Chem 139Summer '14

Unit 2

The outcomes below are the learning outcomes for Unit 2 and the content for the exam. These are transcribed from the text sections’ “Goals,” unless marked by an asterisk. *This is an outcome covered in the text, but not stated in the chapter’s “Goal” section. **This is an outcome not in the textbook.

7/21Outcomes:

  1. Given the formula of a chemical compound (or a name from which the formula may be written), state the number of atoms of each element in the formula unit.
  2. Distinguish among atomic mass, molecular mass, and formula mass.
  3. Calculate the formula (molecular) mass of any compound whose formula is given (or known).
  4. Define the term mole. Identify the number of objects that corresponds to one mole.
  5. Given the number of moles (or units) in any sample, calculate the number of units (or moles) in the sample.
  6. Define molar mass or interpret statements in which the term molar mass is used.
  7. Calculate the molar mass of any substance whose chemical formula is given (or known).
  8. Given any one of the following for a substance with a given (or known) formula, calculate the other two: (a) mass; (b) number of moles; (c) number of formula units, molecules, or atoms.
  9. Calculate the percentage composition of any compound whose formula is given (or known).
  10. Given the mass of any compound with a given (or known) formula, calculate the mass of any element in the sample; or, given the mass of any element in the sample, calculate the mass of the sample or the mass of any other element in the sample.

Reading Assignment:§7.1 through 7.6

Prepared problems:(next class)

7/22Outcomes:

  1. Distinguish between an empirical formula and a molecular formula.
  2. **Explain the equivalence of ionic formula, empirical formula, and formula unit for ionic compounds.
  3. Given data from which the mass of each element in a sample of ca compound can be determined, find the empirical formula of the compound.
  4. Given the molar mass and the empirical formula of a compound, or information from which they can be found, determine the molecular formula of the compound.
  5. Describe five types of evidence detectable by human senses that usually indicate a chemical change.
  6. **Give at least two types of evidence not detectable by human senses, but easily detected by instrumental methods.
  7. Distinguish between an unbalanced and a balanced chemical equation, and explain why (**and when) a chemical equation needs to be balanced.
  8. Given an unbalanced chemical equation, balance it by inspection.

Reading Assignment:§7.7, 7.8; 8.1 – 8.3

Prepared Problems:Ch. 7, #1, 7 (a, d, e), 15, 17, 21 (a,d), 23 (a, c), 27 (a, c), 43 (d, e), 47

Note: If this is unfamiliar content, you are urged to do additional problems.

7/23Outcomes:

  1. Given a balanced chemical equation or information from which it can be written, describe its meaning on the particulate, molar, and macroscopic levels.
  2. Write the equation for the reaction in which a compound is formed by the combination of two or more simpler substances.
  3. Given a compound that is decomposed into simpler substances, either compounds or elements, write the equation for the reaction.
  4. Given the reactants of a single-replacement reaction, write the equation for the reaction. *To recognize that these are usually also “oxidation-reduction” reactions.
  5. Given the reactants in a double-replacement or neutralization (*and evidence that it proceeds), write the equation.
  6. Distinguish among strong electrolytes, weak electrolytes, and non-electrolytes.
  7. Given the formula of a (*soluble) ionic compound (or its name), write the formulas of the ions present when it dissolves in water (*using the format of a chemical reaction).
  8. Explain why the solution of an acid may be a good conductor or a poor conductor of electricity.
  9. Given the formula of a soluble acid (or its name), write the major and minor species present when it is dissolved in water (*using the format of a reversible reaction, if appropriate).

Reading assignment:§8.4 through 8.10; 9.1 – 9.3

Prepared problems:Ch.7, #55, 57, 65; Ch. 8 Practice Exercises 8.1, 8.2, 8.3 (pp.216 - 218) .

7/28 Outcomes:

  1. Distinguish among conventional (**“whole-formula”), total ionic, and net ionic equations.
  2. *Given a conventional reaction equation, recognize spectator ions and be able to write its net ionic equation.
  3. (Goal #6, modified:)* Given a single replacement reaction, be able to write its net ionic equation.
  4. Write the equation for the complete combustion or burning of any compound containing only C and H or only C, H, and O.
  5. *To memorize a small set of solubility “rules” and to be able to look up and interpret others. (See below.)
  6. Predict whether a precipitate will form when known solutions are combined; if a precipitate forms, write the net ionic equation.
  7. Given the product of a precipitation reaction, write the net ionic equation.
  8. Given reactants for a double-replacement reaction that yields a molecular product, write the conventional, total ionic, and net ionic equation.
  9. (§9.9 -9.12) *Given information about specific reactants or products of other double-replacement reactions, write net ionic equations.

