4.5: Atomic Structure
12
Harper College Chemistry Department
[latexpage]
Composition of an Atom
The development of the 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 1 ).
Atoms—and the protons, neutrons and electrons that compose them—are extremely small. For example, a carbon atom weighs less than 2 × 10 −23 g and an electron has a charge of less than 2 × 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” as will be discussed later in this module. Thus, one amu is exactly $\frac{1}{12}$ of the mass of one carbon-12 atom: 1 amu = 1.6605 × 10 −24 g. The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 × 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 1 . An observant student might notice that the sum of an atom’s subatomic particles does not equal the atom’s actual mass: The total mass of six protons, six neutrons and six electrons is 12.0993 amu, slightly larger than 12.00 amu. This “missing” mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.
| Properties of Subatomic Particles | |||||
|---|---|---|---|---|---|
| Name | Location | Charge (C) | Unit Charge | Mass (amu) | Mass (g) |
| electron | outside nucleus | −1.602 × 10 −19 | 1− | 0.00055 | 0.00091 × 10 −24 |
| proton | nucleus | 1.602 × 10 −19 | 1+ | 1.00727 | 1.67262 × 10 −24 |
| neutron | nucleus | 0 | 0 | 1.00866 | 1.67493 × 10 −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. 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.
\[\begin{equation*} \begin{align} \text{atomic number (Z)} &= \text{number of protons} \\ \text{mass number (A)} &= \text{number of protons} + \text{number of neutrons} \\ \text{A} - \text{Z} &= \text{number of neutrons} \\ \end{align} \end{equation*}\]
Atomic charge = number of protons − number of electrons
As will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire a charge by gaining or losing electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an anion. Positively charged atoms called cations are formed when an atom loses one or more electrons. For example, a neutral sodium atom (Z = 11) has 11 protons and 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 protons and electrons, if it gains two electrons it will become an anion with a 2− charge (8 − 10 = 2−).
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 ).
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 as anions and each has a 1− charge and a mass number of 127.
Check Your Learning
Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.
Answer
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 number 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].
Check Your Learning
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?
78 protons; 117 neutrons; charge is 4+
Chemical Symbols
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 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).
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 come from the name in another language (usually Latin). Symbols have one or two letters. 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 formula 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.
| 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(s) 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 seaborgium (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements.
Visit the IUPAC site to learn more about the International Union of Pure and Applied Chemistry and explore its periodic table at the bottom of the web page.
Isotopes
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 4 ). The atomic number is sometimes written as a subscript also on the left of the symbol. However, 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 24 Mg, 25 Mg and 26 Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, 24 Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” 25 Mg 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 24 Mg atom has 12 neutrons in its nucleus, a 25 Mg atom has 13 neutrons and a 26 Mg atom has 14 neutrons.
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 2 H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3 H, is also called tritium and sometimes symbolized T.
| Nuclear Compositions of Atoms of the Very Light Elements | ||||||
|---|---|---|---|---|---|---|
| Element | Symbol | Atomic Number | Number of Protons | Number of Neutrons | Mass (amu) | % Natural Abundance |
| hydrogen | $_{1}^{1}\text{H}$ (protium) | 1 | 1 | 0 | 1.0078 | 99.989 |
| $_{1}^{2}\text{H}$ (deuterium) | 1 | 1 | 1 | 2.0141 | 0.0115 | |
| $_{1}^{3}\text{H}$ (tritium) | 1 | 1 | 2 | 3.01605 | (trace amounts) | |
| helium | $_{2}^{3}\text{He}$ | 2 | 2 | 1 | 3.01603 | 0.00013 |
| $_{2}^{4}\text{He}$ | 2 | 2 | 2 | 4.0026 | 100 | |
| lithium | $_{3}^{6}\text{Li}$ | 3 | 3 | 3 | 6.0151 | 7.59 |
| $_{3}^{7}\text{Li}$ | 3 | 3 | 4 | 7.0160 | 92.41 | |
| beryllium | $_{4}^{9}\text{Be}$ | 4 | 4 | 5 | 9.0122 | 100 |
| boron | $^{10}_{ 5}\text{B}$ | 5 | 5 | 5 | 10.0129 | 19.9 |
| $_{ 5}^{11}\text{B}$ | 5 | 5 | 6 | 11.0093 | 80.1 | |
| carbon | $_{6}^{12}\text{C}$ | 6 | 6 | 6 | 12.0000 | 98.89 |
| $_{6}^{13}\text{C}$ | 6 | 6 | 7 | 13.0034 | 1.11 | |
| $_{6}^{14}\text{C}$ | 6 | 6 | 8 | 14.0032 | (trace amounts) | |
| nitrogen | $_{7}^{14}\text{N}$ | 7 | 7 | 7 | 14.0031 | 99.63 |
| $_{7}^{15}\text{N}$ | 7 | 7 | 8 | 15.0001 | 0.37 | |
| oxygen | $_{8}^{16}\text{O}$ | 8 | 8 | 8 | 15.9949 | 99.757 |
| $_{8}^{17}\text{O}$ | 8 | 8 | 9 | 16.9991 | 0.038 | |
| $_{8}^{18}\text{O}$ | 8 | 8 | 10 | 17.9992 | 0.205 | |
| fluorine | $_{9}^{19}\text{F}$ | 9 | 9 | 10 | 18.9984 | 100 |
| neon | $_{10}^{20}\text{Ne}$ | 10 | 10 | 10 | 19.9924 | 90.48 |
| $_{10}^{21}\text{Ne}$ | 10 | 10 | 11 | 20.9938 | 0.27 | |
| $_{10}^{22}\text{Ne}$ | 10 | 10 | 12 | 21.9914 | 9.25 |
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.
Atomic Mass
The atomic mass of a single atom is approximately equal to its mass number (a whole number) due to the mass of each proton and each neutron contributing approximately one amu each to an atom, while each electron contributes far less. 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 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’s mass multiplied by its fractional abundance.
In case you are wondering
MASS SPECTROMETRYChemistry End of Chapter Exercises
1. In what way are isotopes of a given element always different? In what way(s) are they always the same?
2. Write the symbol for each of the following ions:
(a) the ion with a 1+ charge, atomic number 55, and mass number 133
(b) the ion with 54 electrons, 53 protons, and 74 neutrons
(c) the ion with atomic number 15, mass number 31, and a 3− charge
(d) the ion with 24 electrons, 30 neutrons, and a 3+ charge
3. Write the symbol for each of the following ions:
(a) the ion with a 3+ charge, 28 electrons, and a mass number of 71
(b) the ion with 36 electrons, 35 protons, and 45 neutrons
(c) the ion with 86 electrons, 142 neutrons, and a 4+ charge
(d) the ion with a 2+ charge, atomic number 38, and mass number 87
4. Open the Build an Atom simulation and click on the Atom icon.
(a) Pick any one of the first 10 elements that you would like to build and state its symbol.
(b) Drag protons, neutrons, and electrons onto the atom template to make an atom of your element.
(c) State the numbers of protons, neutrons, and electrons in your atom, as well as the net charge and mass number.
(d) Click on “Net Charge” and “Mass Number,” check your answers to (c), and correct, if needed.
(e) Predict whether your atom will be stable or unstable. State your reasoning.
(f) Check the “Stable/Unstable” box. Was your answer to (e) 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.
5. Open the Build an Atom simulation
(a) 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.
(b) Now add two more electrons to make an ion and give the symbol for the ion you have created.
6. Open the Build an Atom simulation
(a) 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.
(b) Now remove one electron to make an ion and give the symbol for the ion you have created.
7. Determine the number of protons, neutrons, and electrons in the following isotopes that are used in medical diagnoses:
(a) atomic number 9, mass number 18, charge of 1−
(b) atomic number 43, mass number 99, charge of 7+
(c) atomic number 53, atomic mass number 131, charge of 1−
(d) atomic number 81, atomic mass number 201, charge of 1+
(e) Name the elements in parts (a), (b), (c), and (d).
8. 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.
(a) atomic number 26, mass number 58, charge of 2+
(b) atomic number 53, mass number 127, charge of 1−
9. Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:
(a) $^{10}_{ 5}\text{B}$
(b) $^{199}_{80}\text{Hg}$
(c) $^{63}_{29}\text{Cu}$
(d) $^{13}_{6}\text{C}$
(e) $^{77}_{34}\text{Se}$
10. Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes:
(a) $^{7}_{3}\text{Li}$
(b) $^{125}_{52}\text{Te}$
(c) $^{109}_{74}\text{Ag}$
(d) $^{15}_{7}\text{N}$
(e) $^{31}_{15}\text{P}$
Glossary
- anion
- negatively charged atom or molecule (contains more electrons than protons)
- atomic mass
- average mass of atoms of an element, expressed in amu
- atomic mass unit (amu)
- (also, unified atomic mass unit, u, or Dalton, Da) unit of mass equal to 1 12 112 of the mass of a 12 C atom
- atomic number (Z)
- number of protons in the nucleus of an atom
- cation
- positively charged atom or molecule (contains fewer electrons than protons)
- chemical symbol
- one-, two-, or three-letter abbreviation used to represent an element or its atoms
- Dalton (Da)
- alternative unit equivalent to the atomic mass unit
- fundamental unit of charge
- (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 × 10 −19 C
- ion
- electrically charged atom or molecule (contains unequal numbers of protons and electrons)
- isotope
- same number of protons and different number of neutrons
mass number (A)
sum of the numbers of neutrons and protons in the nucleus of an atomunified atomic mass unit (u)
alternative unit equivalent to the atomic mass unit