Chemistry

Utica Community Schools – Semester Two Review

Directions:

  • You may prepare your own hand-written 3x5 note card which will be collected on your exam day. You must bring a pencil and calculator to the exam.
  • You will be given a periodic table.
  • If you complete the entire review sheet by the day of your exam you will be eligible for the curve should one be applied.
  • The answers to this review must be on a separate sheet of lined paper. You will first see the objective of each question followed by the question itself. This is what you are providing the answer to.
  1. Use stoichiometry to solve problems involving the conversion of one substance into another substance and the Limiting Reactant. These can include mole to mole, gram to mole, and liter to mole conversions. Please review the four steps to stoichiometry. Please review how to calculate molar mass as well as the conversion factor from Liters to moles of a gas at STP.
  2. Use the following balanced equation to answer the following problems:

C5H12 + 8O2  5CO2 + 6H2O

i) When 2.5 moles of oxygen gas react completely with pentane, how many moles of

carbon dioxide are produced?

ii) How many grams of pentane gas are needed to completely react with 348.5 grams of

oxygen?

iii) A student reacts 75 grams of pentane with 45 liters of oxygen gas at STP. Which

substance is the limiting reactant?

iv) What volume of water vapor is produced from 30 liters of oxygen gas at 0°C and 1 atm?

  1. Use acid and base chemistry to calculate the pH, determine the acidity of a solution, and explain acid rain.
  2. Calculate the pH for a solution with a hydronium ion concentration of 2.5 x 10-3M.
  3. The concentration of a hydrobromic acid solution is 0.093 moles/liter. What is the pH of the hydrobromic acid solution? Hint: HBrH+1 + Br-1 100% ionization
  4. What is the pH of a 5 x 10-5M hydroxide ion solution? Classify this solution as acid, basic or neutral.
  5. Research the causes of acid rain. What compounds are responsible for acid rain? Where do they primarily come from?
  6. Understand equilibrium and reversible reactions. Know how to write the equilibrium expression and understand the meaning of having a K value greater than one and less than one. Be aware of how the equilibrium will respond to stress.
  7. The Keq of a reaction is 5x10-16. At equilibrium are the reactants or products favored?
  8. If a reaction has an equilibrium constant (Keq) less than 1, what type of reaction is it?
  9. Given the following K values, which will create the highest product concentrations at equilibrium?
  10. 2.1x1023 ii. 6.2x10-14iii. 8.3x104
  11. What is the effect of adding more hydrogen gas to the following equilibrium reaction? Which way will the reaction favor and will more ammonia be produced or less?

N2(g) + 3H2(g) 2NH3(g)

