This Document Contains Chapters 8 to 12 Chapter 8 Atoms and Periodic Properties Contents Atomic Structure Discovered Discovery of the Electron The Nucleus The Bohr Model The Quantum Concept Atomic Spectra Bohr's Theory Quantum Mechanics The Periodic Table A Closer Look: The Rare Earths Metals, Nonmetals, and Semiconductors Overview This chapter is an introduction to the structure of the atom through a historical perspective. Some of the quantitative aspects of electric and magnetic fields, electromagnetic radiation, and mechanics are used by scientists such as Thomson, Millikan, Rutherford, and Bohr as they identify the parts of an atom and determine how the parts are organized. Thomson's discovery of the electron and Rutherford's discovery of the nucleus are discussed, along with Bohr's development of the planetary model of the atom. How Bohr's model accounted for atomic stability and its ability to produce line spectra is analyzed, along with the weaknesses and shortcomings of the model. The wave nature of particles as suggested by de Broglie is considered next, which leads to the modern quantum mechanics model. Overall, students should be impressed with the symmetry found in nature and understand how scientific models are developed as they learn modern atomic theory. Symmetry is also found by organizing the experimentally discovered facts about atomic structure into the major concept represented by the periodic table. The periodic table demonstrates the periodic recurrence of physical and chemical properties of elements. Elements within chemical families are seen to have similar chemical and physical properties because they all have the same number of electrons in their outermost shells. Across a period, elements are seen to have properties of metals with decreasing chemical activity, then properties of semiconductors, then properties of nonmetals with increasing reactivity, and finally the chemically inert noble gases. Metals are recognized as elements with one, two, and sometimes three electrons in the outermost orbital. Nonmetals have five, six, or seven outer orbital electrons. The inertness of the noble gases is related to the stability of the filled outer orbital arrangement. There are many generalizations that can be made from these periodic patterns, making the periodic table an invaluable tool that simplifies the study of matter. Suggestions 1. The law of definite proportions can be introduced early in this chapter, although the topic is not discussed formally until a later chapter. 2. Illustrate the concepts of Thomson's experiment with Crookes tubes connected to a tesla coil. Demonstrate that the cathode rays travel in straight lines with the Maltese cross, that they have mass and velocity with the kinetic energy effect tube. Demonstrate that cathode rays have a negative charge with the deflection effect tube, and that they can be deflected by a magnet. These demonstrations should lead to a clear understanding of Thomson's work and the properties of electrons. 3. Spectra of different elements can be shown by using several gas emission tubes and squares of inexpensive plastic replica diffraction gratings. Have the students hold the plastic diffraction gratings by the edges (to keep finger oils off the plastic surfaces) and look through the grating at a glowing gas emission tube (H, Ne, Ar). They should rotate the grating and look to the right or to the left of the center until they see the spectrum. Radiation from excited atoms is also shown by the popular demonstration of crushing a wintergreen lifesaver between the jaws of a pair of pliers in a dark room. This emission of light is a result of the excitation of sugar molecules by mechanical stress. 4. Apparatuses for demonstrating the photoelectric effect, Thomson's experiment, and Millikan's experiment are available from scientific supply companies. The Millikan apparatus, however, does not lend itself to large-group demonstrations. It is also difficult to use with small groups. Alpha tracks can be shown in a cloud chamber while discussing Rutherford's experiment. 5. Student understanding of the arrangement of electrons in atoms requires that they mentally act on the process of building electron structures as models. If you choose to cover quantum numbers, it may be helpful to use an analogy to an address, such as state, city, street, and street number. Each electron has its own address (set of quantum numbers) just as each student has a particular address. 6. Illustrate metals and nonmetals with samples of gases in tubes; sheets of copper, zinc, lead, and aluminum; and samples of sulfur, carbon, and iodine. Show mixtures of salt and pepper, sand and dirt, or other common mixtures in varying proportions. 7. An electrolysis of water demonstration is an ideal way to begin discussion of compounds and elements. The Brownlee electrolysis apparatus, available from scientific supply houses, is big enough to be viewed by large classes. 8. Illustrate atomic mass units by using a balance to measure the mass of two coins, such as a dime and a quarter. Calculate the ratio of the mass of one coin relative to the other. Measure the mass of a bag containing equal numbers of both coins, then use the ratio to predict the numbers of each coin in the bag. Define atomic mass units and the abbreviation u. 9. Exhibit a large periodic chart of the elements, referring to the chart often during discussion of chemistry. 10. Illustrate similarities in chemical and physical properties of families and periods of the periodic table. Show the physical properties of Li, Na, and K, then show the chemical reactions with water (Caution: Explosive hydrogen gas), testing the water with litmus. Compare the reactions of Mg, Al, S, C, and Si with dilute hydrochloric acid. Review the main shell structures or electron dot symbols to show similarities among families and periods for the first twenty elements. 11. Do the classic demonstration showing the change of physical properties when a mixture of sulfur and iron is converted into a compound. First, demonstrate that a magnet will attract iron, but not sulfur, and a magnet can be used to separate iron from a mixture of iron and sulfur. Second, demonstrate that sulfur is soluble in carbon disulfide (Caution: Explosive vapors) by dissolving a little in a one-fourth test tube of carbon disulfide. Finally, strongly heat a mixture of 4 grams S and 6 grams Fe filings in a test tube until it glows. Break the test tube by immersion in cold water. Compare the appearance and properties of the mixture before and after heating. Discuss if it is still a mixture or a new compound according to the magnet and carbon disulfide tests. 12. Obtain a roll of calculator tape. Draw 100 squares on a length of the tape and write atomic numbers consecutively in the upper left corner of each square. In the center of each square write the chemical symbol for the atomic number. Form a large spiral by taping together, one under the other, the elements of a family that have the similar chemical properties, such as the inert gases. You can hang the spiral in the classroom to make a mobile, then make another one to show how the periodic table is a spread-out spiral. Cut a spiral to the right of the inert gases, then spread and tape the spiral to make a long form of the periodic table. Show how moving two chemical families results in the short form. For Class Discussions 1. According to the quantum mechanics model of the atom, an electron revolves in an orbit around the nucleus if a. it has no energy. b. it can make a standing wave. c. its momentum matches the electrical attraction from the nucleus. d. its wavelength describes an elliptical orbit. 2. Hydrogen, with its one electron, can produce a line spectrum with four visible colors because a. an isotope of hydrogen has four electrons. b. electrons occur naturally with four different colors. c. there are multiple energy levels that an electron can occupy. d. electrons are easily scattered. 3. A photon is emitted from the electronic structure of an atom when an electron a. jumps from a higher to a lower energy level. b. jumps from a lower to a lower higher level. c. reverses its spin by 180°. d. is removed from an atom by a high quantum of energy. 