19. Molecules and Compounds Elements, which are limited in number, combine to form an almost limitless number of compounds. Sometimes elements will combine in surprising ways. For example, the element sodium (Na) is a metal. Chlorine (Cl) is a deadly, greenish gas. But table salt (NaCl) (a compound of these two elements) is not a metal, not deadly, and not a gas. Clearly there is more to the creation of sodium chloride than just mixing atoms of sodium and chlorine together. Something else has occurred. On the other hand, some elements can be mixed together without any change taking place. In these cases the properties of the mixture are the expected combination of the properties of the elements. To begin to understand how compound substances come to be, consider two hydrogen atoms. Each atom has one electron in an s-orbital as depicted in Figure 19.1(a). The electrons are captured by the electrical attraction of the nucleus. The electrons are in an energy well from which they cannot escape without ionization energy being supplied.

d, but will be new orbitals that will be referred to as molecular orbitals. If the electrons in the new orbitals have less energy than they had in the atomic orbitals (the energy difference is radiated away as photons), the electrons will be deeper in an energy well than before and the new structure will be stable. It will remain as a structure until the missing energy is restored by some mechanism. In fact, in our example the new structure does form and is called a hydrogen molecule—symbolized as H2. The subscripted 2 refers to the two atoms that combined to form the new molecule. The hydrogen molecule is a new structure and it has its own characteristic energy levels and its own characteristic discrete emission spectrum. The example of the hydrogen molecule illustrates one of several ways atoms can combine to form molecules or other compound structures. In each case the new orbitals formed have less electron energy. Figure 17.7 diagrams the electron configurations of a number of elements. The noble gases are elements where no electrons are high in the energy well. On the other hand, when an atom exists with one more electron than needed to just fill a shell, the extra electron must go into the next higher shell. Invariably, the electron has relatively high energy. The valence electron in sodium, next after neon in the Periodic Table, is such an example. This electron is easy to remove; that is, sodium has a small ionization energy. In contrast, fluorine is an example of an element that lacks just one electron to fill the first shell. The electrons have low energies (are deep in the well), but not so low as those of neon. If sodium and fluorine are brought together they “react.” In other words, the valence electrons form new orbitals. In this instance the single valence electron of sodium forms an orbital around fluorine, and the electron has essentially been transferred from sodium to fluorine. But in doing so the electron falls deeper into the energy well and the new structure, NaF, is stable. The sodium atom has become a sodium ion (Na+) and the fluorine atom has become a fluorine ion (F–). The + and – in the symbols refer to the net electrical charge of each ion resulting from either the loss or gain of an electron. Atoms have more disorder (entropy) when they are free from one another than when they are bonded

(a)

(b)

Figure 19.1. (a) Two hydrogen atoms with electrons in s-orbitals. (b) When the atoms collide a new molecular orbital is formed which binds the two atoms together as a hydrogen molecule. Imagine our two hydrogen atoms bumping together so that each electron feels the pull from both nuclei. Now imagine, as a result, that new orbital patterns set up in such a way as to surround both nuclei (Fig. 19.1b). The new patterns will not be the atomic orbitals s, p, or

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together so that they must move collectively; their disorder decreases during bonding. However, total disorder never decreases, so something else must happen during bonding to increase the disorder somewhere else. As the electron jumps to a lower energy state, the atom radiates to the surroundings energy that is equivalent to the difference in orbital energies. Because of this release of energy, the entropy increases in the surroundings more than it decreases in the system. The total entropy increases, as it always does in a spontaneous process. We sometimes say that bonds form because the system has found a way to reduce the energy of its electrons. We could also say that bonds form because energy is given off to raise the entropy of the universe.