  1. For the reaction shown above, what would happen to the concentration of reactants if pressure were increased on the system?
  1. Know how to use formulas to get the percent by mass and know how to use percent by mass to get formulas.
  2. Define Percent by Mass, Empirical Formula and Molecular Formula
  3. What is the percent by mass of each element in NH3?
  4. What is the percent by mass of Ca in Ca3P2?
  5. The empirical formula for benzene is CH. If the molecular weight of benzene is approximately 78.0g/mol, what is the molecular formula?
  6. What is the empirical formula of a compound that is 18.8% sodium, 29.0% chlorine and 52.2% oxygen?
  7. Know how to apply the concepts of electrochemistry such as galvanic cells, oxidation states, reactivity, and oxidation/reduction. Know the oxidation number rules.
  8. Write the chemical formula for lithium bromide. What is the oxidation state of bromine in this compound?
  9. Write the chemical formula for strontium sulfide. What is the oxidation state of strontium in this compound?
  10. Define anode and cathode for a galvanic cell. At which electrode does reduction occur in a galvanic cell? In an electrolytic cell? At which electrode does oxidation occur in a galvanic cell? In an electrolytic cell?
  11. Label each reaction below as either being oxidation or reduction:
  12. Cu  Cu+2 + 2e-
  13. Zn+2 + 2e-  Zn
  14. Something that is a great reducing agent is likely to be oxidized. Something that is a great oxidizing agent is likely to be reduced. Looking at the chart below the most reactive metals are the ones likely to be oxidized. The least reactive are likely to be reduced. Use this chart to determine what metals would be best to make an outdoor statue.
  1. Know how to predict the products of single replacement reactions.
  2. 6Na + Cu3(PO4)2  ______(Don’t forget to balance)
  3. 3Li + Al(NO3)3  ______
  4. 3Mg + Fe2O3 ______
  1. Use concepts of solution chemistry to solve colligative property problems.
  2. List three colligativeproperties.
  3. The freezing point of pure water is 0°C. What is the freezing point of a 2.0m solution of KBr? The Kfof water is 1.86°C/m
  4. Why do we put salt on the roads in Michigan during the winter months?
  5. Which solution shown below will have the highest boiling point?
  6. Pure water
  7. 5 moles of NaCl in 100 mL of water
  8. 10 moles of NaCl in 100 mL of water
  9. 5 moles of NaCl in 50 mL of water
  10. Be able to recognize and problem solve for a wide variety of thermochemistry problems.
  11. Are these reactions exothermic or endothermic?
  12. 2H2 + O2 → 2H2O +483.6 kJ
  13. N2 + 2O2 + 34kJ  2NO2
  14. Heat + H2O2 H2O + O2
  15. S + O2 SO2 + Heat
  16. Is the diagram shown below exothermic or endothermic? For an exothermic reaction are the reactants higher or lower than the products? What would happen to the temperature of this reaction mixture?
  1. Is the diagram shown below endothermic or exothermic? Would the ΔH be positive or negative? In an endothermic reaction are the reactants higher or lower than the products.
  1. Answer the questions below about the following temperature versus time graph of water.

i) What state or states of matter are represented at all numbered regions on this graph?

ii) Identify all regions on the graph where average kinetic energy is increasing.

iii) Identify all regions on the graph where only potential energy is increasing.

iv) During phase 2 – what is happening? Why is this process endothermic?

  1. Label the following phase changes as exothermic or endothermic.
  2. Water is boiling, energy is gained by the system. ΔH is positive
  3. Water is condensing, energy is lost to the surroundings. ΔH is negative
  4. Water is freezing, energy is lost to the surroundings. ΔH is negative.
  5. Know how to apply Hess’s Law to solve for heat:
  6. Calculate ∆H for the reaction 4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g), from the following data.

N2 (g) + O2 (g) 2 NO (g) / ∆H = -180.5 kJ
N2 (g) + 3 H2 (g)  2 NH3 (g) / ∆H = -91.8 kJ
2 H2 (g) + O2 (g) 2 H2O (g) / ∆H = -483.6 kJ
  1. Know how to use the enthalpy of a reaction to figure out how much energy is released or absorbed from a certain mass of material. (We called this Heat and the Reaction or Thermostoichiometry). It is similar to stoich in that we convert grams to moles and then use a mole to heat ratio.
  2. How much heat will be released if 25.0 grams of hydrogen react with nitrogen in the following balanced equation: N2 + 3H2  2NH3 ΔH = -91.8 kJ
  3. Understand the concept of specific heat capacity.
  4. Define Specific Heat
  5. Which would require more energy to increase by 1°C? 20 grams of water or 20 grams of aluminum? Note that water’s specific heat is 4.184J/g°C and aluminum is 0.900J/g°C.
  1. Bonding. Know what combinations of elements create ionic bonds and covalent bonds. Know what bonds are the strongest or weakest (single, double, triple). Know how to draw a Lewis Dot diagram for simple covalent molecules. Understand what happens with energy during bond formation and bond breaking.
  2. Identify the type of bonding present in each pair: (Ionic or Covalent)
  3. Ca with Cl
  4. C with H
  5. Carbon monoxide
  6. KBr
  7. SO2
  8. Magnesium iodide
  9. List single, double, and triple bonds in order of increasing strength. In order of increasing length.
  10. Draw the Lewis Dot Diagram for these simple covalent compounds:
  11. Br2
  12. Cl2
  13. O2
  14. N2
  15. CH4
  16. NH3
  17. H2O
  18. HF
  19. C2H5OH
  20. What does a decomposition reaction do? When the bonds are broken, does this release energy or require energy?
  21. What does a synthesis reaction do? When the bonds are formed, does this release energy or require energy?
  22. Fill in the blanks. When bonds are broken, energy is ______. When bonds are formed, energy is ______.
  23. For the reaction given below, list the bonds that are broken. Make a list of the bonds being formed:

