Chem 101 Chapter 2- Atoms and Elements

Chem 101 Chapter 2- Atoms and Elements - The Basic Tools of...

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Unformatted text preview: The Basic Tools of Chemistry 2— Atoms and Elements Stardust A wide array of elements make up planet Earth and every living thing on it. What is science’s view of the cosmic origin of these elements that we take for granted in our environment and in our lives? The “big bang” theory is the generally accepted theory for the origin of the universe. This theory holds that an unimaginably dense, grapefruit-sized sphere of matter exploded about 15 billion years ago, spewing the products of that explosion as a rapidly expanding cloud with a temperature in the range of 1030 K. Within 1 second, the universe was populated with the particles we explore in this chapter: protons, electrons, and neutrons. Within a few more seconds, the universe had cooled by millions and millions of degrees, and protons and neutrons began to combine to form helium nuclei. After only about 8 minutes, scientists believe, the universe was about one-quarter helium and about three-quarters hydrogen. In fact, this is very close to the composition of the universe today, 15 billion years later. But humans, animals, and plants are built mainly from carbon, oxygen, nitrogen, sulfur, phosphorus, iron, and zinc—heavier elements that have only a trace abundance in the universe as a whole. Where do these heavier elements come from? The cloud of hydrogen and helium cooled over a period of thousands of years and condensed into stars like our sun. There hydrogen atoms fuse into more helium atoms and energy streams outward. Every second on the sun, 700 million tons of hydrogen is converted to 695 million tons of helium, and 3.9 1026 joules of energy is evolved. Gradually, over millions of years, a hydrogen-burning star becomes more and more dense and hotter and hotter. The helium atoms initially formed in the star begin to fuse into heavier atoms—first carbon, then oxygen, and then neon, magnesium, silicon, phospho- The supernova of 1987. When a star becomes more and more dense, and hotter and hotter, it can become a “red giant.” The star is unstable and explodes as a “supernova.” One such spectacular event occurred in a distant star in 1987. These explosions are the origin of the heavier elements, such as iron, nickel, and cobalt. 58 Dr. Christopher Burrows, ESA/STSc1 and NASA. Chapter Goals See Chapter Goals Revisited (page 89). Test your knowledge of these goals by taking the exam-prep quiz on the General ChemistryNow CD-ROM or website. 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 Chapter Outline Protons, Electrons, and Neutrons: Development of Atomic Structure Atomic Number and Atomic Mass Isotopes Atomic Weight Atoms and the Mole The Periodic Table An Overview of the Elements, Their Chemistry, and the Periodic Table Essential Elements • Describe atomic structure and define atomic number and mass number. • Understand the nature of isotopes and calculate atomic weight from the isotopic masses and abundances. • Explain the concept of the mole and use molar mass in calculations. • Know the terminology of the periodic table. rus, and argon. The star becomes even hotter and more dense. Hydrogen is forced to the outer reaches of the star, and the star becomes a red giant. Under certain circumstances, the star will explode, and earth-bound observers see it as a supernova. A supernova can be as much as 108 times brighter than the original star. A single supernova is comparable in brightness to the whole of the galaxy in which it is formed! The supernova that appeared in 1987 gave astronomers an opportunity to study what happens in these element factories. It is here that the heavier atoms such as iron form. In fact, it is with iron that nature reaches its zenith of stability. To make heavier and heavier elements requires energy, rather than having energy as an outcome of element synthesis. The elements spewing out of an exploding supernova move through space and gradually condense into planets, of which ours is just one. The mechanism of element formation in stars is reasonably well understood, and much experimental evidence exists to support this theory. However, the way in which these elements are then assembled out of stardust into living organisms on our planet—and perhaps other planets—is not yet understood at all. (See the General ChemistryNow Screen 2.2 Introduction to Atoms, to watch a video on the “big bang” theory.) 1014 1012 1010 Relative abundance 108 106 104 102 0 H He Li Be B C N O F Ne NaMg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Element The abundance of the elements in the solar system from H to Zn. The chart shows a general decline in abundance with increasing mass among the first 30 elements. The decline continues above zinc. Notice that the scale on the vertical axis is logarithmic—that is, it progresses in powers of ten. The abundance of nitrogen, for example, is 1/10,000 (1/104) of the abundance of hydrogen. (All abundances are plotted as the number of atoms per 1012 atoms of H. The fact that the abundances of Li, Be, and B, as well as those of the elements near Fe, do not follow the general decline is a consequence of the way that elements are synthesized in stars.) 59 60 Chapter 2 Atoms and Elements To Review Before You Begin • Names and uses of SI units (Section 1.6) • Solving numerical problems using units (Section 1.8) T Throughout the chapter this icon introduces a list of resources on the General ChemistryNow CD-ROM or website (http://now that will: help you evaluate your knowledge of the material provide homework problems allow you to take an exam-prep quiz provide a personalized Learning Plan targeting resources that address areas you should study he chemical elements are forged in stars. What are the similarities among the elements? What are the differences? What are their physical and chemical properties? How can we tell them apart? This chapter begins our exploration of the chemistry of the elements, the building blocks of the science of chemistry. 2.1—Protons, Electrons, and Neutrons: Development of Atomic Structure Around 1900 a series of experiments done by scientists such as Sir John Joseph Thomson (1856–1940) and Ernest Rutherford (1871–1937) in England established a model of the atom that is still the basis of modern atomic theory. Three subatomic particles make up all atoms: electrically positive protons, electrically neutral neutrons, and electrically negative electrons. The model places the more massive protons and neutrons in a very small nucleus, which contains all the positive charge and almost all the mass of an atom. Electrons, with a much smaller mass than protons or neutrons, surround the nucleus and occupy most of the volume (Figure 2.1). Atoms have no net charge; the positive and negative charges balance. The number of electrons outside the nucleus equals the number of protons within the nucleus. What is the experimental basis of atomic structure? How did the work of Thomson, Rutherford, and others lead to this model? • • • • Electricity Nucleus (protons and neutrons) Electron cloud Figure 2.1 The structure of the atom. All atoms contain a nucleus with one or more protons (positive electric charge) and neutrons (no charge). Electrons (negative electric charge) are arranged in space as a “cloud” around the nucleus. In an electrically neutral atom, the number of electrons equals the number of protons. Note that this figure is not drawn to scale. If the nucleus were really the size depicted here, the electron cloud would extend about 800 feet. The atom is mostly empty space! Electricity is involved in many of the experiments from which the theory of atomic structure was derived. The fact that objects can bear an electric charge was first observed by the ancient Egyptians, who noted that amber, when rubbed with wool or silk, attracted small objects. You can observe the same thing when you rub a balloon on your hair on a dry day—your hair is attracted to the balloon (Figure 2.2a). A bolt of lightning or the shock you get when touching a doorknob results when an electric charge moves from one place to another. Two types of electric charge had been discovered by the time of Benjamin Franklin (1706–1790), the American statesman and inventor. He named them positive ( ) and negative ( ), because they appear as opposites and can neutralize each other. Experiments show that like charges repel each other and unlike charges attract each other. Franklin also concluded that charge is balanced: If a negative charge appears somewhere, a positive charge of the same size must appear somewhere else. The fact that a charge builds up when one substance is rubbed over another implies that the rubbing separates positive and negative charges. By the 19th century it was understood that positive and negative charges are somehow associated with matter—and perhaps with atoms. Radioactivity In 1896 the French physicist Henri Becquerel (1852–1908) discovered that a uranium ore emitted rays that could darken a photographic plate, even though the plate was covered by black paper to protect it from being exposed to light. In 1898 Marie 2.1 Protons, Electrons, and Neutrons: Development of Atomic Structure 61 b particles Photographic film or phosphor screen b particles, attracted to plate Undeflected g rays g rays a particles Lead block shield a particles, attracted to Slit Radioactive element Charged plates plate (a) (b) Figure 2.2 Electricity and radioactivity. (a) If you brush your hair with a balloon, a static electric charge builds up on the surface of the balloon. Experiments show that objects having opposite electric charges attract each other, whereas objects having the same electric charge repel each other. (See the General ChemistryNow Screen 2.4 Electricity and Electric Charge, for an exercise on the effects of charge.) (b) Alpha (a), beta (b), and gamma (g) rays from a radioactive element are separated by passing them between electrically charged plates. Positively charged a particles are attracted to the negative plate, and negatively charged b particles are attracted to the positive plate. (Note that the heavier a particles are deflected less than the lighter b particles.) Gamma rays have no electric charge and pass undeflected between the charged plates. (See the General ChemistryNow Screen 2.5 Evidence of Subatomic Particles, for an exercise on this experiment.) and Pierre Curie (1867–1934) isolated polonium and radium, which also emitted the same kind of rays, and in 1899 they suggested that atoms of certain substances emit these unusual rays when they disintegrate. They named this phenomenon radioactivity, and substances that display this property are said to be radioactive. Early experiments identified three kinds of radiation: alpha (a), beta (b), and gamma (g) rays. These rays behave differently when passed between electrically charged plates (Figure 2.2b). Alpha and b rays are deflected, but g rays pass straight through. This implies that a and b rays are electrically charged particles, because they are attracted or repelled by the charged plates. Even though an a particle was found to have an electric charge (+2) twice as large as that of a b particle ( 1), a particles are deflected less, which implies that a particles must be heavier than b particles. Gamma rays have no detectable charge or mass; they behave like light rays. Marie Curie’s suggestion that atoms disintegrate contradicted ideas put forward in 1803 by John Dalton (1766–1844) that atoms are indivisible. If atoms can break apart, there must be something smaller than an atom; that is, atoms must be composed of even smaller, subatomic particles. Cathode-Ray Tubes and the Characterization of Electrons Further evidence that atoms are composed of smaller particles came from experiments with cathode-ray tubes (Figure 2.3). These are glass tubes from which most of the air has been removed and that contain two metal electrodes. When a sufficiently high voltage is applied to the electrodes, a cathode ray flows from the negative electrode (cathode) to the positive electrode (anode). Experiments showed that cathode rays travel in straight lines, cause gases to glow, can heat metal objects red hot, can be deflected by a magnetic field, and are attracted toward positively charged 62 Slits to focus a narrow beam of rays Chapter 2 Atoms and Elements Electrically charged deflection plates Fluorescent sensitized screen Undeflected electron beam Electrically deflected electron beam 1. 2. Negative electrode Positive electrodes accelerate electrons To vacuum pump 1. A beam of electrons (cathode rays) is accelerated through two focusing slits. 2. When passing through an electric field the beam of electrons is deflected. 3. Magnetic field coil perpendicular to electric field 3. The experiment is arranged so that the electric field causes the beam of electrons to be deflected in one direction. The magnetic field deflects the beam in the opposite direction. Magnetically deflected electron beam 4. 4. By balancing the effects of the electrical and magnetic fields the charge-to-mass ratio of the electron can be determined. Active Figure 2.3 Measuring the electron’s charge-to-mass ratio. This experiment was done by J. J. Thomson in 1896–1897. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. plates. When cathode rays strike a fluorescent screen, light is given off in a series of tiny flashes. We can understand all of these observations if a cathode ray is assumed to be a beam of the negatively charged particles we now know as electrons. You are already familiar with cathode rays. Television pictures and the images on some types of computer monitors are formed by using electrically charged plates to aim cathode rays onto the back of a phosphor screen on which we view the image. Sir Joseph John Thomson (1856–1940) used this principle to prove experimentally the existence of the electron and to study its properties. He applied electric and magnetic fields simultaneously to a beam of cathode rays (Figure 2.3). By balancing the effect of the electric field against that of the magnetic field and using basic laws of electricity and magnetism, Thompson calculated the ratio of the charge to the mass for the particles in the beam. He was not able to determine either charge or mass independently. However, he found the same charge-to-mass ratio in experiments using 20 different metals as cathodes and several different gases. These results suggested that electrons are present in atoms of all elements. It remained for the American physicist Robert Andrews Millikan (1868–1953) to measure the charge on an electron and thereby enable scientists to calculate its mass (Figure 2.4). In his experiment tiny droplets of oil were sprayed into a chamber. As they settled slowly through the air, the droplets were exposed to x-rays, which caused them to acquire an electric charge. Millikan used a small telescope to observe individual droplets. If the electric charge on the plates above and below the droplets was adjusted, the electrostatic attractive force pulling a droplet upward could be balanced by the force of gravity pulling the droplet downward. From the equations describing these forces, Millikan calculated the charge on various droplets. Different droplets had different charges, but Millikan found that each was a whole-number multiple of the same smaller charge, 1.60 10 19 C (where C represents the coulomb, the SI unit of electric charge; Appendix C). Millikan assumed this to be the fundamental unit of charge, the charge on an electron. Because the 2.1 Protons, Electrons, and Neutrons: Development of Atomic Structure 63 Oil atomizer Light source to illuminate drops for viewing X-ray source Oil droplets under observation Oil atomizer Voltage applied to plates Positively charged plate Light source Telescope X-ray source Negatively charged plate 1. A fine mist of oil drops is introduced into one chamber. 2. The droplets fall one by one into the lower chamber under the force of gravity. 3. Gas molecules in the bottom chamber are ionized (split into electrons and a positive fragment) by a beam of x-rays. The electrons adhere to the oil drops, some droplets having one electron, some two, and so on. These negatively charged droplets continue to fall due to gravity. 4. By carefully adjusting the voltage on the plates, the force of gravity on the droplet is exactly counterbalanced by the attraction of the negative droplet to the upper, positively charged plate. Analysis of these forces lead to a value for the charge on the electron. Figure 2.4 Electron Charge. The experiment was done by R. A. Millikan in 1909. (See the General ChemistryNow Screen 2.7 Charge and Mass of the Electron, for an exercise on this experiment.) charge-to-mass ratio of the electron was known, the mass of an electron could be calculated. The currently accepted value for the electron mass is 9.109383 10 28 g, and the electron charge is 1.602176 10 19 C. When describing the properties of fundamental particles, we always express charge relative to the charge on the electron, which is given the value of 1. Additional experiments showed that cathode rays have the same properties as the b particles emitted by radioactive elements. This provided further evidence that the electron is a fundamental particle of matter. Extensive studies with cathode ray tubes in the late nineteenth century provided another dividend. In addition to cathode rays, a second type of radiation was detected. A beam of positively charged particles called canal rays was observed using a specially designed cathode-ray tube with a perforated cathode (Figure 2.5). Anode Cathode rays Positive (Canal) rays Cathode with holes (pierced disk) Like cathode rays, positive rays (or "canal rays") are deflected by electric and magnetic fields but much less so than cathode rays for a given value of the field because positive particles are much heavier than electrons. Electron attracted to anode collides with gas molecule. Electron Gas molecules To vacuum pump 1. Electrons collide with gas molecules in this cathode-ray tube with a perforated cathode. Gas molecule splits into positive ion ( ) and electron ( ). Electrons continue to move to left; positive ion moves to right. Positive ion 3. Some positive particles pass through the holes and form a beam, or "ray." 2. The molecules become positively charged, and these are attracted to the negatively charged, perforated cathode. Figure 2.5 Canal rays. In 1886 E. Goldstein detected a stream of particles traveling in the direction opposite to that of the negatively charged cathode rays. We now know that these particles are positively charged ions, formed by collisions of electrons with gaseous molecules in the cathode-ray tube. (See the General ChemistryNow Screen 2.8 Protons, to view an animation on this experiment.) 64 Chapter 2 Atoms and Elements These particles, which moved in the opposite direction to cathode rays, passed through the holes in the cathode and were detected on the opposite side. Chargeto-mass values for canal rays were much smaller than the corresponding values measured for cathode rays, indicating particles of higher mass. However, the values also varied depending on the nature of the gas in the tube. We now know that canal rays arise through collisions of cathode rays with gaseous atoms within the cathoderay tube, which cause each atom to fragment into a positive ion and an electron. The positive particles are attracted to the negatively charged cathode. See the General ChemistryNow CD-ROM or website: • Screen 2.6 Electrons, for an exercise on cathode rays and an animation on cathode-ray deflection • Screen 2.8 Protons, for an exercise on the properties of nuclei in a canal-ray tube Protons A century after these seminal studies on the structure of the atom, it is easy for us to recognize the proton as the fundamental positively charged particle in an atom. This understanding did not come so easily a hundred years ago, however. This basic fact was not established in one specific experiment or at one specific moment. With the determination that negatively charged electrons were a component of the atom came recognition that positively charged atomic particles must also exist. One hypothesis suggested that there should be a complementary particle to the electron with a corresponding small mass and a 1 charge, but there was no experimental evidence for such a particle. The positive particles detected and studied in early experiments (a particles from radioactive elements and positive ions making up canal rays) were considerably more massive. Ernest Rutherford (1871–1937) probably deserves most of the credit for the discovery of the proton. He carried out experiments in the early 1900s in which various elements were irradiated with a particles. One of his better-known experiments involved the irradiation of metals such as gold, which led to the conclusion that atoms contained a small positively charged nucleus with most of the mass of an atom [ The Nucleus of the Atom, page 65]. At the same time Rutherford was performing similar experiments using gaseous elements, and, in these experiments, he observed the deflection of a particles as a function of atomic mass. From these observations he concluded, in 1911, that “the hydrogen atom has the simplest possible structure with only one unit charge.” However, the formal identification of the proton did not come until almost 10 years later. In experiments in which nitrogen was bombarded with a particles, Rutherford and his collaborators observed highly energetic particles. The values of their charge-to-mass ratio matched those for hydrogen, the positive particle known to have the lowest mass. Unexpectedly they had carried out the first artificial nuclear reaction. Expelling a proton from the nucleus was accepted as definitive evidence of the proton as a nuclear particle. The name “proton” for this particle appears to have been first used by Rutherford in a report in a scientific meeting in 1919. Neutrons Because atoms have no net electric charge, the number of positive protons must equal the number of negative electrons in an atom. Most atoms, however, have 2.1 Protons, Electrons, and Neutrons: Development of Atomic Structure 65 Historical Perspectives Uncovering Atomic Structure The last few years of the 19th century and the first decades of the 20th century were among the most important in the history of science, in part because the structure of the atom was discovered, setting the stage for the explosion of developments in science in the 20th century. The notion that matter was built of atoms and that this structure could be used to explain chemical phenomena was first used by John Dalton (1766–1844). Dalton proposed not only that all matter is made of atoms, but also that all atoms of a given element are identical and that atoms are indivisible and indestructible. Dalton’s ideas were generally accepted within a few years of his proposal, but we know now the last two postulates are not correct. Marie Curie (1867–1934) understood the nature of radioactivity and its implications for the nature of the atom. She was born Marya Sklodovska in Poland. When she later lived in France she was known as Marie, but today she is often referred to as Madame Curie. With her husband Pierre she isolated the previously unknown elements polonium and radium from a uranium-bearing ore. They shared the 1911 Nobel Prize in chemistry for this discovery. One of their daughters, Irène, married Frédéric Joliot, and they shared the 1935 Nobel Prize in chemistry for their discovery of artificial radioactivity. (See the General ChemistryNow Screen 2.5 Evidence of Subatomic Particles, to view an animation on separation of radiation by electric field.) Sir Joseph John Thomson (1856–1940) was Cavendish Professor of Experimental Physics at Cambridge University in England. In 1896 he gave a series of lectures at Princeton University in the United States titled the Discharge of Electricity in Gases. This work on cathode rays led to his discovery of the electron, which he announced at a lecture on the evening of Friday, April 30, 1897. Thomson later published a number of books on the electron and was awarded the Nobel Prize in physics in 1906. (See the General ChemistryNow Screen 2.6 Electrons, to view an animation on cathode-ray deflection.) Ernest Rutherford (1871–1937) was born in New Zealand in 1871 but went to Cambridge University in England to pursue his Ph.D. in physics in 1895. There he worked with J. J. Thomson, and it was at Cambridge that he discovered a and b radiation. At McGill University in Canada in 1899 Rutherford did further experiments to prove that a radiation is composed of helium nuclei and that b radiation consists of electrons. He received the Nobel Prize in chemistry for his work in 1908. His research on the structure of the atom was done after he moved to Manchester University in England. In 1919 he returned to Cambridge University, where he took up the position formerly held by Thomson. In his career, Rutherford guided the work of ten future recipients of the Nobel Prize. Element 104 has been named rutherfordium in his honor. (See the General ChemistryNow Screen 2.10 The Nucleus of the Atom, to view an animation on Rutherford’s a particle experiment.) Photos: (Left and Right) Oesper Collection in the History of Chemistry/University of Cincinnati; (Center Top) E. F. Smith Collection; (Center Bottom) Corbis. masses greater than would be predicted on the basis of only protons and electrons, which suggested that atoms must also contain relatively massive particles with no electric charge. In 1932, the British physicist James Chadwick (1891–1974), a student of Rutherford, presented experimental evidence for the existence of such particles. Chadwick found very penetrating radiation was released when particles from radioactive polonium struck a beryllium target. This radiation was directed at a paraffin wax target, and Chadwick observed protons coming from that target. He reasoned that only a heavy, noncharged particle emanating from the beryllium could have caused this effect. This particle, now known as the neutron, has no electric charge and a mass of 1.674927 10 24 g, slightly greater than the mass of a proton. The Nucleus of the Atom J. J. Thomson had supposed that an atom was a uniform sphere of positively charged matter within which thousands of electrons were embedded. Thomson and his students thought the only question was the number of electrons 66 Nucleus of Beam of a particles gold atoms Atoms in gold foil Chapter 2 Atoms and Elements Electrons occupy space outside nucleus. Undeflected a particles Gold foil Deflected particles Some particles are deflected considerably. A few a particles collide head-on with nuclei and are deflected back toward the source. Most a particles pass straight through or are deflected very little. Source of narrow beam of fast-moving a particles ZnS fluorescent screen Active Figure 2.6 Rutherford’s experiment to determine the structure of the atom. A beam of positively charged a particles was directed at a thin gold foil. A fluorescent screen coated with zinc sulfide (ZnS) was used to detect particles passing through. Most of the particles passed through the foil, but some were deflected from their path. A few were even deflected backward. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. ■ How Small Is an Atom? The radius of the typical atom is between 30 and 300 pm (3 10 11 m to 3 10 10 m). To get a feeling for the incredible smallness of an atom, consider that one teaspoon of water (about 1 cm3) contains about three times as many atoms as the Atlantic Ocean contains teaspoons of water. circulating within this sphere. About 1910, Rutherford decided to test Thomson’s model. Rutherford had discovered earlier that a rays (see Figure 2.2b) consisted of positively charged particles having the same mass as helium atoms. He reasoned that, if Thomson’s atomic model were correct, a beam of such massive particles would be deflected very little as it passed through the atoms in a thin sheet of gold foil. Rutherford, with his associates Hans Geiger (1882–1945) and Ernst Marsden, set up the apparatus diagrammed in Figure 2.6 and observed what happened when a particles hit the foil. Most passed almost straight through, but a few were deflected at large angles, and some came almost straight back! Rutherford later described this unexpected result by saying, “It was about as credible as if you had fired a 15-inch [artillery] shell at a piece of paper and it came back and hit you.” The only way for Rutherford and his colleagues to account for their observations was to propose a new model of the atom, in which all of the positive charge and most of the mass of the atom is concentrated in a very small volume. Rutherford called this tiny core of the atom the nucleus. The electrons occupy the rest of the space in the atom. From their results Rutherford, Geiger, and Marsden calculated that the nucleus of a gold atom had a positive charge in the range of 100 20 and a radius of about 10 12 cm. The currently accepted values are 79 for the charge and about 10 13 cm for the radius. 2.2 Atomic Number and Atomic Mass 67 See the General ChemistryNow CD-ROM or website: • Screen 2.10 The Nucleus of the Atom, for an exercise on an experiment investigating the properties of the nuclei Exercise 2.1—Describing Atoms We know now that the radius of the nucleus is about 0.001 pm, and the radius of an atom is approximately 100 pm. If an atom were a macroscopic object with a radius of 100 m, it would approximately fill a small football stadium. What would be the radius of the nucleus of such an atom? Can you think of an object that is about that size? 2.2—Atomic Number and Atomic Mass Atomic Number All atoms of the same element have the same number of protons in the nucleus. Hydrogen is the simplest element, with one nuclear proton. All helium atoms have two protons, all lithium atoms have three protons, and all beryllium atoms have four protons. The number of protons in the nucleus of an element is its atomic number, generally given the symbol Z. Currently known elements are listed in the periodic table inside the front cover of this book. The integer number at the top of the box for each element is its atomic number. A sodium atom, for example, has an atomic number of 11, so its nucleus contains 11 protons. A uranium atom has 92 nuclear protons and Z 92. ■ The Periodic Table Entry for Copper Copper 29 Atomic number Symbol Atomic weight Cu 63.546 Relative Atomic Mass and the Atomic Mass Unit What is the mass of an atom? Chemists in the 18th and 19th centuries recognized that careful experiments could give relative atomic masses. For example, the mass of an oxygen atom was found to be 1.33 times the mass of a carbon atom, and a calcium atom has 2.5 times the mass of an oxygen atom. Chemistry in the 21st century still uses a system of relative masses. After trying several standards, scientists settled on the current one: A carbon atom having six protons and six neutrons in the nucleus is assigned a mass value of exactly 12.000. An oxygen atom having eight protons and eight neutrons has 1.3333 times the mass of carbon, so it has a relative mass of 16.000. Masses of atoms of other elements have been assigned in a similar manner. Masses of fundamental atomic particles are often expressed in atomic mass units (u). One atomic mass unit, 1 u, is one-twelfth of the mass of an atom of carbon with six protons and six neutrons. Thus, such a carbon atom has a mass of 12.000 u. The atomic mass unit can be related to other units of mass using a conversion factor; that is, 1 u 1.661 10 24 g. Mass Number Protons and neutrons have masses very close to 1 u (Table 2.1). The electron, in contrast, has a mass only about 1/2000 of this value. Because proton and neutron masses are so close to 1 u, the approximate mass of an atom can be estimated if the 68 Chapter 2 Atoms and Elements Table 2.1 Properties of Subatomic Particles* Mass Particle Electron Proton Neutron Grams 9.109383 1.672622 1.674927 10 10 10 28 24 24 Atomic Mass Units 0.0005485799 1.007276 1.008665 Charge 1 1 0 Symbol 0 1e 1 1p 1 0n or e or p or n0 * These values and others in the book are taken from the National Institute of Standards and Technology website at number of neutrons and protons is known. The sum of the number of protons and neutrons for an atom is called its mass number and is given the symbol A. A mass number number of protons number of neutrons For example, a sodium atom, which has 11 protons and 12 neutrons in its nucleus, has a mass number of A 23. The most common atom of uranium has 92 protons and 146 neutrons, and a mass number of A 238. Using this information, we often symbolize atoms with the notation Mass number S A Atomic number S Z X d Element symbol The subscript Z is optional because the element symbol tells us what the atomic number must be. For example, the atoms described previously have the symbols 23 238 23 238 U. In words, we say “sodium-23” or “uranium-238.” 11Na or 92U, or just Na or See the General ChemistryNow CD-ROM or website: • Screen 2.11 Summary of Atomic Composition, for a tutorial on the notation for symbolizing atoms Example 2.1—Atomic Composition Problem What is the composition of an atom of phosphorus with 16 neutrons? What is its mass number? What is the symbol for such an atom? If the atom has an actual mass of 30.9738 u, what is its mass in grams? Strategy All P atoms have the same number of protons, 15, which is given by the atomic number (see the periodic table inside the front cover of this book). The mass number is the sum of the number of protons and neutrons. The mass of the atom in grams can be obtained from the mass in atomic mass units using the conversion factor 1 u 1.661 10 24 g. Solution A phosphorus atom has 15 protons and, because it is electrically neutral, also has 15 electrons. Mass number number of protons number of neutrons 15 16 31 2.3 Isotopes 69 The atom’s complete symbol is Mass of one 31P atom 1g2 31 15P. 130.9738 u2 11.661 10 24 g/u2 5.145 10 23 g Exercise 2.2—Atomic Composition (a) What is the mass number of an iron atom with 30 neutrons? (b) A nickel atom with 32 neutrons has a mass of 59.930788 u. What is its mass in grams? (c) How many protons, neutrons, and electrons are in a 64Zn atom? 2.3—Isotopes In only a few instances (for example, aluminum, fluorine, and phosphorus) do all atoms in a naturally occurring sample of a given element have the same mass. Most elements consist of atoms having several different mass numbers. For example, there are two kinds of boron atoms, one with a mass of about 10 u (10B) and a second with a mass of about 11 u (11B). Atoms of tin can have any of 10 different masses. Atoms with the same atomic number but different mass numbers are called isotopes. All atoms of an element have the same number of protons—five in the case of boron. This means that, to have different masses, isotopes must have different numbers of neutrons. The nucleus of a 10B atom (Z 5) contains five protons and five neutrons, whereas the nucleus of a 11B atom contains five protons and six neutrons. Scientists often refer to a particular isotope by giving its mass number (for example, uranium-238, 238U), but the isotopes of hydrogen are so important that they have special names and symbols. All hydrogen atoms have one proton. When that is the only nuclear particle, the isotope is called protium, or just “hydrogen.” The isotope of hydrogen with one neutron, 2H, is called deuterium, or “heavy hydrogen” 1 (symbol D). The nucleus of radioactive hydrogen-3, 3H, or tritium (symbol T), 1 contains one proton and two neutrons. The substitution of one isotope of an element for another isotope of the same element in a compound sometimes has an interesting effect (Figure 2.7). This is especially true when deuterium is substituted for hydrogen because the mass of deuterium is double that of hydrogen. Solid H2O Liquid H2O Solid D2O Isotope Abundance A sample of water from a stream or lake will consist almost entirely of H2O where the H atoms are the 1H isotope. A few molecules, however, will have deuterium (2H) substituted for 1H. We can predict this outcome because we know that 99.985% of all hydrogen atoms on earth are 1H atoms. That is, the percent abundance of 1H atoms is 99.985%. Percent abundance number of atoms of a given isotope total number of atoms of all isotopes of that element 100% (2.1) The remainder of naturally occurring hydrogen is deuterium, whose abundance is only 0.015% of the total hydrogen atoms. Tritium, the radioactive 3H isotope, does not occur naturally. Charles D. Winters Figure 2.7 Ice made from “heavy water.” Water containing ordinary hydrogen (1H, protium) forms a solid that is less 1 dense (d 0.917 g/cm3 at 0 °C) than liquid H2O (d 0.997 g/cm3 at 25 °C) and so floats in the liquid. (Water is unique in this regard. The solid phase of virtually all other substances sinks in the liquid phase of that substance.) Similarly, “heavy ice” (D2O, deuterium oxide) floats in “heavy water.” D2O-ice is denser than H2O, however, so cubes made of D2O sink in liquid H2O. 70 Chapter 2 Atoms and Elements Consider the two isotopes of boron. The boron-10 isotope has an abundance of 19.91%; the abundance of boron-11 is 80.09%. Thus, if you could count out 10,000 boron atoms from an “average” natural sample, 1991 of them would be boron-10 atoms and 8009 of them would be boron-11 atoms. Example 2.2—Isotopes Problem Silver has two isotopes, one with 60 neutrons (percent abundance 51.839%) and the other with 62 neutrons. What are the mass numbers and symbols of these isotopes? What is the percent abundance of the isotope with 62 neutrons? Strategy Recall that the mass number is the sum of the number of protons and neutrons. The symbol is written as AX, where X is the one or two-letter element symbol. The percent abunZ dances of all isotopes must add up to 100%. Solution Silver has an atomic number of 47, so each silver atom has 47 protons in its nucleus. The two isotopes, therefore, have mass numbers of 107 and 109. ■ Atomic Masses of Some Isotopes Isotope 1, with 47 protons and 60 neutrons A 47 protons 60 neutrons 107 Atom 4 Mass (u) 4.0092603 13.003355 15.994915 57.935348 59.930791 78.918338 80.916291 196.966552 238.050783 He C O Ni Ni Br Br Au U 13 16 58 60 79 81 Isotope 2, with 47 protons and 62 neutrons A 47 protons 62 neutrons 109 The first isotope has a symbol 107Ag and the second is 109 Ag . 47 47 Silver-107 has a percent abundance of 51.839%. Therefore, the percent abundance of silver109 is % Abundance of 109Ag 100.000% 51.839% 48.161% 197 238 Exercise 2.3—Isotopes (a) Argon has three isotopes with 18, 20, and 22 neutrons, respectively. What are the mass numbers and symbols of these three isotopes? (b) Gallium has two isotopes: 69Ga and 71Ga. How many protons and neutrons are in the nuclei of each of these isotopes? If the abundance of 69Ga is 60.1%, what is the abundance of 71Ga? Determining Atomic Mass and Isotope Abundance The masses of isotopes and their percent abundances are determined experimentally using a mass spectrometer (Figure 2.8). A gaseous sample of an element is introduced into the evacuated chamber of the spectrometer, and the molecules or atoms of the sample are converted to charged particles (ions). A beam of these ions is subjected to a magnetic field, which causes the paths of the ions to be deflected. The extent of deflection depends on particle mass: The less massive ions are deflected more, and the more massive ions are deflected less. The ions, now separated by mass, are detected at the end of the chamber. In early experiments, ions were detected using photographic film, but modern instruments measure the electric current in a detector. The darkness of a spot on photographic film, or the amount of 2.3 Isotopes 71 VA PO RIZAT ION IONIZAT ION ACCELERAT ION D EFL EC T IO N D ET EC T IO N Electron gun e e e e e e e e e Magnet Heavy ions are deflected too little. 20Ne A mass spectrum is a plot of the relative abundance of the charged particles versus the ratio of mass/charge (m/z). Relative Abundance 22Ne 21Ne Gas inlet Repeller Electron trap plate Accelerating plates Magnet Light ions are deflected too much. To vacuum pump To mass analyzer 100 80 60 40 20 0 20 21 22 Detector m/z 1. A sample is introduced as a vapor into the ionization chamber. There it is bombarded with highenergy electrons that strip electrons from the atoms or molecules of the sample. 2. The resulting positive particles are accelerated by a series of negatively charged accelerator plates into an analyzing chamber. 3. This chamber is in a magnetic field, which is perpendicular to the direction of the beam of charged particles. The magnetic field causes the beam to curve. The radius of curvature depends on the mass and charge of the particles (as well as the accelerating voltage and strength of the magnetic field). 4. Here particles of 21Ne are focused on the detector, whereas beams of ions of 20Ne and 22Ne (of lighter or heavier mass) experience greater and lesser curvature, respectively, and so fail to be detected. By changing the magnetic field, a beam of charged particles of different mass can be focused on the detector, and a spectrum of masses is observed. Active Figure 2.8 Mass spectrometer. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. current measured, is related to the number of ions of a particular mass and hence to the abundance of the ion. The mass-to-charge ratio for the ions can also be determined from the extent of curvature in the ion’s path to the detector. Knowing that most of the ions within the spectrometer have a 1 charge allows us to derive a value for mass. Chemists using modern instruments can measure isotopic masses with as many as nine significant figures. A Closer Look Atomic Mass and the Mass Defect You might expect that the mass of a deuterium nucleus, 2H, would be the sum of the masses of its constituent particles, a proton and a neutron. 