atoms from one another. However, the ionizing electrons in a spectrometer are never energetic enough to disturb or break up the nucleus or even to remove the most tightly bound electrons of atoms with high atomic numbers. More evidence for the structure of the hydrogen molecule comes from Avogadro’s Hypothesis, which says equal volumes of different gases contain equal numbers of molecules if the temperatures and pressures are the same. A certain volume of helium has a mass of 4 grams, and the same volume of hydrogen has a mass of 2 grams, indicating hydrogen molecules are half as massive as helium molecules. On the other hand, atomic masses show that hydrogen atoms are only one-fourth as massive as helium atoms. Therefore, hydrogen molecules must be made of two atoms of hydrogen. What happens to hydrogen in the mass spectrometer can be represented by the events shown in Figure 19.2. Careful measurements show this to be the only acceptable explanation of the structure of hydrogen and of its analysis in the mass spectrometer. Note that even though hydrogen exists as molecules, it is still an element just as surely as is helium which exists as individual atoms. A number of other elements exist as molecules composed of two atoms: nitrogen, oxygen, fluorine, chlorine, and bromine. In contrast, some molecules are composed of atoms of two or more types. For example,

Molecular Structure When pure hydrogen is placed in the mass spectrometer (see Chapter 15), many of the particles reaching the detector are so slow that they appear twice as heavy as normal hydrogen. A small percentage of these heavy particles is hydrogen atoms, each having an extra neutron, but most of them appear to be two joined hydrogen atoms with one electron missing. There is much evidence for this interpretation. If the electron energy in the ionization stage of the spectrometer is increased, these so-called molecular ions nearly disappear. Apparently, the energetic electrons separate the

Figure 19.2. Behavior of hydrogen in the mass spectrometer. (a) The ionizing electron beam removes an electron from some molecules and splits other molecules in half. (b) Atomic ions are accelerated to higher speeds than molecular ions. (c) Computer display of the times of arrival (masses) of ions.

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an atom of nitrogen attached to an atom of oxygen forms a molecule of nitric oxide. A sample of such molecules would be as invisible as air, but would smell horrible. In fact, nitric oxide is one of the most obnoxious components of smog. Its behavior in the mass spectrometer is indicated in Figure 19.3. Notice the different pattern of peaks compared with that of hydrogen. Each type of molecule forms its own distinctive pattern in the mass spectrometer. This is very useful to people in their efforts to identify molecules. For exam-

ple, Figure 19.4 shows the mass spectrum of a white powder confiscated from a person accused of possessing illegal drugs. The defendant claimed it was merely a painkiller containing aspirin and phenacetin (a nonprescription painkiller). The mass spectrum, however, shows a molecular ion with a mass of 315 amu, much larger than can be found in the materials claimed. This mass and the pattern of the other fragments identify the material as Percodan, a prescription narcotic painkiller that the defendant possessed illegally. This particular

Figure 19.3. (a) Fragments of nitric oxide molecules in the drift tube of a mass spectrometer showing low-mass ions getting ahead of high-mass ions. (b) Computer display of the masses of various fragments.

Figure 19.4. Mass spectrum of Percodan.

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Chemical Reactions

mass spectrum was used as evidence in court and helped lead to a conviction. A mass spectrograph can sometimes give very detailed information. For example, even though sugar molecules are composed of the same types of atoms regardless of the source, the relative heights of the peaks similar to the 316-amu peak in Figure 19.4 allow scientists to determine subtle difference which indicate whether the sugar was made from sugarcane or sugar beets. The peak heights also indicate approximately how many miles from the seacoast the sugar-bearing plant was grown. The mass spectrometer clarifies the relationships among the types of particles that are the building blocks of pure materials. An element is composed of only one type of atom. The atoms can exist as individuals or bonded together as molecules. In contrast to an element, a compound is composed of molecules, each of which is made of at least two types of atoms. Examples of each type are shown in Figure 19.5. In addition to these configurations, the atoms of a pure metal are joined to a number of neighboring metal atoms in solid metallic elements. It is not entirely correct to speak of a lump of metal as a single molecule, even though all of the atoms are connected.

