When Anna and Bill went looking for things related to chemistry (in Chapter 1), they classified nearly everything they saw as matter. Now Anna, who is writing a history paper, wants to know how our modern ideas about the atom arose. She learned in a philosophy class that the ancient Greeks thought the world was made up of four kinds of matter: earth, air, fire, and water. They called them elements. How did we come to understand matter as being composed of atoms of the different elements in the periodic table?

As Anna sets out to learn more, she discovers that the concept of the atom dates back to 450 BC. The Greek philosopher Democritus believed that there was a limit to how far matter could be broken down into smaller pieces. A piece of sand, for example, could be split into smaller pieces, and those pieces into even smaller pieces. But it made no sense to Democritus that a piece of sand could be broken down into smaller pieces forever. Democritus gave the smallest components of matter the name atomos (atom in English), which is a Greek word meaning “unbreakable,” to describe the smallest components of matter.

The concept of the atom did not become popular until centuries later. In the 1700s, chemists realized that they must look for elements by breaking down matter into simpler substances. Once they had something that could not be broken down any further, they called it an element. About 30 different elements had been discovered by the end of the 1700s. Throughout the 1700s, few chemists believed that elements were made of atoms. The experiments of many chemists, however, provided evidence for the atomic nature of matter.

Anna learns that the science and math teacher John Dalton offered the first convincing argument for the existence of atoms, based on well-established experimental results, which he published in a book in 1808. So persuasive were his arguments that the idea came to be called Dalton’s atomic theory. As Anna continues her research, she learns that the acceptance of Dalton’s atomic theory by the scientific community paved the way for our current understanding of the atom—the modern model of the atom. Many additional questions pop into Anna’s head now. How are atoms of different elements, such as carbon and hydrogen, different from one another? Do they have different masses? Are atoms composed of even smaller particles?

Anna isn’t the only one who’s finding that chemistry relates to other subjects. Anna’s roommate Megan, a nutrition major, is quick to realize that the human body is composed of matter, and that everything we are made of is just different combinations of elements. Although we are about 93% carbon, hydrogen, and oxygen, small amounts of many other elements are essential to proper functioning. Many of these are called essential minerals. Minerals are crucial to the growth and production of bones, teeth, hair, blood, nerves, and skin, not to mention the enzymes and hormones that living cells and tissues need to function. The essential minerals that we need in the greatest quantity are calcium, phosphorus, potassium, sodium, chlorine, magnesium, and iron. We get them from the foods we eat and drink. For example, calcium, which makes bones strong, is found in significant amounts in milk products. Potassium, an important regulator of cell fluids, is found in many fruits and vegetables. Iron, a building block for the oxygen-carrying hemoglobin in our blood, comes from fruits, vegetables, and red meat.

Megan notices that many of the minerals are classified as metals on the periodic table. She wonders if we ingest metals in the foods we eat. Generally, no. We take them in as components of compounds in which a metal such as iron exists as an ion. An ion is an atom with an electrical charge. The properties of metal ions are different from the properties of pure metal elements. Ions are important to the chemical processes that occur in our bodies. After learning about ions as nutrients, Megan begins to wonder how ions differ from atoms and why they have different properties.

In this chapter we will learn about the atoms that compose the elements and about the similarities and differences among them. We will answer Anna’s and Megan’s questions and many of your own.

Some essential minerals are involved in many processes. Magnesium, for example, plays important roles in more than 300 chemical processes that occur in the human body.

The common iron supplement iron(II) sulfate contains iron in an ion form.

©McGraw-Hill Education/John Flournoy

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Questions for Consideration

2.1 What evidence suggests that matter is composed of atoms?

2.2 How can the composition of atoms differ?

2.3 How do ions differ from the atoms of elements?

2.4 How can we describe the mass of the atoms of an element?

2.5 How does the periodic table relate to the structure and behavior of atoms?

Math Tools Used in This Chapter

Scientific Notation (Math Toolbox 1.1)

Significant Figures (Math Toolbox 1.2)

2.1 Dalton’s Atomic Theory

In Anna’s research, she learned that before the nineteenth century, many believed that matter was continuous—that is, that it could be divided infinitely (Figure 2.1). That notion began to change in 1808, when John Dalton published his atomic theory. Dalton based his argument on experimental evidence that had begun accumulating about three decades earlier. For example, Antoine Lavoisier’s experimental results led to the law of conservation of mass, which he published in 1787. His experiments showed that no measurable change in mass occurs during a chemical reaction. The mass of the products of a reaction always equals the mass of the reacting substances.

FIGURE 2.1 Even when viewed under an optical microscope (center), the particle composition of the copper wire is not obvious. It seems continuous rather than particulate. As shown by the model (right), it would be necessary to use an electron microscope to observe that copper metal, and all matter, is composed of particles.
©imagebroker/Alamy Stock Photo; (inset): ©Jim Birk

When we observe chemical changes, it may seem in some cases that the law of conservation of mass does not apply. After wood is burned in a bonfire, for example, there is less mass left in the fire pit. When an iron nail rusts, the rusted nail has a greater mass than the original nail. However, if we carry out each of these reactions in a closed container and find the mass of the closed container and its contents before and after the change, the mass does not change. With an open container, the gases produced when wood burns escape into the atmosphere, so they are not included in the final measurement. The gases consumed when a nail rusts are not included in the original mass. The example in Figure 2.2 shows the conservation of mass when a reaction occurs in a closed system. Lavoisier did experiments like these in which he was careful not to let any matter enter or escape.