C3H8 + 5O2 4H2O + 3CO2

  1. Study the diagram below. Answer the questions that follow:

Potential Energy Diagram in the formation of a Hydrogen Molecule

Point R, the lowest point on this diagram, represents the distance at which a bond has reached its lowest potential energy and hence represents the distance at which the bond will form. At R, the bond is formed but notice that to the left and right of the R, the potential energy rises. This means that as bonds are pulled apart, the potential energy rises. As the bond is pushed to close together, the potential rises.

  1. Where is the bond most stable? Before R, at R or after R?
  2. What happens to potential energy as we move away from point R
  3. When atoms bond, they achieve a lower, more stable energy – True or False?
  4. As atoms are pulled apart from bonding, their potential energy increases. True or False?
  1. Be able to use the Lewis Dot Diagrams and VSEPR theory to determine whether or not a molecule is polar or nonpolar.
  2. For every molecule listed below, draw the proper VSEPR shape of the molecule (Linear, Trigonal Planar, Tetrahedral, Bent, trigonal pyramidal.) You should already have done the Lewis Diagram for these molecules.
  3. Cl2
  4. O2
  5. CH4
  6. NH3
  7. H2O
  8. HF
  9. C2H5OH
  10. For every molecule below, state whether the molecule is polar or nonpolar. Look for symmetry – a sign of being NONPOLAR! A lone pair on the central atom generally leads to polarity.
  11. Cl2
  12. O2
  13. CH4
  14. NH3
  15. H2O
  16. HF
  17. C2H5OH
  18. Intermolecular forces: Be able to answer a wide variety of questions regarding intermolecular forces of attraction:
  19. What is a London Dispersion force?
  20. What is a Dipole-Dipole Attraction?
  21. What is Hydrogen Bonding?
  22. Rank the three intermolecular forces from weakest to strongest.
  23. What type of intermolecular force is present on each of the following molecules:
  24. Cl2
  25. O2
  26. CH4
  27. NH3
  28. H2O
  29. HF
  30. Rank these three molecules from highest boiling point to lowest: HCl, CH4, H2O. Why did you choose this order?
  31. What three elements can hydrogen bond with to create the Hydrogen Bonding BETWEEN molecules?
  32. What lines in the diagram below represent intramolecular forces? (Dashed or solid?) Which lines represent intermolecular forces? What type of intermolecular forces is being shown in this diagram?
  1. What happens to the strength of the intermolecular forces as water is freezing? As water is melting?
  2. What type of compounds (IONIC OR COVALENT) are sugar (C6H12O6), ammonia (NH3), water (H2O) and butane (C4H10)? Why is sugar a solid at room temperature yet ammonia is a gas? Why does water boil at a higher temperature than butane?
  3. True or False: Solids have greater intermolecular forces than liquids and gases.
  4. True or False: As you go down the halogen family, intermolecular forces increase.
  5. Explain why fluorine and chlorine are gases but bromine is a liquid and iodine is a solid.
  6. Fill in the blank: The higher the intermolecular forces the ______the melting point and the ______the boiling point.
  7. True or False: The IMF with the highest boiling points is hydrogen bonding.
  8. Arrange the following from lowest melting point to highest melting point: A nonpolar covalent, an ionic compound, a polar covalent compound with hydrogen bonding capability, a polar covalent compound with dipole-dipole attractions.
  1. Understand the role entropy plays in chemical and physical changes.
  2. What is the definition of entropy?
  3. Which changes below increase entropy:
  4. Ice melting
  5. Water decomposes into hydrogen gas and oxygen gas
  6. Steam condenses
  7. Dissolved Oxygen gas escapes ocean water as it warms
  8. Understand the difference between chemical and physical change at the molecular level.
  9. True and False: When water melts, the intermolecular forces are weakened and the water molecules stay intact.