Just how are molecules made from atoms or from other molecules? They are formed during chemical reactions. Some reactions are vigorous and showy, like the reaction of sodium and chlorine or the exploding of fireworks. Other reactions are quiet and slow, such as the chemical reactions that occur in green plants to produce the sugar described in the previous section. Charcoal burning in air is a good example of a chemical reaction. The black charcoal briquettes for barbecuing are nearly pure carbon. Once they get hot enough, they combine with oxygen from the air and apparently disappear. (The small amount of ash that remains is an impurity and does not appear when pure charcoal burns.) Actually, the carbon does not vanish—it becomes carbon dioxide, a transparent gas that goes up with the smoke. To initiate the reaction that burns the carbon, we are required to light a match and provide some extra energy. Often, the starting materials for a chemical reaction are stable molecules or substances for which the electrons are in orbitals that already minimize the electron energy. The additional energy provided by the match gives kinetic energy to the molecules of carbon and oxygen so that they bump into one another and excite the electrons, even to the point of breaking the molecules into ions or atoms. These ions or atoms can now rearrange into new combinations with new electron orbitals that will again lower the electron energy. These new combinations of atoms (CO2 in this case) will be the new molecules that are the products of the chemical reaction. The minimum amount of kinetic energy that molecules must have to initiate the reaction is called the activation energy. Once the reaction begins, it may release enough energy of its own to maintain the activation energy to keep the reaction going. This example shows some common characteristics of chemical reactions: 1. One or more substances are changed to specific amounts of other substances that have different characteristics. 2. Some kinetic energy may be required to initiate the reaction (activation energy). 3. Energy is released by the reaction. (In some chemical reactions, energy is absorbed.) Some similar examples of reactions with oxygen are the burning of hydrogen, iron, and magnesium. Hydrogen is a colorless gas, light enough to make balloons rise, but it is unsafe because it burns. When it burns in oxygen, it produces one of the hottest flames known. The product of the reaction is water (a gas when hot, but a liquid when cooled).

Figure 19.5. Drawings of elements and compounds. Why can’t compounds be atoms?

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Iron, in the form of steel wool, burns well in air; however, it appears as glowing coals rather than an open flame. The product of complete burning is similar to rust. At normal temperatures, this reaction proceeds so slowly that it is almost unnoticeable. Water speeds the rusting reaction enough so that changes can be seen from day to day, as you can discover by leaving steel tools exposed to moisture. Magnesium is a gray metal that is prized for its low density. It helped the original “mag wheels” be lightweight and therefore easy to accelerate. Pure magnesium (not modern mag wheels) burns vigorously in air and produces not only heat but also large amounts of white light. Roadside emergency flares use this reaction, usually with an additive to make the color red as well as bright. The white ash remaining when the reaction is completed is magnesium oxide. Let us summarize these chemical reactions by writing an equation in words for each:

of gas contain equal numbers of molecules at the same temperature and pressure. This provides a basis for stating that there are twice as many hydrogen atoms as oxygen atoms in each water molecule.

Figure 19.6. Simple electrolysis cell with electrodes wired to a battery. Another example of electrolysis is that of common table salt. The reaction proceeds at a reasonable rate only if the salt is melted. The required temperature, 801 °C, is easily attained in the laboratory but is far too hot for kitchen ovens. As the electrolysis proceeds, a pale green gas bubbles from one electrode, and a molten silver metal collects at the other. The gas is chlorine, one of the first poisonous gases used for warfare. Nowadays, many tons of chlorine are produced and used in the manufacture of several familiar products,

