CHAPTER 2

Atoms, Molecules, and Ions

Figure 2.1 Analysis of molecules in an exhaled breath can provide valuable information, leading to early diagnosis
of diseases or detection of environmental exposure to harmful substances. (credit: modification of work by Paul
Flowers)

CHAPTER OUTLINE
2.1 Early Ideas in Atomic Theory
2.2 Evolution of Atomic Theory
2.3 Atomic Structure and Symbolism
2.4 Chemical Formulas
2.5 The Periodic Table
2.6 Ionic and Molecular Compounds
2.7 Chemical Nomenclature

INTRODUCTION Lung diseases and lung cancers are among the world's most devastating illnesses partly due
to delayed detection and diagnosis. Most noninvasive screening procedures aren't reliable, and patients often
resist more accurate methods due to discomfort with the procedures or with the potential danger that the
procedures cause. But what if you could be accurately diagnosed through a simple breath test?

Early detection of biomarkers, substances that indicate an organism’s disease or physiological state, could
allow diagnosis and treatment before a condition becomes serious or irreversible. Recent studies have shown
that your exhaled breath can contain molecules that may be biomarkers for recent exposure to environmental
contaminants or for pathological conditions ranging from asthma to lung cancer. Scientists are working to
develop biomarker “fingerprints” that could be used to diagnose a specific disease based on the amounts and
identities of certain molecules in a patient’s exhaled breath. In Sangeeta Bhatia's lab at MIT, a team used
substances that react specifically inside diseased lung tissue; the products of the reactions will be present as
biomarkers that can be identified through mass spectrometry (an analytical method discussed later in the
chapter). A potential application would allow patients with early symptoms to inhale or ingest a "sensor"
62 2 • Atoms, Molecules, and Ions

substance, and, minutes later, to breathe into a detector for diagnosis. Similar research by scientists such as
Laura López-Sánchez has provided similar processes for lung cancer. An essential concept underlying this goal
is that of a molecule’s identity, which is determined by the numbers and types of atoms it contains, and how
they are bonded together. This chapter will describe some of the fundamental chemical principles related to
the composition of matter, including those central to the concept of molecular identity.

2.1 Early Ideas in Atomic Theory

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• State the postulates of Dalton’s atomic theory
• Use postulates of Dalton’s atomic theory to explain the laws of definite and multiple proportions

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the
scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed
of small, finite particles that they called atomos, a term derived from the Greek word for “indivisible.” They
thought of atoms as moving particles that differed in shape and size, and which could join together. Later,
Aristotle and others came to the conclusion that matter consisted of various combinations of the four
“elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers
thought about atoms and “elements” as philosophical concepts, but apparently never considered performing
experiments to test their ideas.

The Aristotelian view of the composition of matter held sway for over two thousand years, until English
schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter
could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the
microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s
atomic theory.

1. Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element
that can participate in a chemical change.
2. An element consists of only one type of atom, which has a mass that is characteristic of the element and is
the same for all atoms of that element (Figure 2.2). A macroscopic sample of an element contains an
incredibly large number of atoms, all of which have identical chemical properties.

FIGURE 2.2 A pre-1982 copper penny (left) contains approximately 3 1022 copper atoms (several dozen are
represented as brown spheres at the right), each of which has the same chemical properties. (credit:
modification of work by “slgckgc”/Flickr)

3. Atoms of one element differ in properties from atoms of all other elements.
4. A compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a
given compound, the numbers of atoms of each of its elements are always present in the same ratio
(Figure 2.3).

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 2.1 • Early Ideas in Atomic Theory 63

FIGURE 2.3 Copper(II) oxide, a powdery, black compound, results from the combination of two types of
atoms—copper (brown spheres) and oxygen (red spheres)—in a 1:1 ratio. (credit: modification of work by
“Chemicalinterest”/Wikimedia Commons)

5. Atoms are neither created nor destroyed during a chemical change, but are instead rearranged to yield
substances that are different from those present before the change (Figure 2.4).

FIGURE 2.4 When the elements copper (a shiny, red-brown solid, shown here as brown spheres) and oxygen
(a clear and colorless gas, shown here as red spheres) react, their atoms rearrange to form a compound
containing copper and oxygen (a powdery, black solid). (credit copper: modification of work by http://images-of-
elements.com/copper.php)

Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that
you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it
cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms.
And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter
present when matter changes from one type to another will remain constant (the law of conservation of
matter).

EXAMPLE 2.1

Testing Dalton’s Atomic Theory

In the following drawing, the green spheres represent atoms of a certain element. The purple spheres
represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does
the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory?
If so, which one?
64 2 • Atoms, Molecules, and Ions

Solution
The starting materials consist of two green spheres and two purple spheres. The products consist of only one
green sphere and one purple sphere. This violates Dalton’s postulate that atoms are neither created nor
destroyed during a chemical change, but are merely redistributed. (In this case, atoms appear to have been
destroyed.)

Check Your Learning

In the following drawing, the green spheres represent atoms of a certain element. The purple spheres
represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does
the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory?
If so, which one?

Answer:
The starting materials consist of four green spheres and two purple spheres. The products consist of four
green spheres and two purple spheres. This does not violate any of Dalton’s postulates: Atoms are neither
created nor destroyed, but are redistributed in small, whole-number ratios.

Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure
compound contain the same elements in the same proportion by mass. This statement is known as the law of
definite proportions or the law of constant composition. The suggestion that the numbers of atoms of the
elements in a given compound always exist in the same ratio is consistent with these observations. For
example, when different samples of isooctane (a component of gasoline and one of the standards used in the
octane rating system) are analyzed, they are found to have a carbon-to-hydrogen mass ratio of 5.33:1, as
shown in Table 2.1.

Constant Composition of Isooctane

Sample Carbon Hydrogen Mass Ratio

A 14.82 g 2.78 g

B 22.33 g 4.19 g

C 19.40 g 3.64 g

TABLE 2.1

It is worth noting that although all samples of a particular compound have the same mass ratio, the converse is
not true in general. That is, samples that have the same mass ratio are not necessarily the same substance. For
example, there are many compounds other than isooctane that also have a carbon-to-hydrogen mass ratio of
5.33:1.00.

Dalton also used data from Proust, as well as results from his own experiments, to formulate another
interesting law. The law of multiple proportions states that when two elements react to form more than one
compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole

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 2.1 • Early Ideas in Atomic Theory 65

numbers. For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g
chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper.
These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of
these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio.

This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as
the green compound.

This can be explained by atomic theory if the copper-to-chlorine ratio in the brown compound is 1 copper
atom to 2 chlorine atoms, and the ratio in the green compound is 1 copper atom to 1 chlorine atom. The ratio
of chlorine atoms (and thus the ratio of their masses) is therefore 2 to 1 (Figure 2.5).

FIGURE 2.5 Compared to the copper chlorine compound in (a), where copper is represented by brown spheres and
chlorine by green spheres, the copper chlorine compound in (b) has twice as many chlorine atoms per copper atom.
(credit a: modification of work by “Benjah-bmm27”/Wikimedia Commons; credit b: modification of work by
“Walkerma”/Wikimedia Commons)

EXAMPLE 2.2

Laws of Definite and Multiple Proportions

A sample of compound A (a clear, colorless gas) is analyzed and found to contain 4.27 g carbon and 5.69 g
oxygen. A sample of compound B (also a clear, colorless gas) is analyzed and found to contain 5.19 g carbon
and 13.84 g oxygen. Are these data an example of the law of definite proportions, the law of multiple
proportions, or neither? What do these data tell you about substances A and B?

Solution
In compound A, the mass ratio of oxygen to carbon is:

In compound B, the mass ratio of oxygen to carbon is:

The ratio of these ratios is:

66 2 • Atoms, Molecules, and Ions

This supports the law of multiple proportions. This means that A and B are different compounds, with A having
one-half as much oxygen per amount of carbon (or twice as much carbon per amount of oxygen) as B. A
possible pair of compounds that would fit this relationship would be A = CO and B = CO2.

Check Your Learning

A sample of compound X (a clear, colorless, combustible liquid with a noticeable odor) is analyzed and found to
contain 14.13 g carbon and 2.96 g hydrogen. A sample of compound Y (a clear, colorless, combustible liquid
with a noticeable odor that is slightly different from X’s odor) is analyzed and found to contain 19.91 g carbon
and 3.34 g hydrogen. Are these data an example of the law of definite proportions, the law of multiple
proportions, or neither? What do these data tell you about substances X and Y?

