{"id":112,"date":"2017-12-14T21:26:16","date_gmt":"2017-12-14T21:26:16","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/chapter\/atomic-theory\/"},"modified":"2025-06-11T19:43:49","modified_gmt":"2025-06-11T19:43:49","slug":"atomic-theory","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/chapter\/atomic-theory\/","title":{"raw":"3.3 The Atom and Atomic Theory","rendered":"3.3 The Atom and Atomic Theory"},"content":{"raw":"<div class=\"textbox learning-objectives\">\r\n<h3>Learning Objectives<\/h3>\r\nBy the end of this section, you will be able to:\r\n<ul>\r\n \t<li>Outline milestones in the development of modern atomic theory<\/li>\r\n \t<li>Summarize and interpret the results of the experiments of Thomson, Millikan, and Rutherford<\/li>\r\n \t<li>Describe the three subatomic particles that compose atoms<\/li>\r\n \t<li>Define isotopes and give examples for several elements<\/li>\r\n \t<li>Write isotopic symbols for elements and ions.<\/li>\r\n<\/ul>\r\n<\/div>\r\n<p id=\"ball-ch03_s01_p01\" class=\"para editable block\">The smallest piece of an element that maintains the identity of that element is called an <span class=\"margin_term\"><a class=\"glossterm\">atom<\/a><\/span>. Individual atoms are extremely small. It would take about fifty million atoms in a row to make a line that is 1 cm long. The period at the end of a printed sentence has several million atoms in it. Atoms are so small that it is difficult to believe that all matter is made from atoms\u2014but it is.<\/p>\r\n<p id=\"ball-ch03_s01_p02\" class=\"para editable block\">The concept that atoms play a fundamental role in chemistry is formalized by the <span class=\"margin_term\"><a class=\"glossterm\">modern atomic theory<\/a><\/span>, first stated by John Dalton, an English scientist, in 1808. It consists of three parts:<\/p>\r\n\r\n<ol id=\"ball-ch03_s01_l02\" class=\"orderedlist editable block\">\r\n \t<li>All matter is composed of atoms.<\/li>\r\n \t<li>Atoms of the same element are the same; atoms of different elements are different.<\/li>\r\n \t<li>Atoms combine in whole-number ratios to form compounds.<\/li>\r\n<\/ol>\r\n<p id=\"ball-ch03_s01_p03\" class=\"para editable block\">These concepts form the basis of chemistry.<\/p>\r\n\r\n<div id=\"ball-ch03_s01_f01\" class=\"figure large editable block\">\r\n\r\n\u00a0In the two centuries since Dalton developed his ideas, scientists have made significant progress in furthering our understanding of atomic theory. Much of this came from the results of several seminal experiments that revealed the details of the internal structure of atoms. Here, we will discuss some of those key developments, with an emphasis on application of the scientific method, as well as understanding how the experimental evidence was analyzed. While the historical persons and dates behind these experiments can be quite interesting, it is most important to understand the concepts resulting from their work.\r\n\r\n<\/div>\r\n<h2>Evolution of the Atom<\/h2>\r\nIf matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? In the late 1800s, a number of scientists interested in questions like these investigated the electrical discharges that could be produced in low-pressure gases, with the most significant discovery made by English physicist J. J. <strong>Thomson<\/strong> using a <strong>cathode ray<\/strong> tube. This apparatus consisted of a sealed glass tube from which almost all the air had been removed; the tube contained two metal electrodes. When high voltage was applied across the electrodes, a visible beam called a cathode ray appeared between them. This beam was deflected toward the positive charge and away from the negative charge, and was produced in the same way with identical properties when different metals were used for the electrodes. In similar experiments, the ray was simultaneously deflected by an applied magnetic field, and measurements of the extent of deflection and the magnetic field strength allowed Thomson to calculate the charge-to-mass ratio of the cathode ray particles. The results of these measurements indicated that these particles were much lighter than atoms (Figure 1).\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"880\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211009\/CNX_Chem_02_02_CathodeRay1.jpg\" alt=\"Figure A shows a photo of J. J. Thomson working at a desk. Figure B shows a photograph of a cathode ray tube. It is a long, glass tube that is narrow at the left end but expands into a large bulb on the right end. The entire cathode tube is sitting on a wooden stand. Figure C shows the parts of the cathode ray tube. The cathode ray tube consists of a cathode and an anode. The cathode, which has a negative charge, is located in a small bulb of glass on the left side of the cathode ray tube. To the left of the cathode it says \u201cHigh voltage\u201d and indicates a positive and negative charge. The anode, which has a positive charge, is located to the right of the cathode. Two charged plates are located to the right of the anode, and are connected to a battery and two magnets. The magnets are labeled \u201cS\u201d and \u201cN.\u201d A cathode ray is generated from the cathode, travels through the anode and into a wider part of the cathode ray tube, where it travels between a positively charged electrode plate and a negatively charged electrode plate. The ray bends upward and continues to travel until it hits the wide part of the tube on the right. The rightmost end of the tube contains a printed scale that allows one to measure how much the ray was deflected.\" width=\"880\" height=\"616\" \/> Figure 1. (a) J. J. Thomson produced a visible beam in a cathode ray tube. (b) This is an early cathode ray tube, invented in 1897 by Ferdinand Braun. (c) In the cathode ray, the beam (shown in yellow) comes from the cathode and is accelerated past the anode toward a fluorescent scale at the end of the tube. Simultaneous deflections by applied electric and magnetic fields permitted Thomson to calculate the mass-to-charge ratio of the particles composing the cathode ray. (credit a: modification of work by Nobel Foundation; credit b: modification of work by Eugen Nesper; credit c: modification of work by \u201cKurzon\u201d\/Wikimedia Commons)[\/caption]\r\n\r\nBased on his observations, here is what Thomson proposed and why: The particles are attracted by positive (+) charges and repelled by negative (-) charges, so they must be negatively charged (like charges repel and unlike charges attract); they are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms. Although controversial at the time, Thomson\u2019s idea was gradually accepted, and his cathode ray particle is what we now call an <strong>electron<\/strong>, a negatively charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term \u201celectron\u201d was coined in 1891 by Irish physicist George Stoney, from \u201c<em>electr<\/em>ic i<em>on<\/em>.\u201d\r\n<div class=\"textbox\">Click <a href=\"https:\/\/www.aip.org\/history\/electron\/jjsound.htm\" target=\"_blank\" rel=\"noopener noreferrer\">this link to \"JJ Thompson Talks About the Size of the Electron\"<\/a> to hear Thomson describe his discovery in his own voice.<\/div>\r\nScientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson and others, the charge and mass of the negative, subatomic particles\u2014the electrons\u2014were known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson proposed the \u201cplum pudding\u201d model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it, since all atoms are electrically neutral. A competing model had been proposed in 1903 by Hantaro <strong>Nagaoka<\/strong>, who postulated a Saturn-like atom, consisting of a positively charged sphere surrounded by a halo of electrons (Figure 2).\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"878\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211013\/CNX_Chem_02_02_AtomModels1.jpg\" alt=\"Figure A shows a photograph of plum pudding, which is a thick, almost spherical cake containing raisins throughout. To the right, an atom model is round and contains negatively charged electrons embedded within a sphere of positively charged matter. Figure B shows a photograph of the planet Saturn, which has rings. To the right, an atom model is a sphere of positively charged matter encircled by a ring of negatively charged electrons.\" width=\"878\" height=\"266\" \/> Figure 2. (a) Thomson suggested that atoms resembled plum pudding, an English dessert consisting of moist cake with embedded raisins (\u201cplums\u201d). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of electrons surrounding a positive \u201cplanet.\u201d (credit a: modification of work by \u201cMan vyi\u201d\/Wikimedia Commons; credit b: modification of work by \u201cNASA\u201d\/Wikimedia Commons)[\/caption]\r\n\r\nThe next major development in understanding the atom came from Ernest <strong>Rutherford<\/strong>, a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged <strong>alpha particles (\u03b1 particles)<\/strong> that were produced by the radioactive decay of radium; \u03b1 particles consist of two protons and two neutrons (you will learn more about radioactive decay in the chapter on nuclear chemistry). Rutherford and his colleagues Hans <strong>Geiger<\/strong> (later famous for the Geiger counter) and Ernest <strong>Marsden<\/strong> aimed a beam of \u03b1 particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the \u03b1 particles using a luminescent screen that glowed briefly where hit by an \u03b1 particle.\r\n\r\nWhat did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source (Figure 3). Rutherford described finding these results: \u201cIt was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you.\u201d\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"880\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211015\/CNX_Chem_02_02_Rutherford1.jpg\" alt=\"This figure shows a box on the left that contains a radium source of alpha particles which generates a beam of alpha particles. The beam travels through an opening within a ring-shaped luminescent screen which is used to detect scattered alpha particles. A piece of thin gold foil is at the center of the ring formed by the screen. When the beam encounters the gold foil, most of the alpha particles pass straight through it and hit the luminescent screen directly behind the foil. Some of the alpha particles are slightly deflected by the foil and hit the luminescent screen off to the side of the foil. Some alpha particles are significantly deflected and bounce back to hit the front of the screen.\" width=\"880\" height=\"386\" \/> Figure 3. Geiger and Rutherford fired \u03b1 particles at a piece of gold foil and detected where those particles went, as shown in this schematic diagram of their experiment. Most of the particles passed straight through the foil, but a few were deflected slightly and a very small number were significantly deflected.[\/caption]\r\n\r\nHere is what Rutherford deduced: Because most of the fast-moving \u03b1 particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom. Alpha particles are positively charged, so deflections arose when they encountered another positive charge (like charges repel each other). Since like charges repel one another, the few positively charged \u03b1 particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge. Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:\r\n<ol>\r\n \t<li>The volume occupied by an atom must consist of a large amount of empty space.<\/li>\r\n \t<li>A small, relatively heavy, positively charged body, the <strong>nucleus<\/strong>, must be at the center of each atom.<\/li>\r\n<\/ol>\r\n<div class=\"textbox\">View <a href=\"https:\/\/micro.magnet.fsu.edu\/electromag\/java\/rutherford\/\" target=\"_blank\" rel=\"noopener noreferrer\">this simulation of the Rutherford gold foil experiment<\/a>. Adjust the slit width to produce a narrower or broader beam of \u03b1 particles to see how that affects the scattering pattern.<\/div>\r\nThis analysis led Rutherford to propose a model in which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the negatively charged electrons, so that the atom is electrically neutral (Figure 4). After many more experiments, Rutherford also discovered that the nuclei of other elements contain the hydrogen nucleus as a \u201cbuilding block,\u201d and he named this more fundamental particle the <strong>proton<\/strong>, the positively charged, subatomic particle found in the nucleus. With one addition, which you will learn next, this nuclear model of the atom, proposed over a century ago, is still used today.\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"880\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211016\/CNX_Chem_02_02_GoldFoil31.jpg\" alt=\"The left diagram shows a green beam of alpha particles hitting a rectangular piece of gold foil. Some of the alpha particles bounce backwards after hitting the foil. However, most of the particles travel through the foil, with some being deflected as they pass through the foil. A callout box shows a magnified cross section of the gold foil. Most of the alpha particles are not deflected, but pass straight through the foil because they travel between the gold atoms. A very small number of alpha particles are significantly deflected when they hit the nucleus of the gold atoms straight on. A few alpha particles are slightly deflected because they glanced off of the nucleus of a gold atom.\" width=\"880\" height=\"515\" \/> Figure 4. The \u03b1 particles are deflected only when they collide with or pass close to the much heavier, positively charged gold nucleus. Because the nucleus is very small compared to the size of an atom, very few \u03b1 particles are deflected. Most pass through the relatively large region occupied by electrons, which are too light to deflect the rapidly moving particles.[\/caption]\r\n\r\n<div class=\"textbox\">The <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\" target=\"_blank\" rel=\"noopener noreferrer\">Rutherford Scattering simulation<\/a> allows you to investigate the differences between a \u201cplum pudding\u201d atom and a Rutherford atom by firing \u03b1 particles at each type of atom.<\/div>\r\nAnother important finding was the discovery of isotopes. During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a \u201cnew element\u201d produced by the radioactive decay of thorium was initially given the name mesothorium. However, a more detailed analysis showed that mesothorium was chemically identical to radium (another decay product), despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick <strong>Soddy<\/strong> to realize that an element could have types of atoms with different masses that were chemically indistinguishable. These different types are called <strong>isotopes<\/strong>\u2014atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.\r\n\r\nOne puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus. As you might expect, detecting uncharged particles is very challenging, and it was not until 1932 that James <strong>Chadwick<\/strong> found evidence of <strong>neutrons<\/strong>, uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes: They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons. With the discovery of the neutron, the modern model of an atom was established (Figure 5).\r\n\r\n[caption id=\"attachment_4627\" align=\"aligncenter\" width=\"585\"]<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/The-Atom.png\"><img class=\"wp-image-4627\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/2835\/2017\/12\/14212601\/The-Atom-1.png\" alt=\"The Atom\" width=\"585\" height=\"336\" \/><\/a> Figure 5. The structure of an atom. The nucleus of an atom is composed of protons and neutrons. The nucleus is surrounded by an electron cloud.