carbon 1 oxygen → carbon dioxide hydrogen 1 oxygen → water iron 1 oxygen → iron oxide (rust) magnesium 1 oxygen → magnesium oxide The arrow is read yields. The starting materials, termed reactants, appear to the left of the arrows. The resulting materials, or products, are on the right. Although it may appear to deserve no comment, these burning processes never use all of the oxygen in the atmosphere. There seems to be a quota of oxygen for each reaction, and the quota is never exceeded. When one reactant is used up, the unused portion of the other reactant is left over at the end of the reaction. These four reactions comprise a good beginning for our collection of chemical reactions. Two more reactions of a different type tear one substance apart and form two new materials with entirely different characteristics. Such reactions are called electrolyses. They occur when a source of direct electrical current (e.g., an automobile battery) is connected to two bars of metal (e.g., platinum). The bars—or electrodes, as they are called—are then dipped in the liquid that is to be electrolyzed. The reaction proceeds at a rate determined by the amount of electrical current. Figure 19.6 shows a simple electrolysis cell. Figure 19.7 shows a more complex cell that traps any gas that might bubble up from the electrodes, as happens when water is electrolyzed. As the water is consumed, oxygen comes from one electrode and twice its volume of hydrogen comes from the other. This is an example of the Law of Definite Composition and reflects the fact that each water molecule contains twice as many hydrogen atoms as oxygen atoms. Remember that according to Avogadro’s Hypothesis equal volumes

Figure 19.7. Water being electrolyzed in a slightly more elaborate electrolysis cell. such as bleaching and cleaning chemicals. The metal produced in the electrolysis of table salt is sodium, which is just as reactive as chlorine but in different ways. A small lump of sodium dropped into water will react violently, releasing hydrogen from the water at such a rate that the lump turns itself into a miniature jet-propelled boat. The heat released often is often to ignite or explode the hydrogen which has mixed with the oxygen of the air. This is why students in chemistry laboratories should not throw scrap sodium down the drains.

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Figure 19.8. Formulas and drawings of molecules.

Figure 19.9. Drawings of groups of molecules and their representations with chemical formulas. Chemical Formulas

The Law of Definite Composition is obeyed in the sodium chloride electrolysis reaction, but comparison is complicated by the fact that one product is a solid rather than a gas. Nevertheless, for every 35.5 grams of chlorine produced at one electrode, 23 grams of sodium is produced at the other. The proportions of the amounts of the products of a reaction are small whole numbers only when measured in gas volumes or atoms. The numbers are more complex when measured in masses, because the mass of one type of atom is not simply related to the masses of other types of atoms.

Atomic symbols such as H and He are used to denote elements or atoms. Similarly, symbols called chemical formulas can represent compounds or molecules that are composed of these atoms. Numerical subscripts show how many of each type of atom are involved. For example, two atoms of nitrogen combining to form a molecule is represented by N2. Water is constructed of two atoms of hydrogen and one of oxygen, so its formula is H2O. It is possible to write the formulas so that they indicate something about the way the atoms are linked. Water’s formula could correctly be written HOH to

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Figure 19.10. Reaction of hydrogen with the proper amount of oxygen.

Figure 19.11. Reaction of hydrogen with excess oxygen.

show that both hydrogens are connected to the central oxygen. Some examples of molecules and their formulas are drawn in Figure 19.8. You are not expected to know which atom is central or precisely how they are linked, only that chemical formulas written in certain ways can indicate the relative positions of the atoms in a molecule. Suppose we wished to write the symbol for three hydrogen molecules. To avoid confusion we must write 3H2. The symbol 3H would indicate three individual hydrogen atoms, but there are a total of six atoms in three hydrogen molecules. The symbol H6 would indicate that all six were strung together in a single molecule, which is incorrect. By convention, the first number indicates the number of molecules, and the remainder of the formula shows the composition of each molecule. Some examples are diagramed in Figure 19.9. With a knowledge of molecular formulas, it is possible to write equations for the reactions in formulas rather than words:

chemical reaction does not create or destroy atoms; it simply rearranges them into new molecules. This is diagramed for the reaction of hydrogen and oxygen in Figure 19.10, in which the right amount of oxygen was present to complete the reaction. In Figure 19.11, too much oxygen is shown, and some oxygen molecules are left over. Virtually all reactions occurring in nature stop with some reactant left over. Chemical equations do not show excess reactants any more than a recipe shows the number of eggs remaining after one has cooked an omelet. The chemical equation shows the minimum amount of each material needed without arbitrarily splitting molecules. As an example of improper splitting, consider the equation

C 1 O2 → CO2

2H2 1 O2 → 2H2O .