4. Quantum mechanics uses four quantum numbers to describe certain properties of electrons. These properties do not include a. spin. b. what happens in a magnetic field. c. energy. d. mass. 5. According to the Bohr model of the atom, an electron is able to stay in orbit around an atom because of the a. balance between momentum and electrical attraction from the nucleus. b. balance between spin and orbital velocity, which gives the electron high stability. c. rule that electrons could not move from their allowed orbit. d. fact that moving electrons have mass and velocity, and therefore momentum. 6. The quantum mechanics and Bohr models of the atom both agree on a. the significance of the de Broglie wavelength and the circumference of an orbit. b. the importance of momentum in determining the size of an orbit. c. how electrons are able to emit light. d. none of the above. 7. The atomic weight of neon is 20.2. This means a. all atoms of neon have a mass of 20.2 u relative to carbon 12. b. a given atom of neon has a mass of 20.2 g. c. neon occurs as isotopes of different masses and the weighted average is 20.2 u. 8. An atom has 6 protons, 6 electrons, and 6 neutrons so the mass number is a. b. c. d. 9. How many dots would surround the symbol for each alkali metal atom in the electron dot notation for the group IA elements? a. one. b. two. c. six. d. seven. 10. How many dots would surround the symbol for each halogen atom in the electron dot notation for the group VIIA elements? a. one. b. two. c. six. d. seven. 11. Calcium belongs to which chemical family? a. alkali metals b. alkaline earth metals c. halogens d. noble gases 12. The most important factor in the chemical behavior of an atom is the a. number of neutrons. b. structure of the isotopes. c. ratio of neutrons to protons. d. electron structure Answers: 1b, 2c, 3a, 4d, 5a, 6c, 7c, 8c, 9a, 10d, 11b, 12d. Answers to Questions for Thought 1. The cathode rays behaved like negatively charged particles when passing between plates, one of which was positively charged and one negatively charged. By using a combination of electric and magnetic fields, Thomson determined the ratio of charge to mass of the individual particles. Since this ratio was the same, no matter what the conditions under which he measured the ratio, Thomson was convinced that a fundamental particle had been discovered. 2. Rutherford was scattering positively charged alpha particles off thin sheets of metal. Some of the particles were scattered at very large angles, and some particles were indeed scattered back toward the way they came. This behavior could be accounted for only by theorizing that the massive, positively charged alpha particles were being repelled by a massive positive charge concentrated in the center of the metal atoms. Rutherford concluded that the atoms must have a tiny, massive, positively charged nucleus surrounded by the negatively charged electrons. 3. The “solar system” model requires electrons to travel in orbits about the positively charged nucleus and any object moving in a circle is accelerating. Since accelerating electrons emit energy in the form of radiation, this loss of energy would result in the electrons spiraling in to collide with the nucleus. Since atoms are not observed to exhibit this behavior, this model is not wholly correct. 4. If the size of the nucleus was the same size as the thickness of a dime, the atom would be about two football fields wide. 5. The atomic number is the number of protons in the nucleus of an atom of an element. Since the number of protons in the nucleus determines which element the atom is, the atomic number is directly related to the element. If an atom is not ionized (is electrically neutral), the number of electrons must equal the number of protons or the atomic number. 6. The atomic number is 11. The element with an atomic number of 11 is sodium. 7. Atomic weight is a weighted average of the isotopes based on (1) their mass compared to carbon-12, and (2) their relative abundance in nature. 8. The three main points of the Bohr model are that the electrons can revolve around the atom only in specific allowed orbits, the electrons in these orbits do not radiate energy as long as the electrons remain in the orbit, and an electron gains or loses energy only by moving from one allowed orbit to another. 9. The reference level for the potential energy of an electron is zero when the electron is free and removed from an atom. Work must be done on electrons to remove them from an atom, bringing them back up to a potential energy level of zero, so the electrons in an atom have negative energy values. 10. The energy level of an electron is inversely proportional to the square of the energy level, which means that higher energy levels have lower values for their energy. Thus the lowest energy level is closest to the nucleus and the higher energy level is farthest from the nucleus. Since these values are negative, however, the higher energy levels have higher energy (closer to a potential energy of zero). 11. The discrete energy levels, electrons gaining or losing energy by moving from one allowed energy level to another, and the emission and absorption of photons of energy exactly equal to the energy difference of the allowed energy levels are common to both models. Bohr's model was based on the solar system model, which gives exact radii for the electron orbits. It does not take into account the Heisenberg uncertainty principle as the probabilities contained in the quantum mechanical model do. The Bohr model treats electrons exclusively as particles, whereas the quantum mechanical model accounts for the dual nature of particles. 12. The electron associated with the hydrogen atom in the excited state has more energy than the electron associated with the atom in the ground state. Therefore, the atom in the excited state is said to have more energy than the atom in the ground state. Group B Solutions 1. 2. 3. 4. For Further Analysis 1. Answers will vary in this evaluation, but will show understanding of Millikan’s technique. 2. Similarities – isotopes of a particular element all have the same atomic number. Difference – isotopes of a particular element all have different numbers of neutrons in the nucleus, and thus different mass numbers. 3. This requires analyzing and evaluating an interpretation of Thompson’s experiments. The correct analysis should include a discussion of electrons having mass as shown by response to magnetic field, electrons having a negative change, and the mass to charge ratio being the same for all materials. 4. This requires analyzing, exploring beliefs, and clarifying a conceptual understanding. Answers will vary, but should clarify that atomic weight is a weighted average of the isotopes of an element based on abundance and compared to carbon-12. 5. Answer should include discussion of photons, or quanta of energy that is emitted or absorbed. 6. The Bohr model considered the electron to be a particle, and treated it as such with Newtonian mechanics using one quantum number. The Quantum Mechanic model considered the electron to be a wave, and solved a wave equation for four quantum numbers that describe the energy of the electrons. Chapter 9 Chemical Reactions Contents Compounds Elements Chemical Change Valence Electrons and Ions Chemical Bonds Ionic Bonds Covalent Bonds Composition of Compounds A Closer Look: Name that Compound A Closer Look: How to Write a Chemical Formula Chemical Equations A Closer Look: On Balancing Equations Types of Chemical Reactions Combination Reactions Decomposition Reactions Replacement Reactions Ion Exchange Reactions Overview This chapter covers the concepts of chemical change, including elements and compounds, and chemical bonding (ionic and covalent) and the basic types of chemical reactions. The concepts of naming compounds, how to write formulas equations, and how to write and balance chemical equations is presented as “A Closer Look” features. This means the material is optional, and can be selected to fit the objectives of your class. Each section of this chapter can be presented as is or expanded to include more in-depth subject matter if you so desire. Much information is available to persons who know how to read a chemical equation and how to use the periodic chart. It is helpful to begin by reviewing the valence electrons and ionic charges on the monatomic ions of elements in groups I, II, VI, and VII and then writing the ionic formulas. Students should realize that formulas can be easily written by knowing the ion name, symbol, and charge, which can be easily obtained by looking at the group number in the periodic table. Once this is understood, it becomes a matter of following the rules and understanding the basic generalization and models to write balanced chemical equations. Suggestions 1. Illustrate the physical and chemical properties of powdered iron and sulfur before and after heating a mixture in a crucible with a burner to produce an obvious chemical change. 