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3.2: Atomic Structure and Symbols

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  • Page ID
    98536
  • Skills to Develop

    • Write and interpret symbols that depict the atomic number, mass number, and charge of an atom or ion
    • Define the atomic mass unit and average atomic mass
    • Calculate average atomic mass and isotopic abundance

    Despite the success of Dalton's atomic theory in accounting for the mass laws, many scientists throughout the 1800s regarded "atoms" as a convenient fiction, a useful way of keeping track of different elements in different chemical reactions, but not representing a reality. After all, no one had been able to observe these tiny atoms. In fact, it would be almost 100 years after the publication of Dalton's theory that definitive evidence for the real existence of atoms was obtained. In the end, Dalton was proved both correct and incorrect: atoms do indeed exist, but they are more than simple, indivisible spheres. In this section, we will summarize the basic facts known about atoms and the even smaller particles that make them up.

    The Structure of Atoms

    The development of modern atomic theory revealed much about the inner structure of atoms. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10−10 m, whereas the diameter of the nucleus is roughly 10−15 m—about 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 \(\PageIndex{1}\)).

    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.

    Figure \(\PageIndex{1}\): 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 “babyknight”/Wikimedia Commons; credit right: modification of work by Paxson Woelber).

    Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, a carbon atom weighs less than 2 \(\times\) 10−23 g, and an electron has a charge of less than 2 \(\times\) 10−19 C (coulomb). When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu) and the fundamental unit of charge (e). 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 “carbon-12” and will be discussed later in this module.) Thus, one amu is exactly \(1/12\) of the mass of one carbon-12 atom: 1 amu = 1.6605 \(\times\) 10−24 g. (The Dalton (Da) and the unified atomic mass unit (u) 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 \(\times\) 10−19 C.

    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− 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 \(\PageIndex{1}\). (An observant student might notice that the sum of an atom’s subatomic particles does not quite equal the atom’s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than the 12.00 amu of an actual carbon-12 atom. This “missing” mass is known as the mass defect, and results from principles of nuclear chemistry outside our present scope.)

    Table \(\PageIndex{1}\): Properties of Subatomic Particles
    Name Location Charge (C) Unit Charge Mass (amu) Mass (g)
    electron outside nucleus \(−1.602 \times 10^{−19}\) 1− 0.00055 \(0.00091 \times 10^{−24}\)
    proton nucleus \(1.602 \times 10^{−19}\) 1+ 1.00727 \(1.67262 \times 10^{−24}\)
    neutron nucleus 0 0 1.00866 \(1.67493 \times10^{−24}\)

    The number of protons in the nucleus of an atom is its atomic number (Z). 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 mass number (A). The number of neutrons is therefore the difference between the mass number and the atomic number: A – Z = number of neutrons. It turns out that atoms of the same element may have differing numbers of neutrons; these atoms will thus have different masses and are called isotopes.

    \[\begin{align*}
    \ce{atomic\: number\:(Z)\: &= \:number\: of\: protons\\
    mass\: number\:(A)\: &= \:number\: of\: protons + number\: of\: neutrons\\
    A-Z\: &= \:number\: of\: neutrons}
    \end{align*}\]

    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 not equal, the atom is electrically charged and is called an ion. The charge of an atom is defined as follows:

    Atomic charge = number of protons − number of electrons

    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 anion. A positively charged atom is called a cation and is 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 − 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− charge (8 − 10 = 2−).

    Example \(\PageIndex{1}\): Composition of an Atom

    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 \(\PageIndex{2}\)). If left untreated, congenital thyroid problems due to iodine deficiency can result in stunted growth and intellectual disability in children.

    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.

    Figure \(\PageIndex{2}\): (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 “Almazi”/Wikimedia Commons; credit b: modification of work by Mike Mozart)

    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’s population is still at risk of iodine deficiency. The iodine atoms are added to salt as anions, and each has a 1− charge and a mass number of 127. Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.

    Solution

    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 − 53 = 74). Since the iodine is added as a 1− anion, the number of electrons is 54 [53 – (1–) = 54].

    Exercise \(\PageIndex{1}\)

    An ion 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?

    Answer

    78 protons; 117 neutrons; charge is 4+

    Chemical Symbols

    To simplify discussions of chemistry, chemists have come up with a system of symbols to refer to atoms, elements, and compounds. A chemical symbol 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 \(\PageIndex{3}\)). 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).

    A jar is shown with a small amount of liquid mercury in it.

    Figure \(\PageIndex{3}\): The symbol Hg represents the element mercury regardless of the amount; it could represent one atom of mercury or a large amount of mercury. Image used with permission from Wikipedia (user: Materialscientist).

    The symbols for several common elements and their atoms are listed in Table \(\PageIndex{2}\). Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Symbols have one or two letters, for example, H for hydrogen and Cl for chlorine. 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 the periodic table.

