Bauer 2e
Teaching Tips
Chapter 1
Page 8
· Teaching Tip
Ask students to list compounds they encountered today.
Page 9
· Teaching Tip
Ask students to generate a list of examples of matter in the room. In small groups, have them classify the examples as mixtures or pure substances. Then require them to classify the pure substances as elements or compounds. Lastly, have them classify the mixtures as heterogeneous or homogeneous.
Page 11
· Teaching Tip
Provide opportunities for students to draw pictures of matter on a particulate level whenever possible.
Page 12
· Teaching Tip
Given particulate representations, students often mistake diatomic elements for compounds. In order for a molecule to be classified as a compound, there must be two or more atoms of different elements.
Page 15
· Teaching Tip
Rather than teaching students to convert from English to SI units, provide them with practical examples to allow them to think in the SI system. Tell them that a small paper clip has a mass of approximately 1 gram.
Page 15
· Teaching Tip
Emphasize the value of dimensional analysis and that it can be used throughout the course to solve problems.
Page 17
· Teaching Tip
Tell students that a milliliter of water fills a small thimble.
Page 23
· Teaching Tip
Students often do not recognize that pure substances melt and freeze at the same temperature, even a substance with which they have extensive experience, such as water.
Page 31
· Teaching Tip
When discussing scientific inquiry, emphasize an iterative process employing different activities (observation, asking questions, designing experiments, collecting data, developing models, etc.) rather than a series of sequential steps.
Page 32
· Teaching Tip
Many students have trouble distinguishing a prediction from a hypothesis. Use the example with the pennies to point out a prediction based on Anna and Bill’s hypothesis: If Bill and Anna scratched the pennies more, then the reaction would proceed faster.
Page 35
· Teaching Tip
Give the students a list of positive numbers in scientific notation including ones with positive and negative exponents. Ask them to place them in order from smallest to largest. Add zero to the list. Many students will place zero between the numbers with negative and positive exponents. Be sure they understand that the sign on the exponent does not relate to the whether or not the number is more or less than zero.
Page 35
· Teaching Tip
Encourage students to make sure they can perform computations on their calculators using numbers in scientific notation.
Page 35
· Teaching Tip
Point out that in chemistry many of the numbers are actually measurements having a magnitude, unit, and degree of certainty.
Chapter 2
Page 58
· Teaching Tip
A chocolate chip cookie is more familiar to students than plum pudding. Ask the students to explain what the chocolate chips and the cookie dough represent if the cookie resembles Thomson’s model.
Page 69
· Teaching Tip
Students are likely to already be familiar with weighted averages from the experience of computing their grades in a course in which the various course components (exams, quizzes, labs, homework, etc.) are not weighted equally. Make up a problem with grades for a hypothetical student and have the class find the average.
Chapter 3
Page 88
· Teaching Tip
Emphasize that dissolving describes behavior on a macroscopic level and substances that dissolve do not necessarily dissociate. Give the formula for sucrose and ask students if it is ionic or molecular. Ask if sucrose dissolves in water and then if it dissociates in water.
Page 90
· Teaching Tip
Point out to students that the process of sodium chloride dissolving is a physical change, although we represent it with a chemical equation.
Page 91
· Teaching Tip
Students often memorize the charges of many common monatomic ions by their positions in the periodic table without thinking about the number of protons and electrons in each particle or its position relative to the noble gases. Emphasize that monatomic ions in Groups I, II, V, & VII have characteristic charges that give each of them the number of electrons as the nearest noble gas.
Page 95
· Teaching Tip
After students identify the relationship between the names and formulas for most of the oxoanions, recommend that students memorize the common polyatomic ions with the suffix
–ate, such as sulfate, nitrate, carbonate, phosphate, and chlorate, and simply take one oxygen away to form the polyatomic ion with the suffix –ite.
Page 96
· Teaching Tip
Select 10-15 common polyatomic ions for students to memorize. Not having to refer to their notes saves time when writing formulas and naming compounds.
Page 97
· Teaching Tip
Point out the distinction between atom, molecule, ion, and formula unit, and stress the importance of using these terms accurately.
Page 98
· Teaching Tip
Have the students write the formula for magnesium nitrate with and without the parentheses and explain why they are needed.
Page 102
· Teaching Tip
Some students worry about knowing all the possible charges on ions for metals that form more than one ionic compound. Remind them that the Roman numeral in the name tells the charge, and if they are naming, they can deduce the charge of the metal from the charge on the anion.
Page 102
· Teaching Tip
Remind students to include the Roman numeral in the name when it is necessary.
Page 106
· Teaching Tip
To help students memorize the Greek prefixes, remind them of the common ones they are likely to already know. Bicycle – two wheels; tricycle – three wheels; octopus – eight tentacles; etc.
Page 107
Teaching Tip
The treatment of acids and bases here is in the context of nomenclature and is very brief. Acids and bases are presented in detail in Chapter 13.
Page 108
· Teaching Tip
Point out the three different ways to write the formula for acetic acid (HC2H3O2, CH3COOH, and CH3CO2H) and for the acetate ion (C2H3O2-,CH3COO-, and CH3CO2-).
Page 108
· Teaching Tip
Mention that bases commonly are ionic compounds or ammonia, and formulas for molecular substances ending in –OH are known as alcohols.
Page 111
· Teaching Tip
Students often have a great deal of difficulty recognizing when to apply certain naming/formula writing rules. Recommend that they practice with the flowchart (Figure 3.37), but emphasize that they should eventually be able to correctly write names and formulas without it.
Chapter 4
Page 127
· Teaching Tip
Although a mole is analogous to a dozen, the dozen analogy seems to break down for students. Twelve is a relatively small number and students do not see the necessity in using dozens when they can just report a counting number. To help them to understand the utility of the mole, better alternatives are the ream (as in 500 sheets of paper) and the gross (144 items).
Page 127
· Teaching Tip
Students can easily see that a CO2 molecule is composed of 1 carbon atom and 2 oxygen atoms. They get confused however when this is scaled up to moles to say that 1 mole of CO2 is comprised of 1 mole of carbon atoms and 2 moles of oxygen atoms. They often say that 1 + 2 does not equal 1. Emphasize that the C and O atoms are subsets of CO2 molecules much like individual shoes are subsets of pairs of shoes. Ask how many right shoes they have if they have a dozen pairs, and relate it back to atoms and molecules.
Page 127
· Teaching Tip
Instruct the students not to round the molar masses from the periodic table, but rather to use all of the available digits.
Page 128
· Teaching Tip
Students know how to use percentages to compute their grades on assignments. Remind them that a grade of 88% on a test is the ratio between a part (points earned) and the whole (points possible).
Page 129
· Teaching Tip
Point out that students should evaluate their answers to be sure they make conceptual sense. Even students using dimensional analysis may divide when they should multiply and vice versa. For example, if there are 63.55 g in 1 mol of Cu, then 1.27 mol (more than 1 mol) should contain more than 63.55 g. Therefore, if students divide 63.55 g by 1.27 mol, they should recognize that this is not correct. Encourage them to estimate before they calculate and evaluate after.
Page 133
· Teaching Tip
Provide a list of formulas and have students classify them as empirical or molecular. Ask them to write the empirical formula for each one they identified as a molecular formula.
Page 133
· Teaching Tip
Pose these questions to check for understanding: When is an empirical formula and a molecular formula the same? What are examples of compounds with the same empirical and molecular formulas? Ask why the formulas for most ionic compounds are empirical formulas. Can you come up with formulas of ionic compounds that are not empirical formulas?
Page 134
· Teaching Tip
Students tend to refer to all formulas for compounds as molecular, even ionic ones. Emphasize that the term molecule only refers to particles comprised of nonmetallic atoms and that formula units show the simplest whole number ratio of cations (metallic or NH4+ ions) to anions (nonmetallic monatomic or polyatomic ions).
Page 137
· Teaching Tip
Remind students that when converting from grams to moles for hydrogen, oxygen, or any other diatomic element in an empirical or molecular formula problem, they are determining the ratio of atoms, not molecules. This means that the appropriate molar masses for hydrogen and oxygen are 1.008 g/mole and 16.00 g/mole, respectively. You can also remind them that hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, and iodine are diatomic as free or uncombined element, whereas here they have combined with other elements and formed a compound.
Chapter 5
Page 160
· Teaching Tip
Students should see as many chemical reactions (demonstrations and animations) as possible while studying this chapter.
Page 161
· Teaching Tip
Emphasize that coefficients in chemical equations are mathematically analogous to coefficients in algebra. For example, in 2xy the two coefficient doubles both quantities x and y, just as 2MgO in a chemical equation means that there are two magnesium ions and two oxygen ions present in two formula units of magnesium oxide.
Page 163
· Teaching Tip
Before practicing balancing equations, give students the opportunity to practice counting atoms in chemical formulas, such as in (NH4)3PO4, NH4NO3, and Al2(CO3)3, to ensure they are able to correctly interpret subscripts and parentheses.
Page 164
· Teaching Tip
Emphasize to students that they must write correct formulas for all reactants and products before attempting to balance chemical equations.
Page 187
· Teaching Tip
To review the different types of reactions, show the “Reaction Types” animation again.
Page 187
· Teaching Tip
Students often try to write the net ionic equation directly from molecular equation and make errors. Given the molecular equation, encourage them to first write the ionic equation, identify and cancel the spectator ions, and then write the net ionic equation.
Chapter 6
Page 203
· Teaching Tip
Relating the coefficients in balanced chemical equations to ingredients in simple recipes helps students to understand the meaning of the coefficients and the ratios of reactants to products. Use simple recipes such as 2 pieces of bread + 1 piece of bologna + 1 piece of cheese ® 1 bologna sandwich. Later, you can remind students of the recipe analogy when they don’t understand how two moles of hydrogen combines with one mole of oxygen to form two moles of water (2 + 1 ≠ 2).
Page 206
· Teaching Tip
One useful way to help students understand that coefficients represent the mole ratios of the substances in the chemical equation is to scale up the coefficients from molecules to moles. For example, point out the 1:5 ratio between propane and oxygen in the balanced equation for the combustion equation. If there were two molecules of propane, ten molecules of oxygen would be required. If there were 1000 molecules of propane, 5000 molecules of oxygen would be required. If there were 6.02 x 1023 molecules of propane, 5(6.02 x 1023) molecules of oxygen would be required.
Page 206
· Teaching Tip
Emphasize that the coefficients in balanced chemical equations represent mole, not mass, ratios of the substances in the reaction.
Page 217
· Teaching Tip
Perform the Zn and HCl demonstration from Section 6.4 again and tell students the masses of the reactants in each flask. Have students calculate the limiting reactant for each flask. The students may also determine the mass of hydrogen produced in each flask.
Page 219
· Teaching Tip
Emphasize that whenever students calculate the mass of a product from the mass of the limiting reactant, it is the theoretical yield.
Page 220
· Teaching Tip
Remind students that a percentage is the ratio of a part compared with its whole and that the denominator is always larger than the numerator.
Page 224
· Teaching Tip
Ask students if they would rather sit on a metal or wooden bench on a hot day. Although they might not be able to explain why, most students will choose the wooden bench. This is an excellent opportunity to relate the concepts of specific heat and thermal conductivity to their prior experiences. Point out that even if both benches are at the same temperature and more heat is stored in the wooden bench (wood has a higher specific heat than metals), the metal has a greater thermal conductivity, so heat so the heat can be transferred to skin faster, causing a burn.