1 1p However, the mass of 2H is less than the sum of its constituents! Difference in mass, ¢ m mass of product total mass of reactants 2.01355 u 2.015941 u 0.00239 u This “missing mass” is equated to energy, the binding energy for the nucleus. The binding energy can be calculated from Einstein’s equation that relates the mass, 11.007276 u2 1 0n 11.008665 u2 S 2H 12.01355 u2 1 m, to energy, E (E mc2, where c is the velocity of light). Although the mass loss on forming an atomic nucleus from its constituent protons and neutrons seems small, the energy equivalent is enormous. In fact, it is the mass loss from fusing protons into helium nuclei on the sun that provides the energy for life on earth. (See the story, Stardust, on page 58 and see Chapter 23 for more details.) 72 Chapter 2 Atoms and Elements Except for carbon-12, whose mass is defined to be exactly 12u, isotopic masses do not have integer values. However, the isotopic masses are always very close to the mass numbers for the isotope. For example, the mass of an atom of boron-11 (11B, 5 protons and 6 neutrons) is 11.0093 u, and the mass of an atom of iron-58 (58Fe, 26 protons and 32 neutrons) is 57.9333 u. Note also that the masses of individual isotopes are always slightly less than the sum of the masses of the protons, neutrons, and electrons making up the atom. This mass difference, called the “mass defect,” is related to the energy binding the particles of the nucleus together. (See A Closer Look: Atomic Mass and the Mass Defect.) 2.4—Atomic Weight Because every sample of boron has some atoms with a mass of 10.0129 u and others with a mass of 11.0093 u, the average atomic mass must be somewhere between these values. The atomic weight is the average weight of a representative sample of atoms. For boron, the atomic weight is 10.81. In general, the atomic weight of an element can be calculated using the equation a % abundance isotope 1 b 1mass of isotope 12 100 p Atomic weight (2.2) % abundance isotope 2 b 1mass of isotope 22 a 100 ■ Atomic Weight and Units Values of atomic weight are relative to the mass of the carbon-12 isotope and so are unitless numbers. For boron with two isotopes (10B, 19.91% abundant; 11B, 80.09% abundant ), we find Atomic weight a 19.91 b 100 10.0129 a 80.09 b 100 11.0093 10.81 Equation 2.2 gives an average, weighted in terms of the abundance of each isotope for the element. As illustrated by the data in Table 2.2, the atomic weight of an element is always closer to the mass of the more abundant isotope or isotopes. ■ Fractional Abundance The percent abundance of an isotope divided by 100 is called its fractional abundance. Table 2.2 Element Hydrogen Isotope Abundance and Atomic Weight Symbol H D* T† Atomic Weight 1.00794 Mass Number 1 2 3 10.811 20.1797 10 11 20 21 22 Isotopic Mass (u) 1.0078 2.0141 3.0161 10.0129 11.0093 19.9924 20.9938 21.9914 23.9850 24.9858 25.9826 Natural Abundance (%) 99.985 0.015 0 19.91 80.09 90.48 0.27 9.25 78.99 10.00 11.01 Boron Neon B Ne Magnesium Mg 24.305 24 25 26 *D deuterium; †T tritium, radioactive. 2.5 Atoms and the Mole ■ The Periodic Table Entry for Copper Copper 29 Atomic number Symbol Atomic weight 73 The atomic weight of each stable element has been determined experimentally, and these numbers appear in the periodic table inside the front cover of this book. In the periodic table, each element’s box contains the atomic number, the element symbol, and the atomic weight. For unstable (radioactive) elements, the atomic mass or mass number of the most stable isotope is given in parentheses. Cu 63.546 Example 2.3—Calculating Atomic Weight from Isotope Abundance Problem Bromine (used to make silver bromide, the important component of photographic film) has two naturally occurring isotopes. One has a mass of 78.918338 u and an abundance of 50.69%. The other isotope, of mass 80.916291 u, has an abundance of 49.31%. Calculate the atomic weight of bromine. Strategy The atomic weight of any element is the weighted average of the masses of the isotopes in a representative sample. To calculate the atomic weight, multiply the mass of each isotope by its percent abundance divided by 100 (Equation 2.2). (See the General ChemistryNow Screen 2.13 Atomic Mass, for a tutorial on calculating atomic weight from isotope abundance.) Solution Average atomic mass of bromine 150.69/1002178.9183382 79.90 149.31/1002180.9162912 Exercise 2.4—Calculating Atomic Weight Verify that the atomic weight of chlorine is 35.45, given the following information: 35 37 Cl mass Cl mass 34.96885; percent abundance 36.96590; percent abundance 75.77% 24.23% 2.5—Atoms and the Mole One of the most exciting aspects of chemical research is the discovery of some new substance, and part of this process of discovery involves quantitative experiments. When two chemicals react with each other, we want to know how many atoms of each are used so that formulas can be established for the reaction’s products. To do so, we need some method of counting atoms. That is, we must discover a way of connecting the macroscopic world, the world we can see, with the particulate world of atoms, molecules, and ions. The solution to this problem is to define a convenient amount of matter that contains a known number of particles. That chemical unit is the mole. The mole (abbreviated mol ) is the SI base unit for measuring an amount of a substance (see Table 1.2) and is defined as follows: A mole is the amount of a substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 g of the carbon-12 isotope. The key to understanding the concept of the mole is recognizing that one mole always contains the same number of particles, no matter what the substance. One mole of sodium ■ The “Mole” The term “mole” was introduced about 1896 by Wilhelm Ostwald (1853–1932), who derived the term from the Latin word moles, meaning a “heap” or “pile.” 74 Chapter 2 Atoms and Elements Historical Perspectives Amedeo Avogadro and His Number Amedeo Avogadro, Conte di Quaregna (1776–1856) was an Italian nobleman and a lawyer. In about 1800, he turned to science, becoming the first professor of mathematical physics in Italy. Avogadro did not propose the notion of a fixed number of particles in a chemical unit. Rather, the number was named in his honor because he had performed experiments in the 19th century that laid the groundwork for the concept. Just how large is Avogadro’s number? One mole of unpopped popcorn kernels would cover the continental United States to a depth of about 9 miles. One mole of pennies divided equally among every man, woman, and child in the United States would allow each person to pay off the national debt ($5.7 trillion or 5.7 1012 dollars) and there would still be $15 trillion left over! Is Avogadro’s number a unique value like p? No. It is fixed by the definition of the mole as exactly 12 g of carbon-12. If one mole of carbon were defined to have some other mass, then Avogadro’s number would have a different value. Photo: E. F. Smith Collection/Van Pelt Library/ University of Pennsylvania contains the same number of atoms as one mole of iron. How many particles? Many, many experiments over the years have established that number as 1 mole 6.0221415 1023 particles This value is known as Avogadro’s number in honor of Amedeo Avogadro, an Italian lawyer and physicist (1776–1856) who conceived the basic idea (but never determined the number). ■ The Difference Between “Amount” and “Quantity” The terms “amount” and “quantity” are used in a specific sense by chemists. The amount of a substance is the number of moles of that substance. Quantity refers to the mass of the substance. See W. G. Davies and J. W. Moore: Journal of Chemical Education, Vol. 57, p. 303, 1980. See also on the Internet. Molar Mass The mass in grams of one mole of any element (6.0221415 1023 atoms of that element ) is the molar mass of that element. Molar mass is conventionally abbreviated with a capital italicized M and has units of grams per mole (g/mol ). An element’s molar mass is the amount in grams numerically equal to its atomic weight. Using sodium and lead as examples, Molar mass of sodium 1Na2 Molar mass of lead (Pb) mass of 1.000 mol of Na atoms 22.99 g/mol mass of 6.022 1023 Na atoms mass of 1.000 mol of Pb atoms 207.2 g/mol mass of 6.022 1023 Pb atoms Figure 2.9 shows the relative sizes of a mole of some common elements. Although each of these “piles of atoms” has a different volume and different mass, each contains 6.022 1023 atoms. The mole concept is the cornerstone of quantitative chemistry. It is essential to be able to convert from moles to mass and from mass to moles. Dimensional analysis, which is described in Section 1.8 (pages 41–43), shows that this can be done in the following way: 2.5 Atoms and the Mole 75 Figure 2.9 One-mole of common elements. (left to right) Sulfur powder, magnesium chips, tin, and silicon. (above) Copper beads. Copper 63.546 g Charles D. Winters Sulfur 32.066 g Magnesium 24.305 g Tin 118.71 g Silicon 28.086 g MASS · MOLES CONVERSION Moles to Mass Mass to Moles Moles grams 1 mol c molar mass grams Grams 1 mol grams c 1/molar mass moles For example, what mass, in grams, is represented by 0.35 mol of aluminum? Using the molar mass of aluminum (27.0 g/mol ), you can determine that 0.35 mol of Al has a mass of 9.5 g. 0.35 mol Al 27.0 g Al 1 mol Al 9.5 g Al Molar masses are generally known to at least four significant figures. The convention followed in calculations in this book is to use a value of the molar mass with one more significant figure than in any other number in the problem. For example, if you weigh out 16.5 g of carbon, you use 12.01 g/mol for the molar mass of C to find the amount of carbon present. 1 mol C 12.01 g C 16.5 g C 1.37 mol C c Note that four significant figures are used in the molar mass, but there are three in the sample mass. Using one more significant figure for the molar mass means the accuracy of this value will not affect the accuracy of the result. 76 Chapter 2 Atoms and Elements See the General ChemistryNow CD-ROM or website: • Screen 2.14 the Mole, for a tutorial on moles and atoms conversion • Screen 2.15 Moles and Molar Mass of the Elements, for two tutorials on molar mass conversion Charles D. Winters Example 2.4—Mass, Moles, and Atoms Problem Consider two elements in the same vertical column of the periodic table, lead and tin. (a) What mass of lead, in grams, is equivalent to 2.50 mol of lead (Pb, atomic number 82)? (b) What amount of tin, in moles, is represented by 36.5 g of tin (Sn, atomic number How many atoms of tin are in the sample? 50)? Lead. A 150-mL beaker containing 2.50 mol or 518 g of lead. Strategy The molar masses of lead (207.2 g/mol) and tin (118.7 g/mol) are required and can be found in the periodic table inside the front cover of this book. Avogadro’s number is needed to convert the amount of each element to number of atoms. Solution (a) Convert the amount of lead in moles to mass in grams. 2.50 mol Pb 207.2 g 1 mol Pb 518 g Pb Charles D. Winters (b) First convert the mass of tin to the amount in moles. 36.5 g Sn 1 mol Sn 118.7 g Sn 0.308 mol Sn Tin. A sample of tin having a mass of 36.5 g (1.85 1023 atoms). Finally, use Avogadro’s number to find the number of atoms in the sample. 0.308 mol Sn 1.85 6.022 1023 atoms Sn 1 mol Sn 1023 atoms Sn Example 2.5—Mole Calculation Problem The graduated cylinder in the photo contains 32.0 cm3 of mercury. If the density of mercury at 25 °C is 13.534 g/cm3, what amount of mercury, in moles, is in the cylinder? Charles D. Winters Strategy Volume and moles of mercury are not directly connected. You must first use the density of mercury to find the mass of the metal and then use this value with the molar mass of mercury to calculate the amount in moles. Volume 1cm3 2 density 1g/cm3 2 mass of mercury 1g2 Mercury. A graduated cylinder containing 32.0 cm3 of mercury. This is equivalent to 433 g or 2.16 mol of mercury. Amount of mercury 1mol2 mass of mercury 1g2 11/molar mass21mol/g2 2.6 The Periodic Table 77 Solution Combining the volume and density gives the mass of the mercury. 32.0 cm3 13.534 g Hg 1 cm3 433 g Hg Finally, the amount of mercury can be calculated from its mass and molar mass. 433 g Hg 1 mol Hg 200.6 g Hg 2.16 mol Hg Example 2.6—Mass of an Atom Problem What is the average mass of an atom of platinum (Pt)? Strategy The mass of one mole of platinum is 195.08 g. Each mole contains Avogadro’s number of atoms. Solution Here we divide the mass of a mole by the number of objects in that unit. 195.08 g Pt 1 mol Pt 1 mol Pt 6.02214 1023 atoms Pt 3.2394 10 22 g Pt 1 atom Pt Comment Notice that the units “mol Pt” cancel and leave an answer with units of g/atom. Exercise 2.5—Mass/Mole Conversions (a) What is the mass, in grams, of 1.5 mol of silicon? (b) What amount (moles) of sulfur is represented by 454 g? How many atoms? (c) What is the average mass of one sulfur atom? Exercise 2.6—Atoms The density of gold, Au, is 19.32 g/cm3. What is the volume (in cubic centimeters) of a piece of gold that contains 2.6 x 1024 atoms? If the piece of metal is a square with a thickness of 0.10 cm, what is the length (in centimeters) of one side of the piece? 2.6—The Periodic Table The periodic table of elements is one of the most useful tools in chemistry. Not only does it contain a wealth of information, but it can also be used to organize many of the ideas of chemistry. It is important that you become familiar with its main features and terminology. Features of the Periodic Table The main organizational features of the periodic table are the following: • Elements are arranged so that those with similar chemical and physical properties lie in vertical columns called groups or families. The periodic table commonly used in the United States has groups numbered 1 through 8, with each number followed by a letter: A or B. The A groups are often called the main group elements and the B groups are the transition elements. 78 Group 1A Lithium — Li (top) Potassium — K (bottom) Chapter 2 Atoms and Elements Group 2B Zinc — Zn (top) Mercury Hg (bottom) Group 2A Magnesium — Mg Transition Metals Titanium — Ti, Vanadium — V, Chromium — Cr, Manganese — Mn, Iron — Fe, Cobalt — Co, Nickel — Ni, Copper — Cu 1A 1 2 3 4 5 6 7 (3A) (4A) (5A) 2A 3A 4A 5A 6A 7A 8A Li Mg K 3B 4B 5B 6B 7B 8B 1B 2B B C N P S Se Br Ne Al Si Ti V Cr Mn Fe Co Ni Cu Zn Sn Hg Pb (6A) (7A) Group 8A, Noble Gases Neon — Ne Group 4A Carbon — C (top) Lead — Pb (left) Silicon — Si (right) Tin — Sn (bottom) Group 3A Boron — B (top) Aluminum — Al (bottom) Group 5A Nitrogen — N2 (top) Phosphorus — P (bottom) Group 6A Sulfur — S (top) Selenium — Se (bottom) Group 7A Bromine — Br Active Figure 2.10 Some of the 116 known elements. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. Photos: Charles D. Winters 2.6 The Periodic Table 1 2 3 4 5 6 7 Periods 1A 2A 4B 6B 3B 5B 7B 8B 2B 1B 3A 4A 5A 6A 79 • The horizontal rows of the table are called periods, and they are numbered beginning with 1 for the period containing only H and He. For example, sodium, Na, in Group 1A, is the first element in the third period. Mercury, Hg, in Group 2B, is in the sixth period (or sixth row). The periodic table can be divided into several regions according to the properties of the elements. On the table inside the front cover of this book, elements that behave as metals are indicated in purple, those that are nonmetals are indicated in yellow, and elements called metalloids appear in green. Elements gradually become less metallic as one moves from left to right across a period, and the metalloids lie along the metal–nonmetal boundary. Some elements are shown in Figure 2.10. You are probably familiar with many properties of metals from everyday experience (Figure 2.11a). Metals are solids (except for mercury), can conduct electricity, are usually ductile (can be drawn into wires) and malleable (can be rolled into sheets), and can form alloys (solutions of one or more metals in another metal ). Iron (Fe) and aluminum (Al ) are used in automobile parts because of their ductility, malleability, and low cost relative to other metals. Copper (Cu) is used in electric wiring because it conducts electricity better than most other metals. Chromium (Cr) is plated onto automobile parts, not only because its metallic luster makes cars look better but also because chrome-plating protects the underlying metal from reacting with oxygen in the air. The nonmetals lie to the right of a diagonal line that stretches from B to Te in the periodic table and have a wide variety of properties. Some are solids (carbon, sulfur, phosphorus, and iodine). Four elements are gases at room temperature (oxygen, nitrogen, fluorine, and chlorine). One element, bromine, is a liquid at room temperature (Figure 2.11b). With the exception of carbon in the form of graphite, nonmetals do not conduct electricity, which is one of the main features that distinguishes them from metals. Some of the elements next to the diagonal line from boron (B) to tellurium (Te) have properties that make them difficult to classify as metals or nonmetals. Chemists call them metalloids or, sometimes, semimetals (Figure 2.11c). You should 8A 7A Groups or Families ■ Two Ways to Designate Groups One way to designate periodic groups is to number them 1 through 18 from left to right. This method is generally used outside the United States. The system predominant in the United States labels main group elements as Groups 1A–8A and transition elements as Groups 1B–8B. This book uses the A/B system. Magnesium, Mg Copper, Cu Bromine, Br2 Iodine, I2 Forms of silicon Charles D. Winters Aluminum, Al (a) Metals (b) Nonmetals (c) Metalloids Figure 2.11 Representative elements. (a) Magnesium, aluminum, and copper are metals. All can be drawn into wires and conduct electricity. (b) Only 15 or so elements can be classified as nonmetals. Here are orange liquid bromine and purple solid iodine. (c) Only 6 elements are generally classified as metalloids or semimetals. This photograph shows solid silicon in various forms, including a wafer that holds printed electronic circuits. 80 Chapter 2 Atoms and Elements Main Group Metals Transition Metals Metalloids Nonmetals know, however, that chemists often disagree about what a metalloid is as well as which elements fit into this category. We will define a metalloid as an element that has some of the physical characteristics of a metal but some of the chemical characteristics of a nonmetal; we include only B, Si, Ge, As, Sb, and Te in this category. This distinction reflects the ambiguity in the behavior of these elements. Antimony (Sb), for example, conducts electricity as well as many elements that are true metals. Its chemistry, however, resembles that of a nonmetal such as phosphorus. Developing the Periodic Table Although the arrangement of elements in the periodic table can now be understood on the basis of atomic structure [ Chapter 8], the table was originally developed from many, many experimental observations of the chemical and physical properties of elements and is the result of the ideas of a number of chemists in the 18th and 19th centuries. In 1869, at the University of St. Petersburg in Russia, Dmitri Ivanovitch Mendeleev (1834–1907) was pondering the properties of the elements as he wrote a textbook on chemistry. On studying the chemical and physical properties of the elements, he realized that, if the elements were arranged in order of increasing atomic mass, elements with similar properties appeared in a regular pattern. That is, he saw a periodicity or periodic repetition of the properties of elements. Mendeleev organized the known elements into a table by lining them up in a horizontal row in order of increasing atomic mass. Every time he came to an element with properties similar to one already in the row, he started a new row. For example, the elements Li, Be, B, C, N, O, and F were in a row. Sodium was the next element then known; because its properties closely resembled those of Li, Mendeleev started a new row. The columns, then, contained elements such as Li, Na, and K with similar properties. An important feature of Mendeleev’s table—and a mark of his genius—was that he left an empty space in a column when an element was not known but should exist and have properties similar to the element above it in his table. He deduced that these spaces would be filled by undiscovered elements. For example, a space was left between Si (silicon) and Sn (tin) in what is now Group 4A. Based on the progression of properties in this group, Mendeleev was able to predict the properties of this missing element. With the discovery of germanium in 1886, Mendeleev’s prediction was confirmed. In Mendeleev’s table the elements were ordered by increasing mass. A glance at a modern table, however, shows that, on this basis, Ni and Co, Ar and K, and Te and I, should be reversed. Mendeleev assumed the atomic masses known at that time were inaccurate—not a bad assumption based on the analytical methods then in use. In fact, his order was correct and what was wrong was his assumption that element properties were a function of their mass. In 1913 H. G. J. Moseley (1887–1915), a young English scientist working with Ernest Rutherford, corrected Mendeleev’s assumption. Moseley was doing experiments in which he bombarded many different metals with electrons in a cathode-ray tube (Figure 2.3) and examined the x-rays emitted in the process. In seeking some order in his data, he realized that the wavelength of the x-rays emitted by a given element were related in a precise manner to the atomic number of the element. Indeed, chemists quickly recognized that organizing the elements in a table by increasing atomic number corrected the inconsistencies in the Mendeleev table. The law of chemical periodicity is now stated as “the properties of the elements are periodic functions of atomic number.” ■ About the Periodic Table For more information on the periodic table, the central icon of chemistry, we recommend the following: • The American Chemical Society has a description of every element on its website at • J. Emsley: Nature’s Building Blocks—An A–Z Guide to the Elements, New York, Oxford University Press, 2001. • O. Sacks: Uncle Tungsten—Memories of a Chemical Boyhood, New York, Alfred A. Knopf, 2001. ■ Placing H in the Periodic Table Where to place H? Tables often show it in Group 1A even though it is clearly not an alkali metal. However, in its reactions it forms a 1+ ion just like the alkali metals. For this reason, H is often placed in Group 1A. 2.6 The Periodic Table 81 Historical Perspectives Periodic Table In his book Nature’s Building Blocks (p. 527, New York, Oxford University Press, 2001), John Emsley tells us that “As long as chemistry is studied, there will be a periodic table. Even if some day we communicate with another part of the Universe, we can be sure that one thing both cultures will have in common is an ordered system of the elements that will be instantly recognizable by both intelligent life forms.” The person credited with organizing the elements into a periodic table is Dmitri Mendeleev. However, other chemists had long recognized that groups of elements shared similar properties. In 1829 Johann Dobereiner (1780–1849) announced the Law of Triads. He showed that there were groups of three elements (triads), in which 80 70 Atomic volume (mL/mol) 60 50 40 30 20 10 0 0 Ne Ar the middle element had an atomic weight that was the average of the other two. One such triad consisted of Li, Na, and K; another was made up of Cl, Br, and I. Perhaps the first revelation of the periodicity of the elements was published by a French geologist, A. E. Béguyer de Chancourtois (1820–1886), in 1862. He listed the elements on a paper tape, and, according to Emsley, “then wound this, spiral like around a cylinder. The cylinder’s surface was divided into 16 parts, based on the atomic weight of oxygen. De Chancourtois noted that certain triads came together down the cylinder, such as the alkali metals.” He called his model the “telluric screw.” Another attempt at organizing the elements was proposed by John Newlands (1837–1898) in 1864. His “Law of Octaves” proposed that there was a periodic similarity every eight elements, just as the musi- cal scale repeats every eighth note. Unfortunately, his proposal was ridiculed at the time. Julius Lothar Meyer (1830–1895) came closer than any other to discovering the periodic table. He drew a graph of atomic volumes of elements plotted against their atomic weight. This clearly showed a periodic rise and fall in atomic volume on moving across what we now call the periods of the table. Before publishing the paper, Meyer passed it along to a colleague for comment. His colleague was slow to return the paper, and, unfortunately for Meyer, Mendeleev’s paper was published in the interim. Because chemists quickly recognized the importance of Mendeleev’s paper, Meyer was not given the recognition he perhaps deserves. An essay on Mendeleev and his life appears at the beginning of Chapter 8 (pages 332–3). Xe Kr 50 100 150 200 250 Atomic weight (g/mol) Atomic volume plot. Julius Lothar Meyer (1830–1895) illustrated the periodicity of the elements in 1868 by plotting atomic volume against atomic weight. (This plot uses current data.) Source: C. N. Singman: Journal of Chemical Education, Vol. 61, p. 137, 1984. See the General ChemistryNow CD-ROM or website: • Screen 2.16 The Periodic Table, for an exercise on the periodic table organization 82 Chapter 2 Atoms and Elements 2.7—An Overview of the Elements, Their Chemistry, and the Periodic Table The vertical columns, or groups, of the periodic table contain elements having similar chemical and physical properties, and several groups of elements have distinctive names that are useful to know. Group 1A, Alkali Metals: Li, Na, K, Rb, Cs, Fr ■ Alkali and Alkaline The word “alkali” comes from the Arabic language; ancient Arabian chemists discovered that ashes of certain plants, which they called al-qali, gave water solutions that felt slippery and burned the skin. These ashes contain compounds of Group 1A elements that produce alkaline (basic) solutions. Elements in the leftmost column, Group 1A, are known as the alkali metals. All are metals and are solids at room temperature. The metals of Group 1A are all reactive. For example, they react with water to produce hydrogen and alkaline solutions (Figure 2.12). Because of their reactivity, these metals are found in nature only combined in compounds, such as NaCl (Figure 1.7)—never as the free element. Group 2A, Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra The second group in the periodic table, Group 2A, is composed entirely of metals that occur naturally only in compounds (Figure 2.13). Except for beryllium (Be), these elements also react with water to produce alkaline solutions, and most of their oxides (such as lime, CaO) form alkaline solutions; hence, they are known as the alkaline earth metals. Magnesium (Mg) and calcium (Ca) are the seventh and fifth most abundant elements in the earth’s crust, respectively (Table 2.3). Calcium is one of the important elements in teeth and bones, and it occurs naturally in vast limestone deposits. Calcium carbonate (CaCO3) is the chief constituent of limestone and of corals, sea shells, marble, and chalk (see Figure 2.13b). Radium (Ra), the heaviest alkaline earth element, is radioactive and is used to treat some cancers by radiation. Table 2.3 The Ten Most Abundant Elements in the Earth’s Crust Rank 1 2 3 4 5 6 7 8 9 10 * ppm Element Oxygen Silicon Aluminum Iron Calcium Sodium Magnesium Potassium Titanium Hydrogen g per 1000 kg. Abundance (ppm)* 474,000 277,000 82,000 41,000 41,000 23,000 23,000 21,000 5,600 1,520 Group 3A: B, Al, Ga, In, Tl Group 3A contains one element of great importance, aluminum (Figure 2.14). This element and three others (gallium, indium, and thallium) are metals, whereas boron (B) is a metalloid. Aluminum (Al ) is the most abundant metal in the earth’s crust at 8.2% by mass. It is exceeded in abundance only by the nonmetals oxygen and silicon. These three elements are found combined in clays and other common Figure 2.12 Alkali metals. (a) Cutting a bar of sodium with a knife is about like cutting a stick of cold butter. (b) When an alkali metal such as potassium is treated with water, a vigorous reaction occurs, giving an alkaline solution and hydrogen gas, which burns in air. See also Figure 1.7, the reaction of sodium with chlorine. (a) Cutting sodium. (b) Potassium reacts with water. Charles D. Winters 2.7 An Overview of the Elements, Their Chemistry, and the Periodic Table 83 Figure 2.13 Alkaline earth metals. (a) When heated in air, magnesium burns to give magnesium oxide. The white sparks you see in burning fireworks are burning magnesium. (b) Some common calciumcontaining substances: calcite (the clear crystal); a seashell; limestone; and an overthe-counter remedy for excess stomach acid. a, James Cowlin/Image Enterprises, Phoenix, AZ; b, Charles D. Winters (a) Magnesium and strontium in fireworks. (b) Calcium-containing compounds. minerals. Boron occurs in the mineral borax, a compound used as a cleaning agent, antiseptic, and flux for metal work. As a metalloid, boron has a different chemistry than the other elements of Group 3A, all of which are metals. Nonetheless, all form compounds with analogous formulas such as BCl3 and AlCl3, and this similarity marks them as members of the same periodic group. Group 4A: C, Si, Ge, Sn, Pb All of the elements we have described so far, except boron, have been metals. Beginning with Group 4A, however, the groups contain more and more nonmetals. Group 4A includes one nonmetal, carbon (C); two metalloids, silicon (Si) and germanium (Ge), and two metals, tin (Sn) and lead (Pb). Because of the change from nonmetallic to metallic behavior, more variation occurs in the properties of the elements of this group than in most others. Nonetheless, these elements also form 0159/0160 Charles D. Winters (a) Wagons for hauling borax in Death Valley. (b) Aluminum-containing minerals. Figure 2.14 Group 3A elements. (a) Boron is mined as borax, a natural compound used in soap. Borax was mined in Death Valley, California, at the end of the 19th century and was hauled from the mines in wagons drawn by teams of 20 mules. Boron is also a component of borosilicate glass, which is used for laboratory glassware. (b) Aluminum is abundant in the earth’s crust; it is found in all clays and in many minerals and gems. It has many commercial applications as the metal as well as in aluminum sulfate, which is used in water purification. 