FIGURE 2.2 When sodium carbonate contained in the balloon is placed in a solution of hydrochloric acid, a reaction forms carbon dioxide gas. Because the carbon dioxide is not allowed to dissipate into the atmosphere, its mass is included in the final mass in the experiment shown here. How would the result be different if the container were open?
©Jim Birk

Joseph Proust did similar experiments, mostly reacting metals with oxygen. He found that the oxygen content of the products was always fixed at one or two values, rather than showing a range of all possible values. His findings, published between 1797 and 1804, led to the law of definite proportions. This law states that all samples of the same compound always contain the same proportions by mass of the component elements. For example, pure water is always composed of oxygen and hydrogen with a mass ratio of 8 g of oxygen for every 1 g of hydrogen. Conversely, when water is broken down into its elements using electricity, oxygen and hydrogen are formed in this same mass ratio of 8:1.

As you will learn in Chapter 15, mass is not conserved in reactions that involve the nucleus of the atom. In nuclear fusion reactions that occur in the Sun, for example, matter is converted into energy. Nuclear reactions were not known in Lavoisier’s and Dalton’s time.

Dalton reasoned that the law of conservation of mass and the law of definite proportions could be explained only if matter was composed of atoms. The following postulates summarize Dalton’s atomic theory:

1.All matter is composed of exceedingly small, indivisible particles, called atoms.

2.All atoms of a given element are identical both in mass and in chemical properties. However, atoms of different elements have different masses and different chemical properties.

3.Atoms are not created or destroyed in chemical reactions.

4.Atoms combine in simple, fixed, whole-number ratios to form compounds.

A chemical reaction, according to Dalton’s atomic theory, is a rearrangement of atoms into a new combination, resulting in the formation of one or more new chemical substances (Figure 2.3). Although Dalton’s theory has been modified over the past 200 years, it still provides the basis for understanding how atoms are the building blocks of matter. As with most theories, atomic theory has been amended as new evidence has been found. We now know that some of Dalton’s postulates are not quite correct. The first statement, for example, is not strictly true. Atoms are composed of even smaller particles, called subatomic particles. The second statement is not quite accurate either. As you will learn in Section 2.2, atoms of a given element actually can vary in mass.

FIGURE 2.3 Two H₂ molecules combine with one O₂ molecule and rearrange to form two H₂O molecules. A chemical reaction is simply the rearrangement of atoms into new combinations.

In addition to developing atomic theory, John Dalton also studied meteorology and color blindness.

©GeorgiosArt/Getty Images

Page 60Still, Dalton’s central idea has stood the test of time and experimentation—although for nearly two centuries no one had a way to see the atoms that Dalton proposed. Because atoms are so small, they cannot be seen using an optical microscope. However, with the invention of the scanning tunneling microscope (STM) in 1981, scientists can now observe what we interpret to be single atoms on the surface of a material (Figure 2.4).

FIGURE 2.4 This scanning tunneling microscope (STM) image shows atoms on the surface of copper. Because atoms are too small to be seen with an optical microscope, images could not be produced until the invention of the STM in the 1980s.
©Drs. Ali Yazdani & Daniel J. Hornbaker/Science Source

An optical microscope cannot resolve images smaller than the size of the waves making up the light used to examine the object.

2.2 Structure of the Atom

When reading nutritional labels, Megan noticed that a food supplement contains both iron and selenium. She wondered how the atoms of these elements differ. Scientists have also asked this question about atoms in general and found ways to probe their composition. They have studied the smaller particles that actually make up atoms.

Selenium is an essential mineral involved in the function of the thyroid gland.

Subatomic Particles

The existence of the electron, a negatively charged subatomic particle, was demonstrated by J. J. Thomson in 1897. He conducted a series of experiments with cathode-ray tubes (Figure 2.5). In a partially evacuated cathode-ray tube, a voltage is applied by connecting each end of the tube to a battery. Electricity then flows from one end of the tube to the other in the form of a ray. The invisible rays can be observed when they cause certain materials coated on the glass to glow. Thomson found that, in a magnetic or electric field, the rays bent toward a positively charged plate and were deflected away from a negatively charged plate outside the tube. He knew that like electrical charges repel each other and opposite charges attract each other. The bending of the beam toward the positive plate (and away from the negative plate) showed that the beam was composed of negatively charged particles. Thomson showed that the rays had a negative electrical charge no matter what material was used for the source of the rays. This result indicated that the rays were composed of identical, negatively charged particles common to all matter. We call these particles electrons. Thomson was also able to determine the charge-to-mass ratio of the electron from such experiments.

FIGURE 2.5 J. J. Thomson’s experiments with cathode-ray tubes led to the discovery of the electron. (A) In a normal cathode-ray tube, the cathode ray travels in a straight path when there is no external field. (B) When Thomson applied an external electric field to the cathode-ray tube, the cathode ray bent toward the positive plate. It was known that like charges repel each other and opposite charges attract each other. The bending of the beam toward the positive plate (and away from the negative plate) indicated that the beam was composed of negatively charged particles.Page 61

Cathode-ray tubes (CRTs) are the fundamental components of older television picture tubes and computer monitors. The screen contains chemical compounds that glow when struck by fast-moving electrons. Different chemicals that glow different colors provide a color picture.

How did the discovery of electrons make it necessary to change one of Dalton’s postulates?

Building on these results, Robert Millikan experimented with oil droplets in an effort to measure the strength of the negative charge on an electron (Figure 2.6). When he exposed the droplets to radiation, they took on an electrical charge. By measuring the magnitude of the electric field necessary to cause the droplets to hang suspended in air, Millikan determined that the charge on an electron is −1.6022 × 10⁻¹⁹ coulombs (C) (a coulomb is a unit of electrical charge). He then calculated the mass of an electron to be 9.1094 × 10⁻²⁸ g from its charge and the charge-to-mass ratio determined by Thomson. Since even the lightest atoms have a mass greater than 10⁻²⁴ g, an electron contributes only a small part to the mass of an atom. In fact, the mass of an electron is 1836 times less than the mass of one hydrogen atom, the lightest of all the elements. Because of the electron’s small mass, it was originally thought that thousands of them must lie inside a single hydrogen atom. We now know there is only one.