Answer:
In compound X, the mass ratio of carbon to hydrogen is In compound Y, the mass ratio of carbon to

hydrogen is The ratio of these ratios is This small, whole-

number ratio supports the law of multiple proportions. This means that X and Y are different compounds.

2.2 Evolution of Atomic Theory

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• Outline milestones in the development of modern atomic theory
• Summarize and interpret the results of the experiments of Thomson, Millikan, and Rutherford
• Describe the three subatomic particles that compose atoms
• Define isotopes and give examples for several elements

If matter is composed of atoms, what are atoms composed of? Are they the smallest particles, or is 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. Thomson using a cathode ray 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 2.6).

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 2.2 • Evolution of Atomic Theory 67

FIGURE 2.6 (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 “Kurzon”/Wikimedia Commons)

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’s
idea was gradually accepted, and his cathode ray particle is what we now call an electron, a negatively
charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term
“electron” was coined in 1891 by Irish physicist George Stoney, from “electric ion.”

LINK TO LEARNING
Click here (http://openstax.org/l/16JJThomson) to hear Thomson describe his discovery in his own voice.

In 1909, more information about the electron was uncovered by American physicist Robert A. Millikan via his
“oil drop” experiments. Millikan created microscopic oil droplets, which could be electrically charged by
friction as they formed or by using X-rays. These droplets initially fell due to gravity, but their downward
progress could be slowed or even reversed by an electric field lower in the apparatus. By adjusting the electric
field strength and making careful measurements and appropriate calculations, Millikan was able to determine
the charge on individual drops (Figure 2.7).
68 2 • Atoms, Molecules, and Ions

FIGURE 2.7 Millikan’s experiment measured the charge of individual oil drops. The tabulated data are examples of
a few possible values.

Looking at the charge data that Millikan gathered, you may have recognized that the charge of an oil droplet is
always a multiple of a specific charge, 1.6 10−19 C. Millikan concluded that this value must therefore be a
fundamental charge—the charge of a single electron—with his measured charges due to an excess of one
electron (1 times 1.6 10−19 C), two electrons (2 times 1.6 10−19 C), three electrons (3 times 1.6 10−19 C),
and so on, on a given oil droplet. Since the charge of an electron was now known due to Millikan’s research,
and the charge-to-mass ratio was already known due to Thomson’s research (1.759 1011 C/kg), it only
required a simple calculation to determine the mass of the electron as well.

Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of
Thomson, Millikan, and others, the charge and mass of the negative, subatomic particles—the electrons—were
known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson
proposed the “plum pudding” 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 Nagaoka, who postulated a Saturn-like atom,
consisting of a positively charged sphere surrounded by a halo of electrons (Figure 2.8).

FIGURE 2.8 (a) Thomson suggested that atoms resembled plum pudding, an English dessert consisting of moist
cake with embedded raisins (“plums”). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of

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 2.2 • Evolution of Atomic Theory 69

electrons surrounding a positive “planet.” (credit a: modification of work by “Man vyi”/Wikimedia Commons; credit
b: modification of work by “NASA”/Wikimedia Commons)

The next major development in understanding the atom came from Ernest Rutherford, 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 alpha particles (α particles) that were produced by the
radioactive decay of radium; α 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 Geiger (later
famous for the Geiger counter) and Ernest Marsden aimed a beam of α 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 α particles using a luminescent screen that glowed briefly where hit by an α particle.

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 2.9). Rutherford described finding these results: “It 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
1
and it came back and hit you.”

FIGURE 2.9 Geiger and Rutherford fired α 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.

Here is what Rutherford deduced: Because most of the fast-moving α 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 α 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:

1. The volume occupied by an atom must consist of a large amount of empty space.
2. A small, relatively heavy, positively charged body, the nucleus, must be at the center of each atom.

LINK TO LEARNING
View this simulation (http://openstax.org/l/16Rutherford) of the Rutherford gold foil experiment. Adjust the slit
width to produce a narrower or broader beam of α particles to see how that affects the scattering pattern.

1 Ernest Rutherford, “The Development of the Theory of Atomic Structure,” ed. J. A. Ratcliffe, in Background to Modern Science,
eds. Joseph Needham and Walter Pagel, (Cambridge, UK: Cambridge University Press, 1938), 61–74. Accessed September 22, 2014,
https://ia600508.us.archive.org/3/items/backgroundtomode032734mbp/backgroundtomode032734mbp.pdf.
70 2 • Atoms, Molecules, and Ions

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 2.10). After many more experiments, Rutherford also
discovered that the nuclei of other elements contain the hydrogen nucleus as a “building block,” and he named
this more fundamental particle the proton, 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.

FIGURE 2.10 The α 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 α
particles are deflected. Most pass through the relatively large region occupied by electrons, which are too light to
deflect the rapidly moving particles.

LINK TO LEARNING
The Rutherford Scattering simulation (http://openstax.org/l/16PhetScatter) allows you to investigate the
differences between a “plum pudding” atom and a Rutherford atom by firing α particles at each type of atom.

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 “new
element” 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 Soddy to realize that an element could have types of atoms with different masses
that were chemically indistinguishable. These different types are called isotopes—atoms of the same element
that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.

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 Chadwick found evidence of
neutrons, 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

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 2.3 • Atomic Structure and Symbolism 71

neutrons, but they are chemically identical because they have the same number of protons. This will be
explained in more detail later in this chapter.

2.3 Atomic Structure and Symbolism

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• 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

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 2.11).

FIGURE 2.11 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 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 of the
mass of one carbon-12 atom: 1 amu = 1.6605 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 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 2.2. (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.)
72 2 • Atoms, Molecules, and Ions

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

TABLE 2.2

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.

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. 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 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 2.3

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 2.12).

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 2.3 • Atomic Structure and Symbolism 73

FIGURE 2.12 (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 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].

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?

Answer:
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 2.13). 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).
74 2 • Atoms, Molecules, and Ions

FIGURE 2.13 The symbol Hg represents the element mercury regardless of the amount; it could represent one
atom of mercury or a large amount of mercury.

The symbols for several common elements and their atoms are listed in Table 2.3. 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. 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 in Figure 2.26 (also found in Appendix A).

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

TABLE 2.3

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 2.3 • Atomic Structure and Symbolism 75

Element Symbol Element Symbol

hydrogen H tin Sn (from stannum)

iodine I zinc Zn

TABLE 2.3

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 seaborgium (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active
in the discovery of several heavy elements. Element 109 was named in honor of Lise Meitner, who discovered
nuclear fission, a phenomenon that would have world-changing impacts; Meitner also contributed to the
discovery of some major isotopes, discussed immediately below.

LINK TO LEARNING
Visit this site (http://openstax.org/l/16IUPAC) to learn more about IUPAC, the International Union of Pure and
Applied Chemistry, and explore its periodic table.

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 2.14). 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 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.

FIGURE 2.14 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 2.4. 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.
76 2 • Atoms, Molecules, and Ions

Nuclear Compositions of Atoms of the Very Light Elements

Atomic Number of Number of Mass % Natural

Element Symbol
Number Protons Neutrons (amu) Abundance

1 1 0 1.0078 99.989
(protium)

hydrogen 1 1 1 2.0141 0.0115

(deuterium)

1 1 2 3.01605 — (trace)
(tritium)

2 2 1 3.01603 0.00013
helium
2 2 2 4.0026 100

3 3 3 6.0151 7.59
lithium
3 3 4 7.0160 92.41

beryllium 4 4 5 9.0122 100

5 5 5 10.0129 19.9
boron
5 5 6 11.0093 80.1

6 6 6 12.0000 98.89

carbon 6 6 7 13.0034 1.11

6 6 8 14.0032 — (trace)

7 7 7 14.0031 99.63
nitrogen
7 7 8 15.0001 0.37

8 8 8 15.9949 99.757

oxygen 8 8 9 16.9991 0.038

8 8 10 17.9992 0.205

fluorine 9 9 10 18.9984 100

neon 10 10 10 19.9924 90.48

TABLE 2.4

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 2.3 • Atomic Structure and Symbolism 77

Atomic Number of Number of Mass % Natural

Element Symbol
Number Protons Neutrons (amu) Abundance

10 10 11 20.9938 0.27

10 10 12 21.9914 9.25

TABLE 2.4

LINK TO LEARNING
Use this Build an Atom simulator (http://openstax.org/l/16PhetAtomBld) to build atoms of the first 10
elements, see which isotopes exist, check nuclear stability, and gain experience with isotope symbols.

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 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.

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:

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 2.4

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

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
78 2 • Atoms, Molecules, and Ions

isotopes, depending on their origin.)

Check Your Learning

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 2.5

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.

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:

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.