84 Chapter 2 Atoms and Elements Photos: Charles D. Winters (a) Graphite (b) Diamond (c) Buckyballs Figure 2.15 The allotropes of carbon. (a) Graphite consists of layers of carbon atoms. Each carbon atom is linked to three others to form a sheet of six-member, hexagonal rings. (b) In diamond the carbon atoms are also arranged in six-member rings, but the rings are not flat because each C atom is connected tetrahedrally to four other C atoms. (c) A member of the family called buckminsterfullerenes, C60 is an allotrope of carbon. Sixty carbon atoms are arranged in a spherical cage that resembles a hollow soccer ball. Notice that each six-member ring shares an edge with three other six-member rings and three five-member rings. Chemists call this molecule a “buckyball.” C60 is a black powder; it is shown here in the tip of a pointed glass tube. compounds with analogous formulas (such as CO2, SiO2, GeO2, SnO2, and PbO2), so they are assigned to the same group. Carbon is the basis for the great variety of chemical compounds that make up living things. It is found in the earth’s atmosphere as CO2, on the surface of the earth in carbonates like limestone and coral (see Figure 2.13b), and in coal, petroleum, and natural gas—the fossil fuels. One of the most interesting aspects of the chemistry of the nonmetals is that a particular element can often exist in several different and distinct forms, called allotropes, each having its own properties. Carbon has at least three allotropes, the best known of which are graphite and diamond (Figure 2.15). The flat sheets of carbon atoms in graphite (Figure 2.15a) cling only weakly to one another. One layer can slip easily over another, which explains why graphite is soft, is a good lubricant, and is used in pencil lead. (Pencil “lead” is not the element lead, Pb, but rather a composite of clay and graphite that leaves a trail of graphite on the page as you write.) In diamond each carbon atom is connected to four others at the corners of a tetrahedron, and this pattern extends throughout the solid (see Figure 2.15b). This structure causes diamonds to be extremely hard, denser than graphite (d 3.51 g/cm3 for diamond and d 2.22 g/cm3 for graphite), and chemically less reactive. Because diamonds are not only hard but are also excellent conductors of heat, they are used on the tips of metal- and rock-cutting tools. In the late 1980s another form of carbon was identified as a component of black soot, the stuff that collects when carbon-containing materials are burned in a deficiency of oxygen. This substance is made up of molecules with 60 carbon atoms arranged as a spherical “cage” (Figure 2.15c). You may recognize that the surface is made up of five- and six-member rings and resembles a hollow soccer ball. The 2.7 An Overview of the Elements, Their Chemistry, and the Periodic Table 85 shape reminded its discoverers of an architectural dome invented several decades ago by the American philosopher and engineer, R. Buckminster Fuller. The official name of the allotrope is therefore buckminsterfullerene, and chemists often call these molecules “buckyballs.” Oxides of silicon are the basis of many minerals such as clay, quartz, and beautiful gemstones like amethyst (Figure 2.16). Tin and lead have been known for centuries because they are easily smelted from their ores. Tin alloyed with copper makes bronze, which was used in ancient times in utensils and weapons. Lead has been used in water pipes and paint, even though the element is toxic to humans. Charles D. Winters Figure 2.16 Compounds containing Group 5A: N, P, As, Sb, Bi Nitrogen, which occurs naturally in the form of N2 (Figures 2.10 and 2.17), makes up about three-fourths of earth’s atmosphere. It is also incorporated in biochemically important substances such as chlorophyll, proteins, and DNA. Scientists have long sought ways to make compounds from atmospheric nitrogen, a process referred to as “nitrogen fixation.” Nature accomplishes this transformation easily in plants, but severe conditions (high temperatures, for example) must be used in the laboratory and in industry to cause N2 to react with other elements (such as H2 to make ammonia, NH3, which is widely used as a fertilizer). Phosphorus is also essential to life. For example, it is an important constituent in bones and teeth. The element glows in the dark if it is in the air, and its name, based on Greek words meaning “light-bearing,” reflects this property. This element also has several allotropes, the most important being white (Figure 2.10) and red phosphorus. Both forms of phosphorus are used commercially. White phosphorus ignites spontaneously in air, so it is normally stored under water. When it does react with air, it forms P4O10, which can react with water to form phosphoric acid (H3PO4), a compound used in food products such as soft drinks. Red phosphorus also reacts with oxygen in the air and is used in the striking strips on match books. As with Group 4A, we again see nonmetals (N and P), metalloids (As and Sb), and a metal (Bi) in Group 5A. In spite of these variations, all of the members of this group form analogous compounds such as the oxides N2O5, P2O5, and As2O5. silicon. Ordinary clay, sand, and many gemstones are based on compounds of silicon and oxygen. Here clear, colorless quartz and dark purple amethyst lie in a bed of sand. All are made of silicon dioxide, SiO2. The different colors are due to impurities. Group 6A: O, S, Se, Te, Po Oxygen, which constitutes about 20% of earth’s atmosphere and combines readily with most other elements, is found at the top of Group 6A. Most of the energy that H2 N2 O2 O3 Figure 2.17 Elements that exist as diatomic F2 Cl2 Br2 I2 molecules. Seven of the elements in the periodic table exist as diatomic, or two-atom, molecules. Oxygen has an additional allotrope, ozone, with three O atoms in each molecule. 86 Chapter 2 Atoms and Elements Figure 2.18 Sulfur. The most common allotrope of sulfur consists of eightmember, crown-shaped rings. powers life on earth is derived from reactions in which oxygen combines with other substances. Sulfur has been known in elemental form since ancient times as brimstone or “burning stone” (Figure 2.18). Sulfur, selenium, and tellurium are referred to collectively as chalcogens (from the Greek word, khalkos, for copper) because most copper ores contain these elements. Their compounds can be foul-smelling and poisonous; nevertheless, sulfur and selenium are essential components of the human diet. By far the most important compound of sulfur is sulfuric acid (H2SO4), which is manufactured in larger amounts than any other compound. As in Group 5A, the second- and third-period elements have different structures. Like nitrogen, oxygen is a diatomic molecule (see Figure 2.17). Unlike nitrogen, however, oxygen has an allotrope, the well-known ozone, O3. Sulfur, which can be found in nature as a yellow solid, has many allotropes. The most common allotrope consists of eight-member, crown-shaped rings of sulfur atoms (see Figure 2.18). Polonium, a radioactive element, was isolated in 1898 by Marie and Pierre Curie, who separated it from tons of a uranium-containing ore and named it for Madame Curie’s native country, Poland. With Group 6A, once again we observe variations in the properties in a group. Oxygen, sulfur, and selenium are nonmetals, tellurium is a metalloid, and polonium is a metal. Nonetheless, there is a family resemblance in their chemistries. All form oxygen-containing compounds such as SO2, SeO2, and TeO2 and sodiumcontaining compounds such as Na2O, Na2S, Na2Se, and Na2Te. Charles D. Winters Group 7A, Halogens: F, Cl, Br, I, At At the far right of the periodic table are two groups composed entirely of nonmetals. The Group 7A elements—fluorine, chlorine, bromine, and iodine—are nonmetals, all of which exist as diatomic molecules (see Figure 2.17). At room temperature, fluorine (F2) and chlorine (Cl2) are gases. Bromine (Br2) is a liquid and iodine (I2) is a solid, but bromine and iodine vapor are clearly visible over the liquid or solid. The Group 7A elements are among the most reactive of all elements. All combine violently with alkali metals to form salts such as table salt, NaCl (Figure 1.7). The name for this group, the halogens, comes from the Greek words hals, meaning “salt,” and genes, meaning “forming.” The halogens also react with other metals and with most nonmetals to form compounds. Bromine, Br2 Iodine, I2 Group 8A, Noble Gases: He, Ne, Ar, Kr, Xe, Rn The Group 8A elements—helium, neon, argon, krypton, xenon, and radon—are the least reactive elements (Figure 2.19). All are gases, and none is abundant on earth or in earth’s atmosphere. Because of this, they were not discovered until the end of the 19th century. Helium, the second most abundant element in the universe after hydrogen, was detected in the sun in 1868 by analysis of the solar spectrum. (The name of the element comes from the Greek word for the sun, helios.) It was not found on earth until 1895, however. Until 1962, when a compound of xenon was first prepared, it was believed that none of these elements would combine chemically with any other element. The common name noble gases for this group, a term meant to denote their general lack of reactivity, derives from this fact. Charles D. Winters The halogens bromine and iodine. Bromine is a liquid at room temperature and iodine is a solid. However, some of the element exists in the vapor state above the liquid or solid. 2.7 An Overview of the Elements, Their Chemistry, and the Periodic Table 87 Figure 2.19 The noble gases. This kit is sold for detecting the presence of radon gas in the home. Neon gas is used in advertising signs, and helium-filled balloons are popular. For the same reason they are sometimes called the inert gases or, because of their low abundance, the rare gases. The Transition Elements Stretching between Groups 2A and 3A is a series of elements called the transition elements. These fill the B-groups (1B through 8B) in the fourth through the seventh periods in the center of the periodic table. All are metals (see Figure 2.10), and 13 of them are in the top 30 elements in terms of abundance in the earth’s crust. Some, like iron (Fe), are abundant in nature (Table 2.4). Most occur naturally in combination with other elements, but a few—silver (Ag), gold (Au), and platinum (Pt )—are much less reactive and can be found in nature as pure elements. Virtually all of the transition elements have commercial uses. They are used as structural materials (iron, titanium, chromium, copper); in paints (titanium, chromium); in the catalytic converters in automobile exhaust systems (platinum, rhodium); in coins (copper, nickel, zinc); and in batteries (manganese, nickel, cadmium, mercury). A number of the transition elements play important biological roles. For example, iron, a relatively abundant element (see Table 2.3), is the central element in the chemistry of hemoglobin, the oxygen-carrying component of blood. Two rows at the bottom of the table accommodate the lanthanides [the series of elements between the elements lanthanum (Z 57) and hafnium (Z 72)] and the actinides [the series of elements between actinium (Z 89) and rutherfordium (Z 104)]. Some lanthanide compounds are used in color television picture tubes, uranium (Z 92) is the fuel for atomic power plants, and americium (Z 95) is used in smoke detectors. Exercise 2.7—The Periodic Table How many elements are in the third period of the periodic table? Give the name and symbol of each. Tell whether each element in the period is a metal, metalloid, or nonmetal. Charles D. Winters Table 2.4 Abundance of the Ten Most Abundant Transition Elements in the Earth’s Crust Rank 4 9 12 18 19 21 23 24 25 26 * ppm Element Iron Titanium Manganese Zirconium Vanadium Chromium Nickel Zinc Cerium Copper g per 1000 kg. Abundance (ppm)* 41,000 5,600 950 190 160 100 80 75 68 50 88 Table 2.5 Relative Amounts of Elements in the Human Body Element Oxygen Carbon Hydrogen Nitrogen Calcium Phosphorus Potassium, sulfur, chlorine Sodium Magnesium Iron, cobalt, copper, zinc, iodine Selenium, fluorine Percent by Mass 65 18 10 3 1.5 1.2 0.2 0.1 0.05 6 0.05 6 0.01 Chapter 2 Atoms and Elements 2.8—Essential Elements As our knowledge of biochemistry—the chemistry of living systems—increases, we learn more and more about essential elements. These elements are so important to life that a deficiency in any one will result in either death, severe developmental abnormalities, or chronic ailments. No other element can take the place of an essential element. Of the 116 known elements, 11 are predominant in many different biological systems and are present in approximately the same relative amounts (Table 2.5). In humans these 11 elements constitute 99.9% of the total number of atoms present, but 4 of these elements—C, H, N, and O—account for 99% of the total. These elements are found in the basic structure of all biochemical molecules. Additionally, H and O are present in water, a major component of all biological systems. The other 7 elements of the group of 11 elements comprise only 0.9% of the total atoms in the body. These are sodium, potassium, calcium, magnesium, phosphorus, sulfur, and chlorine. These generally occur in the form of ions such as Na , K , Mg2 , Ca2 , Cl , and HPO2 . 