[\/caption]\r\n\r\nThe diameter of an atom is on the order of 10<sup>\u221210<\/sup> m, whereas the diameter of the nucleus is roughly 10<sup>\u221215<\/sup> m\u2014about 100,000 times smaller. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium (Figure 6).\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"1300\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211018\/CNX_Chem_02_03_AtomSize1.jpg\" alt=\"The diagram on the left shows a picture of an atom that is 10 to the negative tenth power meters in diameter. The nucleus is labeled at the center of the atom and is 10 to the negative fifteenth power meters. The central figure shows a photograph of an American football stadium. The figure on the right shows a photograph of a person with a handful of blueberries.\" width=\"1300\" height=\"400\" \/> Figure 6. If an atom could be expanded to the size of a football stadium, the nucleus would be the size of a single blueberry. (credit middle: modification of work by \u201cbabyknight\u201d\/Wikimedia Commons; credit right: modification of work by Paxson Woelber)[\/caption]\r\n\r\nAtoms\u2014and the protons, neutrons, and electrons that compose them\u2014are extremely small. For example, a carbon atom weighs less than 2 \u00d7 10<sup>\u221223<\/sup> g, and an electron has a charge of less than 2 \u00d7 10<sup>\u221219<\/sup> C (coulomb). When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the<strong> atomic mass unit (amu)<\/strong> and the <strong>fundamental unit of charge (e)<\/strong>. The amu was originally defined based on hydrogen, the lightest element, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. (This isotope is known as \u201ccarbon-12\u201d as will be discussed later in this module.) Thus, one amu is exactly [latex]\\frac{1}{12}[\/latex] of the mass of one carbon-12 atom: 1 amu = 1.6605 \u00d7 10<sup>\u221224<\/sup> g. (The <strong>Dalton (Da)<\/strong> and the <strong>unified atomic mass unit (u) <\/strong>are alternative units that are equivalent to the amu.) The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 \u00d7 10<sup>\u221219<\/sup> C.\r\n\r\nA proton has a mass of 1.0073 amu and a charge of 1+. A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero; as its name suggests, it is neutral. The electron has a charge of 1\u2212 and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton. The properties of these fundamental particles are summarized in Table 1. (An observant student might notice that the sum of an atom\u2019s subatomic particles does not equal the atom\u2019s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than 12.00 amu. This \u201cmissing\u201d mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)\r\n<table id=\"fs-idp90857696\" class=\"span-all\" summary=\"This table gives the name, location, charge in C, unit charge, mass in A M U and mass in grams for electrons, protons and neutrons. Electrons are located outside of the nucleus, have a charge of negative 1.602 times 10 to the negative nineteenth power, a unit charge of negative 1, and a mass of 0.00055 A M U or 0.00091 times 10 to the negative twenty-fourth power grams. Protons are located within the nucleus, have a charge of 1.602 times 10 to the negative nineteenth power, have a unit charge of positive 1, and have a mass of 1.0073 A M U or 1.6726 times 10 to the negative twenty-fourth power grams. Neutrons are located within the nucleus, have a charge of 0, have a unit charge of 0, and have a mass of 1.0087 A M U or 1.6749 times 10 to the negative twenty-fourth power grams.\">\r\n<thead>\r\n<tr>\r\n<th colspan=\"6\">Table 1. Properties of Subatomic Particles<\/th>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<th>Name<\/th>\r\n<th>Location<\/th>\r\n<th>Charge (C)<\/th>\r\n<th>Unit Charge<\/th>\r\n<th>Mass (amu)<\/th>\r\n<th>Mass (g)<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr valign=\"top\">\r\n<td>electron<\/td>\r\n<td>outside nucleus<\/td>\r\n<td>\u22121.602 \u00d7 10<sup>\u221219<\/sup><\/td>\r\n<td>1\u2212<\/td>\r\n<td>0.00055<\/td>\r\n<td>0.00091 \u00d7 10<sup>\u221224<\/sup><\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>proton<\/td>\r\n<td>nucleus<\/td>\r\n<td>1.602 \u00d7 10<sup>\u221219<\/sup><\/td>\r\n<td>1+<\/td>\r\n<td>1.00727<\/td>\r\n<td>1.67262 \u00d7 10<sup>\u221224<\/sup><\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>neutron<\/td>\r\n<td>nucleus<\/td>\r\n<td>0<\/td>\r\n<td>0<\/td>\r\n<td>1.00866<\/td>\r\n<td>1.67493 \u00d7 10<sup>\u221224<\/sup><\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\nThe number of protons in the nucleus of an atom is its <strong>atomic number (Z)<\/strong>. This is the defining trait of an element: Its value determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its <strong>mass number (A)<\/strong>. The number of neutrons is therefore the difference between the mass number and the atomic number: A \u2013 Z = number of neutrons.\r\n\r\n[latex]\\begin{array}{ccc}\\hfill \\text{atomic number}\\left(\\text{Z}\\right)&amp; =&amp; \\text{number of protons}\\hfill \\\\ \\hfill \\text{atomic mass}\\left(\\text{A}\\right)&amp; =&amp; \\text{number of protons}+\\text{number of neutrons}\\hfill \\\\ \\hfill \\text{A}-\\text{Z}&amp; =&amp; \\text{number of neutrons}\\hfill \\end{array}[\/latex]\r\n\r\nAtoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. When the numbers of these subatomic particles are <em>not<\/em> equal, the atom is electrically charged and is called an <strong>ion<\/strong>. The charge of an atom is defined as follows:\r\n\r\nAtomic charge = number of protons \u2212 number of electrons\r\n\r\nAs will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an <strong>anion<\/strong>. Positively charged atoms called <strong>cations<\/strong> are formed when an atom loses one or more electrons. For example, a neutral sodium atom (Z = 11) has 11 electrons. If this atom loses one electron, it will become a cation with a 1+ charge (11 \u2212 10 = 1+). A neutral oxygen atom (Z = 8) has eight electrons, and if it gains two electrons it will become an anion with a 2\u2212 charge (8 \u2212 10 = 2\u2212).\r\n<div class=\"textbox examples\">\r\n<h3>Example 1: <strong>Composition of an Atom<\/strong><\/h3>\r\nIodine is an essential trace element in our diet; it is needed to produce thyroid hormone. Insufficient iodine in the diet can lead to the development of a goiter, an enlargement of the thyroid gland (Figure 2).\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"699\"]<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211019\/CNX_Chem_02_03_Iodine1.jpg\" alt=\"Figure A shows a photo of a person who has a very swollen thyroid in his or her neck. Figure B shows a photo of a canister of iodized salt.\" width=\"699\" height=\"310\" \/> Figure 2. (a) Insufficient iodine in the diet can cause an enlargement of the thyroid gland called a goiter. (b) The addition of small amounts of iodine to salt, which prevents the formation of goiters, has helped eliminate this concern in the US where salt consumption is high. (credit a: modification of work by \u201cAlmazi\u201d\/Wikimedia Commons; credit b: modification of work by Mike Mozart)[\/caption]\r\n\r\n&nbsp;\r\n\r\nThe addition of small amounts of iodine to table salt (iodized salt) has essentially eliminated this health concern in the United States, but as much as 40% of the world\u2019s population is still at risk of iodine deficiency. The iodine atoms are added as anions, and each has a 1\u2212 charge and a mass number of 127. Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.\r\n\r\n[reveal-answer q=\"754651\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"754651\"]\r\n\r\nThe atomic number of iodine (53) tells us that a neutral iodine atom contains 53 protons in its nucleus and 53 electrons outside its nucleus. Because the sum of the numbers of protons and neutrons equals the mass number, 127, the number of neutrons is 74 (127 \u2212 53 = 74). Since the iodine is added as a 1\u2212 anion, the number of electrons is 54 [53 \u2212 (\u22121) = 54].\r\n\r\n[\/hidden-answer]\r\n<h4><strong>Check Your Learning<\/strong><\/h4>\r\nAn atom of platinum has a mass number of 195 and contains 74 electrons. How many protons and neutrons does it contain, and what is its charge?\r\n\r\n[reveal-answer q=\"488361\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"488361\"]78 protons; 117 neutrons; charge is 4+[\/hidden-answer]\r\n\r\n<\/div>\r\n<h2>Chemical Symbols<\/h2>\r\n[caption id=\"\" align=\"alignright\" width=\"350\"]<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211020\/CNX_Chem_02_03_SiSymbol1.jpg\" alt=\"A jar labeled \u201cH g\u201d is shown with a small amount of liquid mercury in it.\" width=\"350\" height=\"280\" \/> Figure 7. The symbol Hg represents the element mercury regardless of the amount; it could represent one atom of mercury or a large amount of mercury.[\/caption]\r\n\r\nA <strong>chemical symbol<\/strong> is an abbreviation that we use to indicate an element or an atom of an element. For example, the symbol for mercury is Hg (Figure 7). We use the same symbol to indicate one atom of mercury (microscopic domain) or to label a container of many atoms of the element mercury (macroscopic domain).\r\n\r\nThe symbols for several common elements and their atoms are listed in Table 2. Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Most symbols have one or two letters, but three-letter symbols have been used to describe some elements that have atomic numbers greater than 112.\r\n\r\nTo avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen (O). All known elements and their symbols are in <a title=\"The Periodic Table\" href=\".\/chapter\/the-periodic-table-3\/\" target=\"_blank\" rel=\"noopener\">the periodic table<\/a>.\r\n<table id=\"fs-idm36686800\" class=\"span-all\" summary=\"This table has two columns labeled element and symbol. The first letter of the symbol is always an uppercase letter while the second letter of the symbol is always a lowercase letter. Aluminum has the symbol A L. Bromine has the symbol B R, calcium has the symbol C A, carbon has the symbol C, chlorine has the symbol C L, chromium has the symbol C R, cobalt has the symbol C O, copper has the symbol C U, from cuprum, fluorine has the symbol F, gold has the symbol A U, from aurum, helium has the symbol H E, hydrogen has the symbol H, iodine has the symbol I, iron has the symbol F E, from ferrum, lead has the symbol P B, from plumbum, magnesium has the symbol M G, mercury has the symbol H G from hydrargyrum, nitrogen has the symbol N, oxygen has the symbol O, potassium has the symbol K, from kalium, silicon has the symbol S I, silver has the symbol A G, from argentum, sodium has the symbol N A from natrium, sulfur has the symbol S, tin has the symbol S N from stannum, and zinc has the symbol Z N.\">\r\n<thead>\r\n<tr>\r\n<th colspan=\"4\">Table 2. Some Common Elements and Their Symbols<\/th>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<th>Element<\/th>\r\n<th>Symbol<\/th>\r\n<th>Element<\/th>\r\n<th>Symbol<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr valign=\"top\">\r\n<td>aluminum<\/td>\r\n<td>Al<\/td>\r\n<td>iron<\/td>\r\n<td>Fe (from <em>ferrum<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>bromine<\/td>\r\n<td>Br<\/td>\r\n<td>lead<\/td>\r\n<td>Pb (from <em>plumbum<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>calcium<\/td>\r\n<td>Ca<\/td>\r\n<td>magnesium<\/td>\r\n<td>Mg<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>carbon<\/td>\r\n<td>C<\/td>\r\n<td>mercury<\/td>\r\n<td>Hg (from <em>hydrargyrum<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>chlorine<\/td>\r\n<td>Cl<\/td>\r\n<td>nitrogen<\/td>\r\n<td>N<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>chromium<\/td>\r\n<td>Cr<\/td>\r\n<td>oxygen<\/td>\r\n<td>O<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>cobalt<\/td>\r\n<td>Co<\/td>\r\n<td>potassium<\/td>\r\n<td>K (from <em>kalium<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>copper<\/td>\r\n<td>Cu (from <em>cuprum<\/em>)<\/td>\r\n<td>silicon<\/td>\r\n<td>Si<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>fluorine<\/td>\r\n<td>F<\/td>\r\n<td>silver<\/td>\r\n<td>Ag (from<em> argentum<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>gold<\/td>\r\n<td>Au (from <em>aurum<\/em>)<\/td>\r\n<td>sodium<\/td>\r\n<td>Na (from <em>natrium<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>helium<\/td>\r\n<td>He<\/td>\r\n<td>sulfur<\/td>\r\n<td>S<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>hydrogen<\/td>\r\n<td>H<\/td>\r\n<td>tin<\/td>\r\n<td>Sn (from <em>stannum<\/em>)<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>iodine<\/td>\r\n<td>I<\/td>\r\n<td>zinc<\/td>\r\n<td>Zn<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\nTraditionally, the discoverer (or discoverers) of a new element names the element. However, until the name is recognized by the International Union of Pure and Applied Chemistry (IUPAC), the recommended name of the new element is based on the Latin word(s) for its atomic number. For example, element 106 was called unnilhexium (Unh), element 107 was called unnilseptium (Uns), and element 108 was called unniloctium (Uno) for several years. These elements are now named after scientists (or occasionally locations); for example, element 106 is now known as <em>seaborgium<\/em> (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements.\r\n<div class=\"textbox\">Visit <a href=\"https:\/\/iupac.org\/\" target=\"_blank\" rel=\"noopener\">this site to learn more about IUPAC, the International Union of Pure and Applied Chemistry<\/a>, and explore its periodic table.<\/div>\r\n<h2>Isotopes<\/h2>\r\nThe symbol for a specific isotope of any element is written by placing the mass number as a superscript to the left of the element symbol (Figure 8). The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element\u2019s identity, as does its symbol, it is often omitted. For example, magnesium exists as a mixture of three isotopes, each with an atomic number of 12 and with mass numbers of 24, 25, and 26, respectively. These isotopes can be identified as <sup>24<\/sup>Mg, <sup>25<\/sup>Mg, and <sup>26<\/sup>Mg. These isotope symbols are read as \u201celement, mass number\u201d and can be symbolized consistent with this reading. For instance, <sup>24<\/sup>Mg is read as \u201cmagnesium 24,\u201d and can be written as \u201cmagnesium-24\u201d or \u201cMg-24.\u201d <sup>25<\/sup>Mg is read as \u201cmagnesium 25,\u201d and can be written as \u201cmagnesium-25\u201d or \u201cMg-25.\u201d All magnesium atoms have 12 protons in their nucleus. They differ only because a <sup>24<\/sup>Mg atom has 12 neutrons in its nucleus, a <sup>25<\/sup>Mg atom has 13 neutrons, and a <sup>26<\/sup>Mg has 14 neutrons.\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"650\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211023\/CNX_Chem_02_03_AtomSym1.jpg\" alt=\"This diagram shows the symbol for helium, \u201cH e.\u201d The number to the upper left of the symbol is the mass number, which is 4. The number to the upper right of the symbol is the charge which is positive 2. The number to the lower left of the symbol is the atomic number, which is 2. This number is often omitted. Also shown is \u201cM g\u201d which stands for magnesium It has a mass number of 24, a charge of positive 2, and an atomic number of 12.\" width=\"650\" height=\"144\" \/> Figure 8. The symbol for an atom indicates the element via its usual two-letter symbol, the mass number as a left superscript, the atomic number as a left subscript (sometimes omitted), and the charge as a right superscript.[\/caption]\r\n\r\nInformation about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table 3. Note that in addition to standard names and symbols, the isotopes of hydrogen are often referred to using common names and accompanying symbols. Hydrogen-2, symbolized <sup>2<\/sup>H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized <sup>3<\/sup>H, is also called tritium and sometimes symbolized T.