FIGURE 2.3 The three most abundant isotopes of carbon. Isotopes Carbon -12 of a particular atom have different numbers of 6 Neutro 6 Electrons Isotopes tential to severely damage living cells, producing mutations in Atoms with the same atom nic number(that is, the same num- their genes, and, at high doses, cell death. Consequently,ex osure to radiation is now very carefully controlled and ber of protons) have the same chemical properties and are iated. Scientists who work with radioactivity(basic re- said to belong to the same element. Formally speaking, an element is any substance that cannot be broken down to any searchers as well as applied scientists such as X-ray other substance by ordinary chemical means. However, while technologists)wear radiation-sensitive badges to monitor the all atoms of an element have the same number of protons total amount of radioactivity to which they are exposed. Each they may not all have the same number of neutrons. Atoms of month the badges are collected and scrutinized. Thus, em an element that possess different numbers of neutrons are ployees whose work places them in danger of excessive radio- active exposure are equipped with an"earl ng system as mixtures of different isotopes. Carbon (C), for example, has three isotopes, all containing six protons(figure 2.3) Electrons Over 99% of the carbon found in nature exists as an isotope The positive charges in the nucleus of an atom are counter with six neutrons. Because its total mass is 12 daltons(6 from balanced by negatively charged electrons orbiting at vary- protons plus 6 from neutrons), this isotope is referred to as ing distances around the nucleus. Thus, atoms with the carbon-12, and symbolized 2C. Most of the rest of the natu- same number of protons and electrons are electrically neu- ally occurring carbon is carbon-13 tope with seven tral, having no net ch neutrons. The rarest carbon isotope is carbon-14, with eight Electrons are maintained in their orbits by their attrac neutrons. Unlike the other two isotopes, carbon-14 is unsta- tion to the positively charged nucleus. Sometimes other ble: its nucleus tends to break up into elements with lower forces overcome this attraction and an atom loses one or atomic numbers. This nuclear breakup, which emits a signifi more electrons In other cases, atoms may gain additional cant amount of energy, is called radioactive decay, and iso- electrons. Atoms in which the number of electrons does pes that decay in this fashion are radioactive isotopes not equal the number of protons are known as ions, and however, the rate of decay is constant. This rate is usually called a cation. For example, an atom of sodium(Na)that expressed as the half-life the time it takes for one half of the has lost one electron becomes a sodium ion(Na*), with a atoms in a sample to decay. Carbon-14, for example, has a charge of +1. An atom that has fewer protons than elec half-life of about 5600 years. A sample of carbon containing rons carries a net negative charge and is called an anion. a I gram of carbon-14 today would contain 0.5 gram of car chlorine atom( Ci)that has gained one electron becomes a bon-14 after 5600 years, 0.25 gram 11, 200 years from now, chloride ion( Ch), with a charge of-1 0.125 gram 16, 800 years from now, and so on. By determin ing the ratios of the different isotopes of carbon and other An atom consists of a nucleus of protons and neutrons elements in biological samples and in rocks, scientists ar surrounded by a cloud of electrons. The number of its able to accurately determine when these materials formed electrons largely determines the chemical properties of While there are many useful applications of radioactivity an atom. Atoms that have the same number of protons there are also harmful side effects that must be considered in but different numbers of neutrons are called isotopes any planned use of radioactive substances. Radioactive sub- Isotopes of an atom differ in atomic mass but have stances emit energetic subatomic particles that have the po- similar chemical properties Chapter2 The Nature of Molecules 21
Isotopes Atoms with the same atomic number (that is, the same number of protons) have the same chemical properties and are said to belong to the same element. Formally speaking, an element is any substance that cannot be broken down to any other substance by ordinary chemical means. However, while all atoms of an element have the same number of protons, they may not all have the same number of neutrons. Atoms of an element that possess different numbers of neutrons are called isotopes of that element. Most elements in nature exist as mixtures of different isotopes. Carbon (C), for example, has three isotopes, all containing six protons (figure 2.3). Over 99% of the carbon found in nature exists as an isotope with six neutrons. Because its total mass is 12 daltons (6 from protons plus 6 from neutrons), this isotope is referred to as carbon-12, and symbolized 12C. Most of the rest of the naturally occurring carbon is carbon-13, an isotope with seven neutrons. The rarest carbon isotope is carbon-14, with