experimenting with mixtures, compounds, and elements changes

Experimenting with Mixtures, Compounds, and ElementsChanges to the Teacher’s Edition

andExperimenting with Mixtures, Compounds, and Elements

Changes to the Student Guide

Since publication of the STC–Secondary unit Experimenting with Mixtures, Compounds, and Elements Teacher’s Edition and the Experimenting with Mixtures, Compounds, and Elements Student Guide, chemical formulas and scientific explanations in both books have been revised. In the Teacher’s Edition, these changes affect Tab 4, Lessons 1–9, 11, 12, and 13. In the Student Guide, these changes affect Lessons 1, 3, 4, 6, 8, 9, and 10. Please replace the pages in your texts with the revised pages provided.

This errata set includes the following:

Photocopy and distribute these new instruction pages as needed.

If you have questions about these changes or about the unit in general, call Carolina’s product information staff at 800-227-1150 (8 am–5 pm ET, M–F), or email [email protected].

1206

• FortheExperimenting with Mixtures, Compounds, and ElementsTeacher’s Edition–Tab4revisedpages1-B,1-D,5,6,7,8,13-B,13-E,13-F,13-G,27,31-A,31-B,41,41-A,41-D,57-D,58,67,77-D,89-D,89-E,95,105,128,135-C,145-D,151-A,154,155,161-A,and161-B.

• FortheExperimenting with Mixtures, Compounds, and ElementsStudent Guide–revisedpages5,7,8,27,41,58,67,95,105,and128.

LESSON AT A GLANCE

STC Physical Science Strand: Matter and Change 1–B

INQUIRIES 1.1–1.8

Follow instructions to perform a circuit of inquiries into the nature of pure substances and mixtures.

Make accurate measurements of volume and temperature.

Explain observations using students’ own words and ideas.

Work cooperatively with lab partner.

Practice safe and appropriate laboratory techniques.

Substances vary in their capacity to conduct electricity.

Filtering can be used to separate an insoluble solute in a mixture.

Light and heat energy are evidence of chemical change.

A magnet can separate the magnetic components of a mixture.

Metals vary in their rates of reaction with acids.

Particle size and dispersal determine whether a substance appears to be a heterogenous or homogenous mixture (at the scale of observation used in this unit).

Chemical reactions may produce gas, cause a change in temperature, and/or cause a change in color.

Combining two solutions produces a chemical (and color) change and new products.

Students perform a circuit of eight inquiries, which introduces them to some basic concepts of the properties and behavior of pure substances and mixtures. They record, discuss, and compare their results.

2.0 periods

Content Standard A• Abilities necessary to do scientific inquiry

Content Standard B• Properties and changes of properties in matter

Chemical reaction

Compound

Conductor

Dissolve

Effervescent

Element

Filter

Insoluble

Magnetic

Mixture

Physical change

Product

Reactant

Reactivity

Soluble

STC Physical Science Strand: Matter and Change 1–D

COMMON MISCONCEPTIONS

The inquiries may reveal the following common student misconceptions about matter, pure substances, and mixtures:

• Students may incorrectly assume that only elements (not compounds) are pure substances.

• Students may incorrectly believe that pure substances are transparent, free from additives, and safe to ingest.

• Students may incorrectly assume that all mixtures will appear heterogenous.

• Students may incorrectly give mixtures properties that are independent of their composition.

• Students may confuse compounds with mixtures or elements.

• Students may fail to recognize that elements combine in fixed ratios to form compounds.

• Students may not have observed the production of a gas, a change in temperature, or color change as an example of a chemical reaction.

• Students may confuse dissolving with a chemical reaction.

• Students may use the terms “atoms,” “molecules,” and “electrons” freely with little or no understanding of where they come from or what they mean. The notion of developing concepts based on their observation of scientific phenomena and collection of experimental data may be new for many of them.

The inquiries will reveal misconceptions that relate to other concepts. Some of these are discussed in the “Assessment” section on page 13-A, and many will be addressed in later lessons in which students are able to revisit their earlier ideas and reassess them in the light of the inquiries they have performed.

TEACHER’S NOTES

STC Physical Science Strand: Matter and Change 5

STC Physical Science Strand: Energy in our World 5

SAFETY TIPS Wear safety goggles at all times during the inquiries.

Keep long hair tied back.

Be careful with hydrochloric acid. If acid gets on your clothes or skin, wash it off immediately with lots of cold water.

Tell your teacher immediately about any accidents involving acid.

4

Share your ideas and examples in a class discussion.

5

In this lesson, you will investigate the properties and behaviors of a variety of mixtures and pure substances. Working with another student, you will complete eight inquiries. You will be graded on the seriousness of your efforts, on the carefulness of your observations, and on your cooperation with your lab partner.

6

Listen carefully to the safety instructions given by your teacher. In order to participate in the lab work for this class, you are expected to follow safe laboratory procedures.

7

Each inquiry station has a basic topic. Each pair of students will start at a different inquiry station. At each station you will follow the instructions on the Inquiry Card. These instructions are also in your student guide. When you make observations or think you can explain what you are observing, you should discuss these ideas with your partner. Remember: Exchanging ideas with others is a very important part of science. You will have 4-5 minutes to complete each inquiry and record your observations on Student Sheet 1: Our Ideas About Pure Substances and Mixtures.

8

Your teacher will set up two sets of identical inquiries—A and B. You will be assigned to either A or B. Place an asterisk beside your first inquiry on the student sheet so that you write your answers in the correct spaces.

9

When you have completed each inquiry, put the apparatus back as you found it at the beginning of the experiment.

10

If you have any questions about the procedure, you should ask your teacher now.

STC Physical Science Strand: Energy in our World 5 STC Physical Science Strand: Matter and Change 5

assigned is the number of their first inquiry. Have students find their assigned inquiry on Student Sheet 1 and mark it with an asterisk (*) to reduce confusion.

9. Remind students to return the apparatus at each station to its original condition when they finish the inquiry.

10. Ask students whether they have any questions regarding the procedure. You may ask one group to review briefly what they will be doing for each inquiry.

7. You can also mention the basic topic of each inquiry (“Here you investigate the ability of metals to conduct electricity … the magnetic properties of some substances … the separation of a substance.”).

Explain the following directions for the circuit of inquiries:

• Students should follow the instructions on the card at each inquiry station.

• Students have 4–5 minutes to perform each inquiry. A warning will be issued when they have

2 minutes left, telling them to complete the exercise and clean up (adjust time as necessary).

• Students should leave each station as they found it.

• Students should record their answers in the appropriate spaces on Student Sheet 1 (stress that you will not deduct points for incorrect or incomplete conclusions).

8. Assign a station number (1.1A through 1.8A and 1.1B through 1.8B) to each pair of students. Explain that the number they have been

NOTE Ideally, you will be at this point in the lesson by the end of the first period;

students will have the next two periods to complete the inquiries and to reflect on their experiences.

LESSON

PROCEDURE

1

Set up a circuit with the batteries in holders, a lightbulb, a bulb holder, and the connector wires (see Figure 1.1). Touch the unfastened alligator clips together and observe. Answer the following question on Student Sheet 1: What happens to the bulb when you complete the circuit by touching the alligator clips?

2

Place the copper cylinder in the circuit between the alligator clips, and observe the bulb. Answer the following question: What happens to the lightbulb when the copper cylinder is placed in the circuit?

3

Connect the ends of the small pencil with the alligator clips, and observe the bulb. What happens to the bulb with the pencil in the circuit? Record your observations and ideas.

4

Finally, use the connector wires to put the piece of zinc in the circuit, and observe the bulb. Answer the following question: What happens to the bulb with the zinc in the circuit?

5

What conclusions can you make about the ability of these substances to conduct electricity? Record your ideas.

6

Disconnect the circuit, and return the apparatus to its original condition for the next group.

FINdINg THE CONduCTOR

place each of the objectS to be teSted in the circuit, one at a time.FIgUre 1.1

iNQuirY 1.1

6 STC Unit: Experimenting with Mixtures, Compounds, and Elements

The naTure of MaTTerlesson 1


PROCEDURE FOR THE CIRCUIT OF INQUIRIES

1. Send each pair of students to their first inquiry and have them begin. Circulate among the stations. Use this opportunity to observe student procedures and to listen to and discuss each pair’s observations and conclusions.

2. Allow 4–5 minutes for students to work at each station. When students have 2 minutes left, let them know they should start wrapping up the inquiry. You may wish to ring a bell or buzzer or to set up an automatic timing system. Encourage students to move quickly to the next inquiry.

3. You may need to replenish petri dishes or filter paper for Inquiries 1.2, 1.4, and 1.8.

4. The inquiries may continue into the third period of this lesson. When all of the students have completed the entire circuit, give them a few minutes to complete, review, and revise their answers.

SAFETY TIPSStudents should wear safety goggles at all times.

Students with long hair should keep it tied back.

Care should be taken when handling hydrochloric acid. If acid gets on clothes or skin, students should immediately wash it off with copious amounts of cold water. Spills on benches should be neutralized with baking soda before being wiped up with paper towels. Towels can be thrown into the trash. (You can also use a proprietary spill kit.)

Eyewash facilities should be available in the lab.

Students should immediately report all accidents involving acid to the teacher.

INQUIRY 1.2

FOLD THE FILTER PAPER TWICE IN TWO DIFFERENT DIRECTIONS AS SHOWN.FIGURE 1.2

FILTERING A MIXTURE

PROCEDURE

1

Observe the mixture in the jar labeled Substance A. Does it appear to be more than one substance? Describe it on Student Sheet 1.

2

Place a lab scoop of the mixture in your test tube and add 10 mL of water.

3

Place a stopper on the test tube, and shake the test tube for about 30 seconds. (Be careful not to hit the test tube against your lab table or another hard object.)