\r\n<table id=\"fs-idm87646592\" class=\"span-all\" summary=\"This table has seven columns labeled element, symbol, atomic number, number of protons, number of neutrons, mass in A M U, and percent natural abundance. The symbols for each element each show the mass number in the upper left and the atomic number in the lower left. Therefore hydrogen left superscript 1, left subscript 1, or protium, has a mass number of 1 and an atomic number of 1. Protium has one proton, 0 neutrons, a mass of 1.0078 and a natural abundance percentage of 99.985. Hydrogen left superscript 2, left subscript 1, or deuterium, has an atomic number of 1, 1 proton, 1 neutron, a mass of 2.0141 and a natural abundance percentage of 0.015. Hydrogen left superscript 3, left subscript 1, or tritium, has an atomic number of 11 protons, 2 neutrons, and a mass of 3.01605. No natural abundance percentage is given. Helium left superscript 3, left subscript 2 has an atomic number of 2, 2 protons, 1 neutron, a mass of 3.01603, and a natural abundance percentage of 0.00013. Helium left superscript 4, left subscript 2 has an atomic number of 2, 2 protons, 2 neutrons, a mass of 4.0026 and a natural abundance percentage of 100. Lithium left superscript 6, left subscript 3 has an atomic number of 3, 3 protons, 3 neutrons, a mass of 6.0151, and a natural abundance percentage of 7.42. Lithium left superscript 7, left subscript 3 has an atomic number of 3, 3 protons, 4 neutrons, a mass of 7.0160, and a natural abundance percentage of 92.8. Beryllium left superscript 9, left subscript 4 has an atomic number of 4, 4 protons, 5 neutrons, a mass of 9.0122, and a natural abundance percentage of 100. Boron left superscript 10, left subscript 5 has an atomic number of 5, 5 protons, 5 neutrons and a natural abundance percentage of 19.9. Boron left superscript 11, left subscript 5 has an atomic number of 5, 5 protons, 6 neutrons, a mass of 11.0093 and a natural abundance of 80.1. Carbon left superscript 12, left subscript 6 has an atomic number of 6, 6 protons, 6 neutrons, a mass of 12, and a natural abundance percentage of 98.89. Carbon left superscript 13, left subscript 6 has an atomic number of 6, 6 protons, 7 neutrons, a mass of 13.0033, and a natural abundance percentage of 1.11. Carbon left superscript 14, left subscript 6 has an atomic number of 6, 6 protons, 8 neutrons, and a mass of 14.0032. Its natural abundance percentage is not reported. Nitrogen left superscript 14, left subscript 7 has an atomic number of 7, 7 protons, 7 neutrons, a mass of 14.0031, and a natural abundance percentage of 99.63. Nitrogen left superscript 15, left subscript 7 has an atomic number of 7, 7 protons, 8 neutrons, a mass of 15.0001, and a natural abundance percentage of 0.37. Oxygen left superscript 16, left subscript 8 has an atomic number of 8, 8 protons, 8 neutrons, a mass of 15.9949, and a natural abundance percentage of 99.759. Oxygen left superscript 17, left subscript 8 has an atomic number of 8, 8 protons, 9 neutrons, a mass of 16.9991, and a natural abundance percentage of 0.037. Oxygen left superscript 18, left subscript 8 has an atomic number of 8, 8 protons, 10 neutrons, a mass of 17.9992, and a natural abundance percentage of 0.204. Fluorine left superscript 19, left subscript 9 has an atomic number of 9, 9 protons, 10 neutrons, a mass of 18.9984, and a natural abundance percentage of 100. Neon left superscript 20, left subscript 10 has an atomic number of 10, 10 protons, 10 neutrons, a mass of 19.9924, and a natural abundance percentage of 90.92. Neon left superscript 21, left subscript 10 has an atomic number of 10, 10 protons, 11 neutrons, a mass of 20.994, and a natural abundance percentage of 0.257. Neon left superscript 22, left subscript 10 has an atomic number of 10, 10 protons, 12 neutrons, a mass of 21.9914, and a natural abundance percentage of 8.82.\">\r\n<thead>\r\n<tr>\r\n<th colspan=\"7\">Table 3. Nuclear Compositions of Atoms of the Very Light Elements<\/th>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<th>Element<\/th>\r\n<th>Symbol<\/th>\r\n<th>Atomic Number<\/th>\r\n<th>Number of Protons<\/th>\r\n<th>Number of Neutrons<\/th>\r\n<th>Mass (amu)<\/th>\r\n<th>% Natural Abundance<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"3\">hydrogen<\/td>\r\n<td>[latex]{}_{1}^{1}\\text{H}[\/latex]\r\n(protium)<\/td>\r\n<td>1<\/td>\r\n<td>1<\/td>\r\n<td>0<\/td>\r\n<td>1.0078<\/td>\r\n<td>99.989<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{1}^{2}\\text{H}[\/latex]\r\n(deuterium)<\/td>\r\n<td>1<\/td>\r\n<td>1<\/td>\r\n<td>1<\/td>\r\n<td>2.0141<\/td>\r\n<td>0.0115<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{1}^{3}\\text{H}[\/latex]\r\n(tritium)<\/td>\r\n<td>1<\/td>\r\n<td>1<\/td>\r\n<td>2<\/td>\r\n<td>3.01605<\/td>\r\n<td>\u2014 (trace)<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"2\">helium<\/td>\r\n<td>[latex]{}_{2}^{3}\\text{He}[\/latex]<\/td>\r\n<td>2<\/td>\r\n<td>2<\/td>\r\n<td>1<\/td>\r\n<td>3.01603<\/td>\r\n<td>0.00013<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{2}^{4}\\text{He}[\/latex]<\/td>\r\n<td>2<\/td>\r\n<td>2<\/td>\r\n<td>2<\/td>\r\n<td>4.0026<\/td>\r\n<td>100<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"2\">lithium<\/td>\r\n<td>[latex]{}_{3}^{6}\\text{Li}[\/latex]<\/td>\r\n<td>3<\/td>\r\n<td>3<\/td>\r\n<td>3<\/td>\r\n<td>6.0151<\/td>\r\n<td>7.59<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{3}^{7}\\text{Li}[\/latex]<\/td>\r\n<td>3<\/td>\r\n<td>3<\/td>\r\n<td>4<\/td>\r\n<td>7.0160<\/td>\r\n<td>92.41<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>beryllium<\/td>\r\n<td>[latex]{}_{4}^{9}\\text{Be}[\/latex]<\/td>\r\n<td>4<\/td>\r\n<td>4<\/td>\r\n<td>5<\/td>\r\n<td>9.0122<\/td>\r\n<td>100<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"2\">boron<\/td>\r\n<td>[latex]{}_{5}^{10}\\text{B}[\/latex]<\/td>\r\n<td>5<\/td>\r\n<td>5<\/td>\r\n<td>5<\/td>\r\n<td>10.0129<\/td>\r\n<td>19.9<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{5}^{11}\\text{B}[\/latex]<\/td>\r\n<td>5<\/td>\r\n<td>5<\/td>\r\n<td>6<\/td>\r\n<td>11.0093<\/td>\r\n<td>80.1<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"3\">carbon<\/td>\r\n<td>[latex]{}_{6}^{12}\\text{C}[\/latex]<\/td>\r\n<td>6<\/td>\r\n<td>6<\/td>\r\n<td>6<\/td>\r\n<td>12.0000<\/td>\r\n<td>98.89<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{6}^{13}\\text{C}[\/latex]<\/td>\r\n<td>6<\/td>\r\n<td>6<\/td>\r\n<td>7<\/td>\r\n<td>13.0034<\/td>\r\n<td>1.11<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]{}_{6}^{14}\\text{C}[\/latex]<\/td>\r\n<td>6<\/td>\r\n<td>6<\/td>\r\n<td>8<\/td>\r\n<td>14.0032<\/td>\r\n<td>\u2014 (trace)<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"2\">nitrogen<\/td>\r\n<td>[latex]{}_{7}^{14}\\text{N}[\/latex]<\/td>\r\n<td>7<\/td>\r\n<td>7<\/td>\r\n<td>7<\/td>\r\n<td>14.0031<\/td>\r\n<td>99.63<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{7}^{15}\\text{N}[\/latex]<\/td>\r\n<td>7<\/td>\r\n<td>7<\/td>\r\n<td>8<\/td>\r\n<td>15.0001<\/td>\r\n<td>0.37<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"3\">oxygen<\/td>\r\n<td>[latex]{}_{8}^{16}\\text{O}[\/latex]<\/td>\r\n<td>8<\/td>\r\n<td>8<\/td>\r\n<td>8<\/td>\r\n<td>15.9949<\/td>\r\n<td>99.757<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{8}^{17}\\text{O}[\/latex]<\/td>\r\n<td>8<\/td>\r\n<td>8<\/td>\r\n<td>9<\/td>\r\n<td>16.9991<\/td>\r\n<td>0.038<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{8}^{18}\\text{O}[\/latex]<\/td>\r\n<td>8<\/td>\r\n<td>8<\/td>\r\n<td>10<\/td>\r\n<td>17.9992<\/td>\r\n<td>0.205<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>fluorine<\/td>\r\n<td>[latex]{}_{9}^{19}\\text{F}[\/latex]<\/td>\r\n<td>9<\/td>\r\n<td>9<\/td>\r\n<td>10<\/td>\r\n<td>18.9984<\/td>\r\n<td>100<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td rowspan=\"3\">neon<\/td>\r\n<td>[latex]{}_{10}^{20}\\text{Ne}[\/latex]<\/td>\r\n<td>10<\/td>\r\n<td>10<\/td>\r\n<td>10<\/td>\r\n<td>19.9924<\/td>\r\n<td>90.48<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{10}^{21}\\text{Ne}[\/latex]<\/td>\r\n<td>10<\/td>\r\n<td>10<\/td>\r\n<td>11<\/td>\r\n<td>20.9938<\/td>\r\n<td>0.27<\/td>\r\n<\/tr>\r\n<tr valign=\"middle\">\r\n<td>[latex]{}_{10}^{22}\\text{Ne}[\/latex]<\/td>\r\n<td>10<\/td>\r\n<td>10<\/td>\r\n<td>12<\/td>\r\n<td>21.9914<\/td>\r\n<td>9.25<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\n<div class=\"textbox\"><a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Use this Build an Atom simulator<\/a> to build atoms of the first 10 elements, see which isotopes exist, check nuclear stability, and gain experience with isotope symbols.<\/div>\r\n<h2>Atomic Mass<\/h2>\r\nBecause each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the <strong>atomic mass<\/strong> of a single atom is approximately equal to its mass number (a whole number). However, the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes.\r\n\r\nThe mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope\u2019s mass multiplied by its fractional abundance.\r\n\r\n[latex]\\text{average mass}=\\sum _{i}{\\left(\\text{fractional abundance}\\times \\text{isotopic mass}\\right)}_{i}[\/latex]\r\n\r\nFor example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are <sup>10<\/sup>B with a mass of 10.0129 amu, and the remaining 80.1% are <sup>11<\/sup>B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:\r\n\r\n[latex]\\begin{array}{cc}\\hfill \\text{boron average mass}&amp; =\\left(0.199\\times \\text{10.0129 amu}\\right)+\\left(0.801\\times \\text{11.0093 amu}\\right)\\hfill \\\\ &amp; =\\text{1.99 amu}+\\text{8.82 amu}\\hfill \\\\ &amp; =\\text{10.81 amu}\\hfill \\end{array}[\/latex]\r\n\r\nIt is important to understand that no single boron atom weighs exactly 10.8 amu; 10.8 amu is the average mass of all boron atoms, and individual boron atoms weigh either approximately 10 amu or 11 amu.\r\n<div class=\"textbox examples\">\r\n<h3>Example 2: <strong>Calculation of Average Atomic Mass<\/strong><\/h3>\r\nA meteorite found in central Indiana contains traces of the noble gas neon picked up from the solar wind during the meteorite\u2019s trip through the solar system. Analysis of a sample of the gas showed that it consisted of 91.84% <sup>20<\/sup>Ne (mass 19.9924 amu), 0.47% <sup>21<\/sup>Ne (mass 20.9940 amu), and 7.69% <sup>22<\/sup>Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?\r\n\r\n[reveal-answer q=\"877286\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"877286\"]\r\n\r\n[latex]\\begin{array}{cc}\\hfill \\text{average mass}&amp; =\\left(0.9184\\times \\text{19.9924 amu}\\right)+\\left(0.0047\\times \\text{20.9940 amu}\\right)+\\left(0.0769\\times \\text{21.9914 amu}\\right)\\hfill \\\\ &amp; =\\left(18.36+0.099+1.69\\right)\\text{amu}\\hfill \\\\ &amp; =\\text{20.15 amu}\\hfill \\end{array}[\/latex]\r\n\r\nThe average mass of a neon atom in the solar wind is 20.15 amu. (The average mass of a terrestrial neon atom is 20.1796 amu. This result demonstrates that we may find slight differences in the natural abundance of isotopes, depending on their origin.)\r\n\r\n[\/hidden-answer]\r\n<h4><strong>Check Your Learning<\/strong><\/h4>\r\nA sample of magnesium is found to contain 78.70% of <sup>24<\/sup>Mg atoms (mass 23.98 amu), 10.13% of <sup>25<\/sup>Mg atoms (mass 24.99 amu), and 11.17% of <sup>26<\/sup>Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom.\r\n\r\n[reveal-answer q=\"74864\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"74864\"]24.31 amu\r\n\r\nWe can also do variations of this type of calculation, as shown in the next example.\r\n\r\n[\/hidden-answer]\r\n\r\n<\/div>\r\n<div class=\"textbox\">\r\n\r\n<a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/isotopes-and-atomic-mass\" target=\"_blank\" rel=\"noopener\">Visit the PhET Isotopes and Atomic Mass site<\/a> to make mixtures of the main isotopes of the first 18 elements, gain experience with average atomic mass, and check naturally occurring isotope ratios using the Isotopes and Atomic Mass simulation.\r\n\r\n<\/div>\r\n<div class=\"textbox key-takeaways\">\r\n<h3>Key Concepts and Summary<\/h3>\r\nAlthough no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure. Thomson\u2019s cathode ray tube showed that atoms contain small, negatively charged particles called electrons. Millikan discovered that there is a fundamental electric charge\u2014the charge of an electron. Rutherford\u2019s gold foil experiment showed that atoms have a small, dense, positively charged nucleus; the positively charged particles within the nucleus are called protons. Chadwick discovered that the nucleus also contains neutral particles called neutrons. Soddy demonstrated that atoms of the same element can differ in mass; these are called isotopes.\r\n\r\nAn atom consists of a small, positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons; its diameter is about 100,000 times smaller than that of the atom. The mass of one atom is usually expressed in atomic mass units (amu), which is referred to as the atomic mass. An amu is defined as exactly [latex]\\frac{1}{12}[\/latex] of the mass of a carbon-12 atom and is equal to 1.6605 \u00d7 10<sup>\u221224<\/sup> g.\r\n\r\nProtons are relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu. Neutrons are relatively heavy particles with no charge and a mass of 1.0087 amu. Electrons are light particles with a charge of 1\u2212 and a mass of 0.00055 amu. The number of protons in the nucleus is called the atomic number (Z) and is the property that defines an atom\u2019s elemental identity. The sum of the numbers of protons and neutrons in the nucleus is called the mass number and, expressed in amu, is approximately equal to the mass of the atom. An atom is neutral when it contains equal numbers of electrons and protons.\r\n\r\nIsotopes of an element are atoms with the same atomic number but different mass numbers; isotopes of an element, therefore, differ from each other only in the number of neutrons within the nucleus. When a naturally occurring element is composed of several isotopes, the atomic mass of the element represents the average of the masses of the isotopes involved. A chemical symbol identifies the atoms in a substance using symbols, which are one-, two-, or three-letter abbreviations for the atoms.\r\n<h4>Key Equations<\/h4>\r\n<ul>\r\n \t<li>[latex]\\text{average mass}=\\sum _{i}{\\left(\\text{fractional abundance}\\times \\text{isotopic mass}\\right)}_{i}[\/latex]<\/li>\r\n<\/ul>\r\n<\/div>\r\n<div class=\"textbox exercises\">\r\n<h3>Exercises<\/h3>\r\n<ol>\r\n \t<li>The existence of isotopes violates one of the original ideas of Dalton\u2019s atomic theory. Which one?<\/li>\r\n \t<li>How are electrons and protons similar? How are they different?<\/li>\r\n \t<li>How are protons and neutrons similar? How are they different?<\/li>\r\n \t<li>Predict and test the behavior of \u03b1 particles fired at a \u201cplum pudding\u201d model atom.\r\n<ol style=\"list-style-type: lower-alpha;\">\r\n \t<li>Predict the paths taken by \u03b1 particles that are fired at atoms with a Thomson\u2019s plum pudding model structure. Explain why you expect the \u03b1 particles to take these paths.<\/li>\r\n \t<li>If \u03b1 particles of higher energy than those in (a) are fired at plum pudding atoms, predict how their paths will differ from the lower-energy \u03b1 particle paths. Explain your reasoning.<\/li>\r\n \t<li>Now test your predictions from (a) and (b). Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\" target=\"_blank\" rel=\"noopener noreferrer\">Rutherford Scattering simulation<\/a> and select the \u201cPlum Pudding Atom\u201d tab. Set \u201cAlpha Particles Energy\u201d to \u201cmin,\u201d and select \u201cshow traces.