FIGURE 2.6 In Millikan’s oil-drop experiment, the electric field strength required to suspend an oil droplet depended upon the number of extra electrons on it. These experiments allowed Millikan to determine the charge of a single electron and calculate its mass.

ANIMATION: Cathode-Ray Tube

ANIMATION: Millikan Oil Drop

Proton therapy is a radiation technique used to treat various forms of cancer. During treatment a beam of radiation (protons) is targeted directly at tumor cells. The radiation ionizes molecules in the cells of tumors, causing permanent damage to the DNA of the cells.

The discovery of the electron stimulated many more experiments in search of other subatomic particles. Since atoms are electrically neutral, scientists reasoned that atoms must contain positively charged particles to counter the negatively charged electrons. The positively charged particle, called a proton, has a charge equal in magnitude to the electron but opposite in sign, +1.6022 × 10⁻¹⁹ C. To be electrically neutral, an atom must have equal numbers of protons and electrons. To make it easier to deal with electrical charges in matter, we usually express the charges as a multiple of the charge of an electron or of a proton, instead of in units of coulombs. Expressed in this way, the charge of an electron is 1−, and the charge of a proton is 1+.

The Nuclear Atom

ANIMATION: α-Particle Scattering

ANIMATION: Rutherford’s Gold Foil Experiment

How might protons and electrons be arranged in an atom? Thomson’s model of atomic structure, called the “plum pudding” model, assumed that protons and electrons were evenly distributed throughout the atom (Figure 2.7). Ernest Rutherford designed an experiment to test this model, and his associate, Hans Geiger, carried it out. The experiment involved bombarding thin gold foil with alpha particles. Alpha particles (α) were known at the time as positively charged particles thousands of times greater in mass than electrons. (Today we know them as helium atoms that have lost their electrons.) According to the plum pudding model, none of the alpha particles should have been affected by the dispersed bits of positive and negative charge in the gold atoms. They should have zipped right through the gold foil, and most did. However, some were deflected slightly, and a few actually bounced backwards, as shown in Figure 2.8. The result was quite unexpected. It was as if you fired a bullet at a sheet of tissue paper and it came back and hit you! The deflection of these relatively massive alpha particles suggested that most of the mass of the atom had to be concentrated in a positively charged core, which Rutherford called the nucleus. The electrons, he reasoned, had to be dispersed in the large volume outside of the nucleus. The large electron space was the area penetrated by most of the alpha particles. Only if an alpha particle came close enough to the incredibly dense nucleus would it be deflected from its original path. The alpha particles that hit the nucleus of a gold atom head-on were deflected backwards.

FIGURE 2.7 Thomson’s model suggested that electrons in the atom might be embedded in a sphere of positive charge, like raisins in plum pudding.Page 63FIGURE 2.8 Rutherford’s gold foil experiment led to the nuclear model of the atom. A beam of positively charged alpha (α) particles was directed through a thin layer of gold. A zinc-sulfide screen detected the alpha particles by producing a flash of light upon contact. (A) According to the plum pudding model, all alpha particles should have penetrated straight through the gold atoms. (B) In the experiment, many of the alpha particles penetrated straight through the atom, but some were deflected. (C) The nuclear model of the atom explains the experimental results. The positively charged protons are packed tightly together in the very center of the nucleus. When the nucleus is in the path of an alpha particle, the alpha particle is deflected from its original path.

Although a penny contains about 1 × 10²² atoms, most of the penny is empty space. Why?

Rutherford’s experiment was the basis for the nuclear model of the atom (Figure 2.9), developed in 1907. The model suggested that the nucleus contains the protons and most of the mass of the atom. The electrons exist outside the nucleus in what is often called an “electron cloud.”

FIGURE 2.9 In the nuclear model of the atom, protons (blue spheres) and neutrons (red spheres) are located in a tiny nucleus at the center of the atom. The space outside the nucleus is occupied by the electrons.

The diameter of the nucleus is about 10⁻¹⁴ m, and the diameter of the atom is about 10⁻¹⁰ m. These relative sizes are comparable to a flea in the center of a domed stadium. While the volume of the nucleus is about 10,000 times smaller than the entire atom, the mass of the nucleus accounts for most of the mass of the atom. This is because the nucleus contains heavy particles, such as protons, that are packed closely together.

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Like proton therapy, neutron therapy is a radiation technique used to treat various forms of cancer. However, the scientific evidence is inconclusive about the effectiveness of this form of treatment.

Due to its neutral charge, the neutron is an effective nuclear bombardment particle now used to penetrate and split the nuclei of large atoms in energy-producing fission reactions.

These experiments helped Rutherford and other scientists understand the structure of the atom better, but they could not account for the entire mass of the atom. Most atoms other than the hydrogen atom have masses that are at least twice the sum of the masses of the protons and electrons they contain. For example, calcium contains 20 protons and 20 electrons. Together, their mass is 3.3471 × 10⁻²³ g. Yet a calcium atom has a mass that is nearly twice this value (6.6359 × 10⁻²³ g).

To account for the extra mass, Rutherford hypothesized the neutron. A neutron is an uncharged particle in the nucleus of the atom. Because of the electrical neutrality of the neutron, it was difficult to detect. It was not until 1932 that James Chadwick, a scientist working with Rutherford, did experiments that proved the existence of the neutron. The mass of the neutron was determined to be 1.6749 × 10⁻²⁴ g, slightly greater than the mass of a proton. The properties of the electron, proton, and neutron are summarized in Table 2.1. All of the subatomic particles in the table are important to the nuclear model of the atom as we understand it today (Figure 2.9).

TABLE 2.1 Subatomic Particles

Particle	Mass (g)	Actual Charge (C)	Relative Charge
electron	9.1094 × 10⁻²⁸	−1.6022 × 10⁻¹⁹	1−
proton	1.6726 × 10⁻²⁴	+1.6022 × 10⁻¹⁹	1+
neutron	1.6749 × 10⁻²⁴	0	0

In Chapters 7 and 8 you will learn how electrons relate to chemical reactivity. In Chapter 15, you will learn how neutrons relate to nuclear radioactivity. In this chapter, we will focus on how protons determine the identity of an atom, neutrons help determine its mass, and electrons determine its charge.