Reading Assignment:§9.4 through 9.8; skim 9.9 – 9.12

Prepared problemsCh. 8, #3, 7, 11, 23 (CaF2 is a precipitate), 29; Ch. 9, #3, 5, 11 (why?)

A Summary of Solubility and Ionic Reaction “Rules:”

Solubility in water, hierarchy of decisions (there are some exceptions):

  1. Any group 1A cation or compound with a group 1 cation or ammonium ion is soluble.
  2. Any compound with acetate or nitrate ion is soluble.
  3. Any group 7A anion or a compound with a group 7A anion is soluble, except with Ag, Hg or Pb.
  4. If a compound fits none of the above, its solubility needs to be looked up.

Reactants or products that are solids are written in whole-formula (“conventional”) form, not ionized.

Reactions involving H+ plus OH- as reactants form H2O as product.

Reactants that are weak acids, and products that are weak acids are written in molecular form, not ionized.

7/29 Outcomes

  1. Given a chemical equation, or a reaction for which the equation is known, and the number of moles of one species in the reaction, calculate the number of moles of any other species.
  2. Given a chemical equation, or a reaction for which the equation can be written, and the number of grams or moles of one species in the reaction, find the number of grams or moles of any other species.
  3. Given two of the following, or information from which two of the following may be determined, calculate the third: ideal (theoretical) yield, actual yield, percent yield.
  4. Identify and describe or explain limiting reactants and excess reactants.
  5. *Use the format of an **“i.c.f.” table to identify limiting reactants in mole units. **Recognize the outcome if a negative number of final moles is suggested.

Reading Assignment:§10.1 through 10.4

Prepared Problems:Ch.9, #13, 17, 21 (Ba2+ forms), 25, Practice Exercise 9.12 page 261, chapter-end #29 (PbI2 is insoluble), 43, 45 (CO2(g) is one product), 71 (molecular HF forms).

7/30 Outcomes:

  1. Given a chemical equation, or information from which it may be determined, and initial quantities of two or more reactants, (**use the format of an “i.c.f” table to) (a) identify the limiting reactant, (b) calculate the ideal yield of a specified product assuming complete use of the limiting reactant, and (c) calculate the quantity of the reactant initially in excess that remains unreacted.
  2. *Demonstrate the use of the “smaller amount method” to accomplish the above.
  3. Define and describe electromagnetic radiation.
  4. Distinguish between continuous and line spectra.
  5. Describe the Bohr model of the hydrogen atom.
  6. Explain the meaning of quantized energy levels in an atom and show how these levels relate to the discrete lines in the spectrum of that atom.
  7. Distinguish between ground state and excited state.

Reading Assignment:§10.5, 10.6; 11.1 – 11.2 (Chapter 11 testable in Unit 3).

Prepared Problems:Ch. 10; in all problems clearly show the chemical reaction and

conversion factors you use from the coefficients! #1, 3, 5, 7, 21, 27, 31.

Graded Homework(due at start of class):

Ch.7 #26, 44, 60; Ch.8 #12, 32; Ch.9 #16, 44; Ch.10 #22, 46

8/4 Laboratory Exercise Outcomes

  1. To applyskills of measurement in order to accomplish an experimental task.
  2. To understand and use chemical reaction equation coefficients.
  3. To determine and critically evaluate the ratio of moles in a chemical reaction.
  4. To use energy as a detection method to determine when we have run out of a reactant.
  5. To practice percent by mass and molar mass conversions as a tool in critically evaluating reaction coefficients.

Please come to CC330 at class time. Safety goggles are required and will be provided for you.

Reading Assignment:Handout: “What are Mole Relationships in Chemical Reactions?”

Prelaboratory Assignment:On handout, turn in at the start of class.

Report due:At start of class, 8/6

8/5 Outcomes

  1. To assess and review what we have learned, and practice problem solving.
  2. Success on an hour exam covering the above outcomes, up to but not including Chapter 11.

Prepared Problems:Ch.10, #41, 47, 49; Ch. 11, #5, 7, 16.

Group Sheet 2