eight neutrons. Unlike the other two isotopes, carbon-14 is unstable: its nucleus tends to break up into elements with lower atomic numbers. This nuclear breakup, which emits a significant amount of energy, is called radioactive decay, and isotopes that decay in this fashion are radioactive isotopes. Some radioactive isotopes are more unstable than others and therefore decay more readily. For any given isotope, however, the rate of decay is constant. This rate is usually expressed as the half-life, the time it takes for one half of the atoms in a sample to decay. Carbon-14, for example, has a half-life of about 5600 years. A sample of carbon containing 1 gram of carbon-14 today would contain 0.5 gram of carbon-14 after 5600 years, 0.25 gram 11,200 years from now, 0.125 gram 16,800 years from now, and so on. By determining the ratios of the different isotopes of carbon and other elements in biological samples and in rocks, scientists are able to accurately determine when these materials formed. While there are many useful applications of radioactivity, there are also harmful side effects that must be considered in any planned use of radioactive substances. Radioactive substances emit energetic subatomic particles that have the potential to severely damage living cells, producing mutations in their genes, and, at high doses, cell death. Consequently, exposure to radiation is now very carefully controlled and regulated. Scientists who work with radioactivity (basic researchers as well as applied scientists such as X-ray technologists) wear radiation-sensitive badges to monitor the total amount of radioactivity to which they are exposed. Each month the badges are collected and scrutinized. Thus, employees whose work places them in danger of excessive radioactive exposure are equipped with an “early warning system.” Electrons The positive charges in the nucleus of an atom are counterbalanced by negatively charged electrons orbiting at varying distances around the nucleus. Thus, atoms with the same number of protons and electrons are electrically neutral, having no net charge. Electrons are maintained in their orbits by their attraction to the positively charged nucleus. Sometimes other forces overcome this attraction and an atom loses one or more electrons. In other cases, atoms may gain additional electrons. Atoms in which the number of electrons does not equal the number of protons are known as ions, and they carry a net electrical charge. An atom that has more protons than electrons has a net positive charge and is called a cation. For example, an atom of sodium (Na) that has lost one electron becomes a sodium ion (Na+), with a charge of +1. An atom that has fewer protons than electrons carries a net negative charge and is called an anion. A chlorine atom (Cl) that has gained one electron becomes a chloride ion (Cl–), with a charge of –1. An atom consists of a nucleus of protons and neutrons surrounded by a cloud of electrons. The number of its electrons largely determines the chemical properties of an atom. Atoms that have the same number of protons but different numbers of neutrons are called isotopes. Isotopes of an atom differ in atomic mass but have similar chemical properties. Chapter 2 The Nature of Molecules 21 Carbon-12 6 Protons 6 Neutrons 6 Electrons Carbon-13 6 Protons 7 Neutrons 6 Electrons Carbon-14 6 Protons 8 Neutrons 6 Electrons FIGURE 2.3 The three most abundant isotopes of carbon. Isotopes of a particular atom have different numbers of neutrons
Electrons Determine the Chemical also explains why the isotopes of an element, all of which Behavior of atoms have the same arrangement of electrons, behave the same The key to the chemical behavior of an atom lies in the ar ingement of its electrons in their orbits. It is convenient visualize individual electrons as following discrete circular Energy within the ator orbits around a central nucleus, as in the bohr model of the All atoms possess energy, defined as the ability to do work. atom. However, such a simple picture is not realistic. It is Because electrons are attracted to the positively charged not possible to precisely locate the position of any individual nucleus, it takes work to keep them in orbit, just as it takes electron precisely at any given time. In fact, a particular work to hold a grapefruit in your hand against the pull of electron can be anywhere at a given instant, from close to gravity. The grapefruit is said to possess potential energy the nucleus to infinitely far away from it. the ability to do work, because of its position; if you were However, a particular electron is more likely to be locat to release it, the grapefruit would fall and its energy would ed in some positions than in others. The area around a nu- be reduced. Conversely, if you were to move the grapefruit cleus where an electron is most likely to be found is called to the top of a building, you would increase its potential the orbital of that electron(figure 2. 