CH4O?
  7. Assign oxidation numbers to the following: CaCl2, H2O2, CClF3, Cr2O72-, KNO2
  8. Given the following redox reaction, what species is oxidized? Reduced? Write the two half-reactions.

Mg + HCl MgCl2 + H2

  1. Balance the following redox reaction using the half-reaction method. H3PO2 + H+ + Cr  Cr3+ + P + H2O
  2. Hg22+ + 2e-  2Hg

Cu2+ + 2e-  Cu

  1. Write the balanced overall cell reaction for the given pair of half-reactions.
  2. Write the cell notation.
  3. Calculate E0cell.
  1. A voltaic cell consists of a lead electrode and aluminum electrode. What is the cell potential? Write the balanced overall cell reaction.
  2. CdCd2+Cu2+Cu What is the cell potential of the voltaic cell represented by the cell notation?

Unit 9: Radioactivity and Nuclear Chemistry **May change depending on how far we get.

  • Nuclear chemistry
  • Radioactive decay
  • Alpha decay
  • Beta decay
  • Gamma radiation
  • Positron emission
  • Electron capture
  • Nuclear equation
  • Half-life
  • Fission
  • Fusion

Practice Problems

  1. What is the net change in the nucleus when a nucleus undergoes positron emission? Alpha decay?
  2. Would you rather be exposed to alpha particles or gamma radiation? Why?
  3. What is half-life? A radioactive isotope X has a half-life of 45 minutes. If you start with 10grams of the isotope, how many grams will you have in 4 hours?
  4. What two things must be conserved in a nuclear equation?
  5. Fill in the blanks and identify what kind of decay the following nuclear equations represent.

 ______+

______ +

 ______+

  1. Carbon-14 has a half-life of 5730 years. How much of a 14 gram sample of carbon-14 will remain after 17190 years?
  2. Fluorine-21 has a half-life of 5 seconds. What percentage of the original nuclei would remain after 30 seconds? If you started out with 2 grams of fluorine-21, how many seconds would it take to reduce the sample to 0.5 grams?
  3. The half-life of Iodine-125 is 60 days. After 120 days, there is 131 grams of iodine-125 remains. What was the mass of the original sample?
  4. Write the nuclear equation for the following:

Gold-185 emits an alpha particle.

Thorium-234 undergoes alpha decay.

Potassium-42 emits a beta particle.

Magnesium-32 emits a positron.

Calcium-39 undergoes positron emission

Iodine-125 undergoes electron capture.