Check Your Learning

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

LINK TO LEARNING
Visit this site (http://openstax.org/l/16PhetAtomMass) 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.

As you will learn, isotopes are important in nature and especially in human understanding of science and
medicine. Let's consider just one natural, stable isotope: Oxygen-18, which is noted in the table above and is
referred to as one of the environmental isotopes. It is important in paleoclimatology, for example, because
scientists can use the ratio between Oxygen-18 and Oxygen-16 in an ice core to determine the temperature of

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 2.4 • Chemical Formulas 79

precipitation over time. Oxygen-18 was also critical to the discovery of metabolic pathways and the
mechanisms of enzymes. Mildred Cohn pioneered the usage of these isotopes to act as tracers, so that
researchers could follow their path through reactions and gain a better understanding of what is happening.
One of her first discoveries provided insight into the phosphorylation of glucose that takes place in
mitochondria. And the methods of using isotopes for this research contributed to entire fields of study.

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 2.15), 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
rolling past a magnet is deflected to a lesser extent that that of a small steel ball). 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.

FIGURE 2.15 Analysis of zirconium in a mass spectrometer produces a mass spectrum with peaks showing the
different isotopes of Zr.

LINK TO LEARNING
See an animation (http://openstax.org/l/16MassSpec) that explains mass spectrometry. Watch this video
(http://openstax.org/l/16RSChemistry) from the Royal Society for Chemistry for a brief description of the
rudiments of mass spectrometry.

2.4 Chemical Formulas

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• Symbolize the composition of molecules using molecular formulas and empirical formulas
• Represent the bonding arrangement of atoms within molecules using structural formulas

A molecular formula is a representation of a molecule that uses chemical symbols to indicate the types of
atoms followed by subscripts to show the number of atoms of each type in the molecule. (A subscript is used
only when more than one atom of a given type is present.) Molecular formulas are also used as abbreviations
80 2 • Atoms, Molecules, and Ions

for the names of compounds.

The structural formula for a compound gives the same information as its molecular formula (the types and
numbers of atoms in the molecule) but also shows how the atoms are connected in the molecule. The
structural formula for methane contains symbols for one C atom and four H atoms, indicating the number of
atoms in the molecule (Figure 2.16). The lines represent bonds that hold the atoms together. (A chemical bond
is an attraction between atoms or ions that holds them together in a molecule or a crystal.) We will discuss
chemical bonds and see how to predict the arrangement of atoms in a molecule later. For now, simply know
that the lines are an indication of how the atoms are connected in a molecule. A ball-and-stick model shows
the geometric arrangement of the atoms with atomic sizes not to scale, and a space-filling model shows the
relative sizes of the atoms.

FIGURE 2.16 A methane molecule can be represented as (a) a molecular formula, (b) a structural formula, (c) a
ball-and-stick model, and (d) a space-filling model. Carbon and hydrogen atoms are represented by black and white
spheres, respectively.

Although many elements consist of discrete, individual atoms, some exist as molecules made up of two or
more atoms of the element chemically bonded together. For example, most samples of the elements hydrogen,
oxygen, and nitrogen are composed of molecules that contain two atoms each (called diatomic molecules) and
thus have the molecular formulas H2, O2, and N2, respectively. Other elements commonly found as diatomic
molecules are fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2). The most common form of the
element sulfur is composed of molecules that consist of eight atoms of sulfur; its molecular formula is S8
(Figure 2.17).

FIGURE 2.17 A molecule of sulfur is composed of eight sulfur atoms and is therefore written as S8. It can be
represented as (a) a structural formula, (b) a ball-and-stick model, and (c) a space-filling model. Sulfur atoms are
represented by yellow spheres.

It is important to note that a subscript following a symbol and a number in front of a symbol do not represent
the same thing; for example, H2 and 2H represent distinctly different species. H2 is a molecular formula; it
represents a diatomic molecule of hydrogen, consisting of two atoms of the element that are chemically
bonded together. The expression 2H, on the other hand, indicates two separate hydrogen atoms that are not
combined as a unit. The expression 2H2 represents two molecules of diatomic hydrogen (Figure 2.18).

FIGURE 2.18 The symbols H, 2H, H2, and 2H2 represent very different entities.

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 2.4 • Chemical Formulas 81

Compounds are formed when two or more elements chemically combine, resulting in the formation of bonds.
For example, hydrogen and oxygen can react to form water, and sodium and chlorine can react to form table
salt. We sometimes describe the composition of these compounds with an empirical formula, which indicates
the types of atoms present and the simplest whole-number ratio of the number of atoms (or ions) in the
compound. For example, titanium dioxide (used as pigment in white paint and in the thick, white, blocking
type of sunscreen) has an empirical formula of TiO2. This identifies the elements titanium (Ti) and oxygen (O)
as the constituents of titanium dioxide, and indicates the presence of twice as many atoms of the element
oxygen as atoms of the element titanium (Figure 2.19).

FIGURE 2.19 (a) The white compound titanium dioxide provides effective protection from the sun. (b) A crystal of
titanium dioxide, TiO2, contains titanium and oxygen in a ratio of 1 to 2. The titanium atoms are gray and the oxygen
atoms are red. (credit a: modification of work by “osseous”/Flickr)

As discussed previously, we can describe a compound with a molecular formula, in which the subscripts
indicate the actual numbers of atoms of each element in a molecule of the compound. In many cases, the
molecular formula of a substance is derived from experimental determination of both its empirical formula
and its molecular mass (the sum of atomic masses for all atoms composing the molecule). For example, it can
be determined experimentally that benzene contains two elements, carbon (C) and hydrogen (H), and that for
every carbon atom in benzene, there is one hydrogen atom. Thus, the empirical formula is CH. An
experimental determination of the molecular mass reveals that a molecule of benzene contains six carbon
atoms and six hydrogen atoms, so the molecular formula for benzene is C6H6 (Figure 2.20).

FIGURE 2.20 Benzene, C6H6, is produced during oil refining and has many industrial uses. A benzene molecule can
be represented as (a) a structural formula, (b) a ball-and-stick model, and (c) a space-filling model. (d) Benzene is a
clear liquid. (credit d: modification of work by Sahar Atwa)

If we know a compound’s formula, we can easily determine the empirical formula. (This is somewhat of an
academic exercise; the reverse chronology is generally followed in actual practice.) For example, the molecular
formula for acetic acid, the component that gives vinegar its sharp taste, is C2H4O2. This formula indicates that
a molecule of acetic acid (Figure 2.21) contains two carbon atoms, four hydrogen atoms, and two oxygen atoms.
The ratio of atoms is 2:4:2. Dividing by the lowest common denominator (2) gives the simplest, whole-number
ratio of atoms, 1:2:1, so the empirical formula is CH2O. Note that a molecular formula is always a whole-
number multiple of an empirical formula.
82 2 • Atoms, Molecules, and Ions

FIGURE 2.21 (a) Vinegar contains acetic acid, C2H4O2, which has an empirical formula of CH2O. It can be
represented as (b) a structural formula and (c) as a ball-and-stick model. (credit a: modification of work by
“HomeSpot HQ”/Flickr)

EXAMPLE 2.6

Empirical and Molecular Formulas

Molecules of glucose (blood sugar) contain 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. What are
the molecular and empirical formulas of glucose?

Solution
The molecular formula is C6H12O6 because one molecule actually contains 6 C, 12 H, and 6 O atoms. The
simplest whole-number ratio of C to H to O atoms in glucose is 1:2:1, so the empirical formula is CH2O.

Check Your Learning

A molecule of metaldehyde (a pesticide used for snails and slugs) contains 8 carbon atoms, 16 hydrogen atoms,
and 4 oxygen atoms. What are the molecular and empirical formulas of metaldehyde?

Answer:
Molecular formula, C8H16O4; empirical formula, C2H4O

LINK TO LEARNING
You can explore molecule building (http://openstax.org/l/16molbuilding) using an online simulation.

Portrait of a Chemist
Lee Cronin
What is it that chemists do? According to Lee Cronin (Figure 2.22), chemists make very complicated
molecules by “chopping up” small molecules and “reverse engineering” them. He wonders if we could
“make a really cool universal chemistry set” by what he calls “app-ing” chemistry. Could we “app”
chemistry?

In a 2012 TED talk, Lee describes one fascinating possibility: combining a collection of chemical “inks”
with a 3D printer capable of fabricating a reaction apparatus (tiny test tubes, beakers, and the like) to
fashion a “universal toolkit of chemistry.” This toolkit could be used to create custom-tailored drugs to fight
a new superbug or to “print” medicine personally configured to your genetic makeup, environment, and
health situation. Says Cronin, “What Apple did for music, I’d like to do for the discovery and distribution of
2
prescription drugs.” View his full talk (http://openstax.org/l/16LeeCronin) at the TED website.

