Bauer, Birk, and Marks
Instructor Annotations
Chapter 1
Section 1.1 Compounds
• 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.
Section 1.1 Representations of Matter
• Teaching Tip
Provide opportunities for students to draw pictures of matter on a particulate level
whenever possible.
Section 1.1 Representations of Matter
• 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.
Section 1.1 States of Matter
• Misconception
Many students believe that there is no molecular motion in solids.
Section 1.1 States of Matter
• Misconception
Many students think that condensation on the outside of a glass containing a cold
beverage comes from the liquid inside the glass.
Section 1.2 Physical and Chemical Changes and Properties of Matter
• 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.
Section 1.2 Physical and Chemical Changes and Properties of Matter
• Teaching Tip
Emphasize the value of dimensional analysis and that it can be used throughout the
course to solve problems.
Section 1.2 Volume
• Teaching Tip
Tell students that a milliliter of water fills a small thimble.
Section 1.2 Density
• Demonstration

Materials: 2 index cards, four 250 mL Erlenmeyer flasks, red and blue food dye, hot and
cold water.
Preparation: Add a few drops of red food coloring to two 250 mL flasks. Add hot water
until the flasks are completely filled. Swirl to mix the food coloring. Repeat for the other
two flasks using blue food coloring and cold water.
Presentation: Place the index card over the top of one of the flasks containing the hot
(red) water. Carefully invert the flask and place it on top of one of the flask containing
cold (blue) water. Remove the card and ask the students to observe. Have them predict
what would happen if the flask containing the cold (blue) water was placed on the one
containing the hot (red) water. Repeat the demonstration with the flask with hot water on
the bottom. Ask the students to explain their observations. They should discuss how
the density of water changes with temperature on a macroscopic and particulate level.
Safety: Wear safety goggles in case glassware breaks.
Section 1.2 Density
• Demonstrations
Shakhashiri, Bassam Z., “Density and Miscibility of Liquids,” Chemical Demonstrations:
A Handbook for Teachers of Chemistry, Vol. 3 (The University of Wisconsin Press,
Madison, 1985) pp. 229-233.
Summerlin, L., Borgford, C., & Ealy, J., “The Mysterious Sunken Ice Cube,” Chemical
Demonstrations: A Sourcebook for Teachers, Vol. 2 (American Chemical Society, 1988)
pp. 15-16.
Section 1.2 Temperature
• Demonstration
Materials: Galilean thermometer, large beaker containing ice water.
Presentation: Show students the thermometer. Briefly place the bottom of the device
into an ice bath until a change is observable. Have students describe the changes and
explain how the thermometer works.
Section 1.2 Temperature
• Misconceptions
There are many known misconceptions held by students about heat and temperature:
Heat and temperature are the same thing; temperature is transferred, not heat; and a
temperature decrease is due to the introduction of “cold.”
Section 1.2 Temperature
• 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.
Section 1.2 Physical Changes
• Misconceptions
Many students think that a phase change is a chemical change. A very common
misconception is that water boiling produces hydrogen and oxygen gases. Alternatively,
students who believe that a phase change is a physical change think that such a change
is not evident on a particle level.
Section 1.2 Chemical Changes

• Demonstrations
Summerlin, L., & Ealy, J., “Copper into Gold: The Alchemist’s Dream,” Chemical
Demonstrations: A Sourcebook for Teachers, Vol. 1 (American Chemical Society, 1988)
pp. 137-138.
Summerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: Foam
Production,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 1 (American
Chemical Society, 1988) pp. 101-102.
Section 1.4 Scientific Inquiry
• 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.
Section 1.4 Hypotheses
• 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.
Math Toolbox 1.1 Scientific Notation
• Teaching Tip
Give the students a list of positive numbers in scientific notation including ones with
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.
Math Toolbox 1.1 Scientific Notation
• Teaching Tip
Encourage students to make sure they can perform computations on their calculators
using numbers in scientific notation.
Math Toolbox 1.2 Significant Figures
• Teaching Tip
Point out that in chemistry many of the numbers are actually measurements having a
magnitude, unit, and degree of certainty.
Chapter 2
Section 2.2 Subatomic Particles
• 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.
Section 2.4 Atomic Mass
• Teaching Support Activity
Materials: Baggies each containing 12 navy beans, 6 kidney beans, and 8 lima beans
(dried) for each group of 3-4 students.

Loading page 4...

Directions: Hand out one baggie to each group of students and tell them that the beans
represent the atoms of a fictional element named beanium. Tell the students the mass
of each type of atom (bean) and instruct them to count the number of each. Ask them to
find the percentage abundance of each type isotope of beanium and compute the
average atomic mass (relative atomic mass) of the element.
Discussion: Dalton stated in his atomic theory that all atoms of an element are identical.
Did Dalton know about isotopes? Why or why not?
Section 2.4 Atomic Mass
• 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.
Section 2.5 Classification of Elements
• Demonstration
Bring in samples of metals, nonmetals, and metalloids. Show the samples to the
students and have them describe the properties and classify each element.
Chapter 3
Section 3.1 Ionic and Molecular Compounds
• Demonstration
Shakhashiri, Bassam Z., “Conductivity,” Chemical Demonstrations: A Handbook for
Teachers of Chemistry, Vol. 3 (The University of Wisconsin Press, Madison, 1985) pp.
326-328.
Section 3.1 Ionic and Molecular Compounds
• 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.
Section 3.1 Ionic and Molecular Compounds
• 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.
Section 3.2 Monatomic Ions
• 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.
Section 3.2 Polyatomic Ions

Loading page 5...

• 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.
Section 3.2 Polyatomic Ions
• 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.
Section 3.3 Formulas for Ionic Compounds
• Teaching Tip
Point out the distinction between atom, molecule, ion, and formula unit, and stress the
importance of using these terms accurately.
Section 3.3 Formulas for Ionic Compounds
• Teaching Tip
Have the students write the formula for magnesium nitrate with and without the
parentheses and explain why they are needed.
• Demonstration
Bring in vials of FeCl24H2O and FeCl36H2O (or Fe(NH4)2(SO4)26H2O and Section 3.4
Naming Ionic Compounds
FeNH4(SO4)212H2O) to show students the difference between iron(II) and iron(III)
compounds.
Section 3.4 Naming Ionic Compounds
• 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.
Section 3.4 Naming Ionic Compounds
• Teaching Tip
Remind students to include the Roman numeral in the name when it is necessary.
Section 3.4 Naming Ionic Compounds
• Teaching Support Activity
Materials: Various food, vitamin, nutritional supplement, and beverage labels
Directions: Have the students write chemical formulas for as many ionic compounds as
they can.
Discussion Questions: Did most of the names use the common or Stock system? Which
monatomic ions were common among the compounds? Which polyatomic ions were
common among the compounds? For what names or categories of names were you
unable to write formulas?
Section 3.5 Naming and Writing Formulas for Molecular Compounds
• Teaching Tip

Loading page 6...

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.
Section 3.6 Acids and Bases
• 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.
Section 3.6 Acids and Bases
• Teaching Tip
Point out the three different ways to write the formula for acetic acid (HC2H3O2, CH3-
COOH, and CH3CO2H) and for the acetate ion (C2H3O2−,CH3COO−, and CH3CO2−).
Section 3.6 Acids and Bases
• Teaching Tip
Mention that bases commonly are ionic compounds or ammonia, and formulas for
molecular substances ending in –OH are known as alcohols.
Section 3.7 Predicting Properties and Naming Compounds
• 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
Section 4.1 Mole Quantities
• Demonstration
To help students understand the magnitude of 6.022 x 1023, describe the spending
power of a mole of pennies and the volume of a mole of marshmallows. If a mole of
pennies were evenly distributed to each and every person on Earth and they spent $1
million dollars every hour, 24 hours per day, they would have half of their money unspent
at death. One mole of marshmallows would cover the USA to a depth of 105,000 km
(6500 miles). Next, ask students if they can estimate the volume of one mole of water
molecules. After gathering input, show them a graduated cylinder containing 18 mL of
water. Ask students what this tells them about the number of things in a mole and the
size of a water molecule.
Section 4.1 Mole Quantities
• 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).
Section 4.1 Mole Quantities

Loading page 7...

• 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.
Section 4.1 Mole Quantities
• Teaching Tip
Instruct the students not to round the molar masses from the periodic table, but rather to
use all of the available digits.
Section 4.2 Moles, Masses, and Particles
• 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).
Section 4.2 Moles, Masses, and Particles
• 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.
Section 4.3 Determining Empirical and Molecular Formulas
• 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.
Section 4.3 Empirical and Molecular Formulas
• 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?
Section 4.3 Empirical and Molecular Formulas
• 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).

Loading page 8...

Section 4.3 Empirical Formulas for Compounds Containing More Than Two Elements
• 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
Section 5.1 What is a Chemical Reaction?
• Teaching Tip
Students should see as many chemical reactions (demonstrations and animations) as
possible while studying this chapter.
Section 5.1 What is a Chemical Reaction?
• Misconception
Students often misinterpret chemical equations in thinking that the substances on either
side of the arrow represent a different physical entity in a different location. Emphasize
that the reactants and products are separated not by location, but by time and show the
substances present before and after the chemical reaction has taken place.
Section 5.3 Writing Chemical Equations
• 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.
Section 5.3 Writing Chemical Equations
• 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.
Section 5.3 Writing Chemical Equations
• Teaching Tip
Emphasize to students that they must write correct formulas for all reactants and
products before attempting to balance chemical equations.
Section 5.3 Writing Chemical Equations
• Misconception
Students learn quickly to correctly balance chemical equations but often do not
understand what the equations represent, as evidenced by their inability to draw correct
particulate representations of atoms, molecules, and formula units in a reaction. Be sure
to ask them to draw particulate diagrams of chemical reactions from the balanced
chemical equations. Since coefficients and subscripts serve the same mathematical

Loading page 9...

purpose in counting atoms, students do not distinguish between them in particulate
drawings. For example, students often draw two separate spheres to represent 2C
(correct), but also draw two separate spheres to represent O2 (incorrect).
Section 5.4 Decomposition Reactions
• Demonstration
Summerlin, L., & Ealy, J., “Catalytic Decomposition of Hydrogen Peroxide: foam
Production,” Chemical Demonstrations: A Sourcebook for Teachers, Vol. 1 (American
Chemical Society, 1988) pp. 101-102.
Section 5.4 Decomposition Reactions
• Demonstration
Decomposition of Sodium Hydrogen Carbonate
Materials: Large test tube, 10-15 g of NaHCO3, Bunsen burner, test tube holder or
clamp, wood splint, safety goggles
Procedure: Add the sodium hydrogen carbonate to the test tube, light the burner, and
heat the test tube over the flame. While heating the test tube, place a burning splint into
the test tube. The carbon dioxide produced in the decomposition of sodium hydrogen
carbonate will extinguish the burning splint.
Questions:
What are the products of the reaction?
Why did the flame go out?
Section 5.4 Decomposition Reactions
• Demonstration
Bring in small vials of copper(II) sulfate pentahydrate and anhydrous copper(II) sulfate.
Pass them around to show students how hydrates and anhydrous compounds differ in
appearance.
Section 5.4 Combination Reactions
• Demonstration
Burning Magnesium
Materials: 4-6 cm strip of magnesium ribbon, crucible tongs, watch glass, Bunsen
burner, safety goggles
Procedure: Hold magnesium ribbon with crucible tongs. Caution students not to look
directly at the flame. Ignite the Mg with the burner. After the reaction is complete, place
the magnesium on the watch glass.
Questions:
Is there evidence that a chemical reaction occurred? Describe the evidence.
What are the formulas for the reactants in this demonstration?
What is the formula for the product in this demonstration?
What is the balanced chemical equation for this reaction?
Section 5.4 Combination Reactions
• Demonstration
Making Rust
Materials: Small piece of steel wool, crucible tongs, 9-V battery, safety goggles
Procedure: Hold steel wool with crucible tongs. Touch the terminals of the battery to the
steel wool. Once it ignites, gently blow on the glowing steel wool. Tell the students that
the product of this reaction is iron(III) oxide.

Loading page 10...

Questions:
Is there evidence that a chemical reaction occurred? Describe the evidence.
What are the formulas for the reactants in this demonstration?
What is the balanced chemical equation for this reaction?
Section 5.4 Single-Displacement Reactions
• Demonstration
Metals in HCl
Materials: 100 mL of 1M HCl, small pieces of metals such as magnesium, zinc, copper,
and lead, 4 large test tubes, 1 test tube rack, safety goggles
Procedure: Pour about 20 mL of HCl into each test tube. Drop one metal into each test
tube and observe.
Questions:
Which metals reacted with HCl?
How did you know?
What are the products of the reactions in which the metals replaced hydrogen?
Section 5.4 Single-Displacement Reactions
• Demonstration
Summerlin, L., & Ealy, J., “The Aluminum Cola Can Rip-Off,” Chemical Demonstrations:
A Sourcebook for Teachers, Vol. 1 (American Chemical Society, 1988) pp. 152-153.
Section 5.4 Gas Formation Reactions
• Demonstration
Reaction of CaCO3 with HCl
Materials: About 5 g of CaCO3, large beaker (1 L), 500 mL of 1 M HCl, safety goggles
Procedure: Pour the HCl into the beaker. Add the CaCO3 and observe.
Questions:
What are the products of the double-replacement reaction between calcium carbonate
and hydrochloric acid?
What evidence did you see that carbonic acid decomposed?
Section 5.4 Acid-Base Neutralization Reactions
• Demonstration
Acid-Base Neutralization
Materials: 100 mL of 1 M HCl, 100 mL of 1 M NaOH, thermometer, 2 400-mL beakers,
safety goggles
Procedure: Pour the hydrochloric acid into one beaker and the sodium hydroxide into the
other. Measure and report the temperature of each. Pour the solutions back and forth
between the beakers a few times. Take the temperature of the mixture.
Questions:
What evidence of a chemical reaction was observed?
What are the products of a reaction between hydrochloric acid and aqueous sodium
hydroxide?
What class of reaction is this?
Section 1.4 Combustion Reactions
• Demonstration
Ethanol Cannon

Loading page 11...

Solution Manual For Introduction to Chemistry, 4th Edition

Study Now!

Document Details

Related Documents

Test Bank for Financial Accounting, 11th Edition

Test Bank For Chemistry: An Introduction to General, Organic, and Biological Chemistry 13th Edition Test Bank

Solution Manual for Financial Markets and Institutions, 9th Edition

Test Bank For Essentials of Biology, 3rd Edition

Solution Manual for Financial Accounting, 3rd Edition

Solution Manual for Chemistry A Molecular Approach, 4th Edition

Lecture Notes for Laboratory Manual for Principles of General Chemistry , 10th Edition