    Table \(\PageIndex{2}\): Some Common Elements and Their Symbols
    Element Symbol Element Symbol
    aluminum Al iron Fe (from ferrum)
    bromine Br lead Pb (from plumbum)
    calcium Ca magnesium Mg
    carbon C mercury Hg (from hydrargyrum)
    chlorine Cl nitrogen N
    chromium Cr oxygen O
    cobalt Co potassium K (from kalium)
    copper Cu (from cuprum) silicon Si
    fluorine F silver Ag (from argentum)
    gold Au (from aurum) sodium Na (from natrium)
    helium He sulfur S
    hydrogen H tin Sn (from stannum)
    iodine I zinc Zn

    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 locations; for example, element 106 is now known as seaborgium (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements.

    Isotopes

    Because atoms may exist as isotopes, we also need a symbolic representation for different numbers of neutrons in an atomic nucleus. 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 \(\PageIndex{4}\)). The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element’s 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 24Mg, 25Mg, and 26Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, 24Mg is read as “magnesium 24,” and can also be written as “magnesium-24” or “Mg-24.” 25Mg is read as “magnesium 25,” and can be written as “magnesium-25” or “Mg-25.” All magnesium atoms have 12 protons in their nucleus. They differ only because a 24Mg atom has 12 neutrons in its nucleus, a 25Mg atom has 13 neutrons, and a 26Mg has 14 neutrons.

    This diagram shows the symbol for helium, “H e.” 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 “M g” which stands for magnesium It has a mass number of 24, a charge of positive 2, and an atomic number of 12.

    Figure \(\PageIndex{4}\): 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.

    Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table \(\PageIndex{2}\). 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 2H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3H, is also called tritium and sometimes symbolized T.

    You can use this Build an Atom Simulator to build atoms of the first 10 elements, see which isotopes exist, check nuclear stability, and gain experience with isotope symbols.

    Table \(\PageIndex{2}\): Nuclear Compositions of Atoms of the Very Light Elements
    Element Symbol Atomic Number Number of Protons Number of Neutrons Mass (amu) % Natural Abundance
    hydrogen \(\ce{^1_1H}\)
    (protium)
    1 1 0 1.0078 99.989
    \(\ce{^2_1H}\)
    (deuterium)
    1 1 1 2.0141 0.0115
    \(\ce{^3_1H}\)
    (tritium)
    1 1 2 3.01605 — (trace)
    helium \(\ce{^3_2He}\) 2 2 1 3.01603 0.00013
    \(\ce{^4_2He}\) 2 2 2 4.0026 100
    lithium \(\ce{^6_3Li}\) 3 3 3 6.0151 7.59
    \(\ce{^7_3Li}\) 3 3 4 7.0160 92.41
    beryllium \(\ce{^9_4Be}\) 4 4 5 9.0122 100
    boron \(\ce{^{10}_5B}\) 5 5 5 10.0129 19.9
    \(\ce{^{11}_5B}\) 5 5 6 11.0093 80.1
    carbon \(\ce{^{12}_6C}\) 6 6 6 12.0000 98.89
    \(\ce{^{13}_6C}\) 6 6 7 13.0034 1.11
    \(\ce{^{14}_6C}\) 6 6 8 14.0032 — (trace)
    nitrogen \(\ce{^{14}_7N}\) 7 7 7 14.0031 99.63
    \(\ce{^{15}_7N}\) 7 7 8 15.0001 0.37
    oxygen \(\ce{^{16}_8O}\) 8 8 8 15.9949 99.757
    \(\ce{^{17}_8O}\) 8 8 9 16.9991 0.038
    \(\ce{^{18}_8O}\) 8 8 10 17.9992 0.205
    fluorine \(\ce{^{19}_9F}\) 9 9 10 18.9984 100
    neon \(\ce{^{20}_{10}Ne}\) 10 10 10 19.9924 90.48
    \(\ce{^{21}_{10}Ne}\) 10 10 11 20.9938 0.27
    \(\ce{^{22}_{10}Ne}\) 10 10 12 21.9914 9.25
     

    Atomic Mass

    Because each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the atomic mass 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.

    The mass of an element shown in the 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’s mass multiplied by its fractional abundance.

    \[\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance\times isotopic\: mass})_i\]

    For example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are 10B with a mass of 10.0129 amu, and the remaining 80.1% are 11B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:

    \[\begin{align*}
    \textrm{boron average mass} &=\mathrm{(0.199\times10.0129\: amu)+(0.801\times11.0093\: amu)}\\
    &=\mathrm{1.99\: amu+8.82\: amu}\\
    &=\mathrm{10.81\: amu}
    \end{align*}\]

    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.

    Example \(\PageIndex{2}\): Calculation of Average Atomic Mass

    A meteorite found in central Indiana contains traces of the noble gas neon picked up from the solar wind during the meteorite’s trip through the solar system. Analysis of a sample of the gas showed that it consisted of 91.84% 20Ne (mass 19.9924 amu), 0.47% 21Ne (mass 20.9940 amu), and 7.69% 22Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?