2 Lee Cronin, “Print Your Own Medicine,” Talk presented at TED Global 2012, Edinburgh, Scotland, June 2012.

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 2.4 • Chemical Formulas 83

FIGURE 2.22 Chemist Lee Cronin has been named one of the UK’s 10 most inspirational scientists. The
youngest chair at the University of Glasgow, Lee runs a large research group, collaborates with many scientists
worldwide, has published over 250 papers in top scientific journals, and has given more than 150 invited talks.
His research focuses on complex chemical systems and their potential to transform technology, but also
branches into nanoscience, solar fuels, synthetic biology, and even artificial life and evolution. (credit: image
courtesy of Lee Cronin)

It is important to be aware that it may be possible for the same atoms to be arranged in different ways:
Compounds with the same molecular formula may have different atom-to-atom bonding and therefore
different structures. For example, could there be another compound with the same formula as acetic acid,
C2H4O2? And if so, what would be the structure of its molecules?

If you predict that another compound with the formula C2H4O2 could exist, then you demonstrated good
chemical insight and are correct. Two C atoms, four H atoms, and two O atoms can also be arranged to form a
methyl formate, which is used in manufacturing, as an insecticide, and for quick-drying finishes. Methyl
formate molecules have one of the oxygen atoms between the two carbon atoms, differing from the
arrangement in acetic acid molecules. Acetic acid and methyl formate are examples of isomers—compounds
with the same chemical formula but different molecular structures (Figure 2.23). Note that this small
difference in the arrangement of the atoms has a major effect on their respective chemical properties. You
would certainly not want to use a solution of methyl formate as a substitute for a solution of acetic acid
(vinegar) when you make salad dressing.

FIGURE 2.23 Molecules of (a) acetic acid and methyl formate (b) are structural isomers; they have the same
formula (C2H4O2) but different structures (and therefore different chemical properties).

Many types of isomers exist (Figure 2.24). Acetic acid and methyl formate are structural isomers, compounds
84 2 • Atoms, Molecules, and Ions

in which the molecules differ in how the atoms are connected to each other. There are also various types of
spatial isomers, in which the relative orientations of the atoms in space can be different. For example, the
compound carvone (found in caraway seeds, spearmint, and mandarin orange peels) consists of two isomers
that are mirror images of each other. S-(+)-carvone smells like caraway, and R-(−)-carvone smells like
spearmint.

FIGURE 2.24 Molecules of carvone are spatial isomers; they only differ in the relative orientations of the atoms in
space. (credit bottom left: modification of work by “Miansari66”/Wikimedia Commons; credit bottom right:
modification of work by Forest & Kim Starr)

LINK TO LEARNING
Select this link (http://openstax.org/l/16isomers) to view an explanation of isomers, spatial isomers, and why
they have different smells (select the video titled “Mirror Molecule: Carvone”).

2.5 The Periodic Table

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• State the periodic law and explain the organization of elements in the periodic table
• Predict the general properties of elements based on their location within the periodic table
• Identify metals, nonmetals, and metalloids by their properties and/or location on the periodic table

As early chemists worked to purify ores and discovered more elements, they realized that various elements
could be grouped together by their similar chemical behaviors. One such grouping includes lithium (Li),
sodium (Na), and potassium (K): These elements all are shiny, conduct heat and electricity well, and have
similar chemical properties. A second grouping includes calcium (Ca), strontium (Sr), and barium (Ba), which
also are shiny, good conductors of heat and electricity, and have chemical properties in common. However, the
specific properties of these two groupings are notably different from each other. For example: Li, Na, and K are
much more reactive than are Ca, Sr, and Ba; Li, Na, and K form compounds with oxygen in a ratio of two of their
atoms to one oxygen atom, whereas Ca, Sr, and Ba form compounds with one of their atoms to one oxygen
atom. Fluorine (F), chlorine (Cl), bromine (Br), and iodine (I) also exhibit similar properties to each other, but
these properties are drastically different from those of any of the elements above.

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 2.5 • The Periodic Table 85

Dimitri Mendeleev in Russia (1869) and Lothar Meyer in Germany (1870) independently recognized that there
was a periodic relationship among the properties of the elements known at that time. Both published tables
with the elements arranged according to increasing atomic mass. But Mendeleev went one step further than
Meyer: He used his table to predict the existence of elements that would have the properties similar to
aluminum and silicon, but were yet unknown. The discoveries of gallium (1875) and germanium (1886)
provided great support for Mendeleev’s work. Although Mendeleev and Meyer had a long dispute over priority,
Mendeleev’s contributions to the development of the periodic table are now more widely recognized (Figure
2.25).

FIGURE 2.25 (a) Dimitri Mendeleev is widely credited with creating (b) the first periodic table of the elements.
(credit a: modification of work by Serge Lachinov; credit b: modification of work by “Den fjättrade ankan”/Wikimedia
Commons)

By the twentieth century, it became apparent that the periodic relationship involved atomic numbers rather
than atomic masses. The modern statement of this relationship, the periodic law, is as follows: the properties
of the elements are periodic functions of their atomic numbers. A modern periodic table arranges the
elements in increasing order of their atomic numbers and groups atoms with similar properties in the same
vertical column (Figure 2.26). Each box represents an element and contains its atomic number, symbol,
average atomic mass, and (sometimes) name. The elements are arranged in seven horizontal rows, called
periods or series, and 18 vertical columns, called groups. Groups are labeled at the top of each column. In the
United States, the labels traditionally were numerals with capital letters. However, IUPAC recommends that the
numbers 1 through 18 be used, and these labels are more common. For the table to fit on a single page, parts of
two of the rows, a total of 14 columns, are usually written below the main body of the table.
86 2 • Atoms, Molecules, and Ions

FIGURE 2.26 Elements in the periodic table are organized according to their properties.

Even after the periodic nature of elements and the table itself were widely accepted, gaps remained.
Mendeleev had predicted, and others including Henry Moseley had later confirmed, that there should be
elements below Manganese in Group 7. German chemists Ida Tacke and Walter Noddack set out to find the
elements, a quest being pursued by scientists around the world. Their method was unique in that they did not
only consider the properties of manganese, but also the elements horizontally adjacent to the missing
elements 43 and 75 on the table. Thus, by investigating ores containing minerals of ruthenium (Ru), tungsten
(W), osmium (Os), and so on, they were able to identify naturally occurring elements that helped complete the
table. Rhenium, one of their discoveries, was one of the last natural elements to be discovered and is the last
stable element to be discovered. (Francium, the last natural element to be discovered, was identified by
Marguerite Perey in 1939.)

Many elements differ dramatically in their chemical and physical properties, but some elements are similar in
their behaviors. For example, many elements appear shiny, are malleable (able to be deformed without
breaking) and ductile (can be drawn into wires), and conduct heat and electricity well. Other elements are not
shiny, malleable, or ductile, and are poor conductors of heat and electricity. We can sort the elements into large
classes with common properties: metals (elements that are shiny, malleable, good conductors of heat and
electricity—shaded yellow); nonmetals (elements that appear dull, poor conductors of heat and
electricity—shaded green); and metalloids (elements that conduct heat and electricity moderately well, and
possess some properties of metals and some properties of nonmetals—shaded purple).

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 2.5 • The Periodic Table 87

The elements can also be classified into the main-group elements (or representative elements) in the
3
columns labeled 1, 2, and 13–18; the transition metals in the columns labeled 3–12 ; and inner transition
metals in the two rows at the bottom of the table (the top-row elements are called lanthanides and the bottom-
row elements are actinides; Figure 2.27). The elements can be subdivided further by more specific properties,
such as the composition of the compounds they form. For example, the elements in group 1 (the first column)
form compounds that consist of one atom of the element and one atom of hydrogen. These elements (except
hydrogen) are known as alkali metals, and they all have similar chemical properties. The elements in group 2
(the second column) form compounds consisting of one atom of the element and two atoms of hydrogen: These
are called alkaline earth metals, with similar properties among members of that group. Other groups with
specific names are the pnictogens (group 15), chalcogens (group 16), halogens (group 17), and the noble
gases (group 18, also known as inert gases). The groups can also be referred to by the first element of the
group: For example, the chalcogens can be called the oxygen group or oxygen family. Hydrogen is a unique,
nonmetallic element with properties similar to both group 1 and group 17 elements. For that reason, hydrogen
may be shown at the top of both groups, or by itself.

FIGURE 2.27 The periodic table organizes elements with similar properties into groups.