Isotopes, Atomic Number, and Mass Number

Recall that Anna was wondering how atoms of carbon and hydrogen differ. What is it about an atom of an element that distinguishes it from an atom of some other element? The same scientists who identified the subatomic particles found an answer to this question. They learned, for example, that all hydrogen atoms contain just one proton, and that any atom containing just one proton is a hydrogen atom. In a similar fashion, any atom that contains two protons is a helium atom. An atom with three protons is a lithium atom, and so on. The number of protons in an atom’s nucleus determines the identity of that element. The number of protons in the nucleus of each atom of an element is the atomic number (Z) of that element. On the periodic table in this book, elements are shown with the atomic number just above the element symbol, as shown for the element gold (Au) in Figure 2.10.

FIGURE 2.10 The atomic number of each element is indicated on the periodic table in this book just above the element symbol.

How does the existence of isotopes make it necessary to change one of Dalton’s postulates?

How can we determine the number of electrons and neutrons in an atom? Atoms are electrically neutral. This means that the number of electrons in an atom equals the number of protons—the atomic number. For example, the atomic number of gold (Au) is 79, so an atom of gold has 79 protons and 79 electrons. Not all atoms of an element contain the same number of neutrons, however. An isotope of an element is an atom that contains a specific number of neutrons. Most of the elements consist of more than one isotope with different numbers of neutrons. For example, most hydrogen atoms contain no neutrons. Their nucleus consists only of a single proton. Not all hydrogen atoms are alike, however. Some contain one neutron and some even have two. All three are isotopes of hydrogen. They are still hydrogen atoms because they have one proton, but their nuclei differ in the number of neutrons (Figure 2.11). The isotopes of an element have essentially identical chemical properties, but their physical properties, such as melting point and boiling point, may differ slightly.

FIGURE 2.11 These images show the subatomic particles in the nuclei of hydrogen isotopes. Isotopes of hydrogen vary in the number of neutrons. In these representations, the blue spheres with a plus in them represent protons. The red spheres represent neutrons. How many neutrons are in each isotope? What do the letters and numbers below each diagram represent?

One way to distinguish between isotopes is by their mass number. The mass number (A) of an isotope is the sum of the number of protons (Z) and the number of neutrons (N) in its nucleus:

Mass number = number of protons + number of neutrons

A = Z + N

The three isotopes of hydrogen have different mass numbers (A = 1, 2, and 3) because they have different numbers of neutrons, also called the neutron number (N = 0, 1, and 2). The mass number is not an actual mass; it is a count of the number of particles in the nucleus.

Example 2.1 shows how you can determine the atomic number and mass number for an isotope of an element by looking at an atomic-level representation.

EXAMPLE 2.1Determining Atomic Number and Mass Number

For the atom represented in the following diagram,

(a) Determine the number of protons and neutrons.

(b) Identify the atomic number and the element.

Solution:

(a) The protons and neutrons make up the nucleus. The protons are positively charged and neutrons have no charge. There are five protons and six neutrons.

(b) The atomic number, the number of protons, is 5. From the periodic table we see that the element with atomic number 5 is boron.

Consider This 2.1

If the number of protons were one less than shown in the diagram, and the number of neutrons were one greater, how would the atomic number and mass number change and would this represent the same element or a different element?

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Practice Problem 2.1

For the atom represented in the following diagram,

(a) Determine the number of protons and neutrons.

(b) Identify the atomic number and the element.

Further Practice: Questions 2.29 and 2.30 at the end of the chapter

As shown by Example 2.2, you can determine the number of each type of subatomic particle when given both the atomic number and the mass number.

EXAMPLE 2.2Numbers of Subatomic Particles

The only stable isotope of naturally occurring fluorine has a mass number of 19. How many protons, electrons, and neutrons are in a fluorine atom?

Solution:

The periodic table shows the atomic number of fluorine is 9. Because the atomic number is equal to the number of protons, the number of protons in a fluorine atom is 9. The number of electrons in an atom equals the number of protons, so fluorine has 9 electrons. The number of neutrons can be determined from the known mass number and known number of protons. The mass number, the total number of protons and neutrons, is 19. There are 9 protons, so there must be 10 neutrons to give a mass number of 19.

Consider This 2.2

If, on another planet, there were another stable isotope of fluorine with a mass number of 17, how would the numbers of protons, electrons, and neutrons differ?

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Practice Problem 2.2

A rare isotope of carbon has a mass number of 14. How many protons, electrons, and neutrons are in this isotope of carbon?

Further Practice: Questions 2.33 and 2.34 at the end of the chapter

The mass number and atomic number of an isotope are often represented by the notation called an isotope symbol, as shown in Figure 2.12. In an isotope symbol, X is the element symbol shown on the periodic table, A is the mass number, and Z is the atomic number (number of protons). For example, the isotope symbol for the isotope that contains one proton and no neutrons in its nucleus is , while hydrogen with one proton and one neutron is represented as (see Figure 2.11). The symbol for carbon with six protons and eight neutrons is . Because each element can only have one atomic number Z, sometimes the atomic number is left off the isotope symbol. For example, the isotope of carbon with eight neutrons can also be represented as ¹⁴C. Other common representations use the name of the element followed by the mass number (carbon-14), or the element symbol followed by the mass number (C-14).

FIGURE 2.12 An isotope symbol consists of the element symbol with a subscript to denote the number of protons and a superscript to denote the sum of the number of protons and neutrons, the mass number.

EXAMPLE 2.3Writing Isotope Symbols

Write two representations for the following isotope.