4

What happens to the contents of the test tube? Record your observations.

5

Fold a filter paper (see Figure 1.2) and place it in the funnel. Pour the contents of the test tube through the funnel into the second test tube.

6

Describe the appearance of the substance on the filter paper.

7

What do you think happened to the parts of the original mixture once you added the water and filtered it?

8

How do you think you might get back all the parts of the original mixture?

9

Rinse out the test tubes and funnel and throw away the used filter paper.

10

Replace the apparatus for the next group.



TEACHER’S NOTES

LESSON

PROCEDURE

1

Have your teacher light the candle.

2

What can you see taking place at or near the top of the candle? Write your observations on Student Sheet 1.

3

Place the open end of the beaker over the candle (see Figure 1.3). Let the beaker stay over the candle for a few minutes.

4

What happened after the beaker was placed over the candle? Record your observations.

5

Why do you think the candle reacted the way it did? Record your answer.

6

Restore the apparatus to its original condition for the next group.

THE BURNING CANDLE

INQUIRY 1.3

Upturned beaker

Burning candle

AFTER YOU HAVE RECORDED YOUR OBSERVATIONS OF THE LIT CANDLE, PLACE THE BEAKER OVER THE CANDLE.FIGURE 1.3


THE NATURE OF MATTERLESSON 1



Inquiry 1.5: Adding the AcidUse this inquiry to help students focus on careful and safe laboratory techniques. This inquiry provides students an additional opportunity to compare the properties of different metals. They should clearly recognize differences in the metals’ reactivity with the hydrochloric acid. The more observant students will notice a gas being formed as part of the reaction. Later in the unit, this gas will be tested and identified as hydrogen.

Inquiry 1.6: Comparing the Two MixturesMost students will not immediately think of a solution as a mixture, nor will they likely have any experience distinguishing between homogeneous and heterogeneous mixtures. In this inquiry they should be able to observe the dissolving of the powdered drink mix in the water to form a solution (or homogeneous mixture at this scale of observation). The flour, however, remains in suspension in the water and some eventually settles out. Neither substance is pure, but like most substances the students encounter, both are mixtures. As they progress through the unit, students will identify different types of mixtures and the methods of separating them.

Inquiry 1.7: Reacting a TabletFocus on students’ observations and records of measurement. Examine the language they use to describe the changes that take place. Look for references to the process of dissolving, different phases of matter (bubbles of gas, solid tablet, liquid water), and the concept of chemical reaction or change being related to a change in temperature.

Inquiry 1.8: Mixing the SolutionsStudents will clearly observe the dramatic color change and new solid formed. They may be able to recognize this as a chemical reaction. Later in the unit, when students are using words (and formulas) to describe chemical reactions, you can revisit this inquiry, which provides a straightforward example of a double-replacement reaction:

Lead + Sodium = Lead + Sodium Nitrate Iodide Iodide Nitrate

or

Pb(NO3)2 + 2NaI = PbI2 + 2NaNO3

TEACHER’S NOTES

13–E STC Unit: Experimenting with Mixtures, Compounds, and Elements

OVERVIEW

This lesson focuses on pure substances and mixtures and addresses some of the difficulties students have in defining them. In later lessons, students will investigate particular examples of mixtures, such as aqueous solutions and alloys, in more detail. In Inquiry 2.1, students are given eight different samples of matter—all of which are mixtures. Using a fixed set of apparatus, they are asked to determine whether the samples are pure substances or mixtures. For each sample, they must identify the criteria they used to make their determination. Students discover that although it is relatively easy to visually detect different substances in some very heterogeneous mixtures, it becomes more difficult as the mixtures become more homogeneous. They will find a gradation from the obviously heterogeneous mixtures found in rocks (such as granite) to homogeneous mixtures (such as solutions). Most students will probably try to identify and classify substances as pure based on how they look. They will discuss the limitations of this approach and, over subsequent lessons, re-evaluate it as they investigate the properties of other mixtures.

BACKGROUND

Purity is a very important concept in science. Since the earliest times, people have been concerned with testing the purity of substances (particularly gold and other metals) to determine their value. As a result, assaying, or the testing of metal purity, became an important occupation. In fact, the term “test” comes from the cuplike pot (called testus or cupule) that assayers used to conduct their tests. The assayer and the alchemist were the precursors to modern chemists who, before they could identify substances, had to purify them. Only then could chemists investigate the different physical and chemical properties of pure substances.

Most raw materials are mixtures of different substances. A wide variety of separation techniques are used to obtain pure substances from these mixtures. For example, crude oil contains many different hydrocarbons and other organic compounds and can be

separated into its various components by the process of fractional distillation. All separation techniques exploit the different physical and chemical properties of the substances that make up a mixture.

Students probably have already heard of some of the techniques used to separate mixtures, such as filtration. Students should have studied some of the properties of matter, such as density, solubility, and boiling and freezing points. In this unit, they will use their understanding of characteristic properties to explore the methods used to separate mixtures and identify the components of mixtures.

Table 2.1 lists information about the eight mixtures used in this lesson. (You will need to prepare mixtures E through H.) Students are unlikely to identify all of the components in the mixtures, and they are not required to do so. The emphasis of this inquiry should be on students’ search for evidence that the samples are pure substances or mixtures.

The samples are mixtures in different phases and vary in their level of heterogeneity. They have been selected to encourage students to think about what the concept of a “mixture” encompasses and to express their own ideas about how to determine whether a sample is a mixture or a pure substance. No sample is a pure substance (that is, an element or a compound).

In the context of this unit, the following are useful working definitions:

• A pure substance is matter that contains only one element or compound and has definite chemical and physical properties.

• A mixture is matter that contains two or more substances. The physical and chemical properties of a mixture may vary depending on the relative concentrations of the substances contained within it.

LESSON PURE SUBSTANCE OR MIXTURE?

STC Physical Science Strand: Matter and Change 13–F

TABLE 2.1 MIXTURE INFORMATION

SAMPLE SUBSTANCE COMPOSITION EVIDENCE THAT SUBSTANCES ARE MIXTURES

A GRANITE Feldspar (pink), quartz (white), biotite (black)

Crystals are clearly visible without a loupe.

B SLATE

Composition varies, but usually includes quartz (white), biotite (dark green, brown, or black), chlorite (green), graphite (gray/black), and pyrite (pale yellow)

Very fine grained and usually requires a loupe or microscope to determine the presence of different minerals.

C WHITE SAND Silica (mainly white) and other minerals

Colored grains visible with a magnifier.

D SHAVING CREAM

Mixture of water, stearic acid, triethanol amine, propane gas, butane gas, and fragrance

Bubbles in foam visible with a magnifier.

E OIL, WATER, AND DETERGENT EMULSION

Mixture of vegetable oil, water, and dishwashing detergent

Small droplets of oil may be visible with a magnifier; a microscope is required to see the smallest droplets.

F FOOD COLORING SOLUTION

Blue food coloring (about 10 drops dissolved in 500 mL of water)

Components of the mixture are not visible; if a drop is placed on a piece of paper, the liquid spreads out quickly and then evaporates, leaving a ring of blue food coloring.

G SUGAR AND ZINC OXIDE MIXTURE

Three-quarters confectioners’ sugar and one-quarter zinc oxide powder

Adding water dissolves the sugar but not the zinc oxide; this difference in solubility is easily visible only when water is added to a small amount of the sample spread out in a petri dish against a black background.

H IRON FILINGS AND SULFUR

Equal volumes of fine iron filings and sublimed sulfur powder

Components are visible with a magnifier; iron filings are magnetic; sulfur floats in water; iron sinks.

NOTE You will need to prepare mixtures E, F, G, and H.

13–G STC Unit: Experimenting with Mixtures, Compounds, and Elements


This lesson focuses on purity and mixtures. Both terms have specific definitions in chemistry and wider meanings in common usage. This duality fuels misconceptions that students may have about these concepts. Therefore, one objective of the lesson is to have students discuss their current understanding of these concepts and then, on the basis of their own experiences, reach a scientific understanding of the terms.

Students may have the following common misconceptions about purity:

• Students may incorrectly use the term “pure” in connection with substances that are safe to use or ingest. (For example, students may consider tap water to be pure because it is potable, even though it is a mixture, and even though many students are aware that chlorine is added to water to kill microorganisms.)

• Students may incorrectly use the term “pure” to mean free of additives. (For example, although pure orange juice is visibly a mixture, students may consider it pure because no artificial preservatives or sweeteners have been added.)

• Students may incorrectly relate the term “purity” to terms such as “clean” or “clear.” (The latter is particularly confusing when trying to classify solutions as mixtures.)

• Students may incorrectly consider only elements to be pure substances. (Both elements and compounds are pure substances.)

Students may have the following common misconceptions about mixtures:

• Students readily recognize obviously heterogeneous mixtures but may only reluctantly use the term “mixture” to describe homogeneous mixtures such as solutions or colloids in which the different components are not easily distinguished.

• Students may incorrectly make a distinction between heterogeneous and homogeneous mixtures. (No definite distinction exists between these two types of mixtures; the terms are based on perceptions of scale. Only a pure substance is truly homogeneous at the scale of molecules. You may wish to substitute the terms “poorly mixed” and “well mixed” to help students understand these two concepts.)

• Students may incorrectly give mixtures constant properties that are independent of their composition. (The properties of mixtures are not constant, but depend on the relative quantities of the substances from which they are composed. For example, the properties of concrete depend, in part, on the relative amounts of gravel, sand, and cement used to make it.)

• Students may incorrectly think that compounds, because they are composed of more than one element, are mixtures.

LESSON PURE SUBSTANCE OR MIXTURE?

PROCEDURE

1

Put four lab scoops of rock salt into the plastic cup. Examine it with the magnifying loupe. Write a description of the rock salt on Student Sheet 3.2: Cleaning Rock Salt.