H2 1 O → H2O . This equation implies that oxygen exists as single atoms when it is really composed of molecules, O2. To reflect this fact, we must write

2H2 1 O2 → 2H2O This equation shows the proper number of atoms and molecules on each side of the equation. The equation is said to be balanced when, for each kind of atom, there are the same number of atoms on the left of the arrow as on the right. Once again, this represents the Law of Conservation of Mass. It also shows that an atom of one type cannot be transformed to another type of atom by a chemical reaction. Unfortunately, there is no simple, foolproof recipe

4Fe 1 3O2 → 2Fe2O3 2Mg 1 O2 → 2MgO 2H2O → 2H2 1 O2 2NaCl → 2Na 1 Cl2 These equations obey the Law of Conservation of Mass. That is, there are as many atoms of each element among the reactants as among the products, indicating that a

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for finding the numbers to put in front of each molecular symbol to balance a chemical reaction equation. One simply proceeds by trial and error. With a little practice, one observes that certain patterns begin to repeat themselves. Some chemical formulas can give you a false impression if you are not forewarned. The formula H2O is a symbol for a molecule containing two hydrogen atoms and a single oxygen atom. Likewise, H2 and O2 are molecular formulas, i.e., they represent individual molecules. The equation

1.

2.

3.

2H2 1 O2 → 2H2O

If the formula represents a gas it is a molecular formula. Examples are H2 and O2. (We sometimes convey physical state information by adding (g) for gas, (l) for liquid, and (s) for solid. Examples are H2 (g), H2O (l), and NaCl (s).) If the formula contains only nonmetal elements, it probably is a molecular formula. Examples are CO2 and SF6. If the formula contains a metal element, it is probably a simplest formula, unless it is a gas. Examples are Na and NaCl.

Summary

says that two hydrogen molecules react with one oxygen molecule to form two molecules of water.. Examples of molecular formulas are N2, F2, Cl2, Br2, I2, CO2, and SF6 . The noble gases contain monatomic molecules (single atoms) and have molecular formulas He, Ne, Ar, Kr, Xe, and Rn. The symbol for table salt, NaCl, is not a molecular formula in the above sense because table salt is a crystalline solid substance. The smallest piece of salt that is still recognizable as salt contains trillions of sodium ions and trillions of chloride ions stacked together in a three-dimensional checkerboard pattern. The smallest subunit of this pattern contains one sodium ion and one chloride ion. This subunit is not, strictly speaking, a molecule. Molecules have strong chemical bonds among the atoms within each molecule, but only weak (at best) bonds between different molecules, so each molecule is recognized as a separate unit. Yet in solid table salt, each subunit NaCl has a strong chemical bond to its adjacent subunits. Hence, we say that NaCl is a simplest formula rather than a molecular formula. To indicate two of these subunits, we would write 2NaCl. The symbol Na (sodium) as it applies to a piece of metal is also a simplest formula. It symbolizes the simplest subunit of a piece of sodium metal, the sodium atom, even though the sodium atoms are bonded together to form the piece of metal. Examples of the simplest formulas are Mg, Fe, and MgCl2.