LINK TO LEARNING
Click on this link (https://openstax.org/l/16Periodic) for an interactive periodic table, which you can use to
explore the properties of the elements (includes podcasts and videos of each element). You may also want to
try this one (https://openstax.org/l/16Periodic2) that shows photos of all the elements.

EXAMPLE 2.7

Naming Groups of Elements

Atoms of each of the following elements are essential for life. Give the group name for the following elements:

(a) chlorine

3 Per the IUPAC definition, group 12 elements are not transition metals, though they are often referred to as such. Additional details
on this group’s elements are provided in a chapter on transition metals and coordination chemistry.
88 2 • Atoms, Molecules, and Ions

(b) calcium

(c) sodium

(d) sulfur

Solution
The family names are as follows:

(a) halogen

(b) alkaline earth metal

(c) alkali metal

(d) chalcogen

Check Your Learning

Give the group name for each of the following elements:

(a) krypton

(b) selenium

(c) barium

(d) lithium

Answer:
(a) noble gas; (b) chalcogen; (c) alkaline earth metal; (d) alkali metal

As you will learn in your further study of chemistry, elements in groups often behave in a somewhat similar
manner. This is partly due to the number of electrons in their outer shell and their similar readiness to bond.
These shared properties can have far-ranging implications in nature, science, and medicine. For example,
when Gertrude Elion and George Hitchens were investigating ways to interrupt cell and virus replication to
fight diseases, they utilized the similarity between sulfur and oxygen (both in Group 16) and their capacity to
bond in similar ways. Elion focused on purines, which are key components of DNA and which contain oxygen.
She found that by introducing sulfur-based compounds (called purine analogues) that mimic the structure of
purines, molecules within DNA would bond to the analogues rather than the "regular" DNA purine. With the
normal DNA bonding and structure altered, Elion successfully interrupted cell replication. At its core, the
strategy worked because of the similarity between sulfur and oxygen. Her discovery led directly to important
treatments for leukemia. Overall, Elion's work with George Hitchens not only led to more treatments, but also
changed the entire methodology of drug development. By using specific elements and compounds to target
specific aspects of tumor cells, viruses, and bacteria, they laid the groundwork for many of today's most
common and important medicines, used to help millions of people each year. They were awarded the Nobel
Prize in 1988.

In studying the periodic table, you might have noticed something about the atomic masses of some of the
elements. Element 43 (technetium), element 61 (promethium), and most of the elements with atomic number
84 (polonium) and higher have their atomic mass given in square brackets. This is done for elements that
consist entirely of unstable, radioactive isotopes (you will learn more about radioactivity in the nuclear
chemistry chapter). An average atomic weight cannot be determined for these elements because their
radioisotopes may vary significantly in relative abundance, depending on the source, or may not even exist in
nature. The number in square brackets is the atomic mass number (an approximate atomic mass) of the most
stable isotope of that element.

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 2.6 • Ionic and Molecular Compounds 89

2.6 Ionic and Molecular Compounds

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• Define ionic and molecular (covalent) compounds
• Predict the type of compound formed from elements based on their location within the periodic table
• Determine formulas for simple ionic compounds

In ordinary chemical reactions, the nucleus of each atom (and thus the identity of the element) remains
unchanged. Electrons, however, can be added to atoms by transfer from other atoms, lost by transfer to other
atoms, or shared with other atoms. The transfer and sharing of electrons among atoms govern the chemistry of
the elements. During the formation of some compounds, atoms gain or lose electrons, and form electrically
charged particles called ions (Figure 2.28).

FIGURE 2.28 (a) A sodium atom (Na) has equal numbers of protons and electrons (11) and is uncharged. (b) A
sodium cation (Na+) has lost an electron, so it has one more proton (11) than electrons (10), giving it an overall
positive charge, signified by a superscripted plus sign.

You can use the periodic table to predict whether an atom will form an anion or a cation, and you can often
predict the charge of the resulting ion. Atoms of many main-group metals lose enough electrons to leave them
with the same number of electrons as an atom of the preceding noble gas. To illustrate, an atom of an alkali
metal (group 1) loses one electron and forms a cation with a 1+ charge; an alkaline earth metal (group 2) loses
two electrons and forms a cation with a 2+ charge, and so on. For example, a neutral calcium atom, with 20
protons and 20 electrons, readily loses two electrons. This results in a cation with 20 protons, 18 electrons, and
a 2+ charge. It has the same number of electrons as atoms of the preceding noble gas, argon, and is symbolized
Ca2+. The name of a metal ion is the same as the name of the metal atom from which it forms, so Ca2+ is called
a calcium ion.

When atoms of nonmetal elements form ions, they generally gain enough electrons to give them the same
number of electrons as an atom of the next noble gas in the periodic table. Atoms of group 17 gain one electron
and form anions with a 1− charge; atoms of group 16 gain two electrons and form ions with a 2− charge, and so
on. For example, the neutral bromine atom, with 35 protons and 35 electrons, can gain one electron to provide
it with 36 electrons. This results in an anion with 35 protons, 36 electrons, and a 1− charge. It has the same
number of electrons as atoms of the next noble gas, krypton, and is symbolized Br−. (A discussion of the theory
supporting the favored status of noble gas electron numbers reflected in these predictive rules for ion
formation is provided in a later chapter of this text.)

Note the usefulness of the periodic table in predicting likely ion formation and charge (Figure 2.29). Moving
from the far left to the right on the periodic table, main-group elements tend to form cations with a charge
equal to the group number. That is, group 1 elements form 1+ ions; group 2 elements form 2+ ions, and so on.
Moving from the far right to the left on the periodic table, elements often form anions with a negative charge
90 2 • Atoms, Molecules, and Ions

equal to the number of groups moved left from the noble gases. For example, group 17 elements (one group
left of the noble gases) form 1− ions; group 16 elements (two groups left) form 2− ions, and so on. This trend
can be used as a guide in many cases, but its predictive value decreases when moving toward the center of the
periodic table. In fact, transition metals and some other metals often exhibit variable charges that are not
predictable by their location in the table. For example, copper can form ions with a 1+ or 2+ charge, and iron
can form ions with a 2+ or 3+ charge.

FIGURE 2.29 Some elements exhibit a regular pattern of ionic charge when they form ions.

EXAMPLE 2.8

Composition of Ions
An ion found in some compounds used as antiperspirants contains 13 protons and 10 electrons. What is its
symbol?

Solution
Because the number of protons remains unchanged when an atom forms an ion, the atomic number of the
element must be 13. Knowing this lets us use the periodic table to identify the element as Al (aluminum). The
Al atom has lost three electrons and thus has three more positive charges (13) than it has electrons (10). This is
the aluminum cation, Al3+.

Check Your Learning

Give the symbol and name for the ion with 34 protons and 36 electrons.

Answer:
Se2−, the selenide ion

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 2.6 • Ionic and Molecular Compounds 91

EXAMPLE 2.9

Formation of Ions
Magnesium and nitrogen react to form an ionic compound. Predict which forms an anion, which forms a
cation, and the charges of each ion. Write the symbol for each ion and name them.

Solution
Magnesium’s position in the periodic table (group 2) tells us that it is a metal. Metals form positive ions
(cations). A magnesium atom must lose two electrons to have the same number electrons as an atom of the
previous noble gas, neon. Thus, a magnesium atom will form a cation with two fewer electrons than protons
and a charge of 2+. The symbol for the ion is Mg2+, and it is called a magnesium ion.

Nitrogen’s position in the periodic table (group 15) reveals that it is a nonmetal. Nonmetals form negative ions
(anions). A nitrogen atom must gain three electrons to have the same number of electrons as an atom of the
following noble gas, neon. Thus, a nitrogen atom will form an anion with three more electrons than protons
and a charge of 3−. The symbol for the ion is N3−, and it is called a nitride ion.

Check Your Learning

Aluminum and carbon react to form an ionic compound. Predict which forms an anion, which forms a cation,
and the charges of each ion. Write the symbol for each ion and name them.

Answer:
Al will form a cation with a charge of 3+: Al3+, an aluminum ion. Carbon will form an anion with a charge of 4−:
C4−, a carbide ion.

The ions that we have discussed so far are called monatomic ions, that is, they are ions formed from only one
atom. We also find many polyatomic ions. These ions, which act as discrete units, are electrically charged
molecules (a group of bonded atoms with an overall charge). Some of the more important polyatomic ions are
listed in Table 2.5. Oxyanions are polyatomic ions that contain one or more oxygen atoms. At this point in your
study of chemistry, you should memorize the names, formulas, and charges of the most common polyatomic
ions. Because you will use them repeatedly, they will soon become familiar.