4). Some electron or- energy. Similarly, electrons have potential energy of posi bitals near the nucleus are spherical (s orbitals), while oth- tion. To oppose the attraction of the nucleus and move the ers are dumbbell-shaped (p orbitals). Still other orbitals, electron to a more distant orbital requires an input of en- more distant from the nucleus, may have different shapes ergy and results in an electron with greater potential ener- Regardless of its shape, no orbital may contain more than gy. This is how chlorophyll captures energy from light two electrons during photosynthesis(chapter 10)-the light excites elec- Almost all of the volume of an atom is empty space, be- trons in the chlorophyll. Moving an electron closer to the cause the electrons are quite far from the nucleus relative nucleus has the opposite effect: energy is released, usually to its size. If the nucleus of an atom were the size of an ap- as heat, and the electron ends up with less potential energy ple, the orbit of the nearest electron would be more than (figure 2.5) 600 meters away. Consequently, the nuclei of two atoms a given atom can possess only certain discrete amounts never come close enough in nature to interact with each of energy. Like the potential energy of a grapefruit on a step other. It is for this reason that an atoms electrons, not its of a staircase, the potential energy contributed by the posi protons or neutrons, determine its chemical behavior. This tion of an electron in an atom can have only certain values Orbital for energy level K: Orbitals for energy level L; Composite of rbital (1s) one spherical orbital(2s)and all three dumbbell-shaped orbitals(2p) FIGURE 2. 4 Electron orbitals. The lowest energy level or electron shell, which is nearest the nucleus, is level K It is occupied by a single s orbital, referred to as ls. The next highest energy level, L, is occupied by four orbitals: one s orbital (referred to as the 2s or orbitals(each referred to as a 2p orbital). The four L-level orbitals compactly fill the space around the nucleus, like two pyramids set b 22 Part I The Origin of Living Things
Electrons Determine the Chemical Behavior of Atoms The key to the chemical behavior of an atom lies in the arrangement of its electrons in their orbits. It is convenient to visualize individual electrons as following discrete circular orbits around a central nucleus, as in the Bohr model of the atom. However, such a simple picture is not realistic. It is not possible to precisely locate the position of any individual electron precisely at any given time. In fact, a particular electron can be anywhere at a given instant, from close to the nucleus to infinitely far away from it. However, a particular electron is more likely to be located in some positions than in others. The area around a nucleus where an electron is most likely to be found is called the orbital of that electron (figure 2.4). Some electron orbitals near the nucleus are spherical (s orbitals), while others are dumbbell-shaped (p orbitals). Still other orbitals, more distant from the nucleus, may have different shapes. Regardless of its shape, no orbital may contain more than two electrons. Almost all of the volume of an atom is empty space, because the electrons are quite far from the nucleus relative to its size. If the nucleus of an atom were the size of an apple, the orbit of the nearest electron would be more than 1600 meters away. Consequently, the nuclei of two atoms never come close enough in nature to interact with each other. It is for this reason that an atom’s electrons, not its protons or neutrons, determine its chemical behavior. This also explains why the isotopes of an element, all of which have the same arrangement of electrons, behave the same way chemically. Energy within the Atom All atoms possess energy, defined as the ability to do work. Because electrons are attracted to the positively charged nucleus, it takes work to keep them in orbit, just as it takes work to hold a grapefruit in your hand against the pull of gravity. The grapefruit is said to possess potential energy, the ability to do work, because of its position; if you were to release it, the grapefruit would fall and its energy would be reduced. Conversely, if you were to move the grapefruit to the top of a building, you would increase its potential energy. Similarly, electrons have potential energy of position. To oppose the attraction of the nucleus and move the electron to a more distant orbital requires an input of energy and results in an electron with greater potential energy. This is how chlorophyll captures energy from light during photosynthesis (chapter 10)—the light excites electrons in the chlorophyll. Moving an electron closer to the nucleus has the opposite effect: energy is released, usually as heat, and the electron ends up with less potential energy (figure 2.5). A given atom can possess only certain discrete amounts of energy. Like the potential energy of a grapefruit on a step of a staircase, the potential energy contributed by the position of an electron in an atom can have only certain values. 22 Part I The Origin of Living Things 1s Orbital x x y z Orbital for energy level K: one spherical orbital (1s) 2s Orbital 2p Orbitals Composite of all p orbitals Orbitals for energy level L: one spherical orbital (2s) and three dumbbell-shaped orbitals (2p) z y FIGURE 2.4 Electron orbitals. The lowest energy level or electron shell, which is nearest the nucleus, is level K. It is occupied by a single s orbital, referred to as 1s. The next highest energy level, L, is occupied by four orbitals: one s orbital (referred to as the 2s orbital) and three p orbitals (each referred to as a 2p orbital). The four L-level orbitals compactly fill the space around the nucleus, like two pyramids set baseto-base