Solution:

There are three protons, so the atomic number is 3, corresponding to the element lithium (Li). There are four neutrons, so the mass number is 7, the sum of the protons and neutrons. This isotope can be represented as i, ⁷Li, lithium-7, or Li-7.

Consider This 2.3

If there were one less neutron in the image of the nucleus, how would the isotope representation differ?

Practice Problem 2.3

Write two representations for the following isotope.

Further Practice: Questions 2.35 and 2.36 at the end of the chapter

As shown by Example 2.4, you can determine the number of neutrons in an atom from its isotope symbol by subtracting the atomic number (number of protons) from the mass number (total number of nuclear particles).

EXAMPLE 2.4Interpreting Isotope Symbols

Determine the number of neutrons in each of the following isotopes.

(a) (b) ²³Na (c) hydrogen-3

Solution:

The number of neutrons is equal to the difference between the mass number (A) and the atomic number (Z). If the atomic number is not given in the symbol, look on the periodic table for the number that appears above the element symbol.

(a) 238 − 92 = 146 neutrons

(b) 23 − 11 = 12 neutrons

Consider This 2.4

If the mass numbers were omitted from the isotope symbols provided, could you still determine the number of neutrons in each isotope?

Practice Problem 2.4

Determine the number of neutrons in each of the following isotopes.

(a) (b) ³⁷Cl (c) carbon-13

Further Practice: Questions 2.37 and 2.38 at the end of the chapter

Page 69The isotope of hydrogen that has a mass number of 2 () has a special name, deuterium. Deuterium (sometimes represented by the symbol D) accounts for less than 1% of naturally occurring hydrogen. It is used frequently by chemists to observe specific hydrogen atoms in experiments. When deuterium (D₂) combines with oxygen, deuterated water (D₂O) forms. It behaves chemically just like ordinary water, but it has different physical properties. The photo in Figure 2.13 shows two ice cubes in water. One ice cube is made of H₂O and the other of D₂O. What property is different for these ice cubes?

FIGURE 2.13 Which ice cube contains deuterium (hydrogen-2)? How can you tell?
©Tom Pantages

Since one ice cube sinks and the other does not, we can conclude that they have different densities. Deuterated water (D₂O) is denser than H₂O because deuterium has two particles in its nucleus and hydrogen-1 only has one. In its solid state, the density of D₂O is greater than 1, so it sinks in water. For this reason D₂O is sometimes called heavy water.

Isotopes have been extensively applied in various fields, from archeology to medicine to art. Radioactive isotopes (those that emit radiation) can be used to date objects of archeological or geological significance. Radioactive isotopes can also be used to generate electricity, such as in lightweight portable power packs used on space vehicles or in heart pacemakers. Isotopes find extensive use in medical studies. For example, thyroid problems can be diagnosed by ingestion of radioactive iodine-131. Isotopes can be used to treat cancer as well as to detect it. A higher dose of iodine-131 destroys thyroid tumors. Isotopes don’t always have to be radioactive to be useful. The techniques known as nuclear magnetic resonance or magnetic resonance imaging locate hydrogen-1 atoms for the purpose of determining molecular structures, or for imaging body parts. In a technique known as neutron activation analysis, isotopes can be detected in a nondestructive analysis of objects such as old paintings and other valuable works of art to determine whether the distribution of elements is like those in common use by a supposed artist, thereby authenticating the painting.

2.3 Ions

In the introduction, Megan was interested in the essential minerals that our bodies need. She noted that the minerals we consume in our food are mostly metal elements in their ion (electrically charged) form. She was especially interested in sodium because her doctor told her to reduce it in her diet. Dietary sodium, found in large quantities in snack foods and sports drinks, is also in an ion form. How are ions different from neutral atoms?

In a neutral (uncharged) atom, the number of electrons equals the number of protons, so the overall charge on the atom is zero. This is because the negative charge on each electron is equal in magnitude and opposite in sign to the positive charge on each proton. What happens when the protons and electrons are not equal? Then the overall charge is not zero. The overall charge is positive if the atom contains fewer electrons than protons. It is negative if there are more electrons than protons. When an atom contains more or less electrons than protons, it has a charge and is called an ion.

Many elements exist in nature as ions. The elements sodium and calcium, for example, are generally not found in nature as neutral atoms. Instead, they exist as ions—either combined with other ions in compounds or dissolved in water. Seawater, for example, contains dissolved sodium and calcium ions along with chloride ions. Supplies of elemental sodium and calcium that contain neutral atoms are made from the naturally occurring ionic forms by chemical reactions. Some elements, such as gold, do exist naturally as neutral atoms, but can be made into ions by chemical reactions.

Ions can be classified as cations or anions. A cation is a positively charged ion that contains fewer electrons than the number of protons in the nucleus. The overall charge is represented as a superscript to the right of the element symbol. An example is a calcium ion, Ca²⁺, which contains 20 protons and only 18 electrons. An anion is a negatively charged ion that contains more electrons than the number of protons in the nucleus. An example is Cl⁻, which contains 17 protons and 18 electrons.

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Ions do not form by gaining or losing protons or neutrons. Changing the number of protons would change the identity of the element! This does not occur under conditions other than nuclear reactions.

Although many ions exist naturally, they can also be formed from their neutral atoms. When a magnesium atom loses two electrons, it forms an Mg²⁺ cation. When a nitrogen atom gains three electrons, it forms an N³⁻ anion. Figure 2.14 shows the formation of cations and anions from neutral atoms. The charge on the ion is the net charge that results from unequal numbers of protons and electrons. As shown in Example 2.5, an ion symbol lets us determine the number of protons and electrons in the ion.

FIGURE 2.14 In an ion, the numbers of protons and electrons are not equal. (A) The cation Mg²⁺ forms when a magnesium atom loses two electrons. (B) The anion N³⁻ forms when a nitrogen atom gains three electrons.

EXAMPLE 2.5Ions

How many protons and electrons compose the following ions? Identify each as a cation or anion.