In a chemical reaction, atoms rearrange to form new structures. If the new structures are stable, the valence electrons form new orbitals in which they have less energy than they had in their respective structures before the reaction. The new orbitals bind the new structures together: new chemical bonds are said to have formed. The Law of Definite Proportions is obeyed in chemical reactions: The amounts of products and reactants in a particular chemical reaction always occur in the same definite proportions by volume of gases or by numbers of molecules. In a chemical reaction chemical bonds are made or broken. The structures of molecules formed in chemical reactions can be analyzed by the mass spectrometer. Spontaneous chemical reactions may form more ordered structures, but the energy released will increase the disorder of the surroundings in such a way that the total disorder of the universe is increased. Chemical equations show the arrangement of atoms and molecules at the beginning and the end of chemical reactions. Balanced chemical equations emphasize that atoms are neither created nor destroyed in chemical reactions. Thus, the Law of Conversation of Mass is obeyed. STUDY GUIDE Chapter 19: Molecules and Compounds

The chemical reaction

A. FUNDAMENTAL PRINCIPLES 1. The Law of Definite Proportions is obeyed in chemical reactions: the amounts of products and reactants in a particular chemical reaction always occur in the same definite proportions. The proportions of the amounts of the reaction products are small whole numbers when measured in gas volumes or numbers of molecules. 2. The Law of Conservation of Mass: See Chapter 18.

2Na 1 Cl2 → 2NaCl means that if metallic sodium is burned in chlorine gas, two atoms of sodium metal combine with a molecule of chlorine to form two subunits of the structure we know as table salt. There are three rules of thumb to use in deciding whether a formula is a molecular formula or a simplest formula:

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B. MODELS, IDEAS, QUESTIONS, OR APPLICATIONS 1. What are chemical bonds? 2. What happens to the energy of a system when atoms bond together? 3. What happens to the disorder of a system when atoms bond together? 4. What is a chemical reaction? 5. How are chemical formulas written and balanced?

a crystal or a piece of metal, the simplest formula is the chemical formula of the basic subunit of the structure. In a crystal of salt, the simplest formula is NaCl even though NaCl does not exist as separate molecules. D. FOCUS QUESTIONS 1. In a demonstration, hydrogen is burned in oxygen to form water. For this chemical reaction: a. Write and balance the chemical equation representing the reaction. b. Explain why a match is needed to start the reaction. c. Describe what happens to the energy within the system as the products are formed. What changes in the forms of energy occur? d. Describe what happens to the disorder of the universe as the products are formed. 2. In a demonstration, magnesium is burned in oxygen forming magnesium oxide. For this chemical reaction: a. Write and balance the chemical equation representing the reaction. b. Explain why a match is needed to start the reaction. c. Describe what happens to the energy within the system as the products are formed. What changes in the forms of energy occur? d. Describe what happens to the disorder of the universe as the products are formed.

C. GLOSSARY 1. Activation Energy: The minimum energy that reactants must have to form products. In some chemical reactions, additional evergy must be added from an external source (such as a lighted match) to reach the activation energy. 2. Avogadro’s Hypothesis: See Chapter 18. 3. Chemical Bond: The attachment that results from a binding together of molecules or atoms in definite proportions into new structures with an accompanying release (or, rarely, absorption) of energy. 4. Chemical Reaction: The energy releasing or absorbing process which makes or breaks a chemical bond. 5. Compound: A substance consisting of molecules formed from atoms of at least two different kinds. Water is a compound of hydrogen and oxygen. 6. Electrolysis: A chemical reaction in which the necessary driving energy is provided in the form of an electrical current passing through the reactants. Water may be separated into hydrogen and oxygen by electrolysis. 7. Element: See Chapter 18. 8. Mixture: A physical combination without chemical bonding of compounds and/or elements in indefinite proportions. A mixture of oxygen and hydrogen might have a large number of oxygen molecules and a small number of hydrogen molecules or vice versa. 9. Molecular Formula: Chemical formulas (symbols) representing individual molecules. 10. Molecular Orbital: An electron orbital, different in shape and energy from the atomic electron orbitals (s, p, d, etc.) of atoms, that forms between atoms to form the “glue” that bonds atoms together to form molecules. 11. Molecule: The simplest independent structural unit of a particular element or compound. A molecule of helium is the helium atom. A molecule of water is the structure H2O. 12. Products: The name given to the substances that result from a chemical reaction. 13. Reactants: The name given to the substances before they are combined or separated by a chemical reaction. 14. Simplest Formula: In repeating structures such as

E.