FIGURE 2.5 Atomic energy levels. When an electron absorbs energy, it moves to higher energy levels farther from the nucleus. When ar E Energy Energy nergy Energy electron releases energy, it falls to lower level level level nergy levels closer to the nucleus. Every atom exhibits a ladder of potential energy values rather than a continuous spectrum of possibilities, a discrete set of orbits at particular distances from the nucleus During some chemical reactions, electrons are trans- ferred from one atom to another. In such reactions the loss of an electron is called oxidation, and the gain of an elec tron is called reduction(figure 2.6). It is important to real ize that when an electron is transferred in this way, it keeps Reduction gy of position. In organisms, energy Is FIGURE 2.6 stored in high-energy electrons that are transferred from Oxidation and reduction. Oxidation is the loss of an electron: one atom to another in reactions involving oxidation and reduction is the gain of an electron. eduction Because the amount of energy an electron possesses is related to its distance from the nucleus. electrons that a the same distance from the nucleus have the same energy, even if they occupy different orbitals. Such electrons are said to occupy the same energy level. In a schematic dia gram of an atom(figure 2.7), the nucleus is represented as a small circle and the electron energy levels are drawn as con- centric rings, with the energy level increasing with distance from the nucleus. Be careful not to confuse energy levels, which are drawn as rings to indicate an electron's energy with orbitals, which have a variety of three-dimensional shapes and indicate an electrons most likely location Helium Electrons orbit a nucleus in paths called orbitals. No orbital can contain more than two electrons, but many orbitals may be the same distance from the nucleus and thus, contain electrons of the same energy FIGURE 2.7 Electron energy levels for helium and nitrogen. Gold balls present the electrons. Each concentric circle represents a different distance from the nucleus and thus, a different electron level Chapter2 The Nature of molecul
Every atom exhibits a ladder of potential energy values, rather than a continuous spectrum of possibilities, a discrete set of orbits at particular distances from the nucleus. During some chemical reactions, electrons are transferred from one atom to another. In such reactions, the loss of an electron is called oxidation, and the gain of an electron is called reduction (figure 2.6). It is important to realize that when an electron is transferred in this way, it keeps its energy of position. In organisms, chemical energy is stored in high-energy electrons that are transferred from one atom to another in reactions involving oxidation and reduction. Because the amount of energy an electron possesses is related to its distance from the nucleus, electrons that are the same distance from the nucleus have the same energy, even if they occupy different orbitals. Such electrons are said to occupy the same energy level. In a schematic diagram of an atom (figure 2.7), the nucleus is represented as a small circle and the electron energy levels are drawn as concentric rings, with the energy level increasing with distance from the nucleus. Be careful not to confuse energy levels, which are drawn as rings to indicate an electron’s energy, with orbitals, which have a variety of three-dimensional shapes and indicate an electron’s most likely location. Electrons orbit a nucleus in paths called orbitals. No orbital can contain more than two electrons, but many orbitals may be the same distance from the nucleus and, thus, contain electrons of the same energy. Chapter 2 The Nature of Molecules 23 Energy released Energy level 3 Energy level 2 Energy level 1 – ML K Energy level 1 Energy absorbed Energy level 2 Energy level 3 + – + + + + + + K L M FIGURE 2.5 Atomic energy levels. When an electron absorbs energy, it moves to higher energy levels farther from the nucleus. When an electron releases energy, it falls to lower energy levels closer to the nucleus. FIGURE 2.6 Oxidation and reduction. Oxidation is the loss of an electron; reduction is the gain of an electron. Oxidation Reduction – Helium Nitrogen 7 7n 2 2n K L K Nucleus L M N K Energy level FIGURE 2.7 Electron energy levels for helium and nitrogen. Gold balls represent the electrons. Each concentric circle represents a different distance from the nucleus and, thus, a different electron energy level.