Common Polyatomic Ions

Name Formula Related Acid Formula

ammonium

hydronium

peroxide

hydroxide

acetate acetic acid CH3COOH

cyanide CN− hydrocyanic acid HCN

azide hydrazoic acid HN3

carbonate carbonic acid H2CO3

TABLE 2.5
92 2 • Atoms, Molecules, and Ions

Name Formula Related Acid Formula

bicarbonate

nitrate nitric acid HNO3

nitrite nitrous acid HNO2

sulfate sulfuric acid H2SO4

hydrogen sulfate

sulfite sulfurous acid H2SO3

hydrogen sulfite

phosphate phosphoric acid H3PO4

hydrogen phosphate

dihydrogen phosphate

perchlorate perchloric acid HClO4

chlorate chloric acid HClO3

chlorite chlorous acid HClO2

hypochlorite ClO− hypochlorous acid HClO

chromate chromic acid H2CrO4

dichromate dichromic acid H2Cr2O7

permanganate permanganic acid HMnO4

TABLE 2.5

Note that there is a system for naming some polyatomic ions; -ate and -ite are suffixes designating polyatomic
ions containing more or fewer oxygen atoms. Per- (short for “hyper”) and hypo- (meaning “under”) are
prefixes meaning more oxygen atoms than -ate and fewer oxygen atoms than -ite, respectively. For example,
perchlorate is chlorate is chlorite is and hypochlorite is ClO−. Unfortunately, the
number of oxygen atoms corresponding to a given suffix or prefix is not consistent; for example, nitrate is
while sulfate is This will be covered in more detail in the next module on nomenclature.

The nature of the attractive forces that hold atoms or ions together within a compound is the basis for
classifying chemical bonding. When electrons are transferred and ions form, ionic bonds result. Ionic bonds
are electrostatic forces of attraction, that is, the attractive forces experienced between objects of opposite
electrical charge (in this case, cations and anions). When electrons are “shared” and molecules form, covalent
bonds result. Covalent bonds are the attractive forces between the positively charged nuclei of the bonded
atoms and one or more pairs of electrons that are located between the atoms. Compounds are classified as
ionic or molecular (covalent) on the basis of the bonds present in them.

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 2.6 • Ionic and Molecular Compounds 93

Ionic Compounds
When an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of
atoms that readily gain electrons (a nonmetal), a transfer of electrons usually occurs, producing ions. The
compound formed by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions
of opposite charge present in the compound. For example, when each sodium atom in a sample of sodium
metal (group 1) gives up one electron to form a sodium cation, Na+, and each chlorine atom in a sample of
chlorine gas (group 17) accepts one electron to form a chloride anion, Cl−, the resulting compound, NaCl, is
composed of sodium ions and chloride ions in the ratio of one Na+ ion for each Cl− ion. Similarly, each calcium
atom (group 2) can give up two electrons and transfer one to each of two chlorine atoms to form CaCl2, which is
composed of Ca2+ and Cl− ions in the ratio of one Ca2+ ion to two Cl− ions.

A compound that contains ions and is held together by ionic bonds is called an ionic compound. The periodic
table can help us recognize many of the compounds that are ionic: When a metal is combined with one or more
nonmetals, the compound is usually ionic. This guideline works well for predicting ionic compound formation
for most of the compounds typically encountered in an introductory chemistry course. However, it is not
always true (for example, aluminum chloride, AlCl3, is not ionic).

You can often recognize ionic compounds because of their properties. Ionic compounds are solids that
typically melt at high temperatures and boil at even higher temperatures. For example, sodium chloride melts
at 801 °C and boils at 1413 °C. (As a comparison, the molecular compound water melts at 0 °C and boils at 100
°C.) In solid form, an ionic compound is not electrically conductive because its ions are unable to flow
(“electricity” is the flow of charged particles). When molten, however, it can conduct electricity because its ions
are able to move freely through the liquid (Figure 2.30).

FIGURE 2.30 Sodium chloride melts at 801 °C and conducts electricity when molten. (credit: modification of work
by Mark Blaser and Matt Evans)

LINK TO LEARNING
Watch this video (http://openstax.org/l/16moltensalt) to see a mixture of salts melt and conduct electricity.

In every ionic compound, the total number of positive charges of the cations equals the total number of
negative charges of the anions. Thus, ionic compounds are electrically neutral overall, even though they
contain positive and negative ions. We can use this observation to help us write the formula of an ionic
compound. The formula of an ionic compound must have a ratio of ions such that the numbers of positive and
negative charges are equal.

EXAMPLE 2.10

Predicting the Formula of an Ionic Compound

The gemstone sapphire (Figure 2.31) is mostly a compound of aluminum and oxygen that contains aluminum
94 2 • Atoms, Molecules, and Ions

cations, Al3+, and oxygen anions, O2−. What is the formula of this compound?

FIGURE 2.31 Although pure aluminum oxide is colorless, trace amounts of iron and titanium give blue sapphire its
characteristic color. (credit: modification of work by Stanislav Doronenko)

Solution
Because the ionic compound must be electrically neutral, it must have the same number of positive and
negative charges. Two aluminum ions, each with a charge of 3+, would give us six positive charges, and three
oxide ions, each with a charge of 2−, would give us six negative charges. The formula would be Al2O3.

Check Your Learning

Predict the formula of the ionic compound formed between the sodium cation, Na+, and the sulfide anion, S2−.

Answer:
Na2S

Many ionic compounds contain polyatomic ions (Table 2.5) as the cation, the anion, or both. As with simple
ionic compounds, these compounds must also be electrically neutral, so their formulas can be predicted by
treating the polyatomic ions as discrete units. We use parentheses in a formula to indicate a group of atoms
that behave as a unit. For example, the formula for calcium phosphate, one of the minerals in our bones, is
Ca3(PO4)2. This formula indicates that there are three calcium ions (Ca2+) for every two phosphate
groups. The groups are discrete units, each consisting of one phosphorus atom and four oxygen atoms,
and having an overall charge of 3−. The compound is electrically neutral, and its formula shows a total count of
three Ca, two P, and eight O atoms.

EXAMPLE 2.11

Predicting the Formula of a Compound with a Polyatomic Anion

Baking powder contains calcium dihydrogen phosphate, an ionic compound composed of the ions Ca2+ and
What is the formula of this compound?

Solution
The positive and negative charges must balance, and this ionic compound must be electrically neutral. Thus,
we must have two negative charges to balance the 2+ charge of the calcium ion. This requires a ratio of one
Ca2+ ion to two ions. We designate this by enclosing the formula for the dihydrogen phosphate ion in
parentheses and adding a subscript 2. The formula is Ca(H2PO4)2.

Check Your Learning

Predict the formula of the ionic compound formed between the lithium ion and the peroxide ion, (Hint:
Use the periodic table to predict the sign and the charge on the lithium ion.)

Answer:
Li2O2

Because an ionic compound is not made up of single, discrete molecules, it may not be properly symbolized
using a molecular formula. Instead, ionic compounds must be symbolized by a formula indicating the relative

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 2.6 • Ionic and Molecular Compounds 95

numbers of its constituent ions. For compounds containing only monatomic ions (such as NaCl) and for many
compounds containing polyatomic ions (such as CaSO4), these formulas are just the empirical formulas
introduced earlier in this chapter. However, the formulas for some ionic compounds containing polyatomic
ions are not empirical formulas. For example, the ionic compound sodium oxalate is comprised of Na+ and
ions combined in a 2:1 ratio, and its formula is written as Na2C2O4. The subscripts in this formula are
not the smallest-possible whole numbers, as each can be divided by 2 to yield the empirical formula, NaCO2.
This is not the accepted formula for sodium oxalate, however, as it does not accurately represent the
compound’s polyatomic anion,

Molecular Compounds
Many compounds do not contain ions but instead consist solely of discrete, neutral molecules. These
molecular compounds (covalent compounds) result when atoms share, rather than transfer (gain or lose),
electrons. Covalent bonding is an important and extensive concept in chemistry, and it will be treated in
considerable detail in a later chapter of this text. We can often identify molecular compounds on the basis of
their physical properties. Under normal conditions, molecular compounds often exist as gases, low-boiling
liquids, and low-melting solids, although many important exceptions exist.

Whereas ionic compounds are usually formed when a metal and a nonmetal combine, covalent compounds
are usually formed by a combination of nonmetals. Thus, the periodic table can help us recognize many of the
compounds that are covalent. While we can use the positions of a compound’s elements in the periodic table to
predict whether it is ionic or covalent at this point in our study of chemistry, you should be aware that this is a
very simplistic approach that does not account for a number of interesting exceptions. Shades of gray exist
between ionic and molecular compounds, and you’ll learn more about those later.