(a) Na⁺ (b) O²⁻ (c) Cr³⁺

Solution:

(a) The number of protons is determined from the atomic number. The atomic number of sodium is 11, so there are 11 protons in Na⁺. It has a positive charge because it does not have enough electrons to balance the positive charges from the protons. A 1+ charge shows that there is one less electron than the number of protons, so its total number of electrons must be 10. The Na⁺ ion is a cation because it is positively charged.

(b) The atomic number for oxygen is 8, so there are 8 protons. A negative charge results when there are more electrons than protons. Since the charge is 2−, there are two more electrons than protons, so there are 10 electrons. The O²⁻ ion is an anion because it is negatively charged.

(c) The atomic number of chromium is 24, so there are 24 protons in Cr³⁺. A 3+ charge indicates that there are three fewer electrons than the number of protons, so there are 21 electrons. The Cr³⁺ ion is a cation because it is positively charged.

Consider This 2.5

Which of the three ions has the same number of electrons as a noble gas atom?

Practice Problem 2.5

How many protons and electrons compose the following ions? Identify each as a cation or anion.

(a) F ⁻ (b) Mg²⁺ (c) N³⁻

Further Practice: Questions 2.49 and 2.50 at the end of the chapter

If we know the number of protons and electrons in an ion, we can write an ion symbol for it. If we also know the number of neutrons, we can write a combination isotope-ion symbol by writing an isotope symbol and adding the overall charge as the right-hand superscript. This is shown in Example 2.6.

EXAMPLE 2.6Writing Isotope Symbols for Ions

Write the isotope symbols for ions that have the following numbers of protons, neutrons, and electrons.

(a) 35 protons, 44 neutrons, and 36 electrons

(b) 13 protons, 14 neutrons, and 10 electrons

Solution:

(a) The atomic number (the number of protons) is 35, indicating the element symbol is Br. The mass number is the sum of the protons and neutrons, which is 79. The charge is 1− because there is one more electron than protons. The symbol is .

(b) The atomic number is 13, so the element symbol is Al. The mass number is 27 (13 + 14). The charge is 3+ because there are three fewer electrons than protons. The symbol is .

(c) The atomic number is 47, so the element is silver (Ag). The mass number is 109 (47 + 62). The charge is 1+ because the number of electrons is one less than the number of protons. The symbol is .

Consider This 2.6

If you omitted the left-hand subscripts on your isotope symbol representations, would they still be correct?

Page 72

Practice Problem 2.6

Write the isotope symbols for ions that have the following numbers of protons, neutrons, and electrons.

(a) 16 protons, 18 neutrons, and 18 electrons

(b) 11 protons, 12 neutrons, and 10 electrons

Further Practice: Questions 2.51 and 2.52 at the end of the chapter

2.4 Atomic Mass

How can we describe the mass of an atom of an element? Although single atoms cannot be weighed on a balance, modern techniques such as mass spectrometry (Figure 2.15) can be used to determine individual atomic masses accurately. For example, the mass of a single hydrogen-1 atom is 1.67380 × 10⁻²⁴ g. The mass of a carbon-12 atom is 1.99272 × 10⁻²³ g, about 12 times the mass of a hydrogen-1 atom. Numbers as small as these are difficult to remember and use. It is more convenient to think of a carbon-12 atom as being about 12 times the mass of a hydrogen-1 atom. For this reason, scientists have devised a method for expressing masses of atoms in a more convenient way. They use the atomic mass unit (amu). This mass scale uses carbon-12 (¹²C), the most abundant isotope of carbon, as the standard to which all other atoms are compared. Carbon-12 is assigned an atomic mass of exactly 12 atomic mass units, or 12 amu. One atomic mass unit (amu) is equal to one-twelfth the mass of a carbon-12 atom:

The term relative atomic mass is used by the International Union of Pure and Applied Chemistry (IUPAC), the worldwide governing body that determines official names in chemistry. Relative atomic mass is used to describe the weighted average of the atomic masses of the isotopes that compose an element as it is found in nature. This term is synonymous with the older term atomic weight, which is still commonly used.

Atomic mass units make it easy to compare masses of atoms. For example, if a carbon atom has a mass 12 times that of a hydrogen atom, then the mass of the hydrogen atom is about 1 amu. A hydrogen-2 atom is one-sixth the mass of the carbon-12 atom, so it has a mass of approximately 2 amu. A hydrogen-3 atom is one-fourth the mass of a carbon atom, so its mass is approximately 3 amu.

Note that there are three different, naturally occurring isotopes of hydrogen, each with a different mass. How do we describe the mass of the atoms in a sample of an element if it is composed of isotopes of different masses? From mass spectrometry (Figure 2.15), we can find the mass of each isotope and measure how much of each is present in a sample. This allows us to take a weighted average. The average of the mass of the individual isotopes, taking into account the naturally occurring relative abundance of each, is the relative atomic mass of the element. The relative atomic mass of hydrogen, to four significant figures, is 1.008 amu. The relative atomic mass of hydrogen is closer to that of hydrogen-1 than the other isotopes because the hydrogen-1 isotope is present in the largest abundance (99.99%) in Earth’s crust and atmosphere. The relative atomic mass of carbon is 12.01 amu because carbon-12 is most abundant (98.93%), with smaller amounts of carbon-13 and carbon-14.

FIGURE 2.15 A mass spectrometer measures the masses of individual atoms. Electrons are removed from atoms (or molecules) to produce ions that are accelerated through a magnetic field. The degree of bending in the path of the ions is related to the mass and the charge of the ions. Mass spectrometry is also used to determine the relative amounts of different isotopes in a sample of an element.

Page 73To find the relative amounts of each isotope that compose an element using a mass spectrometer, samples of that element from many parts of the world are analyzed. For example, suppose a sample of neon is injected into a mass spectrometer (Figure 2.15). The detector indicates that there are three isotopes of neon: neon-20, neon-21, and neon-22. The detector also shows that the isotopes are present in different amounts. The percent abundance of each isotope is determined and an average is taken. The percent abundances from many different locations are similar.