2.2 The atoms of living things are among the smallest. Kinds of atoms level can contain no more than eight electrons; the chemi cal behavior of an element reflects how many of the eight There are 92 naturally occurring elements, each with a dif- positions are filled. Elements possessing all eight elec- ferent number of protons and a different arrangement of trons in their outer energy level(two for helium)are electrons. When the nineteenth-century Russian chemist inert, or nonreactive; they include helium(He),neon Dmitri Mendeleev arranged the known elements in a table (Ne), argon(Ar), krypton(Kr), xenon(Xe), and radon according to their atomic mass(figure 2.8), he discovered(Rn). In sharp contrast, elements with seven electrons(one one of the great generalizations in all of science. Mendeleev fewer than the maximum number of eight) in their outer found that the elements in the table exhibited a pattern of energy level, such as fluorine(F), chlorine(CI), and chemical properties that repeated itself in groups of eight el- bromine(Br), are highly reactive. They tend to gain the ements. This periodically repeating pattern lent the table extra electron needed to fill the energy level. Elements name: the periodic table of elements with only one electron in their outer energy level, such as lithium(Li), sodium(Na), and potassium(K), are als The periodic table very reactive; they tend to lose the single electron in their outer level The eight-element periodicity that Mendeleev found is Mendeleev's periodic table thus leads to a useful generali based on the interactions of the electrons in the outer en- zation, the octet rule (Latin octo, "eight")or rule of eight ergy levels of the different elements. These electrons are atoms tend to establish completely full outer energy levels called valence electrons and their interactions are the Most chemical behavior can be predicted quite accurately basis for the differing chemical properties of the elements. from this simple rule, combined with the tendency of at- For most of the atoms important to life, an outer energy oms to balance positive and negative charges B Mg ClA 2728293031 Co Ni Cu Zn Ga Ge As Se Br K Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb TeIXe 767778 8081 Cs Ba La Hf W Re Os Ir Pt Au Hg TI Pb Bi Po AtRn 878889 墨105106107108109110N Fr RaA 6061 626364 666768697071 (Lanthanide series) Ce Pr Nd Pm Sm Eu Gd Dy Ho Er Tm Yb Lu 90919293949596979899100101 I 102 103 Actinide series) Th Pa U Np Pu Am Cm Bk Cf Es Fm Md NoLr FIGURE 2. 8 Periodic table of the elements. In this representation, the frequency of elements that occur in the earth's crust is indicated by the height of the block Elements found in significant amounts in living organisms are shaded in blue. 24 Part I The Origin of Living Things
24 Part I The Origin of Living Things 1 H 1 H 3 Li 4 Be 19 K 12 Mg 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 37 Rb 38 Sr 39 Y 42 Mo 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 21 Sc 40 Zr 22 Ti 23 V 24 Cr 25 Mn 27 Co 28 Ni 29 Cu 30 Zn 36 Kr 5 B 6 C 6 C 8 O 2 He 55 Cs 56 Ba 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 87 Fr 88 Ra 57 La 89 Ac 104 105 106 107 108 109 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 90 Th 91 Pa 92 U (Lanthanide series) (Actinide series) 11 Na 20 Ca 41 Nb 43 Tc 44 Ru 26 Fe 13 Al 31 Ga 32 Ge 14 Si 7 N 15 P 33 As 16 S 35 Br 34 Se 9 F 18 Ar 10 Ne 17 Cl 110 FIGURE 2.8 Periodic table of the elements. In this representation, the frequency of elements that occur in the earth’s crust is indicated by the height of the block. Elements found in significant amounts in living organisms are shaded in blue. Kinds of Atoms There are 92 naturally occurring elements, each with a different number of protons and a different arrangement of electrons. When the nineteenth-century Russian chemist Dmitri Mendeleev arranged the known elements in a table according to their atomic mass (figure 2.8), he discovered one of the great generalizations in all of science. Mendeleev found that the elements in the table exhibited a pattern of chemical properties that repeated itself in groups of eight elements. This periodically repeating pattern lent the table its name: the periodic table of elements. The Periodic Table The eight-element periodicity that Mendeleev found is based on the interactions of the electrons in the outer energy levels of the different elements. These electrons are called valence electrons and their interactions are the basis for the differing chemical properties of the elements. For most of the atoms important to life, an outer energy level can contain no more than eight electrons; the chemical behavior of an element reflects how many of the eight positions are filled. Elements possessing all eight electrons in their outer energy level (two for helium) are inert, or nonreactive; they include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). In sharp contrast, elements with seven electrons (one fewer than the maximum number of eight) in their outer energy level, such as fluorine (F), chlorine (Cl), and bromine (Br), are highly reactive. They tend to gain the extra electron needed to fill the energy level. Elements with only one electron in their outer energy level, such as lithium (Li), sodium (Na), and potassium (K), are also very reactive; they tend to lose the single electron in their outer level. Mendeleev’s periodic table thus leads to a useful generalization, the octet rule (Latin octo, “eight”) or rule of eight: atoms tend to establish completely full outer energy levels. Most chemical behavior can be predicted quite accurately from this simple rule, combined with the tendency of atoms to balance positive and negative charges. 2.2 The atoms of living things are among the smallest