2. Have students make shell diagrams or electron dot symbols for various elements. Then ask them explain how they could attain the more stable electron structures of the noble gas configurations by loss, gain, or sharing of electrons. 3. To illustrate combustion without the presence of oxygen, heat a piece of steel wool and plunge in into a bottle of chlorine (Caution: Rapid oxidation). 4. Show the polarity of water by deflecting a thin stream from a tap with both a negatively and a positively charged rod. A similar stream of a nonpolar compound, such as mineral oil, is not deflected by either charge. 5. Illustrate as many of the types of chemical reactions as possible with demonstrations, explaining the meaning of the generalized form of the reaction in each case. A strip of magnesium can be burned in air, for example, to illustrate a combination reaction. The generalized form of X + Y XY should be described during the discussion. 6. Decomposition reactions can be illustrated by heating calcium carbonate or soda water, then testing for the production of carbon dioxide with limewater. 7. A variety of metal and salt solution replacement reactions can be demonstrated to illustrate replacement reactions, and the results can be used to establish a relative activity series of metals. Include combinations of metals and salt solutions that result in “no reaction” to help establish the activity series. 8. Ion exchange reactions can be illustrated by looking for the formation of a precipitate, a molecular substance, or a gas. Demonstrate several of these reactions before explaining what happened by way of equations. 9. Depending on the science experience and mathematical background of the students, it may be appropriate to illustrate the quantitative interpretations of the mole relationships in a chemical equation. 10. Additional demonstrations: (a) Do the spectacular demonstration of an ammonium dichromate volcano. Place an approximate 5 cm pile of ammonium dichromate on a metal tray. Ignite with a torch or burner flame (or short piece of burning magnesium ribbon). A large volume of green chromic oxide “cinders” will form a conical pile. The burning reaction releases ammonia and oxygen, and the remaining chromic oxide has a weight of about 40 percent less than the original ammonium dichromate. (b) A 110-volt conductivity apparatus with electrodes about 1 cm apart can be used to show that molten NaCl conducts an electric current while solid NaCl does not. Attach the apparatus to a ring stand so the electrodes are in a 2 cm depth of granulated NaCl in a small porcelain crucible. The crucible is in a triangle on a tripod. Show the lamp does not light when the switch is closed. Now melt the NaCl with a Fisher burner under the crucible (requires 3 – 4 minutes), and again close the switch on the conductivity apparatus. Allow the melt to cool and solidify until the lamp goes out. Explain that melted NaCl has mobile ions that can conduct a current. As the rigid lattice structure is established by ionic bonding, the ions are no longer free to move and conduct the current. (c) Burn a strip of Mg (Caution: Use tongs. Do not look directly at burning.) as an example of a combination reaction. (d) As an example of a decomposition reaction, mix concentrated H2SO4 in a small (50 mL or so) beaker half filled with table sugar. The sugar is decomposed to a growing “tower” of carbon. (Caution: Noxious fumes. Use fume hood or perform outside.) (e) As an example for a replacement reaction, place an iron file in copper(II) sulfate solution or Cu wire in silver nitrate solution. (f) As an example of an ion exchange reaction, mix a sodium chloride solution with a solution of silver nitrate. Check for insolubility of products: NaCl + AgNO3 NaNO3 + AgCl. For Class Discussions 1. When involved in a chemical reaction with metals, atoms of a nonmetal tend to a. exchange electrons. b. remain neutral. c. gain electrons. d. lose electrons. 2. Potassium is a metal and oxygen is a nonmetal that react to form K2O, an ionic compound. How many electrons were lost by an atom during the reaction? Each a. potassium atom lost one electron. b. potassium atom lost two electrons. c. oxygen atom lost one electron. d. oxygen atom lost two electrons. 3. A crystal of potassium chloride is made up of a. K and Cl atoms. b. KCl molecules. c. K+ and Cl– ions d. K– and Cl+ ions 4. In a covalent molecule you would find atoms that a. have lost electrons to become ions. b. have gained electrons to become ions. c. are sharing at least one pair of electrons. d. are sharing at least one electron. 5. The coefficients needed to balance the equation Na3PO4 + AgNO3 NaNO3 + Ag3PO4 are a. 1, 1, 1, 1 b. 3, 1, 1, 3 c. 2, 3, 3, 2 d. 1, 3, 3, 1 6. When hydrocarbons and carbohydrates burn with sufficient O2 they a. always give off CO2 and H2O. b. sometimes give off CO2, but never H2O. c. sometimes give off H2O, but never CO2. d. never give off CO2 or H2O. 7. How many atoms are in a molecule of chalk, which is CaCO3? a. 3 b. 5 c. 6 8. How many atoms of oxygen are in a molecule of Al2(CO3)3? a. 3 b. 6 c. 9 d. 14 9. If Iron(II) chromate has a formula of FeCrO4, then potassium chromate must have a formula of a. KCrO4 b. K2CrO4 c. K2(CrO4)2 d. K(CrO4)2 10. Should dihydrogen monoxide, which is responsible for substantial damage to the environment each year, be banned? a. Yes. b. No. Answers: 1a, 2a, 3c, 4c, 5d, 6a, 7b, 8c, 9b, 10b (systematic name for water). Answers to Questions for Thought 1. A chemical change involves a change in the identity and properties of the substance, while a physical change is merely a size, shape, volume, or phase change. Three examples of a chemical change are wood burning, burning gasoline in a car engine, and iron rusting. Three examples of a physical change are ice melting, water boiling, and water condensing on the side of a water glass. 2. (a) A sodium atom and a sodium ion both have the same atomic number, and if the two are the same isotope, they have the same mass number. The atom and the ion will have different numbers of electrons, and consequently, different charges. (b) A sodium ion and a neon atom will have the same number of electrons, but they have different atomic numbers and different net charges. 3. In an ionic bond, one or more electrons are pulled from one atom to fill the outer shell of another atom. The resulting ions that are created are attracted to each other electrostatically. In a covalent bond, electrons are shared between the atoms, and the attraction of each nucleus for the other's electron binds the atoms together. The atoms involved in these two bond types both follow the octet rule. This is done either by sharing electrons to have eight in the outer shell as in a covalent bond or by losing or gaining electrons as in the ionic bond. 4. The octet rule states that atoms attempt to acquire an outer orbital with eight electrons. This rule is a generalization, and not all elements follow this rule. 5. Polyatomic ions are groups of two or more elements bound tightly together which behave much like a single monatomic ion. The hydroxide ion (formula OH) with a net charge of 1 and the cyanide ion (formula CN) with a net charge of 1 are examples. 6. The formula of magnesium hydroxide is Mg(OH)2. The parentheses indicate that the OH pair is a polyatomic ion that occurs twice in this compound. 7. A double bond is a covalent bond in which two pairs of electrons are shared. Similarly, in a triple bond three pairs of electrons are shared in the bond. 8. Conservation of mass states that matter is neither created nor destroyed in a chemical reaction. A chemical equation conforms to this law only if each element appears the same number of times on each side of the equation. 9. Beginning with the unbalanced equation, the number of each kind of atom on each side of the equation must be counted. These numbers are used to determine whether to place coefficients in front of the formulas and what the coefficients should be. The equation should be checked to make sure that the number of each type of atom on each side is equal and that the coefficients are in the lowest whole number ratio. 10. An element higher on the table will replace an element lower on the table in an ionic compound. If the element in the ionic compound is higher on the table than the other element proposed to replace it, no reaction will occur. 11. The formation of a precipitate, a gas, or water must occur for an ion exchange reaction to take place. If this does not occur, no ions have been removed, so no reaction has taken place. 12. (a) The reaction for the combustion of ethyl alcohol is C2H5OH + 3 O2 2 CO2 + 3 H2O. (b) The reaction for the rusting of aluminum is 4 Al + 3 O2 2 Al2O3. Group B Solutions 1. 2. 3. 4. 5. 6. 7. 8. (a) 2 NO + O2 2 NO2 (b) 2 KClO3 2 KCl + 3 O2 (c) 2 NH4Cl + Ca(OH)2 CaCl2 + 2 NH3 + 2 H2O (d) 2 NaNO3 + H2SO4 Na2SO4 + 2 HNO3 (e) PbS + 4 H2O2 PbSO4 + 4 H2O (f) Al2(SO4)3 + 3 BaCl2 2 AlCl3 + 3 BaSO4 9. 10. (a) 2 C3H6(g) + 9 O2(g) 6 CO2(g) + 6 H2O(g) (b) H2SO4(aq) + 2 KOH(aq) K2SO4(aq) + 2 H2O(l) (c) C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) (d) Na3PO4(aq) + 3 AgNO3(aq) 3 NaNO3(aq) + Ag3PO4(aq) (e) 3 NaOH(aq) + Al(NO3)3(aq) 3 NaNO3(aq) + Al(OH)3(aq) (f) 3 Mg(OH)2(aq) + 2 H3PO4(aq) Mg3(PO4)2(aq) + 6 H2O(l) For Further Analysis 1. Similarities – both are recognizable changes in the physical state or chemical composition of matter. Difference – chemical change alters the identity of matter, producing a new substance with different properties; physical change does not alter the identity of matter. 2. You would know for sure that you have a compound and not an element if it is a pure substance that can be broken down into simpler substances by chemical change. 3. If a substance cannot be decomposed it might be an element or it might be a very stable compound. Extended testing might be necessary before you are sure that the substance is an element. 4. Metal atoms from the left side of the periodic table have a few electrons in outer orbitals that can be transferred to nonmetal atoms that are lacking a few electrons to fill the outer orbitals. This sets up ionic bonding reactions as the octet rule is satisfied. Nonmetallic elements, on the other hand, can satisfy the octet rule by sharing electrons through covalent bonds. 5. Word equations describe a chemical change in general, but say nothing about the quantities of the chemicals involved in the change. The advantage of writing a chemical equation with chemical symbols is that more information is provided about the composition of chemicals as well as the quantities involved. The disadvantage to this is that it requires knowledge and understanding of formulas as well as chemical symbols. 6. Answers will vary, but will follow the generalized form of each reaction. 7. (1) write the correct formulas for the reactants and products in a chemical equation. (2) Inventory the number of each kind of atom on each side of the equation. (3) Place coefficients in front of formulas to balance the equation. (4) Take another inventory to determine if the numbers of atoms on both sides are equal. Chapter 10 Water and Solutions Contents Household Water Properties of Water Structure of the Water Molecule The Dissolving Process A Closer Look: Decompression Sickness Solubility Properties of Water Solutions Electrolytes Boiling Point Freezing Point Acids, Bases, and Salts Properties of Acids and Bases Explaining Acid-Base Properties Strong and Weak Acids and Bases The pH Scale Properties of Salts Hard and Soft Water A Closer Look: Acid Rain Overview This chapter provides opportunities for the application of many fundamental chemical principles that were introduced in previous chapters. These principles are centered on solutions —what they are, how they are measured, and how they are used in many everyday activities. This includes the concepts of solubility and insolubility, the physical properties of water, the dissolving process, the properties of acids, bases, and salts, the pH concept, and soft and hard water. The chapter materials can be presented at practically any depth to match the interest and abilities of the class. Student interest in the materials of this chapter is generally high because of the many apparent applications to and explanations of everyday activities. Suggestions 1. A variety of everyday solutions that relate concentrations can be displayed, such as bleaches, vinegar, hydrogen peroxide, beverages, and so forth. Recent news articles that report concentrations of minerals or pollutants in water or in the atmosphere can be analyzed and compared to other concentrations to bring relevancy to the lesson. 2. The concept of a solution is illustrated by mixing 100 mL of absolute alcohol with 100 mL of water. The resulting solution has a volume less than 200 mL, which raises the question of what happened to the missing solution. Use 100 mL of salt mixed with 100 mL of B-Bs for a model to show how smaller particles can get between the spaces of the larger particles. This makes less total volume (a bushel of peas mixed with a bushel of potatoes does not make two bushels of peas and potatoes). Relate the model to the alcohol and water mixture. 3. Stir small amounts of a soluble salt, such as potassium nitrate, in water in ever increasing amounts. Have the class identify when the solution is dilute, when it is concentrated, and when it is saturated. Students soon understand the vagueness of the terms “dilute” and “concentrated,” opening the way for a discussion of different ways to express concentration measurements. 4. Demonstrate how much the boiling point is elevated by adding a pinch of salt to a liter of water. Then relate the findings to the habit of some cooks adding a pinch of salt to water when cooking. See how much salt you must add to cause a meaningful change in the boiling point. Relate findings to text materials. 5. Test the electrical conductivity of various compounds and solutions with a flashlight bulb and battery. Complete the circuit by plunging the two ends of the exposed wires into pure acetic acid placed in a beaker. Then add water to produce ions and test again. Try the same arrangement with powdered sodium hydroxide, then add water and repeat. 6. If a pH meter is available, add equal amounts of hydrochloric acid and acetic acid to two beakers containing 100 mL of water. To give a color variation add several drops of universal indicator to each beaker. With the pH meter, check the pH of bleach and of vinegar. Pour the two substances together very slowly and check the pH to show neutrality. The amounts to produce neutrality should be determined in advance. Demonstrate the pH of distilled water and tap water. 7. Additional demonstrations: Set up the apparatus as illustrated in figure 10.1 for the classic ammonia fountain demonstration. The long glass tube has a tapered end that reaches about two-thirds of the way into a one liter Florence flask. The medicine dropper contains a small amount of water. The beaker contains about 2-L water with about 6 drops of phenolphthalein. Charge the fountain by filling the Florence flask with relatively dry NH3. In a hood, use a gas generator with concentrated NH4OH dropped into solid NaOH pellets. Fill the Florence flask with the dry NH3 by downward displacement of air. Figure 10.1 (Caution: Use a fume hood. Do not inhale ammonia.) Fit the stopper assembly into the flask, then invert over the beaker with phenolphthalein solution. Start the fountain by squeezing the water from the medicine dropper. For Class Discussions 1. Substances that are ionic crystals are made up of a. isotopes. b. ions. c. atoms. d. molecules. 2. The intermolecular force of attraction known as a hydrogen bond comes from a. uneven distribution of charges around a molecule. b. electrons being transferred from one molecule to another. c. the symmetry of even charge distribution in a molecule. d. electrons being shared and reshared in a symmetrical molecule. 3. A solution with a pH of 7 is a. slightly acidic. b. slightly basic. c. neutral. 4. Water is able to dissolve certain ionic compounds because water is made up of a. separated positive and negative ions. b. symmetrical covalent molecules. c. strongly polar molecules. d. slightly acidic molecules. 5. Suppose it is late during a very cold winter and there is a very deep lake that has a thick layer of ice on top. The temperature of the water at the bottom of this lake should be a. 0°C b. 4°C c. 10°C d. unknown 6. Adding a solute, such as salt, will have what effect on the freezing point of water? It a. is lowered. b. is raised. c. remains the same. 7. A strong acid is strong because it is a. completely ionized in water. b. free of dissolved materials. c. mostly insoluble in water. d. made up of nonpolar molecules. 8. Will a solute that raises the boiling point of water also raise the freezing point? a. Yes. b. No. Answers: 1b, 2a, 3c, 4c, 5b, 6a, 7a, 8b. Answers to Questions for Thought 1. A solution is a homogeneous mixture of ions or molecules of two or more substances. Not all mixtures are homogeneous and not all are mixtures of ions or molecules. 2. If the attraction between the ions is greater than the hydration energy the ionic compound is insoluble in water. Ionic compounds whose attraction to the polarized water molecules is greater than those between ions are soluble. 3. The presence of the dissolved salt decreases the number of water molecules at the surface. This occurs since the sodium ions and chlorine ions are taking up some of the space near the surface. Since there are fewer water molecules near the surface to escape, the vapor pressure is lowered. This increases the temperature necessary to obtain the vapor pressure required to boil the water and salt solution. 4. The greatest density of water occurs at a temperature of 4˚ C. A lake covered with ice (0˚ C) would probably have the greatest density water—with a temperature of 4˚ C—at the bottom of the lake. 5. Water at very close to freezing has hydrogen bonds present between many of the molecules. These bonds spread out the water molecules. As the temperature of the water increases above the freezing point (from 0˚ to 4˚ C), the bonds break and allow the molecules to collapse together. 6. A solution of Ca2+ or Mg2+ ions in water is called hard water. Removing the calcium and magnesium ions softens the water. 7. NaCl does not donate protons to form hydronium ions, and it does not accept protons to form hydroxide ions when it is dissolved in water. Since the pH scale is a measure of the hydronium ion concentration, the pH does not change because the hydronium ion concentration does not change. 8. A hydrogen bond is a weak to moderate strength bond between the hydrogen end (+) of a polar molecule and the negative end () of a second polar molecule. The bond forms when the positively charged end of a polar molecule, which is usually a hydrogen atom, is attracted to the negative end of another polar molecule. This negative end is usually an oxygen, fluorine, or nitrogen atom. If the two ends come close enough together, this electrical force of attraction forms a bond. 9. A soap molecule has two parts, a part that is polar, which allows the soap to be soluble in water, and a nonpolar part, which is absorbed into oil. Soap molecules attach onto oil molecules and carry the oil away with rinse water. 10. (a) Acids are associated with the presence of hydronium ions in a water solution. (b) Bases are associated with the presence of hydroxide ions in a water solution. For Further Analysis 1. Possible explanations should involve a discussion of water as a polar molecule and hydrogen bonding. 2. A salt that will dissolve does so because the attraction between the polar water molecules and a charged ion is greater than the attraction between the ions. A salt that will not dissolve has greater attractions between the ions than between the polar water molecules and the ions. 3. The boiling point is defined as occurring in pure water at sea level pressure. 4. Similarities – Acids and bases are chemical opposites dissolved in water and a salt is produced by a neutralization reaction between and acid and a base. Differences – An acid is a proton donor in a water solution; a base is a proton acceptor in a water solution; a salt is an ionic compound produced when acids and bases react. 5. Answers will vary, but should include the basic idea that pH is a power of ten notation that expresses the hydronium ion concentration. The scale ranges from 1 (most acidic) to 14 (most basic), with 7.0 as the pH of a neutral solution. 6. Answers will vary, but should include the basic idea that a solution of Ca2+ or Mg2+ ions in water is called hard water, and removing these ions will soften the water. Chapter 11 Nuclear Reactions Contents Natural Radioactivity Nuclear Equations The Nature of the Nucleus Types of Radioactive Decay Radioactive Decay Series A Closer Look: How is Half-Life Determined? Measurement of Radiation Measurement Methods Radiation Units Radiation Exposure A Closer Look: Nuclear Medicine Nuclear Energy Nuclear Fission Nuclear Power Plants A Closer Look: Three Mile Island and Chernobyl Nuclear Fusion A Closer Look: Nuclear Waste Overview Radioactivity is feared by the general public because the general public does not understand radioactivity. It seems to be something invisible that can move through the air and kill. Just as individuals are afraid of the dark because they do not know what is there, radioactivity is feared because of a lack of knowledge. This does not mean that radioactivity is safe but rather that erroneous and distorted information should not be used as a basis for decisions about nuclear energy. The purpose of this chapter is to develop a more accurate understanding of radioactivity and the issues involved in its uses. As often happens in science, many great discoveries in this area were the result of seemingly unimportant incidents. The accidental exposure of photographic film to a uranium ore led to Becquerel’s discovery of radioactivity. His continued research, along with that of Rutherford and the Curies, opened a new frontier for scientists of all specialties to explore. The nature of the nucleus, the types of radioactive decay, the resulting radioactivity emitted, and the methods of measuring radiation set the stage for the discussion of the uses of nuclear energy. Few students know the identity of alpha and beta particles nor do they understand their origins. The similarities of gamma rays, X rays, radio waves, light, and so forth should be discussed. Only energy and origin separate one from the other since the electromagnetic spectrum is continuous. The binding energy and mass defect are meaningful applications of the relationship of matter to energy, E = mc2. It is generally a surprise for students to learn that it is a physical impossibility for a nuclear power plant to explode as a nuclear bomb. Once this is understood, it is natural then to discuss the consequences of exposure to radioactivity and the controversy surrounding this topic. Finally, the concerns about nuclear wastes are explored. Suggestions 1. The scientific area encompassing radioactivity and its properties is one in which most students are almost completely lacking in knowledge. Their impressions have developed from misconceptions, so as the instructor you should develop their understanding carefully and thoroughly, without bias one way or the other. 2. Hospitals may release to you various X ray photographs that might prove of interest to students. The negatives can be projected with an overhead projector. 3. Most states have a department of disaster or emergency services that have a supply of portable GM counters and other survey meters. You may find that these units are already available within the school, and they are excellent for demonstrations. Check with your radiation safety officer or health physicist. 4. Alpha particles can be detected with certain GM counters, for example, Pancake GM Tube Survey Meter Model #3 from Ludlum Measurements, Inc., Sweetwater TX. Other tubes are also available that can detect the alpha particle with some efficiency. 5. A portable GM counter, along with beta and gamma sources, is most useful in illustrating the properties of the two particles, including their activity, energy, and charge. The counter can also be used to check rocks, watches, pottery, dinnerware, and other substances for radioactivity. The lantern mantle used in unleaded gasoline and propane lanterns contains thorium oxide and provides an excellent source of radioactivity for demonstrations and laboratory measurements. 