An example of a weighted average is your grade point average (GPA). Commonly an A grade is worth 4 points, a B grade is worth 3 points, etc. The GPA is calculated by multiplying 4 by the fraction of course credits that were earned with an A, multiplying 3 by the fraction of course credits that were earned with a B, etc. Then these numbers are added to obtain the GPA, which is the average of the grade points weighted by the fraction of credits that received each grade.

MATH
TOOLBOX
1.2

On the periodic table on the inside front cover of this book, the relative atomic mass of each element is listed below its element symbol. Elements such as hydrogen and carbon have relative atomic mass values very close to the mass of a particular isotope. Others, like silver (Ag), have relative atomic mass values that are averages based on significant contributions from more than one isotope. For example, natural silver has a relative atomic mass of 107.9 amu (Figure 2.16). At first it may seem that natural silver is composed mostly of the isotope ¹⁰⁸Ag. However, silver has been found to consist of two naturally occurring isotopes with different percent abundances: 51.82% ¹⁰⁷Ag and 48.18% ¹⁰⁹Ag. The exact mass of ¹⁰⁷Ag is measured with a mass spectrometer to be 106.9051 amu, and the exact mass of ¹⁰⁹Ag is 108.9048 amu. To obtain the relative atomic mass of silver, we calculate the weighted average using the masses of Ag-107 and Ag-109. To do this, we multiply the mass of each isotope by its relative abundance, expressed in decimal form (percent divided by 100). This gives the mass contribution from each isotope. Summing the mass contributions gives the weighted average, which is the relative atomic mass listed on the periodic table:

Isotope mass × abundance = mass contribution from isotope

¹⁰⁷Ag 106.9051 amu × 0.5182 = 55.40 amu

¹⁰⁹Ag 108.9048 amu × 0.4818 = 52.47 amu

107.87 amu (relative atomic mass of Ag)

FIGURE 2.16 The relative atomic mass of each element is indicated on the periodic table in this book just below the element symbol.

The relative atomic mass of silver listed on the periodic table is 107.9 amu, rounded to four significant figures.

If only two major isotopes comprise an element, we can usually determine which isotope is most abundant by referring to the relative atomic mass of the element on the periodic table. This is shown in Example 2.7.

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EXAMPLE 2.7Relative Atomic Mass

Naturally occurring chlorine consists of ³⁵Cl (34.9689 amu) and ³⁷Cl (36.9659 amu). Which isotope is most abundant?

Solution:

The relative atomic mass of chlorine is 35.45 amu, as shown on the periodic table. Since the relative atomic mass is closer to the mass of ³⁵Cl (34.9689 amu) than the mass of ³⁷Cl (36.9659 amu), the ³⁵Cl isotope must be most abundant. (The actual percentages are 75.77% ³⁵Cl and 24.23% ³⁷Cl.)

Consider This 2.7

If, on a different planet, the relative atomic mass of chlorine were 36.31 amu with the same two isotopes, which isotope would be more abundant?

Practice Problem 2.7

The copper mined from Earth’s crust consists of ⁶³Cu (62.93 amu) and ⁶⁵Cu (64.93 amu). Which isotope is most abundant?

Further Practice: Questions 2.75 and 2.76 at the end of the chapter

Example 2.8 gives you practice calculating the relative atomic mass for an element given the masses and percent abundances of its isotopes.

EXAMPLE 2.8Calculating Relative Atomic Mass

The element magnesium is composed of three isotopes: magnesium-24, magnesium-25, and magnesium-26. The mass and percent abundance of each are listed in the following table. Calculate the relative atomic mass of magnesium, and compare your calculated value to the value in the periodic table.

Isotope	Mass (amu)	Natural Abundance (%)
²⁴Mg	23.985	78.99
²⁵Mg	24.986	10.00
²⁶Mg	25.983	11.01

MATH
TOOLBOX
1.2

Solution:

To determine the mass contribution of each isotope, multiply its mass by its percent abundance in decimal form. To find the relative atomic mass, total the mass contributions of the isotopes:

²⁴Mg 23.985 amu × 0.7899 = 18.95 amu

²⁵Mg 24.986 amu × 0.1000 = 2.499 amu

²⁶Mg 25.983 amu × 0.1101 = 2.861 amu

24.31 amu (calculated relative atomic mass)

The calculated relative atomic mass, 24.31 amu, is the same as that on the periodic table to four significant figures.

Consider This 2.8

What could account for a relative atomic mass of magnesium that was less than 24.31 amu on a different planet?

Practice Problem 2.8

The element lithium is composed of two isotopes, lithium-6 and lithium-7. The mass and percent abundance of each is listed in the following table. Calculate the relative atomic mass of lithium, and then compare your answer to the value listed on the periodic table.

Isotope	Mass (amu)	Natural Abundance (%)
⁶Li	6.01512	7.793
⁷Li	7.01600	92.21

Further Practice: Questions 2.77 and 2.78 at the end of the chapter

The mass of an isotope, as measured by mass spectrometry, is similar in value to its mass number. A magnesium-24 atom, for example, has a mass of 23.985 amu. Its mass number, 24, is a whole number that is close to its atomic mass. RememberPage 75 that the mass number is not an actual mass; it is a count of the number of particles in the nucleus. The mass of an atom must be measured. Because the mass number and the atomic mass of an isotope are always similar in value, for convenience we often approximate the mass of an isotope as a whole number that is also its mass number. The approximate mass of a magnesium-24 atom, therefore, is about 24 amu. Its more exact mass (23.985 amu) is measured experimentally and is provided in published data tables.