Table 2.1 The most common elements on earth and Their Distribution in the human bod A Percent of Percent of Earth's Crust una Elemen by Weight by Weight Importance or Function 65.0 Required for cellular respiration Silicon race Trace Critical component of hemoglobin in the blood Calcium 1.5 Component of bones and teeth; tri gers muscle contraction Sodium 8 Principal positive ion outside cells; important in nerve function Potassium K 2.6 Principal positive ion inside cells; im- portant in nerve function Magnesium Mg Critical component of many energy 0.14 Electron carrier; component of water Manganese Mn Fluorine F 9 Phosphorus Backbone of nucleic acids; important in energy transfer Sulfur 16 0.03 Component of me Chlorine C Principal negative ion outside cells Chromium Coppe Key component of many enzymes Component of all proteins and nucleic Trace Trace Cobalt Trace Ir Zinc Trace Trace Key component of some enzymes Molybdenu M Tr race Key component of many enzymes race Trace Component of thyroid hormone Distribution of the elements earths crust. For example, silicon, aluminum, and iron con- Of the 92 naturally occurring elements on earth, only 11 are stitute 39.2%of the earths crust, but they exist in trace amounts in the human body. On the other hand, carbon at found in organisms in more than trace amounts(0.01% or oms make up 18.5% of the human body but only 0.03%of higher). These 11 elements have atomic numbers less than 21 and. thus have low atomic masses, Table 2.1 lists the the earth,s crust. levels of various elements in the human body; their levels in other organisms are similar. Inspection of this table suggests Ninety-two elements occur naturally on earth; only that the distribution of elements in living systems is by no eleven of them are found in significant amounts in living means accidental. The most common elements inside or organisms. Four of them--oxygen, hydrogen, carbon, ganisms are not the elements that are most abundant in the itrogen-constitute 96. 3% of the weight of your body Chapter 2 The Nature of Molecules 25
Distribution of the Elements Of the 92 naturally occurring elements on earth, only 11 are found in organisms in more than trace amounts (0.01% or higher). These 11 elements have atomic numbers less than 21 and, thus, have low atomic masses. Table 2.1 lists the levels of various elements in the human body; their levels in other organisms are similar. Inspection of this table suggests that the distribution of elements in living systems is by no means accidental. The most common elements inside organisms are not the elements that are most abundant in the earth’s crust. For example, silicon, aluminum, and iron constitute 39.2% of the earth’s crust, but they exist in trace amounts in the human body. On the other hand, carbon atoms make up 18.5% of the human body but only 0.03% of the earth’s crust. Ninety-two elements occur naturally on earth; only eleven of them are found in significant amounts in living organisms. Four of them—oxygen, hydrogen, carbon, nitrogen—constitute 96.3% of the weight of your body. Chapter 2 The Nature of Molecules 25 Table 2.1 The Most Common Elements on Earth and Their Distribution in the Human Body Approximate Percent of Percent of Earth’s Crust Human Body Element Symbol Atomic Number by Weight by Weight Importance or Function Oxygen Silicon Aluminum Iron Calcium Sodium Potassium Magnesium Hydrogen Manganese Fluorine Phosphorus Carbon Sulfur Chlorine Vanadium Chromium Copper Nitrogen Boron Cobalt Zinc Selenium Molybdenum Tin Iodine O Si Al Fe Ca Na K Mg H Mn F P C S Cl V Cr Cu N B Co Zn Se Mo Sn I 8 14 13 26 20 11 19 12 1 25 9 15 6 16 17 23 24 29 7 5 27 30 34 42 50 53 46.6 27.7 6.5 5.0 3.6 2.8 2.6 2.1 0.14 0.1 0.07 0.07 0.03 0.03 0.01 0.01 0.01 0.01 Trace Trace Trace Trace Trace Trace Trace Trace 65.0 Trace Trace Trace 1.5 0.2 0.4 0.1 9.5 Trace Trace 1.0 18.5 0.3 0.2 Trace Trace Trace 3.3 Trace Trace Trace Trace Trace Trace Trace Required for cellular respiration; component of water Critical component of hemoglobin in the blood Component of bones and teeth; triggers muscle contraction Principal positive ion outside cells; important in nerve function Principal positive ion inside cells; important in nerve function Critical component of many energytransferring enzymes Electron carrier; component of water and most organic molecules Backbone of nucleic acids; important in energy transfer Backbone of organic molecules Component of most proteins Principal negative ion outside cells Key component of many enzymes Component of all proteins and nucleic acids Key component of some enzymes Key component of many enzymes Component of thyroid hormone