EXAMPLE 2.12

Predicting the Type of Bonding in Compounds

Predict whether the following compounds are ionic or molecular:

(a) KI, the compound used as a source of iodine in table salt

(b) H2O2, the bleach and disinfectant hydrogen peroxide

(c) CHCl3, the anesthetic chloroform

(d) Li2CO3, a source of lithium in antidepressants

Solution
(a) Potassium (group 1) is a metal, and iodine (group 17) is a nonmetal; KI is predicted to be ionic.

(b) Hydrogen (group 1) is a nonmetal, and oxygen (group 16) is a nonmetal; H2O2 is predicted to be molecular.

(c) Carbon (group 14) is a nonmetal, hydrogen (group 1) is a nonmetal, and chlorine (group 17) is a nonmetal;
CHCl3 is predicted to be molecular.

(d) Lithium (group 1) is a metal, and carbonate is a polyatomic ion; Li2CO3 is predicted to be ionic.

Check Your Learning

Using the periodic table, predict whether the following compounds are ionic or covalent:

(a) SO2

(b) CaF2

(c) N2H4

(d) Al2(SO4)3
96 2 • Atoms, Molecules, and Ions

Answer:
(a) molecular; (b) ionic; (c) molecular; (d) ionic

2.7 Chemical Nomenclature

LEARNING OBJECTIVES
By the end of this section, you will be able to:
• Derive names for common types of inorganic compounds using a systematic approach

Nomenclature, a collection of rules for naming things, is important in science and in many other situations.
This module describes an approach that is used to name simple ionic and molecular compounds, such as
NaCl, CaCO3, and N2O4. The simplest of these are binary compounds, those containing only two elements, but
we will also consider how to name ionic compounds containing polyatomic ions, and one specific, very
important class of compounds known as acids (subsequent chapters in this text will focus on these compounds
in great detail). We will limit our attention here to inorganic compounds, compounds that are composed
principally of elements other than carbon, and will follow the nomenclature guidelines proposed by IUPAC.
The rules for organic compounds, in which carbon is the principle element, will be treated in a later chapter on
organic chemistry.

Ionic Compounds
To name an inorganic compound, we need to consider the answers to several questions. First, is the compound
ionic or molecular? If the compound is ionic, does the metal form ions of only one type (fixed charge) or more
than one type (variable charge)? Are the ions monatomic or polyatomic? If the compound is molecular, does it
contain hydrogen? If so, does it also contain oxygen? From the answers we derive, we place the compound in
an appropriate category and then name it accordingly.

Compounds Containing Only Monatomic Ions

The name of a binary compound containing monatomic ions consists of the name of the cation (the name of
the metal) followed by the name of the anion (the name of the nonmetallic element with its ending replaced by
the suffix –ide). Some examples are given in Table 2.6.

Names of Some Ionic Compounds

NaCl, sodium chloride Na2O, sodium oxide

KBr, potassium bromide CdS, cadmium sulfide

CaI2, calcium iodide Mg3N2, magnesium nitride

CsF, cesium fluoride Ca3P2, calcium phosphide

LiCl, lithium chloride Al4C3, aluminum carbide

TABLE 2.6

Compounds Containing Polyatomic Ions

Compounds containing polyatomic ions are named similarly to those containing only monatomic ions, i.e. by
naming first the cation and then the anion. Examples are shown in Table 2.7.

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 2.7 • Chemical Nomenclature 97

Names of Some Polyatomic Ionic Compounds

KC2H3O2, potassium acetate NH4Cl, ammonium chloride

NaHCO3, sodium bicarbonate CaSO4, calcium sulfate

Al2(CO3)3, aluminum carbonate Mg3(PO4)2, magnesium phosphate

TABLE 2.7

Chemistry in Everyday Life

Ionic Compounds in Your Cabinets
Every day you encounter and use a large number of ionic compounds. Some of these compounds, where
they are found, and what they are used for are listed in Table 2.8. Look at the label or ingredients list on the
various products that you use during the next few days, and see if you run into any of those in this table, or
find other ionic compounds that you could now name or write as a formula.

Everyday Ionic Compounds

Ionic Compound Use

NaCl, sodium chloride ordinary table salt

KI, potassium iodide added to “iodized” salt for thyroid health

NaF, sodium fluoride ingredient in toothpaste

NaHCO3, sodium bicarbonate baking soda; used in cooking (and as antacid)

Na2CO3, sodium carbonate washing soda; used in cleaning agents

NaOCl, sodium hypochlorite active ingredient in household bleach

CaCO3 calcium carbonate ingredient in antacids

Mg(OH)2, magnesium hydroxide ingredient in antacids

Al(OH)3, aluminum hydroxide ingredient in antacids

NaOH, sodium hydroxide lye; used as drain cleaner

K3PO4, potassium phosphate food additive (many purposes)

MgSO4, magnesium sulfate added to purified water

Na2HPO4, sodium hydrogen phosphate anti-caking agent; used in powdered products

TABLE 2.8
98 2 • Atoms, Molecules, and Ions

Ionic Compound Use

Na2SO3, sodium sulfite preservative

TABLE 2.8

Compounds Containing a Metal Ion with a Variable Charge

Most of the transition metals and some main group metals can form two or more cations with different
charges. Compounds of these metals with nonmetals are named with the same method as compounds in the
first category, except the charge of the metal ion is specified by a Roman numeral in parentheses after the
name of the metal. The charge of the metal ion is determined from the formula of the compound and the
charge of the anion. For example, consider binary ionic compounds of iron and chlorine. Iron typically exhibits
a charge of either 2+ or 3+ (see Figure 2.29), and the two corresponding compound formulas are FeCl2 and
FeCl3. The simplest name, “iron chloride,” will, in this case, be ambiguous, as it does not distinguish between
these two compounds. In cases like this, the charge of the metal ion is included as a Roman numeral in
parentheses immediately following the metal name. These two compounds are then unambiguously named
iron(II) chloride and iron(III) chloride, respectively. Other examples are provided in Table 2.9.

Some Ionic Compounds with Variably Charged Metal Ions

Compound Name

FeCl2 iron(II) chloride

FeCl3 iron(III) chloride

Hg2O mercury(I) oxide

HgO mercury(II) oxide

SnF2 tin(II) fluoride

SnF4 tin(IV) fluoride

TABLE 2.9

Out-of-date nomenclature used the suffixes –ic and –ous to designate metals with higher and lower charges,
respectively: Iron(III) chloride, FeCl3, was previously called ferric chloride, and iron(II) chloride, FeCl2, was
known as ferrous chloride. Though this naming convention has been largely abandoned by the scientific
community, it remains in use by some segments of industry. For example, you may see the words stannous
fluoride on a tube of toothpaste. This represents the formula SnF2, which is more properly named tin(II)
fluoride. The other fluoride of tin is SnF4, which was previously called stannic fluoride but is now named
tin(IV) fluoride.

Ionic Hydrates
Ionic compounds that contain water molecules as integral components of their crystals are called hydrates.
The name for an ionic hydrate is derived by adding a term to the name for the anhydrous (meaning “not
hydrated”) compound that indicates the number of water molecules associated with each formula unit of the
compound. The added word begins with a Greek prefix denoting the number of water molecules (see Table
2.10) and ends with “hydrate.” For example, the anhydrous compound copper(II) sulfate also exists as a

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 2.7 • Chemical Nomenclature 99

hydrate containing five water molecules and named copper(II) sulfate pentahydrate. Washing soda is the
common name for a hydrate of sodium carbonate containing 10 water molecules; the systematic name is
sodium carbonate decahydrate.

Formulas for ionic hydrates are written by appending a vertically centered dot, a coefficient representing the
number of water molecules, and the formula for water. The two examples mentioned in the previous paragraph
are represented by the formulas

Nomenclature Prefixes

Number Prefix Number Prefix

1 (sometimes omitted) mono- 6 hexa-

2 di- 7 hepta-

3 tri- 8 octa-

4 tetra- 9 nona-

5 penta- 10 deca-

TABLE 2.10

EXAMPLE 2.13

Naming Ionic Compounds

Name the following ionic compounds

(a) Fe2S3

(b) CuSe

(c) GaN

(d) MgSO4·7H2O

(e) Ti2(SO4)3

Solution
The anions in these compounds have a fixed negative charge (S2−, Se2− , N3−, and and the compounds
must be neutral. Because the total number of positive charges in each compound must equal the total number
of negative charges, the positive ions must be Fe3+, Cu2+, Ga3+, Mg2+, and Ti3+. These charges are used in the
names of the metal ions:

(a) iron(III) sulfide

(b) copper(II) selenide

(c) gallium(III) nitride

(d) magnesium sulfate heptahydrate

(e) titanium(III) sulfate

100 2 • Atoms, Molecules, and Ions

Check Your Learning

Write the formulas of the following ionic compounds:

(a) chromium(III) phosphide

(b) mercury(II) sulfide

(c) manganese(II) phosphate

(d) copper(I) oxide

(e) iron(III) chloride dihydrate

Answer:
(a) CrP; (b) HgS; (c) Mn3(PO4)2; (d) Cu2O; (e) FeCl3·2H2O

Chemistry in Everyday Life

Erin Brockovich and Chromium Contamination
In the early 1990s, legal file clerk Erin Brockovich (Figure 2.32) discovered a high rate of serious illnesses
in the small town of Hinckley, California. Her investigation eventually linked the illnesses to groundwater
contaminated by Cr(VI) used by Pacific Gas & Electric (PG&E) to fight corrosion in a nearby natural gas
pipeline. As dramatized in the film Erin Brockovich (for which Julia Roberts won an Oscar), Erin and lawyer
Edward Masry sued PG&E for contaminating the water near Hinckley in 1993. The settlement they won in
1996—$333 million—was the largest amount ever awarded for a direct-action lawsuit in the US at that time.

FIGURE 2.32 (a) Erin Brockovich found that Cr(VI), used by PG&E, had contaminated the Hinckley, California,
water supply. (b) The Cr(VI) ion is often present in water as the polyatomic ions chromate, (left), and
dichromate, (right).

Chromium compounds are widely used in industry, such as for chrome plating, in dye-making, as
preservatives, and to prevent corrosion in cooling tower water, as occurred near Hinckley. In the
environment, chromium exists primarily in either the Cr(III) or Cr(VI) forms. Cr(III), an ingredient of many
vitamin and nutritional supplements, forms compounds that are not very soluble in water, and it has low
toxicity. But Cr(VI) is much more toxic and forms compounds that are reasonably soluble in water.
Exposure to small amounts of Cr(VI) can lead to damage of the respiratory, gastrointestinal, and immune
systems, as well as the kidneys, liver, blood, and skin.

Despite cleanup efforts, Cr(VI) groundwater contamination remains a problem in Hinckley and other
locations across the globe. A 2010 study by the Environmental Working Group found that of 35 US cities
tested, 31 had higher levels of Cr(VI) in their tap water than the public health goal of 0.02 parts per billion
set by the California Environmental Protection Agency.

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 2.7 • Chemical Nomenclature 101

Molecular (Covalent) Compounds

The bonding characteristics of inorganic molecular compounds are different from ionic compounds, and they
are named using a different system as well. The charges of cations and anions dictate their ratios in ionic
compounds, so specifying the names of the ions provides sufficient information to determine chemical
formulas. However, because covalent bonding allows for significant variation in the combination ratios of the
atoms in a molecule, the names for molecular compounds must explicitly identify these ratios.

Compounds Composed of Two Elements

When two nonmetallic elements form a molecular compound, several combination ratios are often possible.
For example, carbon and oxygen can form the compounds CO and CO2. Since these are different substances
with different properties, they cannot both have the same name (they cannot both be called carbon oxide). To
deal with this situation, we use a naming method that is somewhat similar to that used for ionic compounds,
but with added prefixes to specify the numbers of atoms of each element. The name of the more metallic
element (the one farther to the left and/or bottom of the periodic table) is first, followed by the name of the
more nonmetallic element (the one farther to the right and/or top) with its ending changed to the suffix –ide.
The numbers of atoms of each element are designated by the Greek prefixes shown in Table 2.10.

When only one atom of the first element is present, the prefix mono- is usually deleted from that part. Thus, CO
is named carbon monoxide, and CO2 is called carbon dioxide. When two vowels are adjacent, the a in the Greek
prefix is usually dropped. Some other examples are shown in Table 2.11.

Names of Some Molecular Compounds Composed of Two Elements

Compound Name Compound Name

SO2 sulfur dioxide BCl3 boron trichloride

SO3 sulfur trioxide SF6 sulfur hexafluoride

NO2 nitrogen dioxide PF5 phosphorus pentafluoride

N2O4 dinitrogen tetroxide P4O10 tetraphosphorus decaoxide

N2O5 dinitrogen pentoxide IF7 iodine heptafluoride

TABLE 2.11

There are a few common names that you will encounter as you continue your study of chemistry. For example,
although NO is often called nitric oxide, its proper name is nitrogen monoxide. Similarly, N2O is known as
nitrous oxide even though our rules would specify the name dinitrogen monoxide. (And H2O is usually called
water, not dihydrogen monoxide.) You should commit to memory the common names of compounds as you
encounter them.

EXAMPLE 2.14

Naming Covalent Compounds

Name the following covalent compounds:

(a) SF6

(b) N2O3

(c) Cl2O7
102 2 • Atoms, Molecules, and Ions

(d) P4O6

Solution
Because these compounds consist solely of nonmetals, we use prefixes to designate the number of atoms of
each element:

(a) sulfur hexafluoride

(b) dinitrogen trioxide

(c) dichlorine heptoxide

(d) tetraphosphorus hexoxide

Check Your Learning

Write the formulas for the following compounds:

(a) phosphorus pentachloride

(b) dinitrogen monoxide

(c) iodine heptafluoride

(d) carbon tetrachloride

Answer:
(a) PCl5; (b) N2O; (c) IF7; (d) CCl4

LINK TO LEARNING
The following website (http://openstax.org/l/16chemcompname) provides practice with naming chemical
compounds and writing chemical formulas. You can choose binary, polyatomic, and variable charge ionic
compounds, as well as molecular compounds.

Binary Acids
Some compounds containing hydrogen are members of an important class of substances known as acids. The
chemistry of these compounds is explored in more detail in later chapters of this text, but for now, it will
suffice to note that many acids release hydrogen ions, H+, when dissolved in water. To denote this distinct
chemical property, a mixture of water with an acid is given a name derived from the compound’s name. If the
compound is a binary acid (comprised of hydrogen and one other nonmetallic element):

1. The word “hydrogen” is changed to the prefix hydro-

2. The other nonmetallic element name is modified by adding the suffix -ic
3. The word “acid” is added as a second word

For example, when the gas HCl (hydrogen chloride) is dissolved in water, the solution is called hydrochloric
acid. Several other examples of this nomenclature are shown in Table 2.12.

Names of Some Simple Acids

Name of Gas Name of Acid

HF(g), hydrogen fluoride HF(aq), hydrofluoric acid

HCl(g), hydrogen chloride HCl(aq), hydrochloric acid

TABLE 2.12

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 2.7 • Chemical Nomenclature 103

Name of Gas Name of Acid

HBr(g), hydrogen bromide HBr(aq), hydrobromic acid

HI(g), hydrogen iodide HI(aq), hydroiodic acid

H2S(g), hydrogen sulfide H2S(aq), hydrosulfuric acid

TABLE 2.12

Oxyacids
Many compounds containing three or more elements (such as organic compounds or coordination
compounds) are subject to specialized nomenclature rules that you will learn later. However, we will briefly
discuss the important compounds known as oxyacids, compounds that contain hydrogen, oxygen, and at least
one other element, and are bonded in such a way as to impart acidic properties to the compound (you will
learn the details of this in a later chapter). Typical oxyacids consist of hydrogen combined with a polyatomic,
oxygen-containing ion. To name oxyacids:

1. Omit “hydrogen”
2. Start with the root name of the anion
3. Replace –ate with –ic, or –ite with –ous
4. Add “acid”

For example, consider H2CO3 (which you might be tempted to call “hydrogen carbonate”). To name this
correctly, “hydrogen” is omitted; the –ate of carbonate is replace with –ic; and acid is added—so its name is
carbonic acid. Other examples are given in Table 2.13. There are some exceptions to the general naming
method (e.g., H2SO4 is called sulfuric acid, not sulfic acid, and H2SO3 is sulfurous, not sulfous, acid).

Names of Common Oxyacids

Formula Anion Name Acid Name

HC2H3O2 acetate acetic acid

HNO3 nitrate nitric acid

HNO2 nitrite nitrous acid

HClO4 perchlorate perchloric acid

H2CO3 carbonate carbonic acid

H2SO4 sulfate sulfuric acid

H2SO3 sulfite sulfurous acid

H3PO4 phosphate phosphoric acid

TABLE 2.13