2

Most of the salt used in food is made from rock salt. Discuss these questions with your partner:

A. Would you want to eat this sample?

B. Do you think it is pure?

C. What do you think the contaminants could be?

3

How could you use the remaining apparatus you have been given to obtain only the soluble component of the rock salt? Record your answers to the following questions:

• What are you trying to do?

• What materials will you use?

4

Record the procedure you and your partner devised.

5

Check your ideas with your teacher.

6

Follow your procedure to purify the salt. If you have any problems, consult your teacher.

CLEANING ROCK SALT

INQUIRY 3.2

YOU EAT THIS ROCK. WHAT IS IT AND HOW IS IT PURIFIED?PHOTO: NASA African Monsoon Multidisciplinary Analyses (NAMMA)

SAFETY TIP Do not taste any of the substances used in the laboratory.



INQUIRY 3.2

PROCEDURE

1. Explain that rock salt is an impure form of salt obtained by mining salt deposits. Have students examine a sample of the rock salt by doing Step 1 in the Student Guide. Ask students to describe the rock salt on Student Sheet 3.2.

2. Have students discuss the following questions with their partners:



C. What do you think the contaminants could be? (Focus on insoluble minerals or rock fragments.)

3. Have each pair discuss how they could use the apparatus to obtain only the soluble component of the sample. Students should answer the questions in Steps 2 and 3 on Student Sheet 3.2, and each pair should agree on a procedure.

4. Students should write their procedure under Step 4 on the student sheet.

5. Ask pairs of students about their procedural designs, but do not suggest a “correct” procedure.

6. Have each pair follow its own procedure to purify the salt. Circulate among the pairs, monitoring their progress. If students identify a problem with the procedure, ask them questions that will help them make appropriate modifications.

LESSON SEPARATING SOLUTES

31–A STC Unit: Experimenting with Mixtures, Compounds, and Elements

GETTING STARTED INQUIRY 4.1 Analyzing Inks

OBJECTIVES Observe the solubility of ink from a marker in water.

Use paper chromatography to analyze a mixture of inks.

Separate several solutes from a solution that contains a mixture of solutes.

CONCEPTS A mixture can have properties of one or more of its components or a completely different set of properties.

In chromatography, the characteristic properties of each solute determine the way in which the solute separates from a mixture of solutes in solution.

Solutions can contain more than one solute.

OVERVIEW Students place the tip of a marker in water and observe what happens.

Students place a dot of an ink mixture on chromatography paper, place the tip of the paper in water, and observe the results.

KEY TERMS Absorption Capillary action Chromatography

Adsorption Chromatogram

TIME 1.0 period 1.0 period (including “Introducing Paper Chromatography”)

CORRELATION TO NATIONAL

SCIENCE STANDARDS



LESSON AT A GLANCE


INQUIRY 4.2 Comparing Inks

INQUIRY 4.3 Identifying Inks READING SELECTIONS

Use paper chromatography to compare different mixtures of inks.

Use paper chromatography to identify inks.

Apply paper chromatography techniques to solve a “crime.”

Introducing Paper Chromatography Read an explanation of the use of chromatography to separate a mixture of inks.

The Case of the Unidentified Ink Read about the use of chromatography to analyze and compare the ink from a forged check with that of pens from known sources.

“Separation Science” at the FBI Read about the uses of different types of chromatography in forensics.

Chromatography can be used to analyze and identify solutions that contain several solutes.

The chromatogram a solution, such as ink, from an unknown source can be compared with a solution from a known source for identification.

The dyes in an ink mixture move at different rates up a piece of absorptive paper and leave a unique pattern.

High-pressure liquid chromatography is used to identify substances in a bomb sample.

The class agrees on a procedure to use chromatography to compare four different ink mixtures.

The chromatogram of a solution, such as ink, from an unknown source can be compared with a solution from a known source for identification.

“Introducing Paper Chromatography” explains the process of chromatography.

“The Case of the Unidentified Ink” serves as a basis for Inquiry 4.3.

“‘Separation Science’ at the FBI” provides real-life applications of the chromatography process.

High-pressure liquid chromatography

1.0 period 0.75 period (including “The Case of the Unidentified Ink”)

0.25 period


Content Standard B• Properties and changes of properties in

matter



Content Standard E• Understanding about science and technology

Content Standard G• History of science• Nature of science


EXTENDING YOUR KNOWLEDGE

chromatography its name.) She stands the plate upright, with the samples along the bottom, in a container with a small amount of liquid. As the drops become moist and interact with the coating on the plate, they begin to move up the plate at different rates, depending on the solubility of the components. Once the liquid gets near the top of the plate it is removed from the container. Then, the positions of the drops are compared. All dyes made from the same components will form the same pattern on the plate. So, if the pattern of the crime scene sample matches the pattern of the dye used by banks, another crime has been solved.

BOMBS AND EXPLOSIVESChromatography also comes in handy for analyzing the materials used in bombs and explosive devices. The FBI analyzes samples from all major bombings involving the United States, including the one at the Murrah Federal Building in Oklahoma City and others causing airline crashes. The technique is called high-performance liquid chromatography—HPLC, for short.

The first step is examination under a microscope. “Most bomb samples look pretty much alike. They look like black powder,” says Mount. Even so, this first step is important. The scientists might, for example, be able to sort out small pieces of material from the residue.

The next step is extraction. The chemists place the sample in a solvent such as water. Once in solution, the particles in the sample may, depending on the composition of the sample, separate into smaller particles that carry positive or negative charges.

A small amount of the solution is placed in the HPLC machine. It moves up to the top, where it mixes with another liquid, and is then forced downward under pressure through a narrow glass column that is filled with a porous substance.

What happens in the column is the critical step. “Some of the [particles],” explains Mount, “seem to like it better in the tube than others. They stay longer.”

The speed at which the particles leave the column is recorded by a detector, which then prints out the information. By comparing the time that the particles have stayed in the column with known retention times, Mount and her colleagues are able to distinguish the various types of particles in the test sample.

STILL A LOT TO LEARNDoes it always work? “No,” says Mount. “Sometimes we find nothing. And other times, we find nothing conclusive. It’s also important to note that when it comes to explosive materials, HPLC is only a qualitative analysis technique. It helps us identify what materials are in an unknown powder. It doesn’t provide quantitative information; in other words, we can’t tell how much of each substance is in the powder.”

Kelly Mount loves her work. To prepare for her career, she earned a bachelor’s degree in chemistry and then went on to get a master’s degree in forensic science. Life in the lab is never routine—this means getting called into the lab on weekends or even at night when there is an emergency. Whether the problem is bombs or banks, Mount has the expertise to help the FBI solve its mysteries. ■

1. Envision a crime scene not described in the reading selection in which chromatography could be useful. Describe how chromatography could be used to help solve the crime.

2. Describe the procedure you would use to find out whether carrots get their color from one or many substances.

1.

DISCUSSION QUESTIONS


LESSON CHANGING MIXTURES


GETTING STARTED INQUIRY 5.1 Adding Salt to Ice

INQUIRY 5.2 Adding Salt to Boiling Water

OBJECTIVES Discuss and record different kinds of mixtures.

Measure the effect of different quantities of salt on the melting point of ice.

Measure the effect of different quantities of salt on the boiling point of water.

CONCEPTS A mixture can have properties of one or more of its components or a completely different set of properties.

Salt lowers the melting point (also freezing point) of ice. The decrease in melting point is related to the amount of salt added to the mixture.

Salt raises the boiling point of water.

Changing the concentration of solutes affects the properties of solutions.

OVERVIEW Students discuss different kinds of mixtures.

Students add salt to an ice/water mixture and collect temperature data.

Students add salt to water and collect temperature data as it is heated.

KEY TERMS Component

Mixture

Freezing point

Melting point

Boiling point

TIME 1.0 period 1.0 period 1.0 period (including “Changing Melting and Boiling Points”)

CORRELATION TO NATIONAL

SCIENCE STANDARDS




Content Standard B• Properties and changes of properties in

matter

STC Physical Science Strand: Matte and Change 41–D

In Inquiry 5.2, students discover that the boiling point of water rises when salt is added to water. Boiling point elevation takes place because solute particles interfere with the evaporation of a solvent. In the case of aqueous solutions, solutes that are more volatile than water do not elevate the boiling point. In the case of water, 1 mole of nonvolatile solute particles in 1 kg of water elevates the boiling point by 0.2°C. Students are unlikely to observe an increase in boiling point of more than 1–2°C when they add salt to boiling water.

The phenomena of raising the boiling point (known as boiling point elevation) and lowering the melting point (known as freezing point depression) have several applications. Ethylene glycol, a water-soluble, fairly nonvolatile compound (boiling point, 197°C) is used in antifreeze in car cooling systems; it also allows the coolant to operate at temperatures above the boiling point of water. (However, the pressurized nature of the cooling system plays a more important role.) The freezing point depression effects of salt on ice are used to de-ice roads (see “Changing Melting and Boiling Points” on page 47) and can be used to manufacture ice cream (see “Ice Cream in the Old Days” on pages 56–57).

In Inquiry 5.3, students investigate alloys. Alloys have been described as “solid solutions” with at least one metal as their major component. Their structure is often very complex; their properties are determined not only by their composition but also by the way they are treated when cooling from their molten state. They almost always have a crystalline structure, which is partly determined by this cooling process (see “The Samurai’s Sword” on pages 52–55 for an example). This structure is particularly important in determining their properties, which can be substantially different from those of pure metals.

In fact, most metal objects in common use are alloys. The most prevalent is steel. Many types of steel consist of different alloys of iron and carbon. The relative amounts of iron and carbon within the alloy determine the properties of the different types of steel. Other substances are also added to steel to make alloys of three or more components. Stainless steel, for example, contains chromium, which provides a protective layer of oxide and reduces corrosion. Other examples of alloys are given in Table 5.1 on page 41-E. Students can read more about this topic in “About Alloys” on page 51 and “The Samurai’s Sword.”