4 The 11 essential elements represent 6 of the groups of the periodic table, and all are “light” elements; they have atomic numbers less than 21. Another 17 elements are required by most but not all biological systems. Some may be required by plants, some by animals, and others by only certain plants or animals. With a few exceptions, these elements are generally “heavier” elements, elements having an atomic number greater than 18. They are about evenly divided between metals and nonmetals (or metalloids). Elements in the Human Body Major Elements 99.9% of all atoms (99.5% by mass) C, H, N, O Na, Mg, P, S, Cl Trace Elements 0.1% of all atoms (0.5% by mass) V, Cr, Mo, Mn, Fe, Co, Ni, Cu, Zn B, Si, Se, F, Br, I, As, Sn Sources of Some Biologically Important Elements Element Iron Zinc Copper Calcium Source Brewer’s yeast Eggs Brazil nuts Chicken Oysters Brazil nuts Swiss cheese Whole milk Broccoli Selenium Butter Cider vinegar mg/100 g 17.3 2.3 4.2 2.6 13.7 2.3 925 118 103 0.15 0.09 K, Ca Although many of the metals in this group are required only in trace amounts, they are often an integral part of specific biological molecules—such as hemoglobin (Fe), myoglobin (Fe), and vitamin B12 (Co)—and activate or regulate their functions. Much of the 3 or 4 g of iron in the body is found in hemoglobin, the substance responsible for carrying oxygen to cells. Iron deficiency is marked by fatigue, infections, and mouth inflammation. The average person also contains about 2 g of zinc. A deficiency of this element will be evidenced as loss of appetite, failure to grow, and changes in the skin. The human body has about 75 mg of copper, about one third of which is found in the muscles and the remainder in other tissues. Copper is involved in many biological functions, and a deficiency shows up in a variety of ways: anemia, degeneration of the nervous system, impaired immunity, and defects in hair color and structure. 2.8 Key Equations 89 Chapter Goals Revisited Now that you have studied this chapter, you should ask whether you have met the chapter goals. In particular, you should be able to Describe atomic structure and define atomic number and mass number a. Explain the historical development of the atomic theory and identify some of the scientists who made important contributions (Section 2.1). General ChemistryNow homework: Study Question(s) 65 • • b. Describe electrons, protons, and neutrons, and the general structure of the atom (Section 2.1). General ChemistryNow homework: SQ(s) 6, 12 c. Understand the relative mass scale and the atomic mass unit (Section 2.2). General ChemistryNow homework: SQ(s) 64 See the General ChemistryNow CD-ROM or website to: Assess your understanding with homework questions keyed to each goal Check your readiness for an exam by taking the exam-prep quiz and exploring the resources in the personalized Learning Plan it provides Understand the nature of isotopes and calculate atomic weight from isotope abundances and isotopic masses a. Define isotope and give the mass number and number of neutrons for a specific isotope (Sections 2.2 and 2.3). General ChemistryNow homework: SQ(s) 14 b. Do calculations that relate the atomic weight (atomic mass) of an element and isotopic abundances and masses (Section 2.4). General ChemistryNow homework: SQ(s) 20, 22, 25, 47 Charles D. Winters Explain the concept of the mole and use molar mass in calculations a. Understand that the molar mass of an element is the mass in grams of Avogadro’s number of atoms of that element (Section 2.5). General ChemistryNow homework: SQ(s) 27, 29, 31 b. Know how to use the molar mass of an element and Avogadro’s number in calculations (Section 2.5). General ChemistryNow homework: SQ(s) 33, 57, 67, 77 Know the terminology of the periodic table a. Identify the periodic table locations of groups, periods, metals, metalloids, nonmetals, alkali metals, alkaline earth metals, halogens, noble gases, and the transition elements (Sections 2.6 and 2.7). General ChemistryNow homework: SQ(s) 38, 39, 41, 49 Foods rich in essential elements. b. Recognize similarities and differences in properties of some of the common elements of a group. General ChemistryNow homework: SQ(s) 56 Key Equations Equation 2.1 (page 69) Calculate the percent abundance of an isotope. Percent abundance number of atoms of a given isotope total number of atoms of all isotopes of that element 100% Equation 2.2 (page 72) Calculate the atomic mass (atomic weight ) from isotope abundances and the exact atomic mass of each isotope of an element. Atomic weight a % abundance isotope 1 b 1mass of isotope 12 100 a % abundance isotope 2 b 1mass of isotope 22 100 p 90 Chapter 2 Atoms and Elements Study Questions ▲ denotes more challenging questions. ■ denotes questions available in the Homework and Goals section of the General ChemistryNow CD-ROM or website. Blue numbered questions have answers in Appendix O and fully worked solutions in the Student Solutions Manual. Structures of many of the compounds used in these questions are found on the General ChemistryNow CD-ROM or website in the Models folder. Assess your understanding of this chapter’s topics with additional quizzing and conceptual questions at 10. Give the mass number of (a) a nickel atom with 31 neutrons, (b) a plutonium atom with 150 neutrons, and (c) a tungsten atom with 110 neutrons. 11. Give the complete symbol (AX) for each of the following Z atoms: (a) potassium with 20 neutrons, (b) krypton with 48 neutrons, and (c) cobalt with 33 neutrons. 12. ■ Give the complete symbol (AX) for each of the following Z atoms: (a) fluorine with 10 neutrons, (b) chromium with 28 neutrons, and (c) xenon with 78 neutrons. 13. How many electrons, protons, and neutrons are there in an atom of (a) magnesium-24, 24Mg; (b) tin-119, 119Sn; and (c) thorium-232, 232Th? 14. ■ How many electrons, protons, and neutrons are there in an atom of (a) carbon-13, 13C; (b) copper-63, 63Cu; and (c) bismuth-205, 205Bi? Isotopes (See Example 2.2 and the General Chemistry Now Screen 2.12.) 15. The synthetic radioactive element technetium is used in many medical studies. Give the number of electrons, protons, and neutrons in an atom of technetium-99. 16. Radioactive americium-241 is used in household smoke detectors and in bone mineral analysis. Give the number of electrons, protons, and neutrons in an atom of americium-241. 17. Cobalt has three radioactive isotopes used in medical studies. Atoms of these isotopes have 30, 31, and 33 neutrons, respectively. Give the symbol for each of these isotopes. 18. Which of the following are isotopes of element X, the 9 atomic number for which is 9: 19X, 20X, 18X, and 21X? 9 9 9 Isotope Abundance and Atomic Mass (See Exercises 2.3 and 2.4 and the General ChemistryNow Screen 2.13.) 19. Thallium has two stable isotopes, 203Tl and 205Tl. Knowing that the atomic weight of thallium is 204.4, which isotope is the more abundant of the two? 20. ■ Strontium has four stable isotopes. Strontium-84 has a very low natural abundance, but 86Sr, 87Sr, and 88Sr are all reasonably abundant. Knowing that the atomic weight of strontium is 87.62, which of the more abundant isotopes predominates? 21. Verify that the atomic mass of lithium is 6.94, given the following information: 6 Li, mass 6.015121 u; percent abundance 7.50% 7 Li, mass 7.016003 u; percent abundance 92.50% 22. ■ Verify that the atomic mass of magnesium is 24.31, given the following information: 24 Mg, mass 23.985042 u; percent abundance 78.99% 25 26 Practicing Skills Atoms: Their Composition and Structure (See Example 2.1, Exercise 2.2, and the General ChemistryNow Screen 2.11.) 1. What are the three fundamental particles from which atoms are built? What are their electric charges? Which of these particles constitute the nucleus of an atom? Which is the least massive particle of the three? 2. Around 1910 Rutherford carried out his now-famous alpha-particle scattering experiment. What surprising observation did he make in this experiment and what conclusion did he draw from it? 3. What did the discovery of radioactivity reveal about the structure of atoms? 4. What scientific instrument was used to discover that not all atoms of neon have the same mass? 5. If the nucleus of an atom were the size of a medium-sized orange (say, with a diameter of about 6 cm), what would be the diameter of the atom? 6. ■ If a gold atom has a radius of 145 pm, and you could string gold atoms like beads on a thread, how many atoms would you need to have a necklace 36 cm long? 7. The volcanic eruption of Mount St. Helens in the state of Washington in 1980 produced a considerable quantity of a radioactive element in the gaseous state. The element has atomic number 86. What are the symbol and name of this element? 8. Titanium and thallium have symbols that are easily confused with each other. Give the symbol, atomic number, atomic weight, and group and period number of each element. Are they metals, metalloids, or nonmetals? 9. Give the mass number of each of the following atoms: (a) magnesium with 15 neutrons, (b) titanium with 26 neutrons, and (c) zinc with 32 neutrons. Mg, mass Mg, mass 24.985837 u; percent abundance 25.982593 u; percent abundance 10.00% 11.01% ▲ More challenging ■ In General ChemistryNow Blue-numbered questions answered in Appendix O Study Questions 91 23. Silver (Ag) has two stable isotopes, 107Ag and 109Ag. The isotopic mass of 107Ag is 106.9051 and the isotopic mass of 109 Ag is 108.9047. The atomic weight of Ag, from the periodic table, is 107.868. Estimate the percentage of 107Ag in a sample of the element. (a) 0% (b) 25% (c) 50% (d) 75% 24. Copper exists as two isotopes: 63Cu (62.9298 u) and 65Cu (64.9278 u). What is the approximate percentage of 63Cu in samples of this element? (a) 10% (c) 50% (e) 90% (b) 30% (d) 70% 25. ■ Gallium has two naturally occurring isotopes, 69Ga and 71 Ga, with masses of 68.9257 u and 70.9249 u, respectively. Calculate the percent abundances of these isotopes of gallium. 26. Antimony has two stable isotopes, 121Sb and 123Sb, with masses of 120.9038 u and 122.9042 u, respectively. Calculate the percent abundances of these isotopes of antimony. Atoms and the Mole (See Examples 2.5–2.7 and the General ChemistryNow Screens 2.14 and 2.15.) 27. ■ Calculate the mass, in grams, of the following: (a) 2.5 mol of aluminum (b) 1.25 10 3 mol of iron (c) 0.015 mol of calcium (d) 653 mol of neon 28. Calculate the mass, in grams, of (a) 4.24 mol of gold (b) 15.6 mol of He (c) 0.063 mol of platinum (d) 3.63 10 4 mol of Pu 29. ■ Calculate the amount (moles) represented by each of the following: (a) 127.08 g of Cu (b) 0.012 g of lithium (c) 5.0 mg of americium (d) 6.75 g of Al 30. Calculate the amount (moles) represented by each of the following: (a) 16.0 g of Na (b) 0.876 g of tin (c) 0.0034 g of platinum (d) 0.983 g of Xe 31. ■ You are given 1.0-g samples of He, Fe, Li, Si, and C. Which sample contains the largest number of atoms? Which contains the smallest? 32. You are given 1.0-mol amounts of He, Fe, Li, Si, and C. Which sample has the largest mass? 33. ■ What is the average mass of one copper atom? 34. What is the average mass of one titanium atom? The Periodic Table (See Section 2.6 and Exercise 2.7. See also the Periodic Table Tool on the General ChemistryNow CD-ROM or website.) 35. Give the name and symbol of each of the Group 5A elements. Tell whether each is a metal, nonmetal, or metalloid. 36. Give the name and symbol of each of the fourth-period elements. Tell whether each is a metal, nonmetal, or metalloid. 37. How many periods of the periodic table have 8 elements, how many have 18 elements, and how many have 32 elements? 38. ■ How many elements occur in the seventh period? What is the name given to the majority of these elements and what well-known property characterizes them? 39. ■ Select answers to the questions listed below from the following list elements whose symbols start with the letter C: C, Ca, Cr, Co, Cd, Cl, Cs, Ce, Cm, Cu, and Cf. (You should expect to use some symbols more than once.) (a) Which are nonmetals? (b) Which are main group elements? (c) Which are lanthanides? (d) Which are transition elements? (e) Which are actinides? (f ) Which are gases? 40. Give the name and chemical symbol for the following. (a) a nonmetal in the second period (b) an alkali metal (c) the third-period halogen (d) an element that is a gas at 20° C and 1 atmosphere pressure 41. ■ Classify the following elements as metals, metalloids, or nonmetals: N, Na, Ni, Ne, and Np. 42. Here are symbols for five of the seven elements whose names begin with the letter B: B, Ba, Bk, Bi, and Br. Match each symbol with one of the descriptions below. (a) a radioactive element (b) a liquid at room temperature (c) a metalloid (d) an alkaline earth element (e) a Group 5A element 43. Use the elements in the following list to answer the questions: sodium, silicon, sulfur, scandium, selenium, strontium, silver, and samarium. (Some elements will be entered in more than one category.) (a) Identify those that are metals. (b) Identify those that are main group elements (c) Identify those that are transition metals. 44. Compare the elements silicon (Si) and phosphorus (P) using the following criteria: (a) metal, metalloid, or nonmetal (b) possible conductor of electricity (c) physical state at 25 ° C (solid, liquid, or gas) ■ In General ChemistryNow ▲ More challenging Blue-numbered questions answered in Appendix O 92 Chapter 2 Atoms and Elements General Questions These questions are not designed as to type or location in the chapter. They may combine several concepts. More challenging questions are marked with the icon ▲. 45. Fill in the blanks in the table (one column per element ). Symbol Number of protons Number of neutrons Number of electrons in the neutral atom Name of element Symbol Number of protons Number of neutrons Number of electrons in the neutral atom Name of element 58 Ni ______ ______ 33 S ______ ______ ______ 10 10 ______ ______ ______ 78 117 ______ ______ ______ ______ 30 25 ______ ______ ______ 46 35 ______ 49. ■ The chart shown in the Stardust story (page 58) plots the logarithm of the abundance of elements 1 through 30 in the solar system on a logarithmic scale. (a) What is the most abundant main group metal? (b) What is the most abundant nonmetal? (c) What is the most abundant metalloid? (d) Which of the transition elements is most abundant? (e) Which halogens are included on this plot and which is the most abundant? 50. The molecule buckminsterfullerene, commonly called a “buckyball,” is one of three common allotropes of a familiar element. Identify two other allotropes of this element. 51. Which of the following is impossible? (a) silver foil that is 1.2 10 4 m thick (b) a sample of potassium that contains 1.784 atoms (c) a gold coin of mass 1.23 10 3 kg (d) 3.43 10 27 mol of S8 ______ ______ Cu ______ ______ ______ ______ 65 ______ ______ Kr ______ ______ ______ ______ 86 46. Fill in the blanks in the table (one column per element ). 1024 47. ■ Potassium has three naturally occurring isotopes (39K, 40 K, and 41K), but 40K has a very low natural abundance. Which of the other two isotopes is the more abundant? Briefly explain your answer. 48. Crossword Puzzle: In the 2 2 box shown here, each answer must be correct four ways: horizontally, vertically, diagonally, and by itself. Instead of words, use symbols of elements. When the puzzle is complete, the four spaces will contain the overlapping symbols of ten elements. There is only one correct solution. 1 3 2 4 52. Give the symbol for a metalloid in the third period and then identify a property of this element. 53. Reviewing the periodic table. (a) Name an element in Group 2A. (b) Name an element in the third period. (c) Which element is in the second period in Group 4A? (d) Which element is in the third period in Group 6A? (e) Which halogen is in the fifth period? (f ) Which alkaline earth element is in the third period? (g) Which noble gas element is in the fourth period? (h) Name the nonmetal in Group 6A and the third period. (i) Name a metalloid in the fourth period. 54. Reviewing the periodic table: (a) Name an element in Group 2B. (b) Name an element in the fifth period. (c) Which element is in the sixth period in Group 4A? (d) Which element is in the third period in Group 6A? (e) Which alkali metal is in the third period? (f ) Which noble gas element is in the fifth period? (g) Name the element in Group 6A and the fourth period. Is it a metal, nonmetal, or metalloid? (h) Name a metalloid in Group 5A. 55. The plot on the following page shows the variation in density with atomic number for the first 36 elements. Use this plot to answer the following questions: (a) Which three elements in this series have the highest density? What is their approximate density? Are these elements metals or nonmetals? Horizontal 1–2: Two-letter symbol for a metal used in ancient times 3–4: Two-letter symbol for a metal that burns in air and is found in Group 5A Vertical 1–3: Two-letter symbol for a metalloid 2–4: Two-letter symbol for a metal used in U.S. coins Single squares: all one-letter symbols 1: A colorful nonmetal 2: Colorless gaseous nonmetal 3: An element that makes fireworks green 4: An element that has medicinal uses Diagonal 1–4: Two-letter symbol for an element used in electronics 2–3: Two-letter symbol for a metal used with Zr to make wires for superconducting magnets This puzzle first appeared in Chemical & Engineering News, p. 86, December 14, 1987 (submitted by S. J. Cyvin) and in Chem Matters, October 1988. ▲ More challenging ■ In General ChemistryNow Blue-numbered questions answered in Appendix O Study Questions 93 (b) Which element in the second period has the highest density? Which element in the third period has the highest density? What do these two elements have in common? (c) Some elements have densities so low that they do not show up on the plot. What elements are these? What property do they have in common? 63. Put the following elements in order from smallest to largest mass: (a) 3.79 1024 atoms Fe (e) 9.221 mol Na (b) 19.921 mol H2 (f ) 4.07 1024 atoms Al (c) 8.576 mol C (g) 9.2 mol Cl2 (d) 7.4 mol Si 64. ■ ▲ When a sample of phosphorus burns in air, the compound P4O10 forms. One experiment showed that 0.744 g of phosphorus formed 1.704 g of P4O10. Use this information to determine the ratio of the atomic masses of phosphorus and oxygen (mass P/mass O). If the atomic mass of oxygen is assumed to be 16.000 u, calculate the atomic mass of phosphorus. 65. ■ The data below were collected in a Millikan oil drop experiment. Oil Drop 1 2 3 4 5 Measured Charge on Drop (C) 1.59 11.1 9.54 15.9 6.36 10 10 10 10 10 19 19 19 19 19 2 4 6 8 10 12 14 Atomic number 16 18 20 22 24 26 28 30 32 34 36 0 2 6 8 4 Density (g/cm3) 10 (a) Use these data to calculate the charge on the electron (in coulombs). (b) How many electrons have accumulated on each oil drop? (c) The accepted value of the electron charge is 1.60 10 19 C. Calculate the percent and error for the value determined by the data in the table. 66. ▲ Although carbon-12 is now used as the standard for atomic masses, this has not always been the case. Early attempts at classification used hydrogen as the standard, with the mass of hydrogen being set equal to 1.0000 u. Later attempts defined atomic masses using oxygen (with a mass of 16.0000 u). In each instance, the atomic masses of the other elements were defined relative to these masses. (To answer this question, you need more precise data on current atomic masses: H, 1.00794 u; O, 15.9994 u.) (a) If H 1.0000 u was used as a standard for atomic masses, what would the atomic mass of oxygen be? What would be the value of Avogadro’s number under these circumstances? (b) Assuming the standard is O 16.0000 u, determine the value for the atomic mass of hydrogen and the value of Avogadro’s number. 67. ■ A reagent occasionally used in chemical synthesis is sodium-potassium alloy. (Alloys are mixtures of metals, and Na-K has the interesting property that it is a liquid.) One formulation of the alloy (the one that melts at the lowest temperature) contains 68 atom percent K; that is, out of every 100 atoms, 68 are K and 32 are Na. What is the weight percent of potassium in sodium-potassium alloy? 68. Mass spectrometric analysis showed that there are four isotopes of an unknown element having the following masses and abundances: ■ In General ChemistryNow Blue-numbered questions answered in Appendix O 56. ■ Give two examples of nonmetallic elements that have allotropes. Name those elements and describe the allotropes of each. 57. ■ In each case, decide which represents more mass: (a) 0.5 mol of Na or 0.5 mol of Si (b) 9.0 g of Na or 0.50 mol of Na (c) 10 atoms of Fe or 10 atoms of K 58. A semiconducting material is composed of 52 g of Ga, 9.5 g of Al, and 112 g of As. Which element has the largest number of atoms in the final mixture? 59. You are given 15 g each of yttrium, boron, and copper. Which sample represents the largest number of atoms? 60. Lithium has two stable isotopes: 6Li and 7Li. One of them has an abundance of 92.5%, and the other has an abundance of 7.5%. Knowing that the atomic mass of lithium is 6.941, which is the more abundant isotope? 61. Superman comes from the planet Krypton. If you have 0.00789 g of the gaseous element krypton, how many moles does this represent? How many atoms? 62. The recommended daily allowance (RDA) of iron in your diet is 15 mg. How many moles is this? How many atoms? ▲ More challenging 94 Isotope 1 2 3 4 Mass Number 136 138 140 142 Chapter 2 Atoms and Elements Isotope Mass 135.9090 137.9057 139.9053 141.9090 Abundance (%) 0.193 0.250 88.48 11.07 (a) Based on their relative reactivities, what might you expect to see when barium, another Group 2A element, is placed in water? (b) Give the period in which each element (Mg, Ca, and Ba) is found. What correlation do you think you might find between the reactivity of these elements and their positions in the periodic table? 74. ▲ In an experiment, you need 0.125 mol of sodium metal. Sodium can be cut easily with a knife (Figure 2.12), so if you cut out a block of sodium, what should the volume of the block be in cubic centimeters? If you cut a perfect cube, what is the length of the edge of the cube? (The density of sodium is 0.97 g/cm3.) 75. ▲ Dilithium is the fuel for the Starship Enterprise. Because its density is quite low, however, you need a large space to store a large mass. To estimate the volume required, we shall use the element lithium. If you need 256 mol for an interplanetary trip, what must the volume of the piece of lithium be? If the piece of lithium is a cube, what is the dimension of an edge of the cube? (The density for the element lithium is 0.534 g/cm3 at 20 ° C.) 76. An object is coated with a layer of chromium, 0.015 cm thick. The object has a surface area of 15.3 cm2. How many atoms of chromium are used in the coating? (The density of chromium 7.19 g/cm3.) 77. ■ A cylindrical piece of sodium is 12.00 cm long and has a diameter of 4.5 cm. The density of sodium is 0.971 g/cm3. How many atoms does the piece of sodium contain? (The volume of a cylinder is V p r 2 length.) 78. ▲ Consider an atom of 64Zn. (a) Calculate the density of the nucleus in grams per cubic centimeter, knowing that the nuclear radius is 4.8 10 6 nm and the mass of the 64Zn atom is 1.06 10 22 g. Recall that the volume of a sphere is (4/3)pr 3. (b) Calculate the density of the space occupied by the electrons in the zinc atom, given that the atomic radius is 0.125 nm and the electron mass is 9.11 10 28 g. (c) Having calculated these densities, what statement can you make about the relative densities of the parts of the atom? 79. ▲ Most standard analytical balances can measure accurately to the nearest 0.0001 g. Assume you have weighed out a 2.0000-g sample of carbon. How many atoms are in this sample? Assuming the indicated accuracy of the measurement, what is the largest number of atoms that can be present in the sample? 80. ▲ To estimate the radius of a lead atom: (a) You are given a cube of lead that is 1.000 cm on each side. The density of lead is 11.35 g/cm3. How many atoms of lead are in the sample? (b) Atoms are spherical; therefore, the lead atoms in this sample cannot fill all the available space. As an approximation, assume that 60% of the space of the cube is filled with spherical lead atoms. Calculate the volume Three elements in the periodic table that have atomic weights near these values are lanthanum (La), atomic number 57, atomic weight 139.9055; cerium (Ce), atomic number 58, atomic weight 140.115; and praeseodymium (Pr), atomic number 59, atomic weight 140.9076. Using the data above, calculate the atomic weight and identify the element if possible. Summary and Conceptual Questions The following questions use concepts from the preceding chapter (Chapter 1). 69. Draw a picture showing the approximate positions of all protons, electrons, and neutrons in an atom of helium-4. Make certain that your diagram indicates both the number and position of each type of particle. 70. Draw two boxes, each about 3 cm on a side. In one box, sketch a representation of iron metal. In the other box, sketch a representation of nitrogen gas. How do these drawings differ? 71. ▲ Identify, from the list below, the information needed to calculate the number of atoms in 1 cm3 of iron. Outline the procedure used in this calculation. (a) the structure of solid iron (b) the molar mass of iron (c) Avogadro’s number (d) the density of iron (e) the temperature (f ) iron’s atomic number (g) the number of iron isotopes 72. Consider the plot of relative element abundances on page 58. Is there a relationship between abundance and atomic number? Is there any difference between the relative abundance of an element of even atomic number and the relative abundance of an element of odd atomic number? 73. The photo here depicts what happens when a coil of magnesium ribbon and a few calcium chips are placed in water. Magnesium (left) and calcium (right) in water. ▲ More challenging ■ In General ChemistryNow Blue-numbered questions answered in Appendix O Charles D. Winters Study Questions 95 of one lead atom from this information. From the calculated volume (V ), and the formula V 4/3 1pr 3 2, estimate the radius (r) of a lead atom. 81. A jar contains some number of jelly beans. To find out precisely how many are in the jar you could dump them out and count them. How could you estimate their number without counting each one? (Chemists need to do just this kind of “bean counting” when we work with atoms and molecules. They are too small to count one by one, so we have worked out other methods to “count atoms.”) (See General ChemistryNow Screen 2.18, Chemical Puzzler.) Do you need a live tutor for homework problems? Access vMentor at General ChemistryNow at for one-on-one tutoring from a chemistry expert Charles D. Winters How many jelly beans are in the jar? ▲ More challenging ■ In General ChemistryNow Blue-numbered questions answered in Appendix O ...
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This note was uploaded on 02/18/2009 for the course CHEM 101 taught by Professor Williamson during the Fall '08 term at Texas A&M.

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