is there a smallest piece of matter or can we keep cutting

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Copyright © 2012 Montana Partners This project was largely funded by an ESEA, Title II Part B Mathematics and Science Partnership grant through the Montana Office of Public Instruction.

The Atom Is There a Smallest Piece of Matter or Can We Keep Cutting a Piece in Half Infinitely?

By convention there is color, by convention sweetness, by convention bitterness, but in reality there are atoms and space.

Democritus (460–400 BC)

Engage: How Do You Model Something You Cannot See? A. You’ve learned about atoms in your previous courses. What is an atom? What evidence do

you know that proves that matter is made of atoms?

Before Studying this Unit After Studying this Unit

1. Your teacher will demonstrate a Think Tube. Your task is to determine how the strings

are arranged inside the sealed tube without direct knowledge of how it was constructed. In science, the term model is used to refer to a representation of something else, often something that cannot be seen. In this case, your challenge is to draw a representation—a model—of the inside of the Think Tube. You will have to use observations of the strings outside of the tube to develop your model of its interior.

UNIT

3

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High School Chemistry: An Inquiry Approach

Demonstration 1 Observations Sketch of the Interior of the Tube A characteristic of models is that they can be revised as you obtain more data. Does the next demonstration warrant a revision of your model? Demonstration 2 Observations Sketch of the Interior of the Tube Demonstration 3 Observations Sketch of the Interior of the Tube Demonstration 4 Observations Sketch of the Interior of the Tube What can you conclude about the interior of the tube? Carefully explain how your observations support your model. Check with other students. Does anyone have an alternative model? Use evidence to support or criticize different models.

Thinking About Your Thinking Mental Models

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Explore 1: What Happens to the Amount of Matter When There Is an Observable Change? 2. Obtain three plastic zip lock bags from your teacher and three plastic medicine cups for

you to investigate the amount of matter that exists before and after a change takes place. You are going to perform three tests.

3. Partially fill one of your zip lock bags with some ice, and close the zip lock tightly.

Place the zip lock bag and its contents on a balance and find and record the mass. Let the contents of the bag melt completely and mass it again. Has the mass changed or has it remained constant? Explain.

4. Carefully place one Alka-Seltzer tablet and a plastic medicine cup that has water in it in

a second zip lock bag, seal it and mass the contents without allowing the water to spill from the medicine cup, and record the bag’s mass. Now mix the contents of the bag and allow the Alka-Seltzer and water to react. When the reaction has finished again place the zip lock bag on the scale and record its mass. Has the mass changed or has it remained constant? Explain.

5. Place some iron (III) chloride solution in one medicine cup and some potassium

thiocyanate solution in the second cup. While being very careful to not allow the solutions to spill or mix, place them in a third zip lock bag and seal it, and then find the mass of its contents and record it. Now carefully allow the two solutions to mix and after the reaction has completed, again mass the contents of the bag. Has the mass changed or has it remained the same? Explain.

Explain 1 6. Describe a generalization that can be made about the mass of the contents of the zip

lock bags before and after a change has taken place.

7. Given that the zip lock bags do not allow materials to escape, how could your

generalization above allow you to make predictions about the difference in mass of matter before and after a change in any chemical system?

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Elaborate 1: How can Before and After Masses of a Chemical or Physical Change be Used to Make Predictions? 8. The before and after mass relationship you discovered in the Explore activity is called

The Law of Conservation of Mass. What do you think the word “conservation” refers to?

9. Return to the Alka-Seltzer and water reaction. If this reaction were run in a cup that was

not sealed in a zip lock bag would you expect the final mass to be more or less than the initial mass? Explain your reasoning.

10. If 5.00 grams of Alka-Seltzer were mixed with 5.00 grams of water in an unsealed cup,

and the final mass of the contents of the cup were 7.50 grams, how many grams of gas were produced? Explain you reasoning.

11. Suppose that 25.00 grams of a sodium chloride solution were mixed with 25.00 grams

of a silver nitrate solution in a sealed system and a white solid was observed to form. What would you predict the final mass of the contents to be? Explain your reasoning.

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Explore 2: What is the Relationship Between the Mass of a Substance Before and After Heating? Option A: 12. Place about 1 cm3 of either (a) hydrated copper(II) sulfate or (b) hydrated magnesium

sulfate into a Pyrex test tube. Half of the lab groups in class will be assigned to option (a) and half will be assigned to (b), so you will need to share data with another group. Record your observations in the table below.

13. Light a Bunsen burner according to the instructions provided by your teacher. Adjust

the flame so it burns with maximum efficiency as a two-coned flame. 14. Clamp the test tube near the top with a test tube holder. Turn the tube so that it is

nearly horizontal. Gently tap the tube so that your solid sample spreads out into a thin line that covers at least half of the length of the tube. While keeping the test tube nearly horizontal, point the opening so that it is turned away from you and others in the lab. Gently ease the tube into the flame, keeping it in constant motion to avoid concentrating the heat in any one area. Heat the tube and its contents and observe the changes. When there appears to be no further changes occurring, place the test tube and holder in a test tube rack. Record your observations while the tube cools.

15. When the test tube has cooled, add a drop of water to the solid. Observe what happens,

then add another drop. Hydrated Copper(II) Sulfate Hydrated Magnesium Sulfate

Before Heating

After Heating

After Adding Water

16. Your observations provide evidence of one product of the change that occurred when the hydrated substances were heated. What is that product? Justify your answer with supporting evidence.

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17. Obtain a porcelain crucible and lid. Wash the crucible, rinse it with distilled water, and

heat it for about 4–5 minutes over a Bunsen burner on a wire triangle on a ring stand. Place the lid on the crucible so that it is slightly cocked, with a small opening. Practice using crucible tongs by manipulating the lid and moving the crucible back and forth from the triangle to your designated cooling surface two or three times.

18. Your teacher will assign a hydrated compound and sample size to each group in the class. First weigh your crucible and lid, and then add the assigned mass (±0.1 g) of the assigned hydrated compound to your crucible.

19. Place the crucible, lid, and sample on the wire triangle on the ring stand. Move the lid so that there is a small opening, as you practiced with the empty crucible. Heat the sample slowly for a few minutes, and then increase the heat. Heat for another five minutes, then remove the crucible from the flame and allow it to cool on the designated cooling surface. When the crucible is cool to the touch, its mass can be measured and recorded.

20. Heat the sample again for about five minutes. Allow it to cool, and measure its mass. If the mass after the second heating is ±0.05 g from the mass after the first heating, you have completed this portion of the exercise. If your mass is significantly lower than the mass after the first heating, heat, cool, and determine the mass of the sample one more time.

21. Dispose of your sample and clean your work area as directed by your teacher. Your teacher will provide a method for pooling class data. Record your data and the class data by setting up tables in the space below.

CAUTION! Handle the crucible and lid with tongs only until you are certain that it has cooled. A hot crucible looks like a cooled crucible.

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Option B: 22. Obtain 2 glass Pasteur pipettes. Break off about 1 to 2 cm of the stem of one of the

pipettes so there is a discernable difference between the two. Measure and record the mass of each pipette to the nearest 0.01 gram.

23. Your instructor will assign your lab group to partially fill your tubes ¼, ½, or ¾ full with the hydrated salts. One tube will be used for the copper(II) sulfate hydrate and the other for the magnesium sulfate hydrate. Keep track of which is which! Fill your tubes as assigned and measure and record the mass of the partially filled tubes to the nearest 0.01 gram.

24. Lay the tubes on a wire gauze square that is sitting on a support ring clamped to a ring stand. Slightly bending upward the sides of the wire gauze square will keep the pipettes from rolling off. Light your Bunsen burner and GENTLY heat the two tubes simultaneously for five minutes. Move the flame back and forth slowly beneath the tubes being careful not to overheat any one area or melt the glass pipette. Make and record observations about the changes in appearance of the salts while heating.

25. After five minutes, allow the tubes to cool to touch. Lift the tubes by the stem and re-measure the mass.

26. Return the tubes to the stand and re-heat both tubes for one minute. Allow the tubes to cool and re-measure the mass.

27. Heat the tubes a third time, and a gain cool and re-measure the mass. 28. If the three post-heating masses are all within 0.01 gram of each other, then it can be

assumed that all the water has been driven off. If the three masses are not within 0.01 gram of each other, then continue to heat, cool, and re-measure until they are.

29. After the final mass measurement has been made and the tubes are completely cool the

hydrated salts are completely dehydrated, and are referred to as anhydrous salts. Grasp the middle of the tube in between your thumb and forefinger. One tube at a time, add a couple of drops of water to the large opening rehydrating the anhydrous salt. Record your observations.

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Data table for Option B: Hydrated Copper(II)

Sulfate Hydrated Magnesium Sulfate

Before Heating Qualitative

Quantitative

After Heating Qualitative

Quantitative

1.

2.

3.

After Adding Water Qualitative

30. Dispose of your samples as instructed by your teacher.

31. Your teacher will provide a method for pooling class data. Record both your data and the class data for use in analysis.

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Both Option A and Option B continue here 32. Plot the mass of anhydrous sample (y-axis) versus the mass of the hydrated sample on

the grid below, and plot the mass of water lost versus the mass of the hydrated sample on the second grid.

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Explain 2 33. Draw the line of best fit to your data. Determine the equation of the best-fit lines for

each plot. Write your final equations in a form that allows you to predict (a) the mass of anhydrous sample that will result from heating a given mass of the assigned compound and (b) the mass of the water lost when heating a given mass of the assigned compound.

34. Determine the percentage of water in the hydrated compound from the appropriate line

of best fit.

35. Determine the percentage water in each of the different sample masses.

How does the percentage water obtained by averaging the individual samples compare with the percentage water from the best-fit line in the graph of the mass of the hydrate vs. the mass of the anhydrate? Why does this relationship occur? What are the advantages of using the line of best fit instead of averaging individual samples?

36. Let’s consider two models that may explain your data:

One model depicts matter as made up of tiny bar magnets. When the north pole and south pole of one magnet has stuck to the south pole and north pole other magnets, it has no further ability to stick to other magnets. Anhydrous magnesium sulfate is one magnet in this model, and water is another magnet. Another model depicts matter as made up of different combinations of food coloring. The two types of matter in this investigation, anhydrous magnesium sulfate and water, are like blue and yellow food coloring. The hydrated compound is analogous to green food coloring, a combination of blue and yellow.

(a) You are going to be given three test tubes, one empty, one that has blue coloring in it and one that has yellow coloring in it. Mix the two colors in the third test tube until you have achieved a green color. Do you think every group in the class will produce the same color green? Why or why not?

(b) When finished, compare your green colored test tube to the green test tubes produced

by other groups, by holding them up to a white background in bright light. Is your

Thinking About Your Thinking Proportional Reasoning

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prediction from part (a) supported or not? Why or why not? Which model is a better fit to your data? Explain.

Elaborate 2: How Can a Small-Scale Experiment be Used to Make Large-Scale Predictions? 37. If you were to heat 155 lb of the assigned hydrated compound, how many pounds of

anhydrous compound would be left? How many pounds of water would be lost?

38. How many kilograms of your assigned hydrated compound do you need if you wish to

produce 375 kg of the anhydrous compound by heating the hydrate?



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Explore 3: What is the Mass Relationship in the Reaction of Copper and Sulfur? This exploration describes instructions about performing a guided investigation in a style that is typical of this textbook. However, because of the safety hazards associated with this procedure, we have conducted the investigation for you. Our data appear in the space where you would have recorded your information. Read the instructions and responses, and proceed with the analysis based on our data. 39. Obtain a porcelain crucible and lid. Wash the crucible, rinse it with distilled water, and

heat it for about 4–5 minutes over a Bunsen burner on a wire triangle on a ring stand. Place the lid on the crucible so that it is slightly cocked, with a small opening. After the crucible has been heated long enough to drive off any remaining water, use crucible tongs to move it to the cooling surface designated by your teacher. Allow it to cool to the touch.

40. Determine the mass of your crucible, and record its mass in a table below. Add 0.5 to 9

grams of medium copper shavings or loosely-rolled copper wire to the crucible, and measure the mass of the copper and crucible. Record your qualitative observations about the visible properties of copper.

Mass of crucible: 19.42 g Mass of copper + crucible: 24.46 g The copper wire is metallic-looking and shiny, with a reddish color. The wire is easily bent and retains its shape after bending. Its diameter is about 1 mm. 41. Add powdered sulfur to the container, adding about 2/3 to 3/4 of the mass of the

copper added for each sample. Record the mass of the crucible after the sulfur added, and write your description of the appearance of sulfur.

Mass of sulfur + copper + crucible: 27.99 g The sulfur is a finely-divided bright yellow powder. It emits a faint smell like that of matches. 42. Work in a fume hood as directed by your teacher. Place your Bunsen burner and ring

stand in the hood. Place the lid fully over the crucible. Heat the mixture, slowly at first, gradually increasing the heat until there is no further evidence of sulfur burning around the lid. When there is no further evidence of burning of sulfur, heat the mixture strongly for about 5 more minutes.

43. Allow the crucible and its contents to cool. Do not open the lid during the cooling

process (to prevent reaction with oxygen in the air). 44. When the crucible has cooled, open the lid and observe the contents. If you see any

remaining sulfur, place the lid on the crucible and reheat. If there is no visible evidence of unburned sulfur, determine the mass of the crucible and its contents, and record your qualitative observations of the product compound.

Mass of product compound + crucible: 25.87 g The product still appears like a wire, but now it is grey. 45. Dispose of the product and clean your work area as directed by your teacher. Pool your

data with that of your classmates as directed by your teacher. Setup and complete a data table(s) in the space below.

CAUTION! Handle the crucible and lid with tongs only until you are certain that it has cooled. A hot crucible looks like a cooled crucible.

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Unit 3 The Atom

Student Initials

Mass of crucible

(g)

Mass of Cu +

crucible (g)

Mass of Cu (g)

Mass of S + Cu

+ crucible

(g)

Mass of product

+ crucible

(g)

Mass of product

(g)

Mass of S in

product (g)

DJ + BT

19.42 24.46 27.99 25.87

MH + TF

22.44 23.70 24.54 24.05

SE + LB

22.10 31.00 37.05 33.49

MA + MD

19.37 26.38 31.22 28.34

PP + KS

20.04 26.86 31.63 28.77

JD + MC

19.81 25.94 30.29 27.66

SS + FT

21.39 23.74 25.50 24.40

PS + GS

22.53 26.00 28.50 26.97

AF + PC

21.24 25.93 29.21 27.24

KM + KL

20.67 25.25 28.41 26.53

KK + LR

21.75 26.77 30.38 28.18

BM + JT

20.00 22.60 24.50 23.33

46. Assume that there was more than enough sulfur in the crucible to react with all of the

copper. Any excess sulfur burned in air, forming gaseous products that were vented away. Calculate the mass of sulfur that reacted, and enter this quantity in your data table. Show the setups for all of the MH + TF calculations below.

Explain 3

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47. Compare your qualitative observational data about the appearance of the copper and sulfur before the reaction with the product compound after heating. Is there evidence that the original substances were destroyed and a new substance was formed? Explain.

48. Plot the mass of sulfur in the product (y-axis) versus the mass of copper (x-axis) on the

first grid, and plot the mass of product compound formed (y-axis) versus the mass of copper reacted (x-axis) on the second grid.

Graph 1

Graph 2

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49. Draw the line of best fit to your data. Determine the equation of the best-fit lines. Write

your final equations in a form that allows you to predict (a) the mass of sulfur in the product with any given mass of copper and (b) the mass of product compound formed from any given mass of copper when more than enough sulfur is present.


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50. What do your data indicate about the nature of the reaction of copper and sulfur? Does the fraction of copper and sulfur in the product compound vary or is it a definite, fixed fractional amount? Write your answer to this question and show how your data support your answer.

51. Reconsider the two models of matter proposed earlier in this unit: One model depicts matter as made up of tiny bar magnets. When the north pole and south pole of one magnet has stuck to the south pole and north pole other magnets, it has no further ability to stick to other magnets. Copper is one magnet in this model, and sulfur is another magnet.

Another model to consider depicts matter as made up of different combinations of food coloring. The two types of matter in this investigation—copper and sulfur—are like blue and yellow food coloring. The product compound is analogous to green food coloring, a combination of blue and yellow.

Which model is a better fit to your data? Explain.


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Elaborate 3: How Can We Use the Definite Composition of Compounds to Make Predictions?

52. The relationship you discovered in Explore 1 and 2 is called the Law of Definite

Composition. A student conducted an experiment similar to your copper–sulfur experiment. He heated magnesium metal in the air to allow it to react with oxygen. His data are in the table below.

Mass of Magnesium (g)

Mass of Product (g)

Mass of Oxygen (g)

Percentage Magnesium in Product

Percentage Oxygen

in Product 2.00 3.34

2.40 3.98

2.80 4.64

3.20 5.30

a) Determine the mass of oxygen in the product compound for each trial. Enter it in

the table above.

b) What is the percentage magnesium in the product compound? What is the

percentage oxygen in the product compound? Enter your data in the table above.

c) The student predicted that if he heated 5.00 g of magnesium in air, 8.33 g of

product compound would be formed. Do you agree with his prediction? Explain why a prediction can be made, and explain why you agree or disagree.

53. Analysis of 84.01 g of baking soda indicates that it contains 22.99 g sodium, 1.01 g

hydrogen, 12.01 g carbon, and 48.00 g oxygen. What is the percentage of each element in baking soda?

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Explore 4: What is the Volume Relationship of the Products of the Decomposition of Water? 54. Your teacher will demonstrate the reaction of water that occurs when it is subjected to

an energy input via an electrical current. Record your observations in the table below. Negative Electrode Positive Electrode Qualitative Observations

Measured Volume

Splint Test Results

Explain 4

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55. What do the results of the splint tests indicate about the identity of the gases? 56. What is the volume ratio of the gases identified by the splint tests? 57. An element is a pure substance that cannot be decomposed. A compound is a pure

substance that can be decomposed into other substances. Is water an element or a compound? Explain.

58. The volume ratio of the product gases provides information about the composition of

water. Based on these data, water appears to be made of x parts of substance A and y parts of substance B. What are x, y, A, and B? Show how the experimental data support your conclusion.

59. Again reconsider the two models of matter proposed earlier in this unit: matter is like

tiny magnets and matter is like food coloring. Which model is a better fit to your data about the decomposition of water? Explain.

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Elaborate 4: What do Studies of the Reaction of Two Elements Tell Us About the Nature of Matter? 60. Some elements react with another element to form two or more different compounds.

For example, when oxygen is streamed over carbon, both carbon monoxide and carbon dioxide can be formed, depending on the reaction conditions. Let’s look at the data:

Mass of Carbon

Reacted (g) Mass of Oxygen

Reacted (g) Mass of Product Compound (g)

Low oxygen conditions

3.0 4.0 7.0

Plentiful oxygen conditions

3.0 8.0 11.0

No other reactions occur between carbon and oxygen. When 3.0 g of carbon reacts with

oxygen, either 4.0 g of oxygen reacts or 8.0 g of oxygen reacts. No other masses of oxygen have ever been observed to react with 3.0 g of carbon. Note that the ratio of the mass of oxygen reacted is a whole number:

€

8.0 g oxygen reacted4.0 g oxygen reacted

= 2

Similar data have been observed for many other reactions between two elements. In all

cases, the ratio of masses of one element that reacts with a fixed mass of another element is a ratio of whole numbers, such as 1:1, 2:1, etc. This is called the Law of Multiple Proportions.

61. An experiment indicates that, under varying conditions, 0.44 g, 0.88 g, and 1.76 g of

nitrogen reacts with 1.00 g of oxygen. (a) How many pairwise combinations of masses of nitrogen can be compared? (b) Determine the ratio of masses of nitrogen for each of the possible comparison pairs. (c) Explain how these data conform to the Law of Multiple Proportions.

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62. Two carbon-oxygen compounds are analyzed for their composition by mass. One 83.4-g sample consists of 22.8 g of carbon and the remainder oxygen, and a different 86.7-g sample contains 37.2 g of carbon with the remainder being oxygen. Is the ratio of masses of oxygen that combines with a fixed quantity of carbon a whole number? Use data to justify your answer.

63. Nitrogen and oxygen react to combine in many combinations. One investigation

analyzed two compounds of nitrogen and oxygen, one 30.4% nitrogen with the remainder oxygen, and the other was 46.7% nitrogen and the balance of the mass was oxygen. Show that the ratio of masses of oxygen is a whole number.

64. Again reconsider the two models of matter proposed earlier in this unit: matter is like

tiny magnets and matter is like food coloring. Which model is a better fit based on the Law of Multiple Proportions? Explain.

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Elaborate 5: Why Do Matter Models… Matter? 65. This activity will focus on answering the following question: How can nuts and bolts be

used to model Dalton’s atomic theory (described in Appendix 1)? 66. A number of compounds composed of the elements nitrogen and hydrogen are known

to exist. One is called ammonia and has the formula NH3 (1 nitrogen atom and 3 hydrogen atoms) and another is called hydrazine and has the formula N2H4 (2 nitrogen atoms and 4 hydrogen atoms).

67. Using a bolt to represent an atom of nitrogen and a nut to represent a hydrogen atom,

construct a model of one particle of ammonia and a second model of one particle of hydrazine.

68. What is the ratio of hydrogen atoms to nitrogen atoms in the ammonia particle model? What is the ratio of hydrogen atoms to nitrogen atoms in the hydrazine particle model?

69. What is the ratio of hydrogen atoms per one nitrogen atom in ammonia to the number

of hydrogen atoms per one nitrogen atom in hydrazine? 70. Construct four more particles of ammonia so that you have a total of five. Use an

electronic balance to measure the masses of the following: ammonia particles, nitrogen atoms (bolts) only and hydrogen atoms (nuts) only. Fill in the data table below:

Number of ammonia particles

Total mass of the ammonia particles

Total mass of the nitrogen atoms only (the bolts)

Total mass of the hydrogen atoms only (the nuts)

1

2

3

4

5

71. On the grid below plot the mass of the nitrogen atoms (the bolts) on the x axis and the mass of the hydrogen atoms (the nuts) for the five ammonia particles.

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72. Construct four more particles of hydrazine so that you have a total of five. Use an

electronic balance to measure the masses of the following: hydrazine particles, nitrogen atoms (bolts) only and hydrogen atoms (nuts) only. Fill in the data table below:

Number of hydrazine particles

Total mass of the hydrazine particles

Total mass of the nitrogen atoms only (the bolts)

Total mass of the hydrogen atoms only (the nuts)

1

2

3

4

5

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73. On the grid below plot the mass of the nitrogen atoms (the bolts) on the x axis and the mass of the hydrogen atoms (the nuts) for the five hydrazine particles.

74. Draw a best fit line for each compound, find the slope of each line, and write an

equation for each line.

75. Use the equation to predict the mass of nuts required to make 12 particles of ammonia.

Use the equation to predict the mass of nuts required to make 12 particles of hydrazine.

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Are these masses the same or different? If they are different what has caused the difference?

76. Which one of the three laws: Law of Conservation of Mass, Law of Definite

Proportions, or Law of Multiple Proportions does the results from the above calculation best align with? Explain your reasoning.

77. Find the ratio of the two slopes from your graph. Divide the larger by the smaller.

78. Find the whole number equivalent of the ratio you found in Item 77. Is this whole

number ratio the same or different as the ammonia-to-hydrazine hydrogen atom per one nitrogen atom ratio you found in question 69? What accounts for this similarity or difference?

79. Which of the three laws listed in Item 76 do the results from the ratio calculation in

Items 78 and Item 69 best align with? Explain your reasoning. 80. Return to the question posed in Item 65 and write a conclusion addressing the question.

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Appendix 1: Dalton’s Atomic Theory and Modern Images of Atoms As you study this section, work to achieve this learning goal: • Identify the main features of Dalton’s atomic theory. As early as 400 BC, Greek philosophers had proposed that matter consisted of tiny, indivisible particles, which they called atoms. In 1808 John Dalton, an English chemist and schoolteacher, revived the concept. We now know that the atom consists of even smaller particles. Today, chemists use some of the most sophisticated research methods ever developed as they continue to seek an understanding of how atoms are put together. But it all started with the vision of John Dalton. Dalton knew about the Law of Definite Composition: The percentage by mass of the elements in a compound is always the same. He was also familiar with the Law of Conservation of Mass: In a chemical change, mass is conserved; it is neither created nor destroyed. Dalton’s atomic theory explained these observations. The main features of his theory are as follows (Fig. 3.1):

Figure 3.1 Atoms according to Dalton’s atomic theory. From Cracolice, M. S., & Peters, E. I. (2011). Introductory Chemistry: An Active Learning Approach. Belmont: CA: Brooks/Cole Cengage Learning.

John Dalton (1766-1844). From Cracolice, M. S., & Peters, E. I. (2011). Introductory Chemistry: An Active Learning Approach. Belmont: CA: Brooks/Cole Cengage Learning.

SUMMARY Dalton’s Atomic Theory

1. Each element is made up of tiny, individual particles called atoms.

2. Atoms are indivisible; they cannot be created or destroyed.

3. All atoms of each element are identical in every respect.

4. Atoms of one element are different from atoms of any other element.

5. Atoms of one element may combine with atoms of other elements, usually in the ratio of small whole numbers, to form chemical compounds.

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Dalton’s theory accounts for chemical reactions in this way: Before the reaction, the reacting substances contain a certain number of atoms of different elements. As the reaction proceeds, the atoms are rearranged to form the products. The atoms are neither created nor destroyed, but simply arranged differently. The starting arrangement is destroyed (reactants are destroyed in a chemical change), and a new arrangement is formed (new substances form).

As with many new ideas, Dalton’s theory was not immediately accepted. However, it led to a prediction that must be true if the theory is correct. This is now known as the Law of Multiple Proportions. It states that when two elements combine to form more than one compound, the different weights of one element that combine with the same weight of the other element are in a simple ratio of whole numbers (Fig. 3.2). This is like threading one, two, or three identical nuts onto the same bolt. The mass of the bolt is constant. The mass of two nuts is twice the mass of one; the mass of three nuts is three times the mass of one. The masses of nuts are in a simple ratio of whole numbers, 1:2:3.

The multiple proportion prediction can be confirmed by experiment. Using a theory to predict something unknown and having the prediction confirmed is convincing evidence that the theory is correct. With supporting evidence such as this, Dalton’s atomic theory was accepted.

In the 1980s, a new way of observing atoms was invented. Scientists built a microscope that could see matter at the atomic level. This instrument is called a scanning tunneling microscope (STM). It works much like an old phonograph record player with vinyl records. The needle would detect the ripples in grooves in the vinyl record and electronically translate those ripples into sound. A STM has a needle with a very tiny, sharp tip that detects the bumps in a surface at the atomic level.

Figure 3.2 Explanation of the Law of Multiple Proportions. Carbon and oxygen combine to form more than one compound. Assume, according to the atomic theory, that carbon and oxygen atoms combine in a 1:1 ratio to form one compound and in a 1-carbon-to-2-oxygen-atom ratio to form another compound. If this assumption at the particulate level is true, the mass of the two oxygen atoms in the 1:2 compound is twice the mass of the one oxygen atom in the 1:1 compound. The oxygen mass ratio is 2:1 in the two compounds with the fixed mass of one carbon atom. This analysis at the particulate level predicts a similar result at the macroscopic level, where masses of combining elements can be measured. Going to the laboratory, we find that (a) 1.0 gram of carbon combines with 1.3 grams of oxygen to form carbon monoxide, CO, and (b) 1.0 gram of carbon combines with 2.6 grams of oxygen to form carbon dioxide, CO2. The macroscopic mass ratio of oxygen that combines with 1.0 gram of carbon in the two compounds is 2.6/1.3, which reduces to 2/1, exactly as predicted by the atomic theory at the particulate level. From Cracolice, M. S., & Peters, E. I. (2011). Introductory Chemistry: An Active Learning Approach. Belmont: CA: Brooks/Cole Cengage Learning.

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Figure 3.3 shows the surface of a sheet of nickel. The individual nickel atoms are

distinct from one another, unlike the homogeneous appearance of a macroscopic sample of smooth, polished metal. What we can observe at the macroscopic level is very different from what we would see if we could visualize matter at the particulate level.

As scientists become more familiar with the capabilities of the STM, they found that they could pick up and move individual atoms. The image in Figure 3.4 was created by using a STM to move very-low-temperature xenon atoms on a nickel surface. Such images provide strong evidence of the atomic nature of matter! This evidence, combined with the evidence from the data you’ve been studying in this unit and data from an uncountable number of different types of experiments, have led us to the belief that the data supporting the atomic theory is among the strongest forms of verification of a particulate-level model known to exist.

Figure 3.3 A STM image of the surface of nickel. Nickel atoms are arranged in a regularly repeating pattern called a crystal. Image courtesy of IBM Corporation.

Figure 3.4 A STM image of xenon atoms on a nickel surface. Image courtesy of IBM Corporation.

Unit 3


Unit 3 The Atom

Homework Questions Is Mass Conserved in Chemical Change? 1. If a 1.00 L sample of pure water is heated to boiling and kept boiling until all the liquid

water had been converted to steam, how many grams of steam would be produced?

2. When iron rusts it reacts with oxygen to produce iron oxide. If an iron nail is placed in a beaker of water and allowed to rust will the mass of the nail increase, decrease, or stay the same over time? Explain your reasoning.

3. A 10.00 gram iron nail is placed in a beaker of water an allowed to rust. After some time

its mass is re-measured and found to be 11.24 grams. How many grams of oxygen reacted with the nail? Explain your reasoning.

4. The nail in question 3 is left in the beaker for a month after which time all the water has

evaporated and nothing but a reddish powder remains. Curious, you find the mass of the reddish powder to be 12.86 grams. You add more water to the beaker completely submerging the reddish powder and let it stand. After another month the water has once again evaporated and you re-measure the mass and again find it to be 12.86 grams. What could be a possible explanation for the observed change in mass between the first and second weighing, but no change between the second and third weighing? What is the maximum amount of oxygen that a 10.00 gram iron nail will react with?

Do Compounds have Variable or Definite Composition? 5. In an experiment, a student found that a 10.00 gram sample of a hydrated salt was made

up of 4.60 grams of water and the rest anhydrous salt. A second student found that a 15.00 gram sample was composed of 6.90 grams of water and the rest anhydrous salt. How could the two students use their data to predict the mass of anhydrous salt in a 32.00 gram sample of the hydrated salt? What would the water mass be in the 32.00 gram sample?

6. A group of 5 students isolated the copper in copper(II) chloride and pooled their data.

They plotted the mass of copper isolated (y axis) versus the mass of copper on a graph. They then used the graph to predict the mass of copper in grams that could be isolated from 1.00 pounds of copper(II) chloride. Below is their data; use it to replicate their graph. What did they figure out?

Mass of copper(II) chloride (g) Mass of copper isolated (g) 1.05 0.49 1.15 0.55 2.36 1.09 2.17 1.03 3.21 1.51

7. Does the data in question 7 support or refute the Law of Definite Composition? Does

the same data support or refute the theory that matter is composed of discrete particles? Does the same data support or refute the idea that matter is continuous in nature?

Unit 3



Why are Mass Ratios of Elements in Compounds of the Same Two Elements in the Form of Small Whole Numbers? 8. Sulfur and fluorine form at least two compounds—SF4 and SF6. Explain how these

compounds can be used as an example of the Law of Multiple Proportions. 9. When 10.0 g of chlorine reacts with mercury under varying conditions, either 28.3 g or

56.6 g of mercury is consumed in the reaction. No other combinations occur. Explain these observations in terms of the Law of Multiple Proportions.

10. Sodium oxide and sodium peroxide are two compounds made up of the elements

sodium and oxygen. Sixty-two grams of sodium oxide contains 46 g of sodium and 16 g of oxygen; 78 g of sodium peroxide has 46 g of sodium and 32 g of oxygen. Show how these figures confirm the Law of Multiple Proportions.

11. Two compounds of mercury and chlorine are mercury(I) chloride and mercury(II)

chloride. The amount of mercury(I) chloride that contains 71 g of chlorine has 402 g of mercury; the amount of mercury(II) chloride that has 71 g of chlorine has 201 g of mercury. Show how the Law of Multiple Proportions is illustrated by these quantities.

12. Sulfur forms two molecular compounds with oxygen. Compound A contains 3.811

grams of sulfur combined with 3.810 grams of oxygen. Compound B contains 7.165 grams of sulfur combined with 10.743 grams of oxygen. What are the two ratios of sulfur combined with oxygen in compound A & compound B? What is the multiple proportion (ratio of ratios) of how sulfur combines with oxygen when comparing compound A to compound B?

Is Matter Made of Combinations of Tiny Indestructible Particles? 13. According to Dalton’s atomic theory, can more than one compound be made from

atoms of the same two elements? 14. List the major points in Dalton’s atomic theory. 15. Show that the Law of Definite Composition is explained by Dalton’s atomic theory. 16. How does Dalton’s atomic theory account for the Law of Conservation of Mass? 17. The chemical name for limestone, a compound of calcium, carbon, and oxygen, is

calcium carbonate. When heated, limestone decomposes into solid calcium oxide and gaseous carbon dioxide. From the names of the products, tell where the atoms of each element may be found after the reaction. How does the atomic theory explain this?

18. The brilliance with which magnesium burns makes it ideal for use in flares and

flashbulbs. Compare the mass of magnesium that burns with the mass of magnesium in the magnesium oxide ash that forms. Explain this in terms of atomic theory.

Unit 3


Unit 3 The Atom

Additional Questions 19. Earlier in this unit you heated a hydrated substance to evaporate the water and separate

it from the anhydrous ionic compound. If you found during an experiment that 190.8 grams of hydrated barium chloride produced 36.0 grams of water vapor on heating, how many grams of hydrated barium chloride would you need to start with to produce 100.0 grams of water vapor?

20. Now that you have completed the unit on the atom, what is an atom and what evidence

do you know that proves matter is made up of discrete particles called atoms?

is there a smallest piece of matter or can we keep cutting

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