    Solution

    \[\begin{align*}
    \mathrm{average\: mass} &=\mathrm{(0.9184\times19.9924\: amu)+(0.0047\times20.9940\: amu)+(0.0769\times21.9914\: amu)}\\
    &=\mathrm{(18.36+0.099+1.69)\:amu}\\
    &=\mathrm{20.15\: amu}
    \end{align*}\]

    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.)

    Exercise \(\PageIndex{2}\)

    A sample of magnesium is found to contain 78.70% of 24Mg atoms (mass 23.98 amu), 10.13% of 25Mg atoms (mass 24.99 amu), and 11.17% of 26Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom. 

    Answer

    24.31 amu

    We can also do variations of this type of calculation, as shown in the next example.

    Example \(\PageIndex{3}\): Calculation of Percent Abundance

    Naturally occurring chlorine consists of 35Cl (mass 34.96885 amu) and 37Cl (mass 36.96590 amu), with an average mass of 35.453 amu. What is the percent composition of Cl in terms of these two isotopes?

    Solution

    The average mass of chlorine is the fraction that is 35Cl times the mass of 35Cl plus the fraction that is 37Cl times the mass of 37Cl.

    \[\mathrm{average\: mass=(fraction\: of\: ^{35}Cl\times mass\: of\: ^{35}Cl)+(fraction\: of\: ^{37}Cl\times mass\: of\: ^{37}Cl)}\]

    If we let x represent the fraction that is 35Cl, then the fraction that is 37Cl is represented by 1.00 − x.

    (The fraction that is 35Cl + the fraction that is 37Cl must add up to 1, so the fraction of 37Cl must equal 1.00 − the fraction of 35Cl.)

    Substituting this into the average mass equation, we have:

    \[\begin{align*}
    \mathrm{35.453\: amu} &=(x\times 34.96885\: \ce{amu})+[(1.00-x)\times 36.96590\: \ce{amu}]\\
    35.453 &=34.96885x+36.96590-36.96590x\\
    1.99705x &=1.513\\
    x&=\dfrac{1.513}{1.99705}=0.7576
    \end{align*}\]

    So solving yields: x = 0.7576, which means that 1.00 − 0.7576 = 0.2424. Therefore, chlorine consists of 75.76% 35Cl and 24.24% 37Cl.

    Exercise \(\PageIndex{3}\)

    Naturally occurring copper consists of 63Cu (mass 62.9296 amu) and 65Cu (mass 64.9278 amu), with an average mass of 63.546 amu. What is the percent composition of Cu in terms of these two isotopes?

    Answer

    69.15% Cu-63 and 30.85% Cu-65

    The left diagram shows how a mass spectrometer works, which is primarily a large tube that bends downward at its midpoint. The sample enters on the left side of the tube. A heater heats the sample, causing it to vaporize. The sample is also hit with a beam of electrons as it is being vaporized. Charged particles from the sample, called ions, are then accelerated and pass between two magnets. The magnetic field deflects the lightest ions most. The deflection of the ions is measured by a detector located on the right side of the tube. The graph to the right of the spectrometer shows a mass spectrum of zirconium. The relative abundance, as a percentage from 0 to 100, is graphed on the y axis, and the mass to charge ratio is graphed on the x axis. The sample contains five different isomers of zirconium. Z R 90, which has a mass to charge ratio of 90, is the most abundant isotope at about 51 percent relative abundance. Z R 91 has a mass to charge ratio of 91 and a relative abundance of about 11 percent. Z R 92 has a mass to charge ratio of 92 and a relative abundance of about 18 percent. Z R 94 has a mass to charge ratio of 94 and a relative abundance of about 18 percent. Z R 96, which has a mass to charge ratio of 96, is the least abundant zirconium isotope with a relative abundance of about 2 percent.

      Figure \(\PageIndex{5}\): Analysis of zirconium in a mass spectrometer produces a mass spectrum with peaks showing the different isotopes of Zr.

    The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass spectrometry (MS) is widely used in chemistry, forensics, medicine, environmental science, and many other fields to analyze and help identify the substances in a sample of material. In a typical mass spectrometer (Figure \(\PageIndex{5}\)), the sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms (or molecules) to become electrically charged, typically by losing one or more electrons. These cations then pass through a (variable) electric or magnetic field that deflects each cation’s path to an extent that depends on both its mass and charge (similar to how the path of a large steel ball bearing rolling past a magnet is deflected to a lesser extent that that of a small steel BB). The ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios (a mass spectrum) is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio. Since its initial use during the development of modern atomic theory, MS has evolved to become a powerful tool for chemical analysis in a wide range of applications.

    Video \(\PageIndex{1}\): Watch this video from the Royal Society for Chemistry for a brief description of the rudiments of mass spectrometry.

    Summary

    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 \(1/12\) of the mass of a carbon-12 atom and is equal to 1.6605 \(\times\) 10−24 g.

    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− 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’s 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.

    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- or two-letter abbreviations for the atoms.

    Key Equations

    • \(\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance \times isotopic\: mass})_i\)

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