\u201d Click on the gun to start firing \u03b1 particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Hit the pause button, or \u201cReset All.\u201d Set \u201cAlpha Particles Energy\u201d to \u201cmax,\u201d and start firing \u03b1 particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual paths as shown in the simulation.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Predict and test the behavior of \u03b1 particles fired at a Rutherford atom model.\r\n<ol style=\"list-style-type: lower-alpha;\">\r\n \t<li>Predict the paths taken by \u03b1 particles that are fired at atoms with a Rutherford atom model structure. Explain why you expect the \u03b1 particles to take these paths.<\/li>\r\n \t<li>If \u03b1 particles of higher energy than those in (a) are fired at Rutherford atoms, predict how their paths will differ from the lower-energy \u03b1 particle paths. Explain your reasoning.<\/li>\r\n \t<li>Predict how the paths taken by the \u03b1 particles will differ if they are fired at Rutherford atoms of elements other than gold. What factor do you expect to cause this difference in paths, and why?<\/li>\r\n \t<li>Now test your predictions from (a), (b), and (c). Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\">Rutherford Scattering simulation<\/a> and select the \u201cRutherford Atom\u201d tab. Due to the scale of the simulation, it is best to start with a small nucleus, so select \u201c20\u201d for both protons and neutrons, \u201cmin\u201d for energy, show traces, and then start firing \u03b1 particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Pause or reset, set energy to \u201cmax,\u201d and start firing \u03b1 particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual path as shown in the simulation. Pause or reset, select \u201c40\u201d for both protons and neutrons, \u201cmin\u201d for energy, show traces, and fire away. Does this match your prediction from (c)? If not, explain why the actual path would be that shown in the simulation. Repeat this with larger numbers of protons and neutrons. What generalization can you make regarding the type of atom and effect on the path of \u03b1 particles? Be clear and specific.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>In what way are isotopes of a given element always different? In what way(s) are they always the same?<\/li>\r\n \t<li>Write the symbol for each of the following ions:\r\n<ol>\r\n \t<li>the ion with a 1+ charge, atomic number 55, and mass number 133<\/li>\r\n \t<li>the ion with 54 electrons, 53 protons, and 74 neutrons<\/li>\r\n \t<li>the ion with atomic number 15, mass number 31, and a 3\u2212 charge<\/li>\r\n \t<li>the ion with 24 electrons, 30 neutrons, and a 3+ charge<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Write the symbol for each of the following ions:\r\n<ol>\r\n \t<li>the ion with a 3+ charge, 28 electrons, and a mass number of 71<\/li>\r\n \t<li>the ion with 36 electrons, 35 protons, and 45 neutrons<\/li>\r\n \t<li>the ion with 86 electrons, 142 neutrons, and a 4+ charge<\/li>\r\n \t<li>the ion with a 2+ charge, atomic number 38, and mass number 87<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a> and click on the Atom icon.\r\n<ol>\r\n \t<li>Pick any one of the first 10 elements that you would like to build and state its symbol.<\/li>\r\n \t<li>Drag protons, neutrons, and electrons onto the atom template to make an atom of your element. State the numbers of protons, neutrons, and electrons in your atom, as well as the net charge and mass number.<\/li>\r\n \t<li>Click on \u201cNet Charge\u201d and \u201cMass Number,\u201d check your answers to (b), and correct, if needed.<\/li>\r\n \t<li>Predict whether your atom will be stable or unstable. State your reasoning.<\/li>\r\n \t<li>Check the \u201cStable\/Unstable\u201d box. Was your answer to (d) correct? If not, first predict what you can do to make a stable atom of your element, and then do it and see if it works. Explain your reasoning.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a>\r\n<ol>\r\n \t<li>Drag protons, neutrons, and electrons onto the atom template to make a neutral atom of Oxygen-16 and give the isotope symbol for this atom.<\/li>\r\n \t<li>Now add two more electrons to make an ion and give the symbol for the ion you have created.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a>\r\n<ol>\r\n \t<li>Drag protons, neutrons, and electrons onto the atom template to make a neutral atom of Lithium-6 and give the isotope symbol for this atom.<\/li>\r\n \t<li>Now remove one electron to make an ion and give the symbol for the ion you have created.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Determine the number of protons, neutrons, and electrons in the following isotopes that are used in medical diagnoses:\r\n<ol>\r\n \t<li>atomic number 9, mass number 18, charge of 1\u2212<\/li>\r\n \t<li>atomic number 43, mass number 99, charge of 7+<\/li>\r\n \t<li>atomic number 53, atomic mass number 131, charge of 1\u2212<\/li>\r\n \t<li>atomic number 81, atomic mass number 201, charge of 1+<\/li>\r\n \t<li>Name the elements in parts (a), (b), (c), and (d).<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>The following are properties of isotopes of two elements that are essential in our diet. Determine the number of protons, neutrons and electrons in each and name them.\r\n<ol>\r\n \t<li>atomic number 26, mass number 58, charge of 2+<\/li>\r\n \t<li>atomic number 53, mass number 127, charge of 1\u2212<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:\r\n<ol>\r\n \t<li>[latex]{}_{5}^{10}\\text{B}[\/latex]<\/li>\r\n \t<li>[latex]{}_{80}^{199}\\text{Hg}[\/latex]<\/li>\r\n \t<li>[latex]{}_{29}^{63}\\text{Cu}[\/latex]<\/li>\r\n \t<li>[latex]{}_{6}^{13}\\text{C}[\/latex]<\/li>\r\n \t<li>[latex]{}_{34}^{77}\\text{Se}[\/latex]<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:\r\n<ol>\r\n \t<li>[latex]{}_{3}^{7}\\text{Li}[\/latex]<\/li>\r\n \t<li>[latex]{}_{52}^{125}\\text{Te}[\/latex]<\/li>\r\n \t<li>[latex]{}_{47}^{109}\\text{Ag}[\/latex]<\/li>\r\n \t<li>[latex]{}_{7}^{15}\\text{N}[\/latex]<\/li>\r\n \t<li>[latex]{}_{15}^{31}\\text{P}[\/latex]<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Click on the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/isotopes-and-atomic-mass\" target=\"_blank\" rel=\"noopener\">Isotopes and Common Mass website<\/a> and select the \u201cMix Isotopes\u201d tab, hide the \u201cPercent Composition\u201d and \u201cAverage Atomic Mass\u201d boxes, and then select the element boron\r\n<ol>\r\n \t<li>Write the symbols of the isotopes of boron that are shown as naturally occurring in significant amounts.<\/li>\r\n \t<li>Predict the relative amounts (percentages) of these boron isotopes found in nature. Explain the reasoning behind your choice.<\/li>\r\n \t<li>Add isotopes to the black box to make a mixture that matches your prediction in (b). You may drag isotopes from their bins or click on \u201cMore\u201d and then move the sliders to the appropriate amounts.<\/li>\r\n \t<li>Reveal the \u201cPercent Composition\u201d and \u201cAverage Atomic Mass\u201d boxes. How well does your mixture match with your prediction? If necessary, adjust the isotope amounts to match your prediction.<\/li>\r\n \t<li>Select \u201cNature\u2019s\u201d mix of isotopes and compare it to your prediction. How well does your prediction compare with the naturally occurring mixture? Explain. If necessary, adjust your amounts to make them match \u201cNature\u2019s\u201d amounts as closely as possible.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Repeat Exercise 11 using an element that has three naturally occurring isotopes.<\/li>\r\n \t<li>An element has the following natural abundances and isotopic masses: 90.92% abundance with 19.99 amu, 0.26% abundance with 20.99 amu, and 8.82% abundance with 21.99 amu. Calculate the average atomic mass of this element.<\/li>\r\n \t<li>Average atomic masses listed by IUPAC are based on a study of experimental results. Bromine has two isotopes <sup>79<\/sup>Br and <sup>81<\/sup>Br, whose masses (78.9183 and 80.9163 amu) and abundances (50.69% and 49.31%) were determined in earlier experiments. Calculate the average atomic mass of bromine based on these experiments.<\/li>\r\n \t<li>Variations in average atomic mass may be observed for elements obtained from different sources. Lithium provides an example of this. The isotopic composition of lithium from naturally occurring minerals is 7.5% <sup>6<\/sup>Li and 92.5% <sup>7<\/sup>Li, which have masses of 6.01512 amu and 7.01600 amu, respectively. A commercial source of lithium, recycled from a military source, was 3.75% <sup>6<\/sup>Li (and the rest <sup>7<\/sup>Li). Calculate the average atomic mass values for each of these two sources.<\/li>\r\n \t<li>The average atomic masses of some elements may vary, depending upon the sources of their ores. Naturally occurring boron consists of two isotopes with accurately known masses (<sup>10<\/sup>B, 10.0129 amu and <sup>11<\/sup>B, 11.0931 amu). The actual atomic mass of boron can vary from 10.807 to 10.819, depending on whether the mineral source is from Turkey or the United States. Calculate the percent abundances leading to the two values of the average atomic masses of boron from these two countries.<\/li>\r\n \t<li>The <sup>18<\/sup>O:<sup>16<\/sup>O abundance ratio in some meteorites is greater than that used to calculate the average atomic mass of oxygen on earth. Is the average mass of an oxygen atom in these meteorites greater than, less than, or equal to that of a terrestrial oxygen atom?<\/li>\r\n<\/ol>\r\n[reveal-answer q=\"456814\"]Show Selected Answers[\/reveal-answer]\r\n[hidden-answer a=\"456814\"]\r\n\r\n1. Dalton originally thought that all atoms of a particular element had identical properties, including mass. Thus, the concept of isotopes, in which an element has different masses, was a violation of the original idea. To account for the existence of isotopes, the second postulate of his atomic theory was modified to state that atoms of the same element must have identical chemical properties.\r\n\r\n3. Both are subatomic particles that reside in an atom\u2019s nucleus. Both have approximately the same mass. Protons are positively charged, whereas neutrons are uncharged.\r\n\r\n5. The answers are as follows:\r\n<ol style=\"list-style-type: lower-alpha;\">\r\n \t<li>The Rutherford atom has a small, positively charged nucleus, so most \u03b1 particles will pass through empty space far from the nucleus and be undeflected. Those \u03b1 particles that pass near the nucleus will be deflected from their paths due to positive-positive repulsion. The more directly toward the nucleus the \u03b1 particles are headed, the larger the deflection angle will be.<\/li>\r\n \t<li>Higher-energy \u03b1 particles that pass near the nucleus will still undergo deflection, but the faster they travel, the less the expected angle of deflection.<\/li>\r\n \t<li>If the nucleus is smaller, the positive charge is smaller and the expected deflections are smaller\u2014both in terms of how closely the \u03b1 particles pass by the nucleus undeflected and the angle of deflection. If the nucleus is larger, the positive charge is larger and the expected deflections are larger\u2014more \u03b1 particles will be deflected, and the deflection angles will be larger.<\/li>\r\n \t<li>The paths followed by the \u03b1 particles match the predictions from (a), (b), and (c).<\/li>\r\n<\/ol>\r\n7. (a) <sup>133<\/sup>Cs<sup>+<\/sup>; (b) <sup>127<\/sup>I<sup>\u2212<\/sup>; (c) <sup>31<\/sup>P<sup>3\u2212<\/sup>; (d) <sup>57<\/sup>Co<sup>3+<\/sup>\r\n\r\n9. The answers are as follows:\r\n<ol>\r\n \t<li>Carbon-12, <sup>12<\/sup>C<\/li>\r\n \t<li>This atom contains six protons and six neutrons. There are six electrons in a neutral <sup>12<\/sup>C atom. The net charge of such a neutral atom is zero, and the mass number is 12.<\/li>\r\n \t<li>The preceding answers are correct.<\/li>\r\n \t<li>The atom will be stable since C-12 is a stable isotope of carbon.<\/li>\r\n \t<li>The preceding answer is correct. Other answers for this exercise are possible if a different element of isotope is chosen.<\/li>\r\n<\/ol>\r\n11. The answers are as follows:\r\n<ol>\r\n \t<li>Lithium-6 contains three protons, three neutrons, and three electrons. The isotope symbol is <sup>6<\/sup>Li or [latex]{}_{3}^{6}\\text{Li}[\/latex]<\/li>\r\n \t<li>[latex]{}_{}^{6}{\\text{Li}}^{+}[\/latex] or [latex]{}_{3}^{6}{\\text{Li}}^{+}[\/latex]<\/li>\r\n<\/ol>\r\n13. The answers are as follows:\r\n<ol>\r\n \t<li>iron, 26 protons, 24 electrons, and 32 neutrons<\/li>\r\n \t<li>iodine, 53 protons, 54 electrons, and 74 neutrons<\/li>\r\n<\/ol>\r\n15. The number of protons, electrons, and neutrons for each isotope are as follows:\r\n<ol>\r\n \t<li>3 protons, 3 electrons, 4 neutrons<\/li>\r\n \t<li>52 protons, 52 electrons, 73 neutrons<\/li>\r\n \t<li>47 protons, 47 electrons, 62 neutrons<\/li>\r\n \t<li>7 protons, 7 electrons, 8 neutrons<\/li>\r\n \t<li>15 protons, 15 electrons, 16 neutrons<\/li>\r\n<\/ol>\r\n17. Let us use neon as an example. Since there are three isotopes, there is no way to be sure to accurately predict the abundances to make the total of 20.18 amu average atomic mass. Let us guess that the abundances are 9% Ne-22, 91% Ne-20, and only a trace of Ne-21. The average mass would be 20.18 amu. Checking the nature\u2019s mix of isotopes shows that the abundances are 90.48% Ne-20, 9.25% Ne-22, and 0.27% Ne-21, so our guessed amounts have to be slightly adjusted.\r\n\r\n19. 79.904 amu\r\n\r\n21. Turkey source: 0.2649 (of 10.0129 amu isotope); US source: 0.2537 (of 10.0129 amu isotope)\r\n\r\n[\/hidden-answer]\r\n\r\n<\/div>\r\n<h2>Glossary<\/h2>\r\n<strong>alpha particle (\u03b1 particle): <\/strong>positively charged particle consisting of two protons and two neutrons\r\n\r\n<strong>anion: <\/strong>negatively charged atom or molecule (contains more electrons than protons)\r\n\r\n<strong>atomic mass: <\/strong>average mass of atoms of an element, expressed in amu\r\n\r\n<strong>atomic mass unit (amu): <\/strong>(also, unified atomic mass unit, u, or Dalton, Da) unit of mass equal to [latex]\\frac{1}{12}[\/latex] of the mass of a <sup>12<\/sup>C atom\r\n\r\n<strong>atomic number (Z): <\/strong>number of protons in the nucleus of an atom\r\n\r\n<strong>cation: <\/strong>positively charged atom or molecule (contains fewer electrons than protons)\r\n\r\n<strong>chemical symbol: <\/strong>one-, two-, or three-letter abbreviation used to represent an element or its atoms\r\n\r\n<strong>Dalton (Da): <\/strong>alternative unit equivalent to the atomic mass unit\r\n\r\n<strong>electron: <\/strong>negatively charged, subatomic particle of relatively low mass located outside the nucleus\r\n\r\n<strong>fundamental unit of charge: <\/strong>(also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 \u00d7 10<sup>\u221219<\/sup> C\r\n\r\n<strong>ion: <\/strong>electrically charged atom or molecule (contains unequal numbers of protons and electrons)\r\n\r\n<strong>isotopes: <\/strong>atoms that contain the same number of protons but different numbers of neutrons\r\n\r\n<strong>mass number (A): <\/strong>sum of the numbers of neutrons and protons in the nucleus of an atom\r\n\r\n<strong>neutron: <\/strong>uncharged, subatomic particle located in the nucleus\r\n\r\n<strong>nucleus: <\/strong>massive, positively charged center of an atom made up of protons and neutrons\r\n\r\n<strong>proton: <\/strong>positively charged, subatomic particle located in the nucleus\r\n\r\n<strong>unified atomic mass unit (u): <\/strong>alternative unit equivalent to the atomic mass unit\r\n\r\n&nbsp;","rendered":"<div class=\"textbox learning-objectives\">\n<h3>Learning Objectives<\/h3>\n<p>By the end of this section, you will be able to:<\/p>\n<ul>\n<li>Outline milestones in the development of modern atomic theory<\/li>\n<li>Summarize and interpret the results of the experiments of Thomson, Millikan, and Rutherford<\/li>\n<li>Describe the three subatomic particles that compose atoms<\/li>\n<li>Define isotopes and give examples for several elements<\/li>\n<li>Write isotopic symbols for elements and ions.