6. Additional demonstrations: (a) A cloud chamber can be used to show the tracks produced by emissions from radioactive materials. (b) A Geiger-Muller tube and counter can be used to show background radiation and test everyday materials for radioactivity. Test old, orange-colored pottery since uranium salts were formerly used as a color pigment. (c) Place a metal key on high speed photographic film, then cover the key and film with cutup lantern mantles. Develop the film in about a week or so. The black area around the white key image is a result of exposure to radiation. The lantern mantle contains radioactive thorium oxide. For Class Discussions 1. The property of radioactivity in materials is caused by a. the pull of electrons on and the repulsion of protons in the nucleus. b. instability from struggles between electromagnetic and nuclear forces. c. nuclear instability from an accumulation of isotopes. d. nuclear forces, which squeeze particles out of the nucleus. 2. Radiation from a radioactive material can be any of the following except a. a helium nucleus. b. an electron. c. a proton. d. a gamma ray. 3. A beta particle, sometimes called a beta ray, is a. an electron from the nucleus. b. a photon from the nucleus. c. a proton from the nucleus. d. a photon from the second energy level. 4. Which of the following is most likely to happen to an alpha particle after it is emitted? a. It escapes to space, where it is known as a cosmic ray. b. After being absorbed by matter it is dissipated as radiant energy. c. After joining with subatomic particles it finds itself in a child's balloon. d. It joins others of its kind to make lightning and thunder. 5. Which of the following is most likely to happen to a beta particle after it is emitted? a. It escapes to space, where it is known as a cosmic ray. b. After being absorbed by matter it is dissipated as radiant energy. c. After joining with subatomic particles it finds itself in a child's balloon. d. It joins others of its kind to make lightning and thunder. 6. A nucleus emits a beta particle so the number of nucleons it contains is now a. more. b. less. c. the same. 7. A nucleus emits an alpha particle so the number of nucleons it contains is now a. more. b. less. c. the same. 8. The mass of an atomic nucleus compared to the mass of the individual nucleons making up the nucleus is always a. less b. more c. the same. 9. Enriched uranium has more a. energy. b. radioactivity. c. deuterium d. uranium-235. 10. The maximum binding energy per nucleon occurs in nuclei of a. low mass number. b. high mass number. c. intermediate mass number Answers: 1b, 2c, 3a, 4c, 5d, 6c, 7b, 8a, 9d, 10c. Answers to Questions for Thought 1. Radioactive materials emit invisible radiation. A radioactive material contains atoms that spontaneously emit particles or energy from their nuclei. This often results in the change of identity of these atoms, so radioactive materials often decompose into different materials. Nonradioactive materials exhibit none of this behavior. 2. Radioactive decay is the natural spontaneous disintegration or decomposition of a nucleus. Changing the decay rate is not possible. 3. The three kinds of radiation emitted by radioactive materials are as follows: alpha particles, which are helium nuclei; beta particles, which are electrons ejected at high speed; and gamma radiation, which is a high-energy burst of electromagnetic radiation. Alpha particles eventually acquire electrons and become ordinary helium atoms. Beta particles either remain free electrons or join an ion to become part of an atom. Gamma radiation is absorbed by materials, giving its energy to the material absorbing it. 4. The force known as the strong nuclear force binds protons and neutrons together in a nucleus. This force, over very short distances, is much stronger than the electromagnetic force and overcomes it. 5. A half-life is the time it takes half of the unstable nuclei in a radioactive substance to undergo radioactive decay. Iodine 131 has a half-life of eight days; after forty days, five half-lives have elapsed, and only 1/32 of the original material remains. 6. Isotopes with a longer half-life are less active than isotopes with shorter half-lives because the activity rate determines the half-life of a material. High activity gives short half-lives and more radioactivity. 7. Background radiation is radiation from natural sources that are part of the environment. The normal dose of background radiation is between 100 and 500 millirem per year. 8. Lack of knowledge, as well as the difficulty in obtaining reliable results not due to extraneous factors, fuels the controversy over long-term, low-level radiation exposure. Since cancer can arise from many factors, including a combination of seemingly unrelated actors, it is difficult to say that low-level radiation is responsible for any increased incidence of cancer. 9. The difference between the actual mass of a nucleus and the sum of the masses of individual nucleons making up the nucleus is the mass defect of the nucleus. This difference in mass has been converted to energy, which was released as the more stable nucleus was formed. This energy is equivalent to the binding energy of the nucleus, and it can be calculated by using the mass defect as the mass in Einstein's equation E = mc2. 10. Nuclear fusion is the process in which two nuclei with low mass numbers collide at very high speeds and “stick together” to form a heavier, more stable nucleus. Nuclear fission is the nuclear reaction of splitting a massive nucleus into more stable, less massive nuclei. Both processes release energy because in both cases more stable nuclei are the result. More stable nuclei have larger binding energies, so energy must be released by less stable nuclei to gain stability. Group B Solutions
1. (a) Al: 13 protons; 12 neutrons (c) Sn: 50 protons; 70 neutrons
(b) Tc: 43 protons; 52 neutrons (d) Hg: 80 protons; 120 neutrons
2. (a) (c)
(b) (d)
3. (a) Unstable, because isotopes with an atomic number less than 83 are stable when the ratio of protons to neutrons in the nucleus is 1:1, which is not the case here. (b) There are odd numbers of both protons and neutrons, and the ratio of protons to neutrons is not 1:1, therefore it is unstable. (c) Stable, because there are even numbers of protons and neutrons. In addition, 50 is a particularly stable number of nucleons. (d) Stable, because there are even numbers of protons and neutrons, allowing them to pair up. 4. (a) (b) (c) (d) (e) (f) 5. (a) (b) (c) (d) (e) (f) 6. 7. First, the mass difference needs to be calculated: before the reaction: 7.01435 u after the reaction: 3(1.00728 u) + 4(1.00867 u) = 7.05652 u mass difference: 7.05652 u - 7.01435 u = 0.04217 u For Further Analysis 1. A radioactive isotope has an unstable nucleus that will disintegrate into a simpler nucleus, giving off alpha, beta, or gamma radiation. An isotope that is not radioactive has a stable nucleus that will not disintegrate. 2. All types of radioactive decay are ionizing radiation, and this ionization can produce free polyatomic ions that can interfere with cell functions or disrupt chemical bonds in essential organic molecules such as DNA. 3. Answers will vary, but address that some nuclear arrangements are less stable than others are. 4. Answers will vary, but should illustrate the basic differences between the two models. 5. Figure 11.10 shows that the maximum binding energy per nucleon occurs around mass number 56, and then decreases in both directions. As a result, fission of massive nuclei and fusion of less massive nuclei both release energy. 6. Answers will vary, but should point out that waste fuel rods contain an appreciable amount of usable uranium and plutonium. 7. Similarities – both supply energy through nuclear reactions. Differences – a fission power plant slowly provides energy through fission processes; a fusion power plant provides much energy through fusion processes; a fission reaction can be contained in a nuclear reactor, but a fusion reaction requires confining a hot plasma in a magnetic field. Chapter 12 The Universe Contents The Night Sky Origin of Stars Brightness of Stars Star Temperature Star Types The Life of a Star Galaxies A Closer Look: Extraterrestrials? The Milky Way Galaxy Other Galaxies The Life of a Galaxy A Closer Look: Redshift and Hubble’s Law A Closer Look: Dark Matter Overview It is usually news to some students that our Sun is a star much like the other stars in the universe. You should point out that we know as much about stars as we do because we have a star, our Sun, relatively nearby to study. We are able to make photographs and measurements and obtain instrument readings of the Sun by telescopic observations and flyby space probes. The position of our star in the Hertzsprung-Russell diagram will help students realize that our Sun is but an average star in terms of size, surface temperature, and luminosity. The section on the life of a star can be made more personal to students by discussing the birth, evolution, and death of our Sun and the consequences of these events on Earth. Our galaxy, with its hundreds of billions of stars, also contains many immense patches of gas and dust particles known as nebulas. All nebulas are essentially alike, but bright nebulas are bright because they are close to hot, bright stars, so the nebulas shine by reflections and fluorescence. Dark nebulas are so far from hot, bright stars that they appear dark. Most nebulas have been formed from the debris of exploding stars. New stars are still forming in the regions of many nebulas. This chapter also discusses our galaxy, the Milky Way. Invariably, a few students will confuse our galaxy with our solar system; consequently, you should point out the difference each time either is discussed. Suggestions 1. Compare life among the ancients with our lives today. Why might they have noticed more about the night skies than we do? 2. Point out the ancients’ needs for accurate information about the positions of the stars, the Moon, and the Sun compared to our needs for such information. 3. You might have available crude models of the simple instruments the ancients used to make accurate measurements of the positions of the Sun, the stars, and other astronomical objects. To illustrate stellar parallax, draw a spot on the board. Have the students close one eye, hold a finger of an extended arm on the spot, and move their heads in a circular motion to simulate the revolution of Earth about the Sun. 4. Have an art major illustrate his or her interpretation of the birth of stars. Others may give their renditions of the various star classifications. 5. Pulsars can be related to the flashing lights of lighthouses, ambulances, or police vehicles. Recordings of sounds made by various pulsars can be obtained from radio- telescope observations. 6. Plastic balls of various sizes and colors can be prepared to represent the different classifications of stars. The sequence involved in the life of a star can be illustrated very vividly with the models. 7. Have students observe the Milky Way if you are in a location where it can be seen at night. Point out that they will be looking along the long axis of our galaxy. 8. Have students determine the length of a light year and then the dimensions of our galaxy in miles and kilometers. 9. The big bang theory of the expanding universe is well illustrated by a balloon with spots painted on it. An ant on the balloon's surface would see all spots recede as the balloon expands. For Class Discussions 1. The radius of Earth's orbit around the Sun is a. a pair of seconds. b. a light year. c. an astronomical unit. d. an arc second. 2. More massive stars generate a. higher temperatures. b. lower temperatures. c. average temperatures. d. temperatures that are neither higher nor lower than less massive stars. 3. More massive stars on the main sequence tend to be a. brighter and bluish. b. white, with an average brightness. c. yellow, with an average brightness. d. fainter, with an reddish color. 4. Where a star is on the main sequence and what happens to it next depends on its a. composition. b. radius. c. temperature. d. mass. 5. Comparing the following list of star types, the oldest star is probably a(n) a. main sequence star. b. red giant. c. white dwarf. 6. Only a more massive main sequence star will have an opportunity to become a(n) a. white dwarf. b. supernova. c. red giant. 7. The star known as the Cepheid variable is important because they a. are very bright and easily identified. b. can be used as cosmic distance mileposts. c. help determine where a main sequence star is in the H-R diagram. d. serve as lighthouses, warning of cosmic hazards. 8. Our Milky Wag galaxy a. is about 33,000 light years across. b. has a radius of approximately 1 or 2 light years. c. contains about 95% of all the know stars in the universe. d. has a diameter of about 100,000 light years 9. According to the accepted model of the life cycle of a star, our Sun will next become a a. red giant b. white dwarf c. supernova d. black hole 10. A black hole might form when a. a white dwarf cools, no longer emitting light. b. the core of a super-massive star collapses. c. a neutron star begins to spin, creating an extremely powerful magnetic field. d. a supernova implodes. Answers: 1c, 2a, 3a, 4d, 5c, 6b, 7b, 8d, 9a, 10b. Answers to Questions for Thought 1. A light year is a unit of distance defined as the distance light travels during one year. This distance is about 9.5 1012 km. 2. Astronomical distances are so large that measuring them in units of miles or kilometers would result in huge numbers. In addition, the large distances cause the standard units of length to have little meaning because there are no referent points of comparison. So light-years are used to measure distances in terms of time, and parsecs measure distances in terms of angles. 3. As the gas forming the protostar falls together, the gravitational potential energy of the gas is decreased. This energy is converted first to kinetic energy of the gas molecules. This kinetic energy is then converted to heat when the gas molecules collide as the gases grow denser. 4. Located at the center of a star and extending about one quarter of the distance to the surface of the star is the core. This region contains about half the mass of the star at a density of about twelve times that of solid lead and at a pressure of about 300 billion atmospheres. The temperature is about 15 million degrees Celsius. Above the core extending to about nine-tenths the way to the surface is the radiation zone. It has about the density of water and absorbs and re-radiates energy from the core. The convection zone is in the final tenth of the way to the star’s surface and is not very dense—only about 1 percent the density of water. Gases at the boundary between this zone and the radiation zone are heated, rise to the surface, and radiate away the energy that was originally released in the core of the star. The cooler gases then contract and fall to begin the process again. The temperature at the surface is about 5500°C. 5. The more massive a star, the higher the pressure at the core, and the faster the fusion reactions take place. Therefore, in general, even though more massive stars have more fuel, they use it at such a rapid rate that the massive stars “burn out” more quickly than less massive stars. A star that has a mass of one twenty-fifth that of the Sun would last longer than both the Sun and a star sixty times as massive as the Sun. 6. Apparent magnitude is a measure of how bright a star appears in the sky. Since different stars of the same size might be at different distances from Earth, they might have different apparent magnitudes. To be able to compare the brightness of stars to one another directly, the absolute magnitude scale is used. This scale gives the brightness of a star if it were located 10 parsecs away from Earth. 7. The color and temperature of an incandescent object have long been understood to be closely related. This relationship is due to the electrons surrounding an atom being excited by higher temperatures to higher energy levels. When the electrons return to the ground state energy level, they emit light. The more energy that the electron has to lose to return to the ground state, the higher energy light it has to emit. Higher energy light is shorter frequency light, or “bluer” light. So hotter objects glow closer to the blue end of the spectrum and cooler objects emit light that is closer to the red end of the spectrum. Comparing the intensities of blue light and red light emitted by a star gives a good measure of the temperature. The spectral classification scheme is based upon the temperature of the star. Originally it was based upon the intensities of the hydrogen lines in the star’s spectra. The current system has a set of temperature ranges assigned to each classification. Type O is 30,000 to 80,000 K and is bluish. These stars are short-lived. Type B is 10,000 to 30,000 K and again is bluish. The range of 7,500 to 10,000 K is type A, which is still bluish. The white type F has a range of 6,000 to 7,500 K, and the yellow stars like our type G Sun has a temperature of 5,000 to 6,000 K. The relatively cool type K stars are 3,500 to 5,000 K and are orange-red in color. The type M stars have a temperature of 2,000 to 3,500 K and are reddish. 8. The Hertzsprung-Russell diagram is a plot of the temperatures of stars indicated by spectral type versus the brightness of the stars indicated by absolute magnitude. The diagram shows that most stars fall close to a narrow diagonal band called the main sequence. The position of a star on this diagram suggests the star’s mass and age. 9. The main sequence on the H-R diagram is the narrow diagonal band that runs from the bright blue stars down to the dimmer red stars. Most stars fall on or near this diagonal band. The mass of a main sequence star is the prime factor in determining if a star is hot and blue, cooler and yellow, or cold and red. 10. A star begins as a large cloud of gas that is gradually attracted together by gravity into a protostar. The protostar collapses down until the temperature and pressure at the center become high enough to sustain nuclear fusion. This fusion heats the core even more and the pressure at the core increases. Eventually equilibrium between the outward force of the pressure and the inward force of gravity is established after about fifty million years. The star is now a part of the main sequence and burns its fuel steadily. Eventually, the hydrogen in the center of the star begins to run out. The fusion reaction slows down and the equilibrium between the pressure and gravity is upset. The star again begins to collapse until the hydrogen remaining outside the core begins to fuse. This causes the outer layers of the star to expand, and the star becomes a red giant. The star is no longer on the main sequence because its properties have changed. The helium core is heated by the fusing hydrogen surrounding it and eventually, after its temperature has risen enough, begins to fuse as well. The star moves back toward the main sequence when this stage is reached. After millions of years of helium fusion in the core and hydrogen fusion in the outer layers, the star heats up and again expands to a red giant. After this expansion, the center may cool enough so that the atoms in the core become neutral rather than part of an ionized plasma. These neutral atoms absorb heat from the surrounding star and the star expands and cools. As it cools it begins to collapse back, heating the star again. This occurs several times and eventually blows the outer layers of the star away, forming a planetary nebula. The carbon core and helium fusion layer surrounding it gravitationally collapses to form a white dwarf. 11. A violent flare up of a star for a short time is called a nova. Many novas occur because they are part of a binary star system. An old white dwarf star can pull matter off a younger companion star until enough matter builds up on the surface of the white dwarf for a fusion explosion to ignite. A supernova is the rebound explosion from the catastrophic collapse of a massive star that has run out of fuel. The collapse is caused by the loss of the supporting pressure from the fusion in the interior of the star as the fuel runs out. 12. (a) A white dwarf is formed if the star is about the mass of the Sun or less. The low mass of the star allows its center to convert from hot plasma to neutral matter. This absorbs heat from the surrounding star until the surrounding material undergoes an expansion-cooling and contraction-heating process. Eventually the matter surrounding the core is blown away in a violent explosion. (b) A red giant star forms as the hydrogen in the core runs out and the star begins to collapse. When the hydrogen in the outer layers gets hot enough to fuse, the star expands to a red giant. A star can become a red giant twice in its life if helium in the core begins to fuse after it collapses from its initial red giant stage. The heat from this new energy source again expands the star. (c) When a star has used up all of its energy sources, it collapses in a catastrophic explosion called a supernova. If the remains of the compressed core after the supernova are between about 1.4 and 3 solar masses, the remaining matter is compressed so much that the nuclei of all the atoms collapse, forcing protons and electrons together into neutrons and forming the core of a small, superdense neutron star. Such a neutron star has a center core of pure neutrons. (d) A black hole is the final stage of a supermassive star (with a mass greater than 3 solar masses) when all the fuel in the center has been used up. The core undergoes a gravitational collapse that has nothing to stop it. The force of gravity overwhelms all other forces and even prevents light from escaping. (e) A supernova is the result of the collapse of a star more massive than the Sun when all its fuel has been exhausted. The energy released in this collapse causes the outer layers of material to rebound outward in a violent explosion. 13. The pressure from the high temperatures at the center of a star trying to expand the core balances the force of gravity attempting to collapse the star. This equilibrium can last millions of years because the pressure is a result of the high temperatures from the fusion process. Since there is so much matter in a star, it takes millions of years to fuse together. 14. Massive stars can fuse elements to create heavier elements up to iron. If the stars are sufficiently massive to undergo a supernova, the energy released in the supernova explosion can fuse atoms to make even heavier elements. 15. Helium nuclei are more massive than hydrogen nuclei, and therefore when traveling at the same speed as hydrogen atoms they have a higher kinetic energy. Since temperature is a measure of the energy of the atoms, the temperature is higher. In addition, a helium nucleus has a higher positive charge than a hydrogen nucleus. This higher positive charge creates a larger repulsive force between helium nuclei than hydrogen nuclei, requiring even higher speeds for the nuclei to collide and fuse. These higher speeds mean higher kinetic energies of the atoms, which mean higher temperatures. 16. A protostar becomes a star when the fusion process is started in the core. If this does not happen, the protostar could result in a cold gas giant such as Jupiter or Saturn. 17. A red giant star is a bright, low-temperature star that is very large. A red giant star forms as the hydrogen in the core runs out and the star begins to collapse. When the hydrogen in the outer layers gets hot enough to fuse, the star expands to a red giant. A star can become a red giant twice in its life if helium in the core begins to fuse after it collapses from its initial red giant stage. The heat from this new energy source again expands the star. A red giant can be brighter than it was as a main sequence star because even though it has a lower surface temperature and is not as luminous, it has much more surface area radiating light, making the star very bright. 18. Since there is not enough mass to create the high pressures and temperatures required to fuse the heavier elements, the fusion process stops at helium fusion. 19. The expansion of the universe is the expansion of space itself. The galaxies are not growing larger, they are moving with space as it expands. 20. The big bang theory is supported by the presence of present-day microwave background radiation from outer space, the current data on the expansion of the universe, the relative abundance of hydrogen and the products of fusion reactions, and measurements of background microwave radiation by satellites. For Further Analysis 1. If this year is 2012, light leaving the star 520 years ago would have left in 1492 which is about when Columbus sailed the ocean blue. 2. In general, a massive star is brighter, hotter, and has shorter a live than average-sized stars such as the Sun. This means the massive star will go through the various stages of its life in a shorter period. 3. Answers will vary, but should include the understanding that absolute magnitude compensates for distance, while apparent magnitude represents what you see. 4. The Hertzsprung-Russell diagram is a temperature-brightness graph that shows groups of stars. These groups represent stars in various stages of their lives. 5. Answers will vary. Instructor Manual for Integrated Science Bill W. Tillery, Eldon D. Enger , Frederick C. Ross 9780073512259
