Day 1

Matter

  1. Describe the density of the three states of matter of water. Which state of water is most dense? How do you know?
  2. What is the difference between a chemical and physical change?
  3. What is the difference between chemical and physical properties?
  4. Is rusting a chemical or physical change?
  5. What is necessary for an iron nail to rust?
  6. Does an iron nail that has rust on the surface weigh more or less than the original nail? Why is this?
  7. Describe the difference between homogeneous and heterogeneous mixtures.
  8. How do you know that a bowl of cereal is heterogeneous and not homogeneous?
  9. What is the conservation of mass?
  10. Create a list of 5 physical properties and 5 chemical properties
  11. Create a list of 5 chemical changes and 5 physical changes.

Atomic Structure

  1. List three parts of the atom and describe their location. Which one has the least mass?
  2. Which part of the atom determines its identity?
  3. Which part of the atom determines the chemical reactivity of the element?
  4. Write the complete nuclear symbols for the following. Be sure to label the mass number, atomic number and charge
  5. Oxygen-16carbon-14uranium-238potassium-40 ion
  6. Write the ground state electron configurations for the following atoms
  7. Lithiumneonsulfurarsenic

Periodic Table/Trends

  1. What is an oxidation number? What is another name for oxidation number?
  2. Write the oxidation numbers for the following elements
  3. Calcium, Potassium, Aluminum, Chlorine, Neon, Nitrogen, Carbon, Oxygen
  4. Describe the trends for the following properties
  5. Atomic Radius
  6. Ionic Radius
  7. Ionization Energy
  8. Electronegativity
  9. Define electronegativity.
  10. Describe the chemical properties of the following families on the periodic table
  11. Alkali metals, alkaline earth metals, halogens, noble gases
  12. Describe physical properties of the following
  13. Metals
  14. Non-metals
  15. Metalloids
  16. What is the periodic table considered to be “periodic?”

Day 2

Bonding, Naming Compounds and Writing Formulas

  1. Differentiate between ionic and covalent bonds.
  2. What is the difference between a molecule and formula unit.
  3. Write the symbols for the following ions. Include the oxidation numbers
  4. Potassium ionoxidesilver (I)iron (III)
  5. Write formulas for the following compounds
  6. Sodium bromide
  7. Ammonium nitrate
  8. Sodium chromate
  9. Carbon disulfide
  10. Dinitrogenpentoxide
  11. What is the common name for the following compounds
  12. H2ONH3HCl
  13. Write names for the following formulas
  14. LiOH
  15. BaSO4
  16. (NH4)3PO4
  17. NCl3
  18. SF6
  19. Complete the following
  20. Covalent compound : atoms :: ionic compound : ______
  21. Formula unit : ionic compound :: ______: covalent compound
  22. Non-metal plus non-metal : covalent :: ______: ionic
  23. ionic : donate :: covalent : ______
  24. Draw the Lewis or electron dot structures for the following. How many lone pairs of electrons to each have?
  25. NaBrOSiAr
  26. Draw Lewis structures and label the geometry for the following molecules
  27. CH4NH3H2OHCl
  28. How many bonds and lone pairs of electrons surround the central atom for the following molecular geometries
  29. Tetrahedral
  30. Trigonal pyramidal
  31. Bent
  32. Covalent bonds exist when two ______are ______between two atoms. A double bond is when ______electrons are shared and is represented by ______lines.

Equations and Reactions

  1. Describe what coefficients and subscripts tell you
  2. Which of the two do you use to balance an equation?
  3. Why can you never alter a subscript once the correct formula has been written?
  4. Write a balanced equation that shows the production of solid sodium chloride (NaCl) from elemental sodium and diatomic chlorine (Cl2). What type of reaction is this?
  5. If heat is produced in this reaction, which side of the equation would you write the word “heat?”
  6. Write a balanced equation that shows the formation of water from its elements. Don’t forget about “HONClBrIF.”

Day 3

Mole stuff and stoichiometry

  1. Describe how you find the molar mass of a substance.
  2. Determine the mass percent of carbon, hydrogen and oxygen in sucrose (C12H22O11).
  3. Determine the molar mass of the following
  4. NaClH2OC12H22O11(NH4)2SO4
  5. How many molecules are in 4 moles of glucose (C6H12O6)?
  6. Which of the following would have the greatest mass? You must do some legwork for this one!
  7. 10 g of NH310 mol H210 molecules of CO210 L O2 at STP
  8. C3H8 + 5 O2  3 CO2 + 4 H2O
  9. How many grams of carbon dioxide are produced when 1 mole of propane (C3H8) reacts?
  10. How many moles of water are made when 6 moles of carbon dioxide are produced?
  11. How many grams of oxygen are needed to produce 18 grams of carbon dioxide?