<\/li>\n<\/ul>\n<\/div>\n<p id=\"ball-ch03_s01_p01\" class=\"para editable block\">The smallest piece of an element that maintains the identity of that element is called an <span class=\"margin_term\"><a class=\"glossterm\">atom<\/a><\/span>. Individual atoms are extremely small. It would take about fifty million atoms in a row to make a line that is 1 cm long. The period at the end of a printed sentence has several million atoms in it. Atoms are so small that it is difficult to believe that all matter is made from atoms\u2014but it is.<\/p>\n<p id=\"ball-ch03_s01_p02\" class=\"para editable block\">The concept that atoms play a fundamental role in chemistry is formalized by the <span class=\"margin_term\"><a class=\"glossterm\">modern atomic theory<\/a><\/span>, first stated by John Dalton, an English scientist, in 1808. It consists of three parts:<\/p>\n<ol id=\"ball-ch03_s01_l02\" class=\"orderedlist editable block\">\n<li>All matter is composed of atoms.<\/li>\n<li>Atoms of the same element are the same; atoms of different elements are different.<\/li>\n<li>Atoms combine in whole-number ratios to form compounds.<\/li>\n<\/ol>\n<p id=\"ball-ch03_s01_p03\" class=\"para editable block\">These concepts form the basis of chemistry.<\/p>\n<div id=\"ball-ch03_s01_f01\" class=\"figure large editable block\">\n<p>\u00a0In the two centuries since Dalton developed his ideas, scientists have made significant progress in furthering our understanding of atomic theory. Much of this came from the results of several seminal experiments that revealed the details of the internal structure of atoms. Here, we will discuss some of those key developments, with an emphasis on application of the scientific method, as well as understanding how the experimental evidence was analyzed. While the historical persons and dates behind these experiments can be quite interesting, it is most important to understand the concepts resulting from their work.<\/p>\n<\/div>\n<h2>Evolution of the Atom<\/h2>\n<p>If matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? In the late 1800s, a number of scientists interested in questions like these investigated the electrical discharges that could be produced in low-pressure gases, with the most significant discovery made by English physicist J. J. <strong>Thomson<\/strong> using a <strong>cathode ray<\/strong> tube. This apparatus consisted of a sealed glass tube from which almost all the air had been removed; the tube contained two metal electrodes. When high voltage was applied across the electrodes, a visible beam called a cathode ray appeared between them. This beam was deflected toward the positive charge and away from the negative charge, and was produced in the same way with identical properties when different metals were used for the electrodes. In similar experiments, the ray was simultaneously deflected by an applied magnetic field, and measurements of the extent of deflection and the magnetic field strength allowed Thomson to calculate the charge-to-mass ratio of the cathode ray particles. The results of these measurements indicated that these particles were much lighter than atoms (Figure 1).<\/p>\n<div style=\"width: 890px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211009\/CNX_Chem_02_02_CathodeRay1.jpg\" alt=\"Figure A shows a photo of J. J. Thomson working at a desk. Figure B shows a photograph of a cathode ray tube. It is a long, glass tube that is narrow at the left end but expands into a large bulb on the right end. The entire cathode tube is sitting on a wooden stand. Figure C shows the parts of the cathode ray tube. The cathode ray tube consists of a cathode and an anode. The cathode, which has a negative charge, is located in a small bulb of glass on the left side of the cathode ray tube. To the left of the cathode it says \u201cHigh voltage\u201d and indicates a positive and negative charge. The anode, which has a positive charge, is located to the right of the cathode. Two charged plates are located to the right of the anode, and are connected to a battery and two magnets. The magnets are labeled \u201cS\u201d and \u201cN.\u201d A cathode ray is generated from the cathode, travels through the anode and into a wider part of the cathode ray tube, where it travels between a positively charged electrode plate and a negatively charged electrode plate. The ray bends upward and continues to travel until it hits the wide part of the tube on the right. The rightmost end of the tube contains a printed scale that allows one to measure how much the ray was deflected.\" width=\"880\" height=\"616\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 1. (a) J. J. Thomson produced a visible beam in a cathode ray tube. (b) This is an early cathode ray tube, invented in 1897 by Ferdinand Braun. (c) In the cathode ray, the beam (shown in yellow) comes from the cathode and is accelerated past the anode toward a fluorescent scale at the end of the tube. Simultaneous deflections by applied electric and magnetic fields permitted Thomson to calculate the mass-to-charge ratio of the particles composing the cathode ray. (credit a: modification of work by Nobel Foundation; credit b: modification of work by Eugen Nesper; credit c: modification of work by \u201cKurzon\u201d\/Wikimedia Commons)<\/p>\n<\/div>\n<p>Based on his observations, here is what Thomson proposed and why: The particles are attracted by positive (+) charges and repelled by negative (-) charges, so they must be negatively charged (like charges repel and unlike charges attract); they are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms. Although controversial at the time, Thomson\u2019s idea was gradually accepted, and his cathode ray particle is what we now call an <strong>electron<\/strong>, a negatively charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term \u201celectron\u201d was coined in 1891 by Irish physicist George Stoney, from \u201c<em>electr<\/em>ic i<em>on<\/em>.\u201d<\/p>\n<div class=\"textbox\">Click <a href=\"https:\/\/www.aip.org\/history\/electron\/jjsound.htm\" target=\"_blank\" rel=\"noopener noreferrer\">this link to &#8220;JJ Thompson Talks About the Size of the Electron&#8221;<\/a> to hear Thomson describe his discovery in his own voice.<\/div>\n<p>Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson and others, the charge and mass of the negative, subatomic particles\u2014the electrons\u2014were known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson proposed the \u201cplum pudding\u201d model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it, since all atoms are electrically neutral. A competing model had been proposed in 1903 by Hantaro <strong>Nagaoka<\/strong>, who postulated a Saturn-like atom, consisting of a positively charged sphere surrounded by a halo of electrons (Figure 2).<\/p>\n<div style=\"width: 888px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211013\/CNX_Chem_02_02_AtomModels1.jpg\" alt=\"Figure A shows a photograph of plum pudding, which is a thick, almost spherical cake containing raisins throughout. To the right, an atom model is round and contains negatively charged electrons embedded within a sphere of positively charged matter. Figure B shows a photograph of the planet Saturn, which has rings. To the right, an atom model is a sphere of positively charged matter encircled by a ring of negatively charged electrons.\" width=\"878\" height=\"266\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 2. (a) Thomson suggested that atoms resembled plum pudding, an English dessert consisting of moist cake with embedded raisins (\u201cplums\u201d). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of electrons surrounding a positive \u201cplanet.\u201d (credit a: modification of work by \u201cMan vyi\u201d\/Wikimedia Commons; credit b: modification of work by \u201cNASA\u201d\/Wikimedia Commons)<\/p>\n<\/div>\n<p>The next major development in understanding the atom came from Ernest <strong>Rutherford<\/strong>, a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged <strong>alpha particles (\u03b1 particles)<\/strong> that were produced by the radioactive decay of radium; \u03b1 particles consist of two protons and two neutrons (you will learn more about radioactive decay in the chapter on nuclear chemistry). Rutherford and his colleagues Hans <strong>Geiger<\/strong> (later famous for the Geiger counter) and Ernest <strong>Marsden<\/strong> aimed a beam of \u03b1 particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the \u03b1 particles using a luminescent screen that glowed briefly where hit by an \u03b1 particle.<\/p>\n<p>What did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source (Figure 3). Rutherford described finding these results: \u201cIt was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you.\u201d<\/p>\n<div style=\"width: 890px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211015\/CNX_Chem_02_02_Rutherford1.jpg\" alt=\"This figure shows a box on the left that contains a radium source of alpha particles which generates a beam of alpha particles. The beam travels through an opening within a ring-shaped luminescent screen which is used to detect scattered alpha particles. A piece of thin gold foil is at the center of the ring formed by the screen. When the beam encounters the gold foil, most of the alpha particles pass straight through it and hit the luminescent screen directly behind the foil. Some of the alpha particles are slightly deflected by the foil and hit the luminescent screen off to the side of the foil. Some alpha particles are significantly deflected and bounce back to hit the front of the screen.\" width=\"880\" height=\"386\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 3. Geiger and Rutherford fired \u03b1 particles at a piece of gold foil and detected where those particles went, as shown in this schematic diagram of their experiment. Most of the particles passed straight through the foil, but a few were deflected slightly and a very small number were significantly deflected.<\/p>\n<\/div>\n<p>Here is what Rutherford deduced: Because most of the fast-moving \u03b1 particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom. Alpha particles are positively charged, so deflections arose when they encountered another positive charge (like charges repel each other). Since like charges repel one another, the few positively charged \u03b1 particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge. Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:<\/p>\n<ol>\n<li>The volume occupied by an atom must consist of a large amount of empty space.<\/li>\n<li>A small, relatively heavy, positively charged body, the <strong>nucleus<\/strong>, must be at the center of each atom.<\/li>\n<\/ol>\n<div class=\"textbox\">View <a href=\"https:\/\/micro.magnet.fsu.edu\/electromag\/java\/rutherford\/\" target=\"_blank\" rel=\"noopener noreferrer\">this simulation of the Rutherford gold foil experiment<\/a>. Adjust the slit width to produce a narrower or broader beam of \u03b1 particles to see how that affects the scattering pattern.<\/div>\n<p>This analysis led Rutherford to propose a model in which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the negatively charged electrons, so that the atom is electrically neutral (Figure 4). After many more experiments, Rutherford also discovered that the nuclei of other elements contain the hydrogen nucleus as a \u201cbuilding block,\u201d and he named this more fundamental particle the <strong>proton<\/strong>, the positively charged, subatomic particle found in the nucleus. With one addition, which you will learn next, this nuclear model of the atom, proposed over a century ago, is still used today.<\/p>\n<div style=\"width: 890px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211016\/CNX_Chem_02_02_GoldFoil31.jpg\" alt=\"The left diagram shows a green beam of alpha particles hitting a rectangular piece of gold foil. Some of the alpha particles bounce backwards after hitting the foil. However, most of the particles travel through the foil, with some being deflected as they pass through the foil. A callout box shows a magnified cross section of the gold foil. Most of the alpha particles are not deflected, but pass straight through the foil because they travel between the gold atoms. A very small number of alpha particles are significantly deflected when they hit the nucleus of the gold atoms straight on. A few alpha particles are slightly deflected because they glanced off of the nucleus of a gold atom.\" width=\"880\" height=\"515\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 4. The \u03b1 particles are deflected only when they collide with or pass close to the much heavier, positively charged gold nucleus. Because the nucleus is very small compared to the size of an atom, very few \u03b1 particles are deflected. Most pass through the relatively large region occupied by electrons, which are too light to deflect the rapidly moving particles.<\/p>\n<\/div>\n<div class=\"textbox\">The <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\" target=\"_blank\" rel=\"noopener noreferrer\">Rutherford Scattering simulation<\/a> allows you to investigate the differences between a \u201cplum pudding\u201d atom and a Rutherford atom by firing \u03b1 particles at each type of atom.<\/div>\n<p>Another important finding was the discovery of isotopes. During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a \u201cnew element\u201d produced by the radioactive decay of thorium was initially given the name mesothorium. However, a more detailed analysis showed that mesothorium was chemically identical to radium (another decay product), despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick <strong>Soddy<\/strong> to realize that an element could have types of atoms with different masses that were chemically indistinguishable. These different types are called <strong>isotopes<\/strong>\u2014atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.<\/p>\n<p>One puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus. As you might expect, detecting uncharged particles is very challenging, and it was not until 1932 that James <strong>Chadwick<\/strong> found evidence of <strong>neutrons<\/strong>, uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes: They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons. With the discovery of the neutron, the modern model of an atom was established (Figure 5).<\/p>\n<div id=\"attachment_4627\" style=\"width: 595px\" class=\"wp-caption aligncenter\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/The-Atom.png\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-4627\" class=\"wp-image-4627\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/2835\/2017\/12\/14212601\/The-Atom-1.png\" alt=\"The Atom\" width=\"585\" height=\"336\" \/><\/a><\/p>\n<p id=\"caption-attachment-4627\" class=\"wp-caption-text\">Figure 5. The structure of an atom. The nucleus of an atom is composed of protons and neutrons. The nucleus is surrounded by an electron cloud.<\/p>\n<\/div>\n<p>The diameter of an atom is on the order of 10<sup>\u221210<\/sup> m, whereas the diameter of the nucleus is roughly 10<sup>\u221215<\/sup> m\u2014about 100,000 times smaller. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium (Figure 6).<\/p>\n<div style=\"width: 1310px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211018\/CNX_Chem_02_03_AtomSize1.jpg\" alt=\"The diagram on the left shows a picture of an atom that is 10 to the negative tenth power meters in diameter. The nucleus is labeled at the center of the atom and is 10 to the negative fifteenth power meters. The central figure shows a photograph of an American football stadium. The figure on the right shows a photograph of a person with a handful of blueberries.\" width=\"1300\" height=\"400\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 6. If an atom could be expanded to the size of a football stadium, the nucleus would be the size of a single blueberry. (credit middle: modification of work by \u201cbabyknight\u201d\/Wikimedia Commons; credit right: modification of work by Paxson Woelber)<\/p>\n<\/div>\n<p>Atoms\u2014and the protons, neutrons, and electrons that compose them\u2014are extremely small. For example, a carbon atom weighs less than 2 \u00d7 10<sup>\u221223<\/sup> g, and an electron has a charge of less than 2 \u00d7 10<sup>\u221219<\/sup> C (coulomb). When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the<strong> atomic mass unit (amu)<\/strong> and the <strong>fundamental unit of charge (e)<\/strong>. The amu was originally defined based on hydrogen, the lightest element, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. (This isotope is known as \u201ccarbon-12\u201d as will be discussed later in this module.) Thus, one amu is exactly [latex]\\frac{1}{12}[\/latex] of the mass of one carbon-12 atom: 1 amu = 1.6605 \u00d7 10<sup>\u221224<\/sup> g. (The <strong>Dalton (Da)<\/strong> and the <strong>unified atomic mass unit (u) <\/strong>are alternative units that are equivalent to the amu.) The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 \u00d7 10<sup>\u221219<\/sup> C.<\/p>\n<p>A proton has a mass of 1.0073 amu and a charge of 1+. A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero; as its name suggests, it is neutral. The electron has a charge of 1\u2212 and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton. The properties of these fundamental particles are summarized in Table 1. (An observant student might notice that the sum of an atom\u2019s subatomic particles does not equal the atom\u2019s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than 12.00 amu. This \u201cmissing\u201d mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)<\/p>\n<table id=\"fs-idp90857696\" class=\"span-all\" summary=\"This table gives the name, location, charge in C, unit charge, mass in A M U and mass in grams for electrons, protons and neutrons. Electrons are located outside of the nucleus, have a charge of negative 1.602 times 10 to the negative nineteenth power, a unit charge of negative 1, and a mass of 0.00055 A M U or 0.00091 times 10 to the negative twenty-fourth power grams. Protons are located within the nucleus, have a charge of 1.602 times 10 to the negative nineteenth power, have a unit charge of positive 1, and have a mass of 1.0073 A M U or 1.6726 times 10 to the negative twenty-fourth power grams. Neutrons are located within the nucleus, have a charge of 0, have a unit charge of 0, and have a mass of 1.0087 A M U or 1.6749 times 10 to the negative twenty-fourth power grams.\">\n<thead>\n<tr>\n<th colspan=\"6\">Table 1. Properties of Subatomic Particles<\/th>\n<\/tr>\n<tr valign=\"top\">\n<th>Name<\/th>\n<th>Location<\/th>\n<th>Charge (C)<\/th>\n<th>Unit Charge<\/th>\n<th>Mass (amu)<\/th>\n<th>Mass (g)<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr valign=\"top\">\n<td>electron<\/td>\n<td>outside nucleus<\/td>\n<td>\u22121.602 \u00d7 10<sup>\u221219<\/sup><\/td>\n<td>1\u2212<\/td>\n<td>0.00055<\/td>\n<td>0.00091 \u00d7 10<sup>\u221224<\/sup><\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>proton<\/td>\n<td>nucleus<\/td>\n<td>1.602 \u00d7 10<sup>\u221219<\/sup><\/td>\n<td>1+<\/td>\n<td>1.00727<\/td>\n<td>1.67262 \u00d7 10<sup>\u221224<\/sup><\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>neutron<\/td>\n<td>nucleus<\/td>\n<td>0<\/td>\n<td>0<\/td>\n<td>1.00866<\/td>\n<td>1.67493 \u00d7 10<sup>\u221224<\/sup><\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<p>The number of protons in the nucleus of an atom is its <strong>atomic number (Z)<\/strong>. This is the defining trait of an element: Its value determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its <strong>mass number (A)<\/strong>. The number of neutrons is therefore the difference between the mass number and the atomic number: A \u2013 Z = number of neutrons.<\/p>\n<p>[latex]\\begin{array}{ccc}\\hfill \\text{atomic number}\\left(\\text{Z}\\right)& =& \\text{number of protons}\\hfill \\\\ \\hfill \\text{atomic mass}\\left(\\text{A}\\right)& =& \\text{number of protons}+\\text{number of neutrons}\\hfill \\\\ \\hfill \\text{A}-\\text{Z}& =& \\text{number of neutrons}\\hfill \\end{array}[\/latex]<\/p>\n<p>Atoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. When the numbers of these subatomic particles are <em>not<\/em> equal, the atom is electrically charged and is called an <strong>ion<\/strong>. The charge of an atom is defined as follows:<\/p>\n<p>Atomic charge = number of protons \u2212 number of electrons<\/p>\n<p>As will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an <strong>anion<\/strong>. Positively charged atoms called <strong>cations<\/strong> are formed when an atom loses one or more electrons. For example, a neutral sodium atom (Z = 11) has 11 electrons. If this atom loses one electron, it will become a cation with a 1+ charge (11 \u2212 10 = 1+). A neutral oxygen atom (Z = 8) has eight electrons, and if it gains two electrons it will become an anion with a 2\u2212 charge (8 \u2212 10 = 2\u2212).<\/p>\n<div class=\"textbox examples\">\n<h3>Example 1: <strong>Composition of an Atom<\/strong><\/h3>\n<p>Iodine is an essential trace element in our diet; it is needed to produce thyroid hormone. Insufficient iodine in the diet can lead to the development of a goiter, an enlargement of the thyroid gland (Figure 2).<\/p>\n<div style=\"width: 709px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211019\/CNX_Chem_02_03_Iodine1.jpg\" alt=\"Figure A shows a photo of a person who has a very swollen thyroid in his or her neck. Figure B shows a photo of a canister of iodized salt.\" width=\"699\" height=\"310\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 2. (a) Insufficient iodine in the diet can cause an enlargement of the thyroid gland called a goiter. (b) The addition of small amounts of iodine to salt, which prevents the formation of goiters, has helped eliminate this concern in the US where salt consumption is high. (credit a: modification of work by \u201cAlmazi\u201d\/Wikimedia Commons; credit b: modification of work by Mike Mozart)<\/p>\n<\/div>\n<p>&nbsp;<\/p>\n<p>The addition of small amounts of iodine to table salt (iodized salt) has essentially eliminated this health concern in the United States, but as much as 40% of the world\u2019s population is still at risk of iodine deficiency. The iodine atoms are added as anions, and each has a 1\u2212 charge and a mass number of 127. Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q754651\">Show Answer<\/span><\/p>\n<div id=\"q754651\" class=\"hidden-answer\" style=\"display: none\">\n<p>The atomic number of iodine (53) tells us that a neutral iodine atom contains 53 protons in its nucleus and 53 electrons outside its nucleus. Because the sum of the numbers of protons and neutrons equals the mass number, 127, the number of neutrons is 74 (127 \u2212 53 = 74). Since the iodine is added as a 1\u2212 anion, the number of electrons is 54 [53 \u2212 (\u22121) = 54].<\/p>\n<\/div>\n<\/div>\n<h4><strong>Check Your Learning<\/strong><\/h4>\n<p>An atom of platinum has a mass number of 195 and contains 74 electrons. How many protons and neutrons does it contain, and what is its charge?<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q488361\">Show Answer<\/span><\/p>\n<div id=\"q488361\" class=\"hidden-answer\" style=\"display: none\">78 protons; 117 neutrons; charge is 4+<\/div>\n<\/div>\n<\/div>\n<h2>Chemical Symbols<\/h2>\n<div style=\"width: 360px\" class=\"wp-caption alignright\"><img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211020\/CNX_Chem_02_03_SiSymbol1.jpg\" alt=\"A jar labeled \u201cH g\u201d is shown with a small amount of liquid mercury in it.\" width=\"350\" height=\"280\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 7. The symbol Hg represents the element mercury regardless of the amount; it could represent one atom of mercury or a large amount of mercury.<\/p>\n<\/div>\n<p>A <strong>chemical symbol<\/strong> is an abbreviation that we use to indicate an element or an atom of an element. For example, the symbol for mercury is Hg (Figure 7). We use the same symbol to indicate one atom of mercury (microscopic domain) or to label a container of many atoms of the element mercury (macroscopic domain).<\/p>\n<p>The symbols for several common elements and their atoms are listed in Table 2. Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Most symbols have one or two letters, but three-letter symbols have been used to describe some elements that have atomic numbers greater than 112.<\/p>\n<p>To avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen (O). All known elements and their symbols are in <a title=\"The Periodic Table\" href=\".\/chapter\/the-periodic-table-3\/\" target=\"_blank\" rel=\"noopener\">the periodic table<\/a>.<\/p>\n<table id=\"fs-idm36686800\" class=\"span-all\" summary=\"This table has two columns labeled element and symbol. The first letter of the symbol is always an uppercase letter while the second letter of the symbol is always a lowercase letter. Aluminum has the symbol A L. Bromine has the symbol B R, calcium has the symbol C A, carbon has the symbol C, chlorine has the symbol C L, chromium has the symbol C R, cobalt has the symbol C O, copper has the symbol C U, from cuprum, fluorine has the symbol F, gold has the symbol A U, from aurum, helium has the symbol H E, hydrogen has the symbol H, iodine has the symbol I, iron has the symbol F E, from ferrum, lead has the symbol P B, from plumbum, magnesium has the symbol M G, mercury has the symbol H G from hydrargyrum, nitrogen has the symbol N, oxygen has the symbol O, potassium has the symbol K, from kalium, silicon has the symbol S I, silver has the symbol A G, from argentum, sodium has the symbol N A from natrium, sulfur has the symbol S, tin has the symbol S N from stannum, and zinc has the symbol Z N.\">\n<thead>\n<tr>\n<th colspan=\"4\">Table 2. Some Common Elements and Their Symbols<\/th>\n<\/tr>\n<tr valign=\"top\">\n<th>Element<\/th>\n<th>Symbol<\/th>\n<th>Element<\/th>\n<th>Symbol<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr valign=\"top\">\n<td>aluminum<\/td>\n<td>Al<\/td>\n<td>iron<\/td>\n<td>Fe (from <em>ferrum<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>bromine<\/td>\n<td>Br<\/td>\n<td>lead<\/td>\n<td>Pb (from <em>plumbum<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>calcium<\/td>\n<td>Ca<\/td>\n<td>magnesium<\/td>\n<td>Mg<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>carbon<\/td>\n<td>C<\/td>\n<td>mercury<\/td>\n<td>Hg (from <em>hydrargyrum<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>chlorine<\/td>\n<td>Cl<\/td>\n<td>nitrogen<\/td>\n<td>N<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>chromium<\/td>\n<td>Cr<\/td>\n<td>oxygen<\/td>\n<td>O<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>cobalt<\/td>\n<td>Co<\/td>\n<td>potassium<\/td>\n<td>K (from <em>kalium<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>copper<\/td>\n<td>Cu (from <em>cuprum<\/em>)<\/td>\n<td>silicon<\/td>\n<td>Si<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>fluorine<\/td>\n<td>F<\/td>\n<td>silver<\/td>\n<td>Ag (from<em> argentum<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>gold<\/td>\n<td>Au (from <em>aurum<\/em>)<\/td>\n<td>sodium<\/td>\n<td>Na (from <em>natrium<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>helium<\/td>\n<td>He<\/td>\n<td>sulfur<\/td>\n<td>S<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>hydrogen<\/td>\n<td>H<\/td>\n<td>tin<\/td>\n<td>Sn (from <em>stannum<\/em>)<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>iodine<\/td>\n<td>I<\/td>\n<td>zinc<\/td>\n<td>Zn<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<p>Traditionally, the discoverer (or discoverers) of a new element names the element. However, until the name is recognized by the International Union of Pure and Applied Chemistry (IUPAC), the recommended name of the new element is based on the Latin word(s) for its atomic number. For example, element 106 was called unnilhexium (Unh), element 107 was called unnilseptium (Uns), and element 108 was called unniloctium (Uno) for several years. These elements are now named after scientists (or occasionally locations); for example, element 106 is now known as <em>seaborgium<\/em> (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements.<\/p>\n<div class=\"textbox\">Visit <a href=\"https:\/\/iupac.org\/\" target=\"_blank\" rel=\"noopener\">this site to learn more about IUPAC, the International Union of Pure and Applied Chemistry<\/a>, and explore its periodic table.<\/div>\n<h2>Isotopes<\/h2>\n<p>The symbol for a specific isotope of any element is written by placing the mass number as a superscript to the left of the element symbol (Figure 8). The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element\u2019s identity, as does its symbol, it is often omitted. For example, magnesium exists as a mixture of three isotopes, each with an atomic number of 12 and with mass numbers of 24, 25, and 26, respectively. These isotopes can be identified as <sup>24<\/sup>Mg, <sup>25<\/sup>Mg, and <sup>26<\/sup>Mg. These isotope symbols are read as \u201celement, mass number\u201d and can be symbolized consistent with this reading. For instance, <sup>24<\/sup>Mg is read as \u201cmagnesium 24,\u201d and can be written as \u201cmagnesium-24\u201d or \u201cMg-24.\u201d <sup>25<\/sup>Mg is read as \u201cmagnesium 25,\u201d and can be written as \u201cmagnesium-25\u201d or \u201cMg-25.\u201d All magnesium atoms have 12 protons in their nucleus. They differ only because a <sup>24<\/sup>Mg atom has 12 neutrons in its nucleus, a <sup>25<\/sup>Mg atom has 13 neutrons, and a <sup>26<\/sup>Mg has 14 neutrons.<\/p>\n<div style=\"width: 660px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211023\/CNX_Chem_02_03_AtomSym1.jpg\" alt=\"This diagram shows the symbol for helium, \u201cH e.\u201d The number to the upper left of the symbol is the mass number, which is 4. The number to the upper right of the symbol is the charge which is positive 2. The number to the lower left of the symbol is the atomic number, which is 2. This number is often omitted. Also shown is \u201cM g\u201d which stands for magnesium It has a mass number of 24, a charge of positive 2, and an atomic number of 12.\" width=\"650\" height=\"144\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 8. The symbol for an atom indicates the element via its usual two-letter symbol, the mass number as a left superscript, the atomic number as a left subscript (sometimes omitted), and the charge as a right superscript.<\/p>\n<\/div>\n<p>Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table 3. Note that in addition to standard names and symbols, the isotopes of hydrogen are often referred to using common names and accompanying symbols. Hydrogen-2, symbolized <sup>2<\/sup>H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized <sup>3<\/sup>H, is also called tritium and sometimes symbolized T.<\/p>\n<table id=\"fs-idm87646592\" class=\"span-all\" summary=\"This table has seven columns labeled element, symbol, atomic number, number of protons, number of neutrons, mass in A M U, and percent natural abundance. The symbols for each element each show the mass number in the upper left and the atomic number in the lower left. Therefore hydrogen left superscript 1, left subscript 1, or protium, has a mass number of 1 and an atomic number of 1. Protium has one proton, 0 neutrons, a mass of 1.0078 and a natural abundance percentage of 99.985. Hydrogen left superscript 2, left subscript 1, or deuterium, has an atomic number of 1, 1 proton, 1 neutron, a mass of 2.0141 and a natural abundance percentage of 0.015. Hydrogen left superscript 3, left subscript 1, or tritium, has an atomic number of 11 protons, 2 neutrons, and a mass of 3.01605. No natural abundance percentage is given. Helium left superscript 3, left subscript 2 has an atomic number of 2, 2 protons, 1 neutron, a mass of 3.01603, and a natural abundance percentage of 0.00013. Helium left superscript 4, left subscript 2 has an atomic number of 2, 2 protons, 2 neutrons, a mass of 4.0026 and a natural abundance percentage of 100. Lithium left superscript 6, left subscript 3 has an atomic number of 3, 3 protons, 3 neutrons, a mass of 6.0151, and a natural abundance percentage of 7.42. Lithium left superscript 7, left subscript 3 has an atomic number of 3, 3 protons, 4 neutrons, a mass of 7.0160, and a natural abundance percentage of 92.8. Beryllium left superscript 9, left subscript 4 has an atomic number of 4, 4 protons, 5 neutrons, a mass of 9.0122, and a natural abundance percentage of 100. Boron left superscript 10, left subscript 5 has an atomic number of 5, 5 protons, 5 neutrons and a natural abundance percentage of 19.9. Boron left superscript 11, left subscript 5 has an atomic number of 5, 5 protons, 6 neutrons, a mass of 11.0093 and a natural abundance of 80.1. Carbon left superscript 12, left subscript 6 has an atomic number of 6, 6 protons, 6 neutrons, a mass of 12, and a natural abundance percentage of 98.89. Carbon left superscript 13, left subscript 6 has an atomic number of 6, 6 protons, 7 neutrons, a mass of 13.0033, and a natural abundance percentage of 1.11. Carbon left superscript 14, left subscript 6 has an atomic number of 6, 6 protons, 8 neutrons, and a mass of 14.0032. Its natural abundance percentage is not reported. Nitrogen left superscript 14, left subscript 7 has an atomic number of 7, 7 protons, 7 neutrons, a mass of 14.0031, and a natural abundance percentage of 99.63. Nitrogen left superscript 15, left subscript 7 has an atomic number of 7, 7 protons, 8 neutrons, a mass of 15.0001, and a natural abundance percentage of 0.37. Oxygen left superscript 16, left subscript 8 has an atomic number of 8, 8 protons, 8 neutrons, a mass of 15.9949, and a natural abundance percentage of 99.759. Oxygen left superscript 17, left subscript 8 has an atomic number of 8, 8 protons, 9 neutrons, a mass of 16.9991, and a natural abundance percentage of 0.037. Oxygen left superscript 18, left subscript 8 has an atomic number of 8, 8 protons, 10 neutrons, a mass of 17.9992, and a natural abundance percentage of 0.204. Fluorine left superscript 19, left subscript 9 has an atomic number of 9, 9 protons, 10 neutrons, a mass of 18.9984, and a natural abundance percentage of 100. Neon left superscript 20, left subscript 10 has an atomic number of 10, 10 protons, 10 neutrons, a mass of 19.9924, and a natural abundance percentage of 90.92. Neon left superscript 21, left subscript 10 has an atomic number of 10, 10 protons, 11 neutrons, a mass of 20.994, and a natural abundance percentage of 0.257. Neon left superscript 22, left subscript 10 has an atomic number of 10, 10 protons, 12 neutrons, a mass of 21.9914, and a natural abundance percentage of 8.82.\">\n<thead>\n<tr>\n<th colspan=\"7\">Table 3. Nuclear Compositions of Atoms of the Very Light Elements<\/th>\n<\/tr>\n<tr valign=\"top\">\n<th>Element<\/th>\n<th>Symbol<\/th>\n<th>Atomic Number<\/th>\n<th>Number of Protons<\/th>\n<th>Number of Neutrons<\/th>\n<th>Mass (amu)<\/th>\n<th>% Natural Abundance<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr valign=\"middle\">\n<td rowspan=\"3\">hydrogen<\/td>\n<td>[latex]{}_{1}^{1}\\text{H}[\/latex]<br \/>\n(protium)<\/td>\n<td>1<\/td>\n<td>1<\/td>\n<td>0<\/td>\n<td>1.0078<\/td>\n<td>99.989<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{1}^{2}\\text{H}[\/latex]<br \/>\n(deuterium)<\/td>\n<td>1<\/td>\n<td>1<\/td>\n<td>1<\/td>\n<td>2.0141<\/td>\n<td>0.0115<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{1}^{3}\\text{H}[\/latex]<br \/>\n(tritium)<\/td>\n<td>1<\/td>\n<td>1<\/td>\n<td>2<\/td>\n<td>3.01605<\/td>\n<td>\u2014 (trace)<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"2\">helium<\/td>\n<td>[latex]{}_{2}^{3}\\text{He}[\/latex]<\/td>\n<td>2<\/td>\n<td>2<\/td>\n<td>1<\/td>\n<td>3.01603<\/td>\n<td>0.00013<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{2}^{4}\\text{He}[\/latex]<\/td>\n<td>2<\/td>\n<td>2<\/td>\n<td>2<\/td>\n<td>4.0026<\/td>\n<td>100<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"2\">lithium<\/td>\n<td>[latex]{}_{3}^{6}\\text{Li}[\/latex]<\/td>\n<td>3<\/td>\n<td>3<\/td>\n<td>3<\/td>\n<td>6.0151<\/td>\n<td>7.59<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{3}^{7}\\text{Li}[\/latex]<\/td>\n<td>3<\/td>\n<td>3<\/td>\n<td>4<\/td>\n<td>7.0160<\/td>\n<td>92.41<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>beryllium<\/td>\n<td>[latex]{}_{4}^{9}\\text{Be}[\/latex]<\/td>\n<td>4<\/td>\n<td>4<\/td>\n<td>5<\/td>\n<td>9.0122<\/td>\n<td>100<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"2\">boron<\/td>\n<td>[latex]{}_{5}^{10}\\text{B}[\/latex]<\/td>\n<td>5<\/td>\n<td>5<\/td>\n<td>5<\/td>\n<td>10.0129<\/td>\n<td>19.9<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{5}^{11}\\text{B}[\/latex]<\/td>\n<td>5<\/td>\n<td>5<\/td>\n<td>6<\/td>\n<td>11.0093<\/td>\n<td>80.1<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"3\">carbon<\/td>\n<td>[latex]{}_{6}^{12}\\text{C}[\/latex]<\/td>\n<td>6<\/td>\n<td>6<\/td>\n<td>6<\/td>\n<td>12.0000<\/td>\n<td>98.89<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{6}^{13}\\text{C}[\/latex]<\/td>\n<td>6<\/td>\n<td>6<\/td>\n<td>7<\/td>\n<td>13.0034<\/td>\n<td>1.11<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]{}_{6}^{14}\\text{C}[\/latex]<\/td>\n<td>6<\/td>\n<td>6<\/td>\n<td>8<\/td>\n<td>14.0032<\/td>\n<td>\u2014 (trace)<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"2\">nitrogen<\/td>\n<td>[latex]{}_{7}^{14}\\text{N}[\/latex]<\/td>\n<td>7<\/td>\n<td>7<\/td>\n<td>7<\/td>\n<td>14.0031<\/td>\n<td>99.63<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{7}^{15}\\text{N}[\/latex]<\/td>\n<td>7<\/td>\n<td>7<\/td>\n<td>8<\/td>\n<td>15.0001<\/td>\n<td>0.37<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"3\">oxygen<\/td>\n<td>[latex]{}_{8}^{16}\\text{O}[\/latex]<\/td>\n<td>8<\/td>\n<td>8<\/td>\n<td>8<\/td>\n<td>15.9949<\/td>\n<td>99.757<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{8}^{17}\\text{O}[\/latex]<\/td>\n<td>8<\/td>\n<td>8<\/td>\n<td>9<\/td>\n<td>16.9991<\/td>\n<td>0.038<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{8}^{18}\\text{O}[\/latex]<\/td>\n<td>8<\/td>\n<td>8<\/td>\n<td>10<\/td>\n<td>17.9992<\/td>\n<td>0.205<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>fluorine<\/td>\n<td>[latex]{}_{9}^{19}\\text{F}[\/latex]<\/td>\n<td>9<\/td>\n<td>9<\/td>\n<td>10<\/td>\n<td>18.9984<\/td>\n<td>100<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td rowspan=\"3\">neon<\/td>\n<td>[latex]{}_{10}^{20}\\text{Ne}[\/latex]<\/td>\n<td>10<\/td>\n<td>10<\/td>\n<td>10<\/td>\n<td>19.9924<\/td>\n<td>90.48<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{10}^{21}\\text{Ne}[\/latex]<\/td>\n<td>10<\/td>\n<td>10<\/td>\n<td>11<\/td>\n<td>20.9938<\/td>\n<td>0.27<\/td>\n<\/tr>\n<tr valign=\"middle\">\n<td>[latex]{}_{10}^{22}\\text{Ne}[\/latex]<\/td>\n<td>10<\/td>\n<td>10<\/td>\n<td>12<\/td>\n<td>21.9914<\/td>\n<td>9.25<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<div class=\"textbox\"><a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Use this Build an Atom simulator<\/a> to build atoms of the first 10 elements, see which isotopes exist, check nuclear stability, and gain experience with isotope symbols.<\/div>\n<h2>Atomic Mass<\/h2>\n<p>Because each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the <strong>atomic mass<\/strong> of a single atom is approximately equal to its mass number (a whole number). However, the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes.<\/p>\n<p>The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope\u2019s mass multiplied by its fractional abundance.<\/p>\n<p>[latex]\\text{average mass}=\\sum _{i}{\\left(\\text{fractional abundance}\\times \\text{isotopic mass}\\right)}_{i}[\/latex]<\/p>\n<p>For example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are <sup>10<\/sup>B with a mass of 10.0129 amu, and the remaining 80.1% are <sup>11<\/sup>B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:<\/p>\n<p>[latex]\\begin{array}{cc}\\hfill \\text{boron average mass}& =\\left(0.199\\times \\text{10.0129 amu}\\right)+\\left(0.801\\times \\text{11.0093 amu}\\right)\\hfill \\\\ & =\\text{1.99 amu}+\\text{8.82 amu}\\hfill \\\\ & =\\text{10.81 amu}\\hfill \\end{array}[\/latex]<\/p>\n<p>It is important to understand that no single boron atom weighs exactly 10.8 amu; 10.8 amu is the average mass of all boron atoms, and individual boron atoms weigh either approximately 10 amu or 11 amu.<\/p>\n<div class=\"textbox examples\">\n<h3>Example 2: <strong>Calculation of Average Atomic Mass<\/strong><\/h3>\n<p>A meteorite found in central Indiana contains traces of the noble gas neon picked up from the solar wind during the meteorite\u2019s trip through the solar system. Analysis of a sample of the gas showed that it consisted of 91.84% <sup>20<\/sup>Ne (mass 19.9924 amu), 0.47% <sup>21<\/sup>Ne (mass 20.9940 amu), and 7.69% <sup>22<\/sup>Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q877286\">Show Answer<\/span><\/p>\n<div id=\"q877286\" class=\"hidden-answer\" style=\"display: none\">\n<p>[latex]\\begin{array}{cc}\\hfill \\text{average mass}& =\\left(0.9184\\times \\text{19.9924 amu}\\right)+\\left(0.0047\\times \\text{20.9940 amu}\\right)+\\left(0.0769\\times \\text{21.9914 amu}\\right)\\hfill \\\\ & =\\left(18.36+0.099+1.69\\right)\\text{amu}\\hfill \\\\ & =\\text{20.15 amu}\\hfill \\end{array}[\/latex]<\/p>\n<p>The average mass of a neon atom in the solar wind is 20.15 amu. (The average mass of a terrestrial neon atom is 20.1796 amu. This result demonstrates that we may find slight differences in the natural abundance of isotopes, depending on their origin.)<\/p>\n<\/div>\n<\/div>\n<h4><strong>Check Your Learning<\/strong><\/h4>\n<p>A sample of magnesium is found to contain 78.70% of <sup>24<\/sup>Mg atoms (mass 23.98 amu), 10.13% of <sup>25<\/sup>Mg atoms (mass 24.99 amu), and 11.17% of <sup>26<\/sup>Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom.<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q74864\">Show Answer<\/span><\/p>\n<div id=\"q74864\" class=\"hidden-answer\" style=\"display: none\">24.31 amu<\/p>\n<p>We can also do variations of this type of calculation, as shown in the next example.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<div class=\"textbox\">\n<p><a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/isotopes-and-atomic-mass\" target=\"_blank\" rel=\"noopener\">Visit the PhET Isotopes and Atomic Mass site<\/a> to make mixtures of the main isotopes of the first 18 elements, gain experience with average atomic mass, and check naturally occurring isotope ratios using the Isotopes and Atomic Mass simulation.<\/p>\n<\/div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Concepts and Summary<\/h3>\n<p>Although no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure. Thomson\u2019s cathode ray tube showed that atoms contain small, negatively charged particles called electrons. Millikan discovered that there is a fundamental electric charge\u2014the charge of an electron. Rutherford\u2019s gold foil experiment showed that atoms have a small, dense, positively charged nucleus; the positively charged particles within the nucleus are called protons. Chadwick discovered that the nucleus also contains neutral particles called neutrons. Soddy demonstrated that atoms of the same element can differ in mass; these are called isotopes.<\/p>\n<p>An atom consists of a small, positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons; its diameter is about 100,000 times smaller than that of the atom. The mass of one atom is usually expressed in atomic mass units (amu), which is referred to as the atomic mass. An amu is defined as exactly [latex]\\frac{1}{12}[\/latex] of the mass of a carbon-12 atom and is equal to 1.6605 \u00d7 10<sup>\u221224<\/sup> g.<\/p>\n<p>Protons are relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu. Neutrons are relatively heavy particles with no charge and a mass of 1.0087 amu. Electrons are light particles with a charge of 1\u2212 and a mass of 0.00055 amu. The number of protons in the nucleus is called the atomic number (Z) and is the property that defines an atom\u2019s elemental identity. The sum of the numbers of protons and neutrons in the nucleus is called the mass number and, expressed in amu, is approximately equal to the mass of the atom. An atom is neutral when it contains equal numbers of electrons and protons.<\/p>\n<p>Isotopes of an element are atoms with the same atomic number but different mass numbers; isotopes of an element, therefore, differ from each other only in the number of neutrons within the nucleus. When a naturally occurring element is composed of several isotopes, the atomic mass of the element represents the average of the masses of the isotopes involved. A chemical symbol identifies the atoms in a substance using symbols, which are one-, two-, or three-letter abbreviations for the atoms.<\/p>\n<h4>Key Equations<\/h4>\n<ul>\n<li>[latex]\\text{average mass}=\\sum _{i}{\\left(\\text{fractional abundance}\\times \\text{isotopic mass}\\right)}_{i}[\/latex]<\/li>\n<\/ul>\n<\/div>\n<div class=\"textbox exercises\">\n<h3>Exercises<\/h3>\n<ol>\n<li>The existence of isotopes violates one of the original ideas of Dalton\u2019s atomic theory. Which one?<\/li>\n<li>How are electrons and protons similar? How are they different?<\/li>\n<li>How are protons and neutrons similar? How are they different?<\/li>\n<li>Predict and test the behavior of \u03b1 particles fired at a \u201cplum pudding\u201d model atom.\n<ol style=\"list-style-type: lower-alpha;\">\n<li>Predict the paths taken by \u03b1 particles that are fired at atoms with a Thomson\u2019s plum pudding model structure. Explain why you expect the \u03b1 particles to take these paths.<\/li>\n<li>If \u03b1 particles of higher energy than those in (a) are fired at plum pudding atoms, predict how their paths will differ from the lower-energy \u03b1 particle paths. Explain your reasoning.<\/li>\n<li>Now test your predictions from (a) and (b). Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\" target=\"_blank\" rel=\"noopener noreferrer\">Rutherford Scattering simulation<\/a> and select the \u201cPlum Pudding Atom\u201d tab. Set \u201cAlpha Particles Energy\u201d to \u201cmin,\u201d and select \u201cshow traces.\u201d Click on the gun to start firing \u03b1 particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Hit the pause button, or \u201cReset All.\u201d Set \u201cAlpha Particles Energy\u201d to \u201cmax,\u201d and start firing \u03b1 particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual paths as shown in the simulation.<\/li>\n<\/ol>\n<\/li>\n<li>Predict and test the behavior of \u03b1 particles fired at a Rutherford atom model.\n<ol style=\"list-style-type: lower-alpha;\">\n<li>Predict the paths taken by \u03b1 particles that are fired at atoms with a Rutherford atom model structure. Explain why you expect the \u03b1 particles to take these paths.<\/li>\n<li>If \u03b1 particles of higher energy than those in (a) are fired at Rutherford atoms, predict how their paths will differ from the lower-energy \u03b1 particle paths. Explain your reasoning.<\/li>\n<li>Predict how the paths taken by the \u03b1 particles will differ if they are fired at Rutherford atoms of elements other than gold. What factor do you expect to cause this difference in paths, and why?<\/li>\n<li>Now test your predictions from (a), (b), and (c). Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/rutherford-scattering\">Rutherford Scattering simulation<\/a> and select the \u201cRutherford Atom\u201d tab. Due to the scale of the simulation, it is best to start with a small nucleus, so select \u201c20\u201d for both protons and neutrons, \u201cmin\u201d for energy, show traces, and then start firing \u03b1 particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Pause or reset, set energy to \u201cmax,\u201d and start firing \u03b1 particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual path as shown in the simulation. Pause or reset, select \u201c40\u201d for both protons and neutrons, \u201cmin\u201d for energy, show traces, and fire away. Does this match your prediction from (c)? If not, explain why the actual path would be that shown in the simulation. Repeat this with larger numbers of protons and neutrons. What generalization can you make regarding the type of atom and effect on the path of \u03b1 particles? Be clear and specific.<\/li>\n<\/ol>\n<\/li>\n<li>In what way are isotopes of a given element always different? In what way(s) are they always the same?<\/li>\n<li>Write the symbol for each of the following ions:\n<ol>\n<li>the ion with a 1+ charge, atomic number 55, and mass number 133<\/li>\n<li>the ion with 54 electrons, 53 protons, and 74 neutrons<\/li>\n<li>the ion with atomic number 15, mass number 31, and a 3\u2212 charge<\/li>\n<li>the ion with 24 electrons, 30 neutrons, and a 3+ charge<\/li>\n<\/ol>\n<\/li>\n<li>Write the symbol for each of the following ions:\n<ol>\n<li>the ion with a 3+ charge, 28 electrons, and a mass number of 71<\/li>\n<li>the ion with 36 electrons, 35 protons, and 45 neutrons<\/li>\n<li>the ion with 86 electrons, 142 neutrons, and a 4+ charge<\/li>\n<li>the ion with a 2+ charge, atomic number 38, and mass number 87<\/li>\n<\/ol>\n<\/li>\n<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a> and click on the Atom icon.\n<ol>\n<li>Pick any one of the first 10 elements that you would like to build and state its symbol.<\/li>\n<li>Drag protons, neutrons, and electrons onto the atom template to make an atom of your element. State the numbers of protons, neutrons, and electrons in your atom, as well as the net charge and mass number.<\/li>\n<li>Click on \u201cNet Charge\u201d and \u201cMass Number,\u201d check your answers to (b), and correct, if needed.<\/li>\n<li>Predict whether your atom will be stable or unstable. State your reasoning.<\/li>\n<li>Check the \u201cStable\/Unstable\u201d box. Was your answer to (d) correct? If not, first predict what you can do to make a stable atom of your element, and then do it and see if it works. Explain your reasoning.<\/li>\n<\/ol>\n<\/li>\n<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a>\n<ol>\n<li>Drag protons, neutrons, and electrons onto the atom template to make a neutral atom of Oxygen-16 and give the isotope symbol for this atom.<\/li>\n<li>Now add two more electrons to make an ion and give the symbol for the ion you have created.<\/li>\n<\/ol>\n<\/li>\n<li>Open the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/build-an-atom\" target=\"_blank\" rel=\"noopener\">Build an Atom simulation<\/a>\n<ol>\n<li>Drag protons, neutrons, and electrons onto the atom template to make a neutral atom of Lithium-6 and give the isotope symbol for this atom.<\/li>\n<li>Now remove one electron to make an ion and give the symbol for the ion you have created.<\/li>\n<\/ol>\n<\/li>\n<li>Determine the number of protons, neutrons, and electrons in the following isotopes that are used in medical diagnoses:\n<ol>\n<li>atomic number 9, mass number 18, charge of 1\u2212<\/li>\n<li>atomic number 43, mass number 99, charge of 7+<\/li>\n<li>atomic number 53, atomic mass number 131, charge of 1\u2212<\/li>\n<li>atomic number 81, atomic mass number 201, charge of 1+<\/li>\n<li>Name the elements in parts (a), (b), (c), and (d).<\/li>\n<\/ol>\n<\/li>\n<li>The following are properties of isotopes of two elements that are essential in our diet. Determine the number of protons, neutrons and electrons in each and name them.\n<ol>\n<li>atomic number 26, mass number 58, charge of 2+<\/li>\n<li>atomic number 53, mass number 127, charge of 1\u2212<\/li>\n<\/ol>\n<\/li>\n<li>Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:\n<ol>\n<li>[latex]{}_{5}^{10}\\text{B}[\/latex]<\/li>\n<li>[latex]{}_{80}^{199}\\text{Hg}[\/latex]<\/li>\n<li>[latex]{}_{29}^{63}\\text{Cu}[\/latex]<\/li>\n<li>[latex]{}_{6}^{13}\\text{C}[\/latex]<\/li>\n<li>[latex]{}_{34}^{77}\\text{Se}[\/latex]<\/li>\n<\/ol>\n<\/li>\n<li>Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:\n<ol>\n<li>[latex]{}_{3}^{7}\\text{Li}[\/latex]<\/li>\n<li>[latex]{}_{52}^{125}\\text{Te}[\/latex]<\/li>\n<li>[latex]{}_{47}^{109}\\text{Ag}[\/latex]<\/li>\n<li>[latex]{}_{7}^{15}\\text{N}[\/latex]<\/li>\n<li>[latex]{}_{15}^{31}\\text{P}[\/latex]<\/li>\n<\/ol>\n<\/li>\n<li>Click on the <a href=\"https:\/\/phet.colorado.edu\/en\/simulation\/isotopes-and-atomic-mass\" target=\"_blank\" rel=\"noopener\">Isotopes and Common Mass website<\/a> and select the \u201cMix Isotopes\u201d tab, hide the \u201cPercent Composition\u201d and \u201cAverage Atomic Mass\u201d boxes, and then select the element boron\n<ol>\n<li>Write the symbols of the isotopes of boron that are shown as naturally occurring in significant amounts.<\/li>\n<li>Predict the relative amounts (percentages) of these boron isotopes found in nature. Explain the reasoning behind your choice.<\/li>\n<li>Add isotopes to the black box to make a mixture that matches your prediction in (b). You may drag isotopes from their bins or click on \u201cMore\u201d and then move the sliders to the appropriate amounts.<\/li>\n<li>Reveal the \u201cPercent Composition\u201d and \u201cAverage Atomic Mass\u201d boxes. How well does your mixture match with your prediction? If necessary, adjust the isotope amounts to match your prediction.<\/li>\n<li>Select \u201cNature\u2019s\u201d mix of isotopes and compare it to your prediction. How well does your prediction compare with the naturally occurring mixture? Explain. If necessary, adjust your amounts to make them match \u201cNature\u2019s\u201d amounts as closely as possible.<\/li>\n<\/ol>\n<\/li>\n<li>Repeat Exercise 11 using an element that has three naturally occurring isotopes.<\/li>\n<li>An element has the following natural abundances and isotopic masses: 90.92% abundance with 19.99 amu, 0.26% abundance with 20.99 amu, and 8.82% abundance with 21.99 amu. Calculate the average atomic mass of this element.<\/li>\n<li>Average atomic masses listed by IUPAC are based on a study of experimental results. Bromine has two isotopes <sup>79<\/sup>Br and <sup>81<\/sup>Br, whose masses (78.9183 and 80.9163 amu) and abundances (50.69% and 49.31%) were determined in earlier experiments. Calculate the average atomic mass of bromine based on these experiments.<\/li>\n<li>Variations in average atomic mass may be observed for elements obtained from different sources. Lithium provides an example of this. The isotopic composition of lithium from naturally occurring minerals is 7.5% <sup>6<\/sup>Li and 92.5% <sup>7<\/sup>Li, which have masses of 6.01512 amu and 7.01600 amu, respectively. A commercial source of lithium, recycled from a military source, was 3.75% <sup>6<\/sup>Li (and the rest <sup>7<\/sup>Li). Calculate the average atomic mass values for each of these two sources.<\/li>\n<li>The average atomic masses of some elements may vary, depending upon the sources of their ores. Naturally occurring boron consists of two isotopes with accurately known masses (<sup>10<\/sup>B, 10.0129 amu and <sup>11<\/sup>B, 11.0931 amu). The actual atomic mass of boron can vary from 10.807 to 10.819, depending on whether the mineral source is from Turkey or the United States. Calculate the percent abundances leading to the two values of the average atomic masses of boron from these two countries.<\/li>\n<li>The <sup>18<\/sup>O:<sup>16<\/sup>O abundance ratio in some meteorites is greater than that used to calculate the average atomic mass of oxygen on earth. Is the average mass of an oxygen atom in these meteorites greater than, less than, or equal to that of a terrestrial oxygen atom?<\/li>\n<\/ol>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q456814\">Show Selected Answers<\/span><\/p>\n<div id=\"q456814\" class=\"hidden-answer\" style=\"display: none\">\n<p>1. Dalton originally thought that all atoms of a particular element had identical properties, including mass. Thus, the concept of isotopes, in which an element has different masses, was a violation of the original idea. To account for the existence of isotopes, the second postulate of his atomic theory was modified to state that atoms of the same element must have identical chemical properties.<\/p>\n<p>3. Both are subatomic particles that reside in an atom\u2019s nucleus. Both have approximately the same mass. Protons are positively charged, whereas neutrons are uncharged.<\/p>\n<p>5. The answers are as follows:<\/p>\n<ol style=\"list-style-type: lower-alpha;\">\n<li>The Rutherford atom has a small, positively charged nucleus, so most \u03b1 particles will pass through empty space far from the nucleus and be undeflected. Those \u03b1 particles that pass near the nucleus will be deflected from their paths due to positive-positive repulsion. The more directly toward the nucleus the \u03b1 particles are headed, the larger the deflection angle will be.<\/li>\n<li>Higher-energy \u03b1 particles that pass near the nucleus will still undergo deflection, but the faster they travel, the less the expected angle of deflection.<\/li>\n<li>If the nucleus is smaller, the positive charge is smaller and the expected deflections are smaller\u2014both in terms of how closely the \u03b1 particles pass by the nucleus undeflected and the angle of deflection. If the nucleus is larger, the positive charge is larger and the expected deflections are larger\u2014more \u03b1 particles will be deflected, and the deflection angles will be larger.<\/li>\n<li>The paths followed by the \u03b1 particles match the predictions from (a), (b), and (c).<\/li>\n<\/ol>\n<p>7. (a) <sup>133<\/sup>Cs<sup>+<\/sup>; (b) <sup>127<\/sup>I<sup>\u2212<\/sup>; (c) <sup>31<\/sup>P<sup>3\u2212<\/sup>; (d) <sup>57<\/sup>Co<sup>3+<\/sup><\/p>\n<p>9. The answers are as follows:<\/p>\n<ol>\n<li>Carbon-12, <sup>12<\/sup>C<\/li>\n<li>This atom contains six protons and six neutrons. There are six electrons in a neutral <sup>12<\/sup>C atom. The net charge of such a neutral atom is zero, and the mass number is 12.<\/li>\n<li>The preceding answers are correct.<\/li>\n<li>The atom will be stable since C-12 is a stable isotope of carbon.<\/li>\n<li>The preceding answer is correct. Other answers for this exercise are possible if a different element of isotope is chosen.<\/li>\n<\/ol>\n<p>11. The answers are as follows:<\/p>\n<ol>\n<li>Lithium-6 contains three protons, three neutrons, and three electrons. The isotope symbol is <sup>6<\/sup>Li or [latex]{}_{3}^{6}\\text{Li}[\/latex]<\/li>\n<li>[latex]{}_{}^{6}{\\text{Li}}^{+}[\/latex] or [latex]{}_{3}^{6}{\\text{Li}}^{+}[\/latex]<\/li>\n<\/ol>\n<p>13. The answers are as follows:<\/p>\n<ol>\n<li>iron, 26 protons, 24 electrons, and 32 neutrons<\/li>\n<li>iodine, 53 protons, 54 electrons, and 74 neutrons<\/li>\n<\/ol>\n<p>15. The number of protons, electrons, and neutrons for each isotope are as follows:<\/p>\n<ol>\n<li>3 protons, 3 electrons, 4 neutrons<\/li>\n<li>52 protons, 52 electrons, 73 neutrons<\/li>\n<li>47 protons, 47 electrons, 62 neutrons<\/li>\n<li>7 protons, 7 electrons, 8 neutrons<\/li>\n<li>15 protons, 15 electrons, 16 neutrons<\/li>\n<\/ol>\n<p>17. Let us use neon as an example. Since there are three isotopes, there is no way to be sure to accurately predict the abundances to make the total of 20.18 amu average atomic mass. Let us guess that the abundances are 9% Ne-22, 91% Ne-20, and only a trace of Ne-21. The average mass would be 20.18 amu. Checking the nature\u2019s mix of isotopes shows that the abundances are 90.48% Ne-20, 9.25% Ne-22, and 0.27% Ne-21, so our guessed amounts have to be slightly adjusted.<\/p>\n<p>19. 79.904 amu<\/p>\n<p>21. Turkey source: 0.2649 (of 10.0129 amu isotope); US source: 0.2537 (of 10.0129 amu isotope)<\/p>\n<\/div>\n<\/div>\n<\/div>\n<h2>Glossary<\/h2>\n<p><strong>alpha particle (\u03b1 particle): <\/strong>positively charged particle consisting of two protons and two neutrons<\/p>\n<p><strong>anion: <\/strong>negatively charged atom or molecule (contains more electrons than protons)<\/p>\n<p><strong>atomic mass: <\/strong>average mass of atoms of an element, expressed in amu<\/p>\n<p><strong>atomic mass unit (amu): <\/strong>(also, unified atomic mass unit, u, or Dalton, Da) unit of mass equal to [latex]\\frac{1}{12}[\/latex] of the mass of a <sup>12<\/sup>C atom<\/p>\n<p><strong>atomic number (Z): <\/strong>number of protons in the nucleus of an atom<\/p>\n<p><strong>cation: <\/strong>positively charged atom or molecule (contains fewer electrons than protons)<\/p>\n<p><strong>chemical symbol: <\/strong>one-, two-, or three-letter abbreviation used to represent an element or its atoms<\/p>\n<p><strong>Dalton (Da): <\/strong>alternative unit equivalent to the atomic mass unit<\/p>\n<p><strong>electron: <\/strong>negatively charged, subatomic particle of relatively low mass located outside the nucleus<\/p>\n<p><strong>fundamental unit of charge: <\/strong>(also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 \u00d7 10<sup>\u221219<\/sup> C<\/p>\n<p><strong>ion: <\/strong>electrically charged atom or molecule (contains unequal numbers of protons and electrons)<\/p>\n<p><strong>isotopes: <\/strong>atoms that contain the same number of protons but different numbers of neutrons<\/p>\n<p><strong>mass number (A): <\/strong>sum of the numbers of neutrons and protons in the nucleus of an atom<\/p>\n<p><strong>neutron: <\/strong>uncharged, subatomic particle located in the nucleus<\/p>\n<p><strong>nucleus: <\/strong>massive, positively charged center of an atom made up of protons and neutrons<\/p>\n<p><strong>proton: <\/strong>positively charged, subatomic particle located in the nucleus<\/p>\n<p><strong>unified atomic mass unit (u): <\/strong>alternative unit equivalent to the atomic mass unit<\/p>\n<p>&nbsp;<\/p>\n\n\t\t\t <section class=\"citations-section\" role=\"contentinfo\">\n\t\t\t <h3>Candela Citations<\/h3>\n\t\t\t\t\t <div>\n\t\t\t\t\t\t <div id=\"citation-list-112\">\n\t\t\t\t\t\t\t <div class=\"licensing\"><div class=\"license-attribution-dropdown-subheading\">CC licensed content, Shared previously<\/div><ul class=\"citation-list\"><li>Introductory Chemistry- 1st Canadian Edition . <strong>Authored by<\/strong>: Jessie A. Key and David W. Ball. <strong>Provided by<\/strong>: BCCampus. <strong>Located at<\/strong>: <a target=\"_blank\" href=\"https:\/\/opentextbc.ca\/introductorychemistry\/\">https:\/\/opentextbc.ca\/introductorychemistry\/<\/a>. <strong>License<\/strong>: <em><a target=\"_blank\" rel=\"license\" href=\"https:\/\/creativecommons.org\/licenses\/by-nc-sa\/4.0\/\">CC BY-NC-SA: Attribution-NonCommercial-ShareAlike<\/a><\/em>. <strong>License Terms<\/strong>: Download this book for free at http:\/\/open.bccampus.ca<\/li><\/ul><\/div>\n\t\t\t\t\t\t <\/div>\n\t\t\t\t\t <\/div>\n\t\t\t <\/section>","protected":false},"author":23485,"menu_order":3,"template":"","meta":{"_candela_citation":"[{\"type\":\"cc\",\"description\":\"Introductory Chemistry- 1st Canadian Edition \",\"author\":\"Jessie A. Key and David W. 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