The alloys used in Inquiry 5.3 are solders (used to join together metal materials); they were selected on the basis of their low melting points and ready availability. The use of a particular solder depends, in part, on its melting point. Table 5.2 on page 41-E lists the composition, melting points, and uses of the three solders investigated in this experiment.

The properties of mixtures other than solutions and alloys will probably come up during class discussions. Many of these mixtures are composite materials that consist of at least two discrete substances bound together.


What happens when sodium sulfate is added to the water? Sodium sulfate, unlike water, is a strong electrolyte, readily producing sodium and sulfate ions when added to water. When a potential difference (voltage) is applied across the electrodes, the sodium ions (Na+) migrate to the cathode and the sulfate ions (SO

42-) migrate to the anode. This process increases

the current flow around the circuit. When hydrogen ions (H+) and sodium ions arrive at the cathode, the sodium ions remain in solution (because they are more electropositive than hydrogen) and the hydrogen ions are reduced, producing hydrogen gas.

As hydrogen ions are removed from the water, more water dissociates to maintain the ionic product at its constant value. The ionic product of water is the product of the molar concentration of hydrogen ions and the molar concentration of hydroxide ions (OH-) present in water, and is constant at a specific temperature. When the sulfate ions arrive at the anode, they do not lose electrons as easily as the water, so oxygen is liberated. The net effect of adding the sodium sulfate is an acceleration of the electrolysis of water.

Other reactions also take place in the electrolytic cell. The cell used in Inquiry 6.1 has stainless steel electrodes. These are not very reactive, but after long periods of electrolysis, gelatinous hydroxides of iron will appear in the beaker and electrodes. For this reason, you should dispose of the sodium sulfate solution and wash the container and electrodes at the end of each period.

THE ELECTROLYTIC CELL USED IN INQUIRY 6.1 FIGURE 6.1

THE ELECTROLYTIC

Test tubes full of sodium sulfate solution

Stainless steel electrodes

Sodium sulfate solution

Students observe that twice the volume of hydrogen collects over the cathode as oxygen over the anode. They can test these gases with a burning splint or, if they are familiar with the test for oxygen, they can test for oxygen with a glowing splint. The burning splint test for hydrogen works best if the test tube containing hydrogen is kept inverted (open end down). This prevents the hydrogen gas, which is lighter than air, from escaping upward. The glowing splint test for oxygen works best if the test tube containing oxygen is sealed with a thumb and held upright.

When exposed to a burning splint, hydrogen burns with a squeaky pop. A glowing splint will relight when exposed to oxygen. A burning splint burns brighter, for a few seconds, in a tube containing oxygen compared with a tube containing air. The difference in the way hydrogen and oxygen behave is adequate evidence that two different gases have been collected, but these tests cannot be considered a definitive identification of oxygen.

Do not expect loud pops or sustained bright burning with the small quantities of gas collected during this experiment. The mass, and therefore the volume, of each gas liberated at a specific temperature is determined by the amount of electricity passed through the electrolytic cell. This quantity is a function of the size of the current passed through the cell and the length of time the cell is operated. It takes about 15 minutes to collect a small tube of hydrogen and twice as long to collect the same volume of oxygen. As the

LESSON


students do not have time to complete the inquiry in one lesson.

lesson


The surface of The earTh is four-fifThs waTer. whaT is waTer made from?PHOTO: NASA

INTRODUCTION

In your previous work, you looked at the characteristic properties of pure

substances. Then you investigated how those properties can differ from the

properties of mixtures. Now you will focus on two groups of pure substances known as elements and compounds. In this lesson, you will examine the composition of the pure substance you have encountered most often during the course of this unit—water. You know that water has several characteristic properties that can identify it as a single substance rather than a mixture.

These properties include its appearance, density, melting

and boiling points, and ability to dissolve a wide range of solutes.

You will investigate what happens to water when electricity is passed through

it. Sometimes, passing electricity through a liquid can give you clues about the composition

of the liquid. If water is a pure substance, why try to find out its composition? Do the inquiry and find out what happens.

Breaking Down a CompounD

PREPARATION

1. Make one copy of Student Sheet 6.1: Electrolysis of Water for each student and make a transparency of TE Figure 6.1. This figure is provided on the Teacher’s Tools CD.

2. Place the plastic boxes containing the materials around the lab so groups can share the batteries (two groups share two batteries).

3. Between each set of plastic boxes, place two 6-V batteries, connected as shown in Figure 6.2.

NOTE Test tubes of hydrogen and oxygen may be needed to confirm

student results or to demonstrate the properties of oxygen and hydrogen if

SAFETY TIP Because hydrogen is extremely explosive, you should prepare it only in small quantities (enough to fill a few test tubes at a time).

NOTE Lesson 7 requires some advance preparation. Be sure to refer to page 65-B

for more information.

4. Prepare three test tubes containing oxygen and three containing hydrogen. You may prepare the oxygen and hydrogen by electrolysis or chemical reaction. For the latter, consult a chemistry textbook.

LESSON


THE SURFACE OF THE EARTH IS 72 PERCENT WATER. WHAT IS WATER MADE FROM?PHOTO: NASA

INTRODUCTION









BREAKING DOWN A COMPOUND


BUILDING YOUR UNDERSTANDING

Water is a compound made up of two elements—hydrogen and oxygen. The characteristic properties of these elements are different from those of water. However, hydrogen and oxygen have some common properties. They are both colorless, odorless gases, and they both readily react with other elements, making them “reactive” elements. But in many ways they are very different from each other.

Hydrogen has the lowest density of all the elements. It is very reactive, which is one reason why it is present in only very small quantities in air. It reacts with oxygen. You reacted it with oxygen when it burned with a squeaky pop. What do you think was made in that chemical reaction?

It may come as a surprise to you to discover that hydrogen is the most common element in the universe. The sun and other stars are mainly hydrogen gas. Hydrogen is found in many compounds. For example, sulfuric acid (used in car batteries) and hydrogen peroxide (used in hair dyes) contain hydrogen.

Oxygen reacts with other substances. Oxygen is needed for burning to take place. Things burn well in oxygen, producing hotter flames. For example, what happened to the glowing splint when it was put into a tube of almost pure oxygen? Some welding and metal-cutting equipment use flammable gases and pure oxygen to produce the high temperatures needed to melt metal.

Oxygen also reacts slowly with many substances. Many compounds containing oxygen are called oxides. You may be familiar with two oxides that are gases—carbon dioxide and sulfur dioxide—but most oxides are solids. In fact, oxygen is the most common element in the Earth’s crust, but most of it is combined with other elements to form minerals that make up rocks. n

WATER IS A COMPOUND FORMED WHEN INFLAMMABLE HYDROGEN REACTS WITH OXYGEN. HERE IT IS BEING USED TO PUT OUT A FIRE. LIKE ALL COMPOUNDS, THE PROPERTIES OF WATER ARE VERY DIFFERENT THAN THOSE OF THE ELEMENTS FROM WHICH IT IS COMPOSED. PHOTO: U.S. Air Force photo by Senior Master Sgt. David H. Lipp



In 1869, Russian chemist Dmitry Ivanovich Mendeleyev classified the 63 elements known at the time in a “Periodic Table of the Elements,” in which he related atomic mass to physical and chemical properties. Mendeleyev was able to use this table to predict the properties of some yet-to-be-discovered elements. He left gaps between some elements in the table to be filled when new elements were discovered. During his lifetime, gallium, scandium, and germanium were discovered. They perfectly matched the predictions Mendeleyev had made on the basis of his periodic table. The table was subsequently modified in 1913 by English chemist Henry Moseley. Armed with knowledge of atomic structure, Moseley related properties of elements to atomic number rather than to atomic mass.

In the periodic table (see Figure 7.1, page 77-E), the horizontal rows are called periods and the vertical columns are called groups. Several approaches exist for labeling these periods and groups. For example, some scientists number the groups as 1–18 across the top of the table. Others favor the traditional approach, using Roman numerals (I through VII) and 0 to number the groups. The latter approach is used here. The Roman numerals correspond to the number of electrons in the outer shell of the elements in a group (with the exception of group 0 elements, which have complete outer shells). The pattern of the table is most easily followed for the first six elements. The sequence is then complicated by a block of metallic elements including the transition elements, the lanthanides and actinides. These last two groups are also called the inner transition elements.

The following general patterns can be observed in the periodic table:

• Metals are on the left side of the table; nonmetals are on the right.

• The most reactive metals are at the bottom of group I. These metals have larger atoms, and they more easily lose electrons in the outer shells, making them more reactive.

• The most reactive nonmetals have the smallest atoms and are found at the top of group VII. All nonmetals (excluding the noble gases) have incomplete outer shells; therefore, they readily accept electrons.

• Elements within a group combine with particular elements in similar ratios. For example, group I elements, the alkali metals (such as potassium and sodium), always form chlorides at the ratio of 1:1.

• Group number is related to the number of valence electrons—a measure of the number of atoms that an element can combine with (for example, hydrogen has 1 valence electron and oxygen has 2 valence electrons, which accounts for the formula of water, H2O). In groups I–IV, the group number is the same as the number of valence electrons. In groups V–VII, the number of valence electrons is 8 minus the number of the group. For example, group V elements have 3 valence electrons (8–5).

The periodic table is continually expanded as new elements are discovered or synthesized. Some elements are synthesized by bombarding existing elements with high-energy particles. For example, meitnerium (Mt), element 109, was synthesized by bombarding bismuth-69 (69 is the isotope—the sum of the neutrons and protons in the element’s nucleus—not the atomic number) with iron-108. Elements such as meitnerium exist as very short-lived radioactive isotopes.

The elements students investigate in Inquiry 7.1 are inexpensive and safe to use in the classroom when the directions for proper handling are followed. Most are metals, reflecting the dominance of elements with metallic properties in the periodic table. Some of the elements in this inquiry are represented only by the Element Cards provided on the Teacher’s Tools CD because they are rare or potentially dangerous for students to handle. Many of the elements investigated in this lesson are found in the compounds and the reactions observed in the unit. They are also discussed in the reading selections throughout the unit. Information about uranium is included in SG Figure 7.2 on page 82 to provide an example of a radioactive element.

TE Figure 7.1 shows the position in the periodic table of the elements that appear in Table 1 on Student Sheet 7.1a: Examining and Grouping Elements. SG Figures 7.1, 7.2, and 7.3 on pages 81–83 and the Element Cards provide information on chemical symbols, physical and chemical properties, compounds, use, and other facts (for example, occurrence and toxicity). Students are not expected to learn all of the chemical symbols at this stage. (You could have students learn the symbols of the elements studied in this lesson for a homework assignment.) Symbols are included on the Element Cards so students can identify elements when they refer to the periodic table.


To the right of the dividing line, the nonmetals can be conveniently split into two groups: the p-block elements (so called because their outermost electrons go into p subshells; refer to an advanced chemistry text for more information on the electronic configuration of elements) and group 0, the noble gases. The noble gases were originally called the inert gases because it was once thought that they did not react with other elements to produce compounds. However, several compounds of these gases (mainly fluorides and oxides) have now been identified.

Table 8.1 summarizes the properties of metals and nonmetals. As with most classifications, there are exceptions to some of the criteria in each group. Expect students to emphasize appearance, conductivity, and other physical properties when they classify metals and nonmetals.

The physical properties of metals can be explained in terms of their structure. Roughly, electrons exist in discrete

energy levels called “shells.” In the most common model of metal structure, the electrons in the outer shell of each atom move randomly through a lattice of cations (positive ions). This negative “cloud” of electrons attracts the metal cations, binding them in a very efficiently packed crystal structure. This tight “metallic” bonding explains the high boiling point and high density of most metals. It also explains the high electrical conductivity of metals—their outer-shell electrons can move freely throughout the cloud. When voltage is applied across a metal, the movement of electrons is no longer random, because electrons are repelled by the negative electrode and attracted by the positive, creating a flow of electrons. Thermal conductivity can also be explained in terms of electrons with high kinetic energy in regions of high temperature moving toward cooler regions and transferring their energy to other electrons throughout the sample of metal.

TABLE 8.1 PROPERTIES OF METALS AND NONMETALS

PROPERTY NONMETALS METALS

Phase at room temperature

Mainly gases, a few solids, and one liquid

(bromide)

Solids (except mercury, which is liquid)

Density Low High

AppearanceUsually dull or gaseous (except silicon, which is

very shiny)

Shiny (may need to scratch the surface to remove

an oxide layer)

Electrical conductivity

Poor (except graphite, an allotrope

of carbon); increases with increasing temperature

Good; decreases with increasing temperature

Thermal conductivity Poor Good

Flexibility and ductility Poor (brittle) Usually good

Oxides Solid or gaseous, acidic or neutral

Solids that are basic or amphoteric*

Chlorides Solid or liquid covalent chlorides

Crystalline solids with ionic bonds

* Amphoteric oxides are oxides that can behave as bases or acids, depending on conditions. Aluminum, zinc, and lead have amphoteric properties.

COMBINING ELEMENTSLESSON

89–E STC Unit: Experimenting with Mixtures, Compounds, and Elements

When atoms of elements react with one another, they achieve stable outer electron shells by either losing or gaining electrons. Metals tend to react by losing electrons, whereas nonmetals tend to react by gaining electrons. Nonmetals often react by combining with other nonmetallic atoms. (These atoms can be from the same element or a different one.) By sharing electrons, nonmetals complete their outer shells and form strong covalent bonds. Several nonmetallic elements exist in diatomic form, for example, nitrogen (N

2), oxygen (O

2),

and chlorine (Cl2). Although the covalent bonds that

hold the atoms together as molecules (N2, O

2, Cl

2) are

very strong, the forces between the molecules are weak. Thus, many of the nonmetals exist as gases, liquids, or “soft” solids (for example, phosphorus).

The pattern of covalent bonding between atoms of the same element can vary. These variations produce different forms of an element, known as allotropes. Different patterns of bonding produce very different characteristic properties. The best example is carbon. In one form, it is soft, grayish-black graphite, and in another, it is hard, transparent, colorless diamond. The electrons in the covalently bonded molecules of nonmetals are firmly held and are less mobile than the electrons of metals. These strong bonds account for the poor thermal and electrical conductivity of nonmetals.

Reacting Iron and Oxygen After identifying metals and nonmetals, students examine the properties of metals in more detail over the next few lessons. In Inquiry 8.2, students examine the synthesis reaction between iron and oxygen. In Lesson 10, they study metals as reactants and order them according to their reactivity with acid. In Lesson 11, students apply their knowledge of reactivity to the problem of corrosion and its prevention.

The reaction between iron and oxygen is rapid, and if finely divided iron in the form of clean steel wool is used, the reaction is readily observable over a period of a few minutes at room temperature. In Inquiry 8.2, steel wool is rinsed with vinegar (dilute acetic acid)

and then placed in a test tube, which is inverted over water. Movement of water up the tube (compared with a control containing no steel wool) is immediately visible as the iron in the steel wool reacts with the oxygen in the air. The reaction is complete within about 15 minutes. The water will have moved about one-fifth of the way up the tube, having consumed the oxygen in the air. The steel wool, which is a shiny metallic gray when rinsed in vinegar, will have turned dark gray to black. The reaction does not produce the red-brown oxide of iron associated with rust—iron (III) oxide (Fe

2O

3); in the

absence of excess oxygen, black iron (II) oxide (FeO) is produced. If the tube is removed from the water and allowed to stand open to the air over a few days, the steel wool will oxidize further, producing iron (III) oxide. (More details on the rusting process are provided in “Background,” Lesson 11.)

The reaction that takes place in Inquiry 8.2 can be represented as a simple word equation:

iron + oxygen ➝ iron oxide

The reaction is exothermic; an investigation of the exothermic nature of this reaction is provided as Extension 3.


9

Immediately show the water level by filling in the diagram in Table 1. Write a description of the damp steel wool in the third column of the table.

10

Watch the apparatus carefully. Do you notice anything happening in the tubes?

11

After about 15 minutes, look at the level of the water in each tube. Record your observations in the appropriate places in Table 1.

12

Describe the appearance of the steel wool.

13

Air is about 21 percent oxygen. Keeping the tubes in position, remove the card and use the ruler to measure how far up the test tube the water has moved. Answer these questions on Student Sheet 8.2:

•Has the water level changed in either tube? If so, can you explain the change?

•What can you conclude from your observations?

14

Clean the test tubes and the beaker. Make sure you remove the steel wool from the test tube and place it in the trash. Return the cleaned apparatus and the index card to the plastic box.

REFLECTING ON WHAT YOU’VE DONE

Discuss your results and compare them with those of other pairs.

Answer the questions in Steps 4–7 on Student Sheet 8.2.

All chemical reactions have reactants (inputs) and products (outputs) and can be written as simple word equations. For example, you know that hydrogen combines with oxygen to form water. A simple word equation to describe this equation is as follows:

hydrogen + oxygen = water

On Student Sheet 8.2, complete Step 8 by writing a word equation for the reaction that took place in the test tube. Label the reactants and the products in the equation. This is an example of the synthesis of a new product from a chemical reaction. In Lesson 10, you will investigate how metals react with acids to form new products.

1

2

3



1. Discuss students’ results by using their answers to the questions in Steps 2 and 3 on Student Sheet 8.2. Focus on the following points:

• The changes in the steel wool

• The difference in the water level in each tube at the end of the inquiry

• Students’ explanations of these differences.

(Ask students how they determined that these two elements were involved in the reaction.)

• The substance produced was iron oxide. (Although in the absence of excess oxygen, this is iron (II) oxide, it is less confusing to use the term “iron oxide” when constructing a simple word equation.) Encourage students to use components of the words “iron” and “oxygen” in their answer.

• The iron and oxygen are reactants, and the iron oxide is the product.

3. The chemical reaction can be expressed as:

iron + oxygen ➝ iron oxide

Write the word equation on the board. Explain that the arrow in the equation indicates the direction of the reaction. To clarify the role of reactants and products, write “reactants” and “product” above the equation:

reactants product iron + oxygen ➝ iron oxide

Explain that this reaction is called a synthesis reaction because a new substance (the compound iron oxide) is made when the elements iron and oxygen combine. Tell students that in Lesson 10 they will compare how different metals react with acid.

2. Try to obtain the following answers from students for Steps 4–7 on Student Sheet 8.2:

• Two test tubes were used in order to compare what happened to the water level with and without steel wool. The empty tube served as a control.

• The two elements that reacted together were iron and oxygen.



EXPLORATION ACTIVITY COMPOUNDS

This is a list of elements, their familiar compounds, and the chemical formulas for those compounds:

• Hydrogen compounds—acids, hydrides, and hydroxides

Hydrochloric acid (HCl) Boric acid (H3BO3) Carbonic acid (H2CO3) Sulfuric acid (H2SO4) Nitric acid (HNO3) Hydrogen peroxide (H2O2) Methane (CH4) Ammonia (NH3) Butane (C4H10)

• Carbon compounds or organic compounds, which make up a separate branch of chemistry

Carbon dioxide (CO2), a major greenhouse gas

Carbon tetrachloride (CCl4), formerly a dry-cleaning compound

Freon (CCl2F2) formerly in air conditioners

• Oxygen compounds—oxides Silicon dioxide (SiO2) Magnesium oxide (MgO) Iron (II) oxide (Fe2O3) Nitrous oxide (N2O) Sulfur dioxide (SO2)

• Sodium, magnesium, phosphorus, and calcium compounds—abundant in earth’s crust

Sodium hydroxide (NaOH), a drain cleaner

Sodium bicarbonate (Na2S2O3), baking soda

Sodium thiosulfate (NaS2O3), a fixing agent in photo printing

Magnesium chloride (MgCl2), abundant in seawater

Magnesium oxide (MgO), 35% of earth’s crust

Hydrogen phosphate (H3PO4), in soft drinks, fertilizers, and solvents

Calcium oxide (CaO), lime Calcium carbonate (CaCO3), limestone Calcium sulfate (CaSO4), in gypsum

and plaster of paris Calcium phosphate Ca3(PO4)2,

the primary inorganic substance in bones and teeth

• Sulfur compounds Hydrogen sulfide (H2S),

has a “rotten egg” odor Sulfur dioxide (SO2), a refrigerant,

bleaching agent, and disinfectant Potassium sulfate (K2SO4), in fertilizer,

pigments, and dyes


LESSON


cids and bases were known and used by alchemists before their chemical structures were identified. Substances such as vinegar and citrus fruit juice were identified as sour tasting and described as acids.

The term “acid” is taken from the Latin “acidus” for “sour.” French chemist Antoine Lavoisier identified the gas produced by the reaction of acids with metals. Other substances known for their bitter taste and slippery feel were called “alkalis.” The word “alkali” came from the Arabic word for “ashes” because these substances were obtained by soaking ashes from wood fires in water.

In the late 1800s, Swedish chemist Svante Arrhenius was testing the ability of certain solutions to conduct electricity. (The invention of a storage battery to produce electricity led to many electrolysis experiments.) He theorized that some solutions, called electrolytes, decomposed into charged particles, or ions, and produced electricity when dissolved in water. Acids, he suggested, split up when placed in water and produced hydrogen ions with a single positive charge (H+). He described bases as substances that dissolved in water and formed negatively charged hydroxide ions (OH—).

Some common acids and their chemical formulas are hydrochloric (HCl), nitric (HNO3), sulfuric (H

2SO

4),

carbonic (H2CO

3), and acetic (C

2H

4O

2). All carbonated beverages contain carbonic acid, and many have small

quantities of phosphoric acids. Citrus fruits include ascorbic acid, or vitamin C (C6H

8O

8). Many household

cleaning products contain hydrochloric acid. You will notice that all these compounds contain hydrogen, which produces the Arrhenius H+ ion in water.

Bases are equally common in everyday life and use. They include sodium hydroxide (NaOH) (often called lye and sometimes used as a drain cleaner), ammonium hydroxide (NH

4OH), baking soda (NaHCO

3),

substances contained in ashes (K2CO

3), and soaps. When solutions of bases and acids are mixed, water and

a new substance are produced. This reaction is called a neutralization reaction because the acid (H+) and the alkali (OH—) neutralize one another to form H

2O. The new substances formed in neutralization reactions are

called “salts.” One of these salts, NaCl, the product of the reaction between hydrochloric acid and sodium hydroxide, is what we call “table salt,” or sodium chloride. This is only one of many substances called salts.

Arrhenius’s definitions of acids and bases were modified in the 20th century as the understanding of atomic structure grew. Two scientists working independently reached the same conclusion on the structure of acids and bases. Chemists T.M. Lowry from England and J.N. Brønsted of Denmark each developed new definitions of acids and bases. They defined an acid as a substance that can donate a proton, the H+ ion. They defined a base as a substance that can accept a proton. This definition of acid-base reactions, known as the Brønsted-Lowry Theory of acids and bases, is widely accepted by scientists today. The Brønsted-Lowry Theory illustrates the gradual progression of explanations in science. New understandings of the structure of matter have brought new explanations for well-known behaviors of matter. n

1. How do acids and bases differ in their chemical and physical properties?

2. Why do scientists change their explanations of things over time?

1.


READING SELECTION EXTENDING YOUR KNOWLEDGE


CHEMICAL REACTIONS INVOLVING METALSLESSON 10

COUNTERING CORROSIONLESSON

135–C STC Unit: Experimenting with Mixtures, Compounds, and Elements

OVERVIEW

Students have investigated several chemical reactions involving metals. They have compared the reactivities of metals and related their differences to the process of corrosion. In this lesson, they discover that their knowledge of chemical reactions can be applied to prevent corrosion. The lesson begins with a discussion that expands into a brainstorming session about methods of preventing rust. Students use a standard set of materials to design an inquiry to compare the effectiveness of several rust-prevention treatments. They also design a results table. Students maintain the experiment for three to four days before recording their results and reporting their findings to the rest of the class. Students then use the summary of class results to explain what happened for each rust treatment.

BACKGROUND

In this lesson, students investigate the corrosion of iron. Iron was chosen because it is, in the form of steel alloys, economically the most important metal. Rusting costs the U.S. economy tens of billions of dollars annually in the replacement and treatment of rusted components (from nails to bridges) and in the cost of rust-prevention measures.

The main cause of rusting is the exposure of iron to oxygen. The overall reaction in natural rusting is different from the one that took place in Inquiry 8.2 because it takes place in an excess of oxygen. This reaction is shown in the following equation:

4Fe + 3O2 ➝ 2Fe2O3

In this process, the iron becomes oxidized, losing electrons to oxygen and forming iron (III) oxide. But why do iron objects in very dry climates rust slowly? What role does water play in the rusting process? For iron to rust, two reactions must occur. In the first, iron is oxidized to iron (II) ions; in the second, the electrons released react with oxygen and water. These reactions are shown in the following two equations:

Fe(s) ➝ Fe2+(aq) + 2e−

O2(g) + 4e− + 2H2O ➝ 4OH−(aq)

For these reactions to take place, the electrons must move from the place where the iron dissolves to the oxygen, and, to complete this circuit, ions must move in the opposite direction. Therefore, electrolytes present in water speed up this process, which is why salt (an electrolyte) in water causes the rapid rusting of unprotected iron and steel. In the second reaction, Fe2+ is oxidized further to Fe3+ by exposure to more oxygen:

2Fe2+(aq) + 2OH−(aq) + O2(g) ➝ Fe2O3(s) + H2O(l)

This explains why black iron (II) oxide (FeO) was produced in Inquiry 8.2. If the steel wool used in the experiment had been exposed to an excess of oxygen, reddish-brown iron (III) oxide (Fe

2O

3) would have

formed. Rust is a hydrated form of iron (III) oxide, so water plays an additional role in the hydration of the iron (III) oxide. Most rust is a mixture of these two iron oxide compounds and is given the combined formula of Fe

3O

4.

The rate at which iron and steel (as well as some other metals) corrode can be slowed in several ways. The most common approach is to coat the metal to prevent it from coming into contact with water and air. The metal can be coated with oil, grease, porcelain (enamel—not to be confused with enamel paints), paint, or plastic. It can also be coated with another, less reactive metal or a more reactive one that readily forms a protective oxide layer. When metal objects are stored, materials that absorb water, such as silica gel, can be stored along with them to reduce corrosion.



Students may have misconceptions about changes in the mass of matter during chemical reactions, including the following:

• Students may incorrectly think that matter “gets lighter” in chemical reactions in which there is an apparent loss of matter. (This misconception commonly arises when dealing with reactions that have at least one gaseous product, for example, the combustion of organic substances.)

• Students may incorrectly think that matter can disappear during a chemical reaction.

• Students may incorrectly think that the appearance of products is the result of the creation of “new matter” rather than of the rearrangement of the bonds between the atoms of the reactants. They may therefore consider the reactants to be independent of the products of a chemical reaction.

• Students may incorrectly equate an apparent loss of mass during a chemical reaction with the destruction of matter.

By comparing the same chemical reaction under two different circumstances, students will have the opportunity to discover that the loss of mass in the open system is the result of escaping gas, not the disappearance or destruction of matter. This concept is reinforced by the audible release of the gas when the cap on the bottle is loosened and by the subsequent loss of mass in the closed system. Students should then be able to conclude that matter is not destroyed and does not disappear during chemical reactions.

READING SELECTION

“The Mass of Matter” provides students with an historical account of the development of the principle of the conservation of mass (later to become a law) by French chemist Antoine Lavoisier and his wife Marie-Anne, who was also proficient in scientific investigation. Students may be interested in doing further research on Antoine Lavoisier and his untimely demise (due to his unfortunate choice of standing with the Royalists during the French Revolution).

MASS AND CHEMICAL REACTIONSLESSON


HOMEWORK

Period 1Lesson 13 is an assessment for the unit. Hand out Student Sheet 12 and have students review it.

Period 2Have students read “The Mass of Matter” on pages 152–153 and answer the accompanying questions.

EXTENSIONS

Science/Social Studies1. Earth is sometimes described as a closed system when it comes to matter but as an open system as far as energy is concerned. Have students write a paragraph about each of the ideas expressed in this statement. They should explain whether they agree or disagree with each statement.

Science/Math2. Have students apply their knowledge of the conservation of mass to the following problems:

• When electrolysis is used to decompose 5.0 g of water into its constituent elements, 0.55 g of hydrogen are produced. What is the mass of oxygen released? Explain how you obtained your answer.

• Exactly 56.0 g of iron will react with exactly 32.0 g of sulfur. If 1.4 g of iron were heated with 10.0 g of sulfur, how much iron (II) sulfide would be made? How much sulfur would be left unreacted?

NOTE It is suggested that students present their work on the Exploration Activity before taking the final assessment. Students

should be completing their cubes and practicing their presentations with their partners as well as reviewing for the assessment. Students may require additional homework or class time after their presentations to review the work they have done during the unit.

LESSON


lesson


INTRODUCTION

This lesson is designed to assess how much you have learned while working on the unit Experimenting with Mixtures, Compounds, and Elements. The assessment consists of two parts: a performance assessment (Inquiry 13.1) and a written assessment.

Final assessment For Experimenting with Mixtures, Compounds, and Elements

in this lesson, you will conduct an inquiry that requires you to follow instructions carefully. you will also have to make and record accurate measurements and observations.PHOTO: © Terry G. McCrea, Smithsonian Institution

PREPARATION

1. Make one copy of Student Sheet 13.1: Performance Assessment and one copy of Student Sheet 13: Written Assessment for each student.

2. Fill 16 2-oz jars with copper (II) sulfate. Label the jars “A.”

3. Fill 16 125-mL bottles with calcium hydroxide solution (limewater) and label the bottles “B.”

4. Place the materials on students’ desks before the class enters the room.

5. Set out a container for collecting liquid waste.

NOTE Check with your environmental health coordinator to determine whether copper waste can be discharged to

your school’s wastewater system. If so, flush wastes down the drain with copious amounts of water. If not, follow the disposal procedures established by your school district.


OBJECTIVES FOR THIS LESSON


MATERIALS FOR LESSON 13For you 1 copy of Student Sheet 13.1:

Performance Assessment 1 copy of Student Sheet 13:

Written Assessment 1 pair of safety gogglesFor you and your partner 1 jar of solid A 1 bottle of clear solution B 1 250-mL beaker 1 lab scoop 3 test tubes 1 metric ruler 1 black permanent marker 3 labels Access to an electronic

balance Access to water

Use your knowledge and skills to complete an assessment of what you have learned during the unit Experimenting with Mixtures, Compounds, and Elements.

MATERIALS FOR LESSON 13For the teacher

1 pair of safety goggles*

For each student 1 copy of Student

Sheet 13.1: Performance Assessment*

1 copy of Student Sheet 13: Written Assessment*

1 pair of safety goggles*

For each pair of students 1 jar (2 oz) of copper (II)

sulfate, labeled A 1 bottle (125 mL) of

limewater (calcium hydroxide solution), labeled B

1 beaker, 250 mL 1 lab scoop 3 test tubes, 16 × 125 mm 1 metric ruler* (marked in

millimeters) 1 black permanent marker 3 labels

For the class 4 electronic balances* 1 waste container* Access to water* *Needed but not supplied

FINAL ASSESSMENTLESSON


ASSESSMENT

1. Table 1 on Inquiry Master 13 is a suggested scoring rubric for Table 1 on Student Sheet 13.1.

2. Table 2 on Inquiry Master 13 provides answers for the remaining questions on Student Sheet 13.1.

3. Table 3 on Inquiry Master 13 lists the answers for the written assessment.

Any three of the following properties: blue, solid, crystalline, soluble

A solid was formed and the solution became cloudy.

A new substance was formed.

Answers based on student’s results.

Answers No or 0 g, or gives an explanation that the error could be a result of the student’s own mistake or the precision of the balance.

Answers based on student’s results.

Answers No or 0 g, or gives an explanation that the error could be a result of the student’s own mistake or the precision of the balance.

Explanation correct.

Explanation includes the correct contextual use of the words provided, up to a maximum of any six.

2A–2C

3A

3B

4A

4B

4C

QUESTION ANSWER

TABLE 13.2 ANSWERS FOR QUESTIONS 2 THROUGH 4

Mass 1

Mass 2 = Mass 1 ± 0.1 g

Mass 3 = Mass 1 ± 0.1 g

—

Correct calculation

Correct calculation

Before Mixing Contents of Tube A and Tube B

After Mixing Contents of Tube A and Tube B

After Mixing Contents of Tube A and Tube C

MASS OF THE THREE TUBES AND SUBSTANCE(S) (g)

CHANGE IN MASS (g)

TABLE 13.1 ANSWERS FOR TABLE 1 ON STUDENT SHEET 13.1


1

2A

2B

3A

3B

3C

4A

4B

5

6A

6B

6C

6D

6E

7A

7B

7C

8A

8B

8C

C

D

A

C

B

A

D

D

D

C

B

D

B

B

A

D

B

B

B

D

QUESTION ANSWER

TABLE 13.3 ANSWER KEY FOR STUDENT SHEET 13


SAFETY TIPS Wear safety goggles at all times during the inquiries.

Keep long hair tied back.

Be careful with hydrochloric acid. If acid gets on your clothes or skin, wash it off immediately with lots of cold water.

Tell your teacher immediately about any accidents involving acid.

4

Share your ideas and examples in a class discussion.

5

In this lesson, you will investigate the properties and behaviors of a variety of mixtures and pure substances. Working with another student, you will complete eight inquiries. You will be graded on the seriousness of your efforts, on the carefulness of your observations, and on your cooperation with your lab partner.

6

Listen carefully to the safety instructions given by your teacher. In order to participate in the lab work for this class, you are expected to follow safe laboratory procedures.

7

Each inquiry station has a basic topic. Each pair of students will start at a different inquiry station. At each station you will follow the instructions on the Inquiry Card. These instructions are also in your student guide. When you make observations or think you can explain what you are observing, you should discuss these ideas with your partner. Remember: Exchanging ideas with others is a very important part of science. You will have 4-5 minutes to complete each inquiry and record your observations on Student Sheet 1: Our Ideas About Pure Substances and Mixtures.

8

Your teacher will set up two sets of identical inquiries—A and B. You will be assigned to either A or B. Place an asterisk beside your first inquiry on the student sheet so that you write your answers in the correct spaces.

9

When you have completed each inquiry, put the apparatus back as you found it at the beginning of the experiment.

10

If you have any questions about the procedure, you should ask your teacher now.


INQUIRY 1.2

FOLD THE FILTER PAPER TWICE IN TWO DIFFERENT DIRECTIONS AS SHOWN.FIGURE 1.2

FILTERING A MIXTURE

PROCEDURE

1

Observe the mixture in the jar labeled Substance A. Does it appear to be more than one substance? Describe it on Student Sheet 1.

2

Place a lab scoop of the mixture in your test tube and add 10 mL of water.

3

Place a stopper on the test tube, and shake the test tube for about 30 seconds. (Be careful not to hit the test tube against your lab table or another hard object.)

4

What happens to the contents of the test tube? Record your observations.

5

Fold a filter paper (see Figure 1.2) and place it in the funnel. Pour the contents of the test tube through the funnel into the second test tube.

6

Describe the appearance of the substance on the filter paper.

7

What do you think happened to the parts of the original mixture once you added the water and filtered it?

8

How do you think you might get back all the parts of the original mixture?

9

Rinse out the test tubes and funnel and throw away the used filter paper.

10

Replace the apparatus for the next group.


PROCEDURE

1

Have your teacher light the candle.

2

What can you see taking place at or near the top of the candle? Write your observations on Student Sheet 1.

3

Place the open end of the beaker over the candle (see Figure 1.3). Let the beaker stay over the candle for a few minutes.

4

What happened after the beaker was placed over the candle? Record your observations.

5

Why do you think the candle reacted the way it did? Record your answer.

6

Restore the apparatus to its original condition for the next group.

THE BURNING CANDLE

INQUIRY 1.3

Upturned beaker

Burning candle

AFTER YOU HAVE RECORDED YOUR OBSERVATIONS OF THE LIT CANDLE, PLACE THE BEAKER OVER THE CANDLE.FIGURE 1.3


THE NATURE OF MATTERLESSON 1

PROCEDURE

1

Put four lab scoops of rock salt into the plastic cup. Examine it with the magnifying loupe. Write a description of the rock salt on Student Sheet 3.2: Cleaning Rock Salt.

2

Most of the salt used in food is made from rock salt. Discuss these questions with your partner:



C. What do you think the contaminants could be?

3

How could you use the remaining apparatus you have been given to obtain only the soluble component of the rock salt? Record your answers to the following questions:

• What are you trying to do?

• What materials will you use?

4

Record the procedure you and your partner devised.

5

Check your ideas with your teacher.

6

Follow your procedure to purify the salt. If you have any problems, consult your teacher.

CLEANING ROCK SALT

INQUIRY 3.2

YOU EAT THIS ROCK. WHAT IS IT AND HOW IS IT PURIFIED?PHOTO: NASA African Monsoon Multidisciplinary Analyses (NAMMA)

SAFETY TIP Do not taste any of the substances used in the laboratory.


EXTENDING YOUR KNOWLEDGE

chromatography its name.) She stands the plate upright, with the samples along the bottom, in a container with a small amount of liquid. As the drops become moist and interact with the coating on the plate, they begin to move up the plate at different rates, depending on the solubility of the components. Once the liquid gets near the top of the plate it is removed from the container. Then, the positions of the drops are compared. All dyes made from the same components will form the same pattern on the plate. So, if the pattern of the crime scene sample matches the pattern of the dye used by banks, another crime has been solved.

BOMBS AND EXPLOSIVESChromatography also comes in handy for analyzing the materials used in bombs and explosive devices. The FBI analyzes samples from all major bombings involving the United States, including the one at the Murrah Federal Building in Oklahoma City and others causing airline crashes. The technique is called high-performance liquid chromatography—HPLC, for short.

The first step is examination under a microscope. “Most bomb samples look pretty much alike. They look like black powder,” says Mount. Even so, this first step is important. The scientists might, for example, be able to sort out small pieces of material from the residue.

The next step is extraction. The chemists place the sample in a solvent such as water. Once in solution, the particles in the sample may, depending on the composition of the sample, separate into smaller particles that carry positive or negative charges.