EXERCISES 19.1. Notice the peak at 1 amu in Figure 19.2c. What ion caused this peak? What ion caused the peak at 2 amu in Figure 19.2c? 19.2. In Figure 19.3b which peak corresponds to the O1+ ion? (Use Table 17.1 to find the mass of an oxygen atom.) What ions caused the other peaks? 19.3. What molecule was put in the mass spec-

Figure 19.12. Masses of fragments of an unknown molecule.

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F2 CO2 Cu

Ar BeO CH4

19.7. Label the groups of molecules in Figure 19.13 the way they are labeled in Figure 19.9. 19.8. How many atoms of hydrogen are there in each molecule of propane that has the formula H3CCH2CH3? Use this formula to draw a diagram of the molecule of propane similar to the diagrams in Figure 19.8. 19.9. How many atoms of nitrogen are there in three molecules of laughing gas, N2O? (Notice that we could have written 3N2O to represent the three molecules.) How many atoms of oxygen are there in 5Al2O3? 19.10. Which of the following equations are balanced? (a) 2C + O2 → 2CO (b) H2 + F2 → HF (c) Mg + F2 → MgF2 (d) 2H2O2 → H2O + O2

Figure 19.13. Groups of molecules to be labeled in Exercise 19.7. trometer to cause the peaks shown in Figure 19.12?

19.11. Suppose someone asked you to finance a promising silver extraction process based on the equation 2Ag2O → 8Ag + O2

19.4. What fragments would the mass spectrometer make from water molecules that are composed of two hydrogen atoms attached to an oxygen atom?

Would you consider him a competent person or a crackpot?

19.5. Listed below are some chemical reactions with which we hope you are familiar. Classify each as explosive, fast, slow, or nonoccurring (under natural conditions). Where you can, describe the characteristics of the products. Remember burning yields carbon dioxide from materials that contain carbon and water from things that contain hydrogen. (a) Rusting of cans along the highway . (b) Burning of gasoline vapor in air when ignited (gasoline is made of carbon and hydrogen). (c) Burning of diamonds in air when ignited (diamond is pure carbon and behaves similarly to charcoal in chemical reactions). (d) Burning of diamonds without ignition (will diamonds “rust”?). (e) Burning large amounts of gunpowder. (f) Tarnishing of silver. (g) Fading of clothes dyes in sunlight. (h) Decay of the insecticide DDT in the environment.

19.12. What are the important differences between a chemical reaction, such as the electrolysis of molten salt, and a physical process, such as the crushing of salt crystals? 19.13. Classify each of the following as pure element, pure compound, or mixture as best you can from your experience and the information in this unit: 7-Up orange juice air nickel CCl4, carbon tetrachloride tin vitamin A diamond tungsten metal chlorine 19.14. What compound would (in analogy to Fig. 19.3) give mass spectrometer peaks at 1, 12, 14, 26, and 27 amu? 19.15. How many atoms of fluorine are in 3SF6?

19.6. Label each of the following formulas as representing an element or a compound:

19.16. Which of the following equations are bal-

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anced? (a) (b) (c) (d)

Ca + O2 → CaO NaOH + HCl → H2O + NaCl 4Al + 3O2 → 2Al2O3 NO2 + H2O → HNO3

19.17. Which of the following is not a compound? (a) sulfur dioxide (b) nitrous oxide (c) water molecule (d) oxygen molecule (e) rust 19.18. Which equation is balanced correctly? (a) H2 + O2 → 2H2O (b) H2 + O2 → H2O (c) Ca + Cl2 → CaCl2 (d) S + O2 → SO4 (e) N2 + O2 → NO

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