Gases

  1. List the postulates of the kinetic molecular theory.
  2. How does the speed of gas molecules relate to the temperature of the gas?
  3. How does the speed of the gas molecules relate to the pressure of a gas?
  4. Why would doubling the temperature of a gas double its pressure?
  5. How does the number of collisions between gas molecules and its container relate to pressure?
  6. How does the volume of gas change when the pressure increases? Assume temperature remains constant.
  7. CH4 + 2 O2  CO2 + 2 H2O

What volume of carbon dioxide can be produced from the combustion of 0.6 L of methane (CH4) at STP?

  1. Use the following formula to solve the following (P1V1)/T1 = (P2V2)/T2
  2. What is the new pressure of 50 mL of gas that is compressed to 25 mL when the original pressure was 5.0 ATM. Assume temperature is constant.
  3. 10 L of a gas at 500 K and pressure P are compressed to 5 L and cooled to 250 K. What is the new pressure in terms of P? You do not use a calculator for this.
  4. Set up this problem but do not solve it. A sample of gas occupies 2.0 L at 760 torr. If the temperature is kept constant, what would the volume be if the pressure was raised to 1520 torr?

Nuclear Chemistry

  1. List three types of radioactive decay
  2. Write the nuclear symbols for the three types of radioactive decay.
  3. Show the complete nuclear equation for the alpha decay of uranium-238.
  4. Arrange the three types of radiation in order of increasing energy.
  5. Radioactive isotopes have unstable nuclei. When energy is emitted by the nuclei how does the stability of the daughter isotope change?
  6. Uranium-235 is the isotope of uranium used to make uranium nuclear bombs. What does the 235 represent?
  7. Very heavy isotopes of an element are usually highly unstable. We know this because they have extremely short half-lives. Why do you think we know so little about the chemistry of elements above uranium on the periodic table?

Day 4

Solutions

  1. What is the formula for molarity?
  2. What does molarity measure?
  3. Use the formula M = mol/L to answer the following
  4. What is the molarity of a solution made by dissolving 2 moles of sugar in enough water to make 4 liters of solution
  5. Determine the molarity of a solution made by dissolving 5.8 grams of NaCl in enough water to make 0.5 L of solution.
  6. Water is a very polar molecule. What does this mean?
  7. Why is water considered to be the “universal” solvent?
  8. What does the phrase “like dissolves like” mean?
  9. Differentiate between solvent, solute and solution.
  10. There are three ways to increase the rate or speed at which a solute dissolves in a solvent. Think about dissolving sugar in ice tea. What are three things you could do to increase the rate of dissolution?
  11. What is a “test” you could perform on a solution to see if it was saturated, unsaturated or supersaturated?
  12. What would the results of the previous test be for each case?
  13. What is an electrolyte?
  14. What is a non-electrolyte?

Acids and Bases

  1. Describe the pH for acids, bases and neutral solutions.
  2. How would the pH of a dilute basic solution and concentrated basic solution compare?
  3. What is the definition of an Arrhenius Acid and Base?
  4. What does an Arrhenius Acid produce in a solution?
  5. What is the definition of a Bronsted-Lowry Acid and Base?
  6. What does a Bronsted-Lowry acid do?
  7. How do acids taste?
  8. How do bases taste?
  9. Should you ever taste a solution in the lab to determine if it is acid or base?
  10. What should you use to determine if a solution is acidic or basic in the lab? List several options.

Thermochemistry

  1. What is calorimetry?
  2. What is the definition of a calorie?
  3. How does a calorie compare to a food calorie?
  4. Define endothermic and exothermic processes.
  5. Use the equation h = mCΔT to solve the following problems.
  6. How many joules of heat are needed to heat 100 grams of aluminum from 20 OC to 100 OC if aluminum has a specific heat of 0.90 J/g oC?
  7. Water has a specific heat of 4.184 J/g OC. How much heat is required to raise 2300 grams of water from 25oC to boiling?
  8. If energy is released in an exothermic process, do the products have more or less energy than the reactants?
  9. If energy is absorbed in an endothermic process, do the products have more or less energy than the reactants?
  10. ΔH is the symbol used to describe the heat change of a chemical reaction. What does negative value for ΔH indicate? Which side of the equation would the value for heat be placed on if it had a positive ΔH?
  11. How does the heat content of the surroundings change for endothermic processes?