A small amount of the solution is placed in the HPLC machine. It moves up to the top, where it mixes with another liquid, and is then forced downward under pressure through a narrow glass column that is filled with a porous substance.

What happens in the column is the critical step. “Some of the [particles],” explains Mount, “seem to like it better in the tube than others. They stay longer.”

The speed at which the particles leave the column is recorded by a detector, which then prints out the information. By comparing the time that the particles have stayed in the column with known retention times, Mount and her colleagues are able to distinguish the various types of particles in the test sample.

STILL A LOT TO LEARNDoes it always work? “No,” says Mount. “Sometimes we find nothing. And other times, we find nothing conclusive. It’s also important to note that when it comes to explosive materials, HPLC is only a qualitative analysis technique. It helps us identify what materials are in an unknown powder. It doesn’t provide quantitative information; in other words, we can’t tell how much of each substance is in the powder.”

Kelly Mount loves her work. To prepare for her career, she earned a bachelor’s degree in chemistry and then went on to get a master’s degree in forensic science. Life in the lab is never routine—this means getting called into the lab on weekends or even at night when there is an emergency. Whether the problem is bombs or banks, Mount has the expertise to help the FBI solve its mysteries. n

1. Envision a crime scene not described in the reading selection in which chromatography could be useful. Describe how chromatography could be used to help solve the crime.

2. Describe the procedure you would use to find out whether carrots get their color from one or many substances.

1.



LESSON


THE SURFACE OF THE EARTH IS 72 PERCENT WATER. WHAT IS WATER MADE FROM?PHOTO: NASA

INTRODUCTION









BREAKING DOWN A COMPOUND

BUILDING YOUR UNDERSTANDING

Water is a compound made up of two elements—hydrogen and oxygen. The characteristic properties of these elements are different from those of water. However, hydrogen and oxygen have some common properties. They are both colorless, odorless gases, and they both readily react with other elements, making them “reactive” elements. But in many ways they are very different from each other.

Hydrogen has the lowest density of all the elements. It is very reactive, which is one reason why it is present in only very small quantities in air. It reacts with oxygen. You reacted it with oxygen when it burned with a squeaky pop. What do you think was made in that chemical reaction?

It may come as a surprise to you to discover that hydrogen is the most common element in the universe. The sun and other stars are mainly hydrogen gas. Hydrogen is found in many compounds. For example, sulfuric acid (used in car batteries) and hydrogen peroxide (used in hair dyes) contain hydrogen.

Oxygen reacts with other substances. Oxygen is needed for burning to take place. Things burn well in oxygen, producing hotter flames. For example, what happened to the glowing splint when it was put into a tube of almost pure oxygen? Some welding and metal-cutting equipment use flammable gases and pure oxygen to produce the high temperatures needed to melt metal.

Oxygen also reacts slowly with many substances. Many compounds containing oxygen are called oxides. You may be familiar with two oxides that are gases—carbon dioxide and sulfur dioxide—but most oxides are solids. In fact, oxygen is the most common element in the Earth’s crust, but most of it is combined with other elements to form minerals that make up rocks. n

WATER IS A COMPOUND FORMED WHEN INFLAMMABLE HYDROGEN REACTS WITH OXYGEN. HERE IT IS BEING USED TO PUT OUT A FIRE. LIKE ALL COMPOUNDS, THE PROPERTIES OF WATER ARE VERY DIFFERENT THAN THOSE OF THE ELEMENTS FROM WHICH IT IS COMPOSED. PHOTO: U.S. Air Force photo by Senior Master Sgt. David H. Lipp


9

Immediately show the water level by filling in the diagram in Table 1. Write a description of the damp steel wool in the third column of the table.

10

Watch the apparatus carefully. Do you notice anything happening in the tubes?

11

After about 15 minutes, look at the level of the water in each tube. Record your observations in the appropriate places in Table 1.

12

Describe the appearance of the steel wool.

13

Air is about 21 percent oxygen. Keeping the tubes in position, remove the card and use the ruler to measure how far up the test tube the water has moved. Answer these questions on Student Sheet 8.2:

•Has the water level changed in either tube? If so, can you explain the change?

•What can you conclude from your observations?

14

Clean the test tubes and the beaker. Make sure you remove the steel wool from the test tube and place it in the trash. Return the cleaned apparatus and the index card to the plastic box.


Discuss your results and compare them with those of other pairs.

Answer the questions in Steps 4–7 on Student Sheet 8.2.

All chemical reactions have reactants (inputs) and products (outputs) and can be written as simple word equations. For example, you know that hydrogen combines with oxygen to form water. A simple word equation to describe this equation is as follows:

hydrogen + oxygen = water

On Student Sheet 8.2, complete Step 8 by writing a word equation for the reaction that took place in the test tube. Label the reactants and the products in the equation. This is an example of the synthesis of a new product from a chemical reaction. In Lesson 10, you will investigate how metals react with acids to form new products.

1

2

3



EXPLORATION ACTIVITY COMPOUNDSThis is a list of elements, their familiar compounds, and the chemical formulas for those compounds:

• Hydrogen compounds—acids, hydrides, and hydroxides

Hydrochloric acid (HCl) Boric acid (H3BO3) Carbonic acid (H2CO3) Sulfuric acid (H2SO4) Nitric acid (HNO3) Hydrogen peroxide (H2O2) Methane (CH4) Ammonia (NH3) Butane (C4H10)

• Carbon compounds or organic compounds, which make up a separate branch of chemistry

Carbon dioxide (CO2), a major greenhouse gas

Carbon tetrachloride (CCl4), formerly a dry-cleaning compound

Freon (CCl2F2) formerly in air conditioners

• Oxygen compounds—oxides Silicon dioxide (SiO2) Magnesium oxide (MgO) Iron (II) oxide (Fe2O3) Nitrous oxide (N2O) Sulfur dioxide (SO2)

• Sodium, magnesium, phosphorus, and calcium compounds—abundant in earth’s crust

Sodium hydroxide (NaOH), a drain cleaner

Sodium bicarbonate (Na2S2O3), baking soda

Sodium thiosulfate (NaS2O3), a fixing agent in photo printing

Magnesium chloride (MgCl2), abundant in seawater

Magnesium oxide (MgO), 35% of earth’s crust

Hydrogen phosphate (H3PO4), in soft drinks, fertilizers, and solvents

Calcium oxide (CaO), lime Calcium carbonate (CaCO3), limestone Calcium sulfate (CaSO4), in gypsum

and plaster of paris Calcium phosphate Ca3(PO4)2,

the primary inorganic substance in bones and teeth

• Sulfur compounds Hydrogen sulfide (H2S),

has a “rotten egg” odor Sulfur dioxide (SO2), a refrigerant,

bleaching agent, and disinfectant Potassium sulfate (K2SO4), in fertilizer,

pigments, and dyes


cids and bases were known and used by alchemists before their chemical structures were identified. Substances such as vinegar and citrus fruit juice were identified as sour tasting and described as acids.

The term “acid” is taken from the Latin “acidus” for “sour.” French chemist Antoine Lavoisier identified the gas produced by the reaction of acids with metals. Other substances known for their bitter taste and slippery feel were called “alkalis.” The word “alkali” came from the Arabic word for “ashes” because these substances were obtained by soaking ashes from wood fires in water.

In the late 1800s, Swedish chemist Svante Arrhenius was testing the ability of certain solutions to conduct electricity. (The invention of a storage battery to produce electricity led to many electrolysis experiments.) He theorized that some solutions, called electrolytes, decomposed into charged particles, or ions, and produced electricity when dissolved in water. Acids, he suggested, split up when placed in water and produced hydrogen ions with a single positive charge (H+). He described bases as substances that dissolved in water and formed negatively charged hydroxide ions (OH—).

Some common acids and their chemical formulas are hydrochloric (HCl), nitric (HNO3), sulfuric (H

2SO

4),

carbonic (H2CO

3), and acetic (C

2H

4O

2). All carbonated beverages contain carbonic acid, and many have small

quantities of phosphoric acids. Citrus fruits include ascorbic acid, or vitamin C (C6H

8O

8). Many household

cleaning products contain hydrochloric acid. You will notice that all these compounds contain hydrogen, which produces the Arrhenius H+ ion in water.

Bases are equally common in everyday life and use. They include sodium hydroxide (NaOH) (often called lye and sometimes used as a drain cleaner), ammonium hydroxide (NH

4OH), baking soda (NaHCO

3),

substances contained in ashes (K2CO

3), and soaps. When solutions of bases and acids are mixed, water and

a new substance are produced. This reaction is called a neutralization reaction because the acid (H+) and the alkali (OH—) neutralize one another to form H

2O. The new substances formed in neutralization reactions are

called “salts.” One of these salts, NaCl, the product of the reaction between hydrochloric acid and sodium hydroxide, is what we call “table salt,” or sodium chloride. This is only one of many substances called salts.

Arrhenius’s definitions of acids and bases were modified in the 20th century as the understanding of atomic structure grew. Two scientists working independently reached the same conclusion on the structure of acids and bases. Chemists T.M. Lowry from England and J.N. Brønsted of Denmark each developed new definitions of acids and bases. They defined an acid as a substance that can donate a proton, the H+ ion. They defined a base as a substance that can accept a proton. This definition of acid-base reactions, known as the Brønsted-Lowry Theory of acids and bases, is widely accepted by scientists today. The Brønsted-Lowry Theory illustrates the gradual progression of explanations in science. New understandings of the structure of matter have brought new explanations for well-known behaviors of matter. n

1. How do acids and bases differ in their chemical and physical properties?

2. Why do scientists change their explanations of things over time?

1.


READING SELECTION EXTENDING YOUR KNOWLEDGE


CHEMICAL REACTIONS INVOLVING METALSLESSON 10

experimenting with mixtures, compounds, and elements changes

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