We use relative atomic mass when we refer to an “average atom” of an element, or a group of atoms. For example, the mass of 100 magnesium atoms can be determined by multiplying 100 atoms by the relative atomic mass of magnesium, 24.31 amu per atom, to get a total mass of 2431 amu. What is the mass of 1000 lithium atoms?

2.5 The Periodic Table

As Anna continued her historical research, she read about the Russian scientist Dmitri Mendeleev. He developed and published the basic arrangement of the periodic table between 1869 and 1871 (Figure 2.17), decades before the discovery of protons. Mendeleev did not arrange the elements in order of atomic number. He arranged the 63 known elements in order of increasing relative atomic mass, and he grouped elements with similar properties into columns and rows so that the properties of the elements varied in a regular pattern (periodically). Once he arranged all of the elements that were known, he predicted the existence and properties of three elements that were unknown at the time: gallium (Ga), scandium (Sc), and germanium (Ge). He also placed two pairs of elements (Co/Ni and Te/I) in an order that did not match their relative atomic masses. Using what he thought was the correct order based on the properties of these elements, Mendeleev thought that their calculated relative atomic mass values must be in error. We now know that the periodicity of the properties of the elements occurs in the order of atomic number (number of protons), not relative atomic mass.

FIGURE 2.17 Mendeleev based his original periodic table on the elements that were known at the time. He used it to predict the existence of other, then unknown, elements. The numbers accompanying the element symbols are relative atomic mass values that were known at the time.

Classification of Elements

The modern version of the periodic table, in which all the known elements are arranged in columns and rows to emphasize periodic properties, is shown in Figure 2.18 and on the Periodic Table. In the periodic table, the elements are in the order of their atomic number (number of protons) from smallest to largest. There are many ways to use the classification system the periodic table provides. One way is to look at elements in the same vertical column. Because they have similar properties, we call them a family or a group. Each group has a Roman numeral (I through VIII) and a letter (A or B) associated with it. A newer system (shown in parentheses after the Roman numeral/letter combination) uses only Arabic numbers (1 through 18) to designate each group.

Page 76

FIGURE 2.18 The periodic table helps us to classify elements in a variety of ways.

Metalloids such as silicon and germanium are commonly used in semiconductor devices.

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A horizontal row of elements in the periodic table is a period. Elements in the same period have properties that tend to vary in a regular fashion. Periods are labeled with numbers (1 through 7). Figure 2.19 shows a picture of elements in period 3 and in group VA (15). Note the gradual variation in the physical appearance of these elements.

FIGURE 2.19 The physical appearance of elements close to each other in the same group or period are similar. Their properties vary gradually from one end of the group or period to the other.
©McGraw-Hill Education/Stephen Frisch

Look for the stair-step line that begins at boron and moves down the periodic table to astatine. This line separates the metals (to its left) from the nonmetals (to its right). (The properties of metals and nonmetals were discussed in Chapter 1.) A metalloid, or semimetal, is an element that has physical properties resembling a metal but chemical reactivity more like that of a nonmetal. The metalloids, identified by green boxes, lie along the stair-step line. Some metals, nonmetals, and metalloids are shown in Figure 2.20.

FIGURE 2.20 The metals, on the left side of the periodic table, have a shiny appearance. They are malleable and ductile (can be pulled into thin wires). They conduct heat and electricity. The nonmetals, on the right side of the periodic table, are gases, liquids, or powder or crystalline solids. They are insulators (nonconductors of heat and electricity). The metalloids have properties between the metals and the nonmetals.
©McGraw-Hill Education/Stephen Frisch

Any element in one of the eight groups labeled with the letter A is a main-group element, or a representative element. An element in any of the 10 groups labeled with the letter B is a transition metal. The 14 members of the lanthanide and actinide series of elements are usually placed on separate lines at the bottom of the table to conserve space, but they belong immediately after elements La (Z = 57) and Ac (Z = 89), respectively. Any lanthanide or actinide element is an inner-transition metal.

Several of the groups have descriptive names that are frequently used instead of their group numbers. Any member of group IA (1), except hydrogen, is an alkali metal. Any element in group IIA (2) is an alkaline earth metal. An element in group VIIA (17) is a halogen. On the far right are the group VIIIA (18) elements. An element in this group is called a noble gas. These names remind us that elements in a group display similar properties. The alkali metals, for example, are said to be reactive. This means they react readily with many other elements and compounds, including water (Figure 2.21). The alkaline earth metals are less reactive than the alkali metals, but are more reactive than most of the transition metals.

FIGURE 2.21 Potassium, an alkali metal, reacts violently with water, producing flammable hydrogen gas.
©McGraw-Hill Education/Stephen Frisch

ANIMATION: Properties of Alkali and Alkaline Earth Materials

Page 78

The halogens are similar in several ways. One is that they occur as elements in the form of diatomic molecules. A diatomic molecule is a molecule consisting of two atoms. When we represent the halogens as elements, we use a subscript to show that they are diatomic molecules: F₂, Cl₂, Br₂, I₂. Other elements that exist naturally as diatomic molecules are hydrogen (H₂), oxygen (O₂), and nitrogen (N₂). Elements that occur as diatomic molecules are shown in Figure 2.22.

FIGURE 2.22 Seven elements occur as diatomic molecules.

The noble gases are unique in that they exist naturally in their elemental form as single atoms. They are said to be inert; that is, they do not react chemically with other elements or compounds. The noble gases were not known until the late 1800s, although many of them rank as highly abundant elements in the universe and in Earth’s atmosphere. As you will learn later, the unique stability of the noble gases helps us understand the chemical reactivity of other elements.

EXAMPLE 2.9Classification of Elements

Classify each of the following elements by group number, group name (if applicable), and period, and as a metal, nonmetal, or metalloid.

(a) sodium (b) silicon (c) bromine (d) copper

Solution:

(a) Na is in group IA (1), the alkali metal group, and in period 3, and is a metal.

(b) Si is in group IVA (14) and in period 3, and is a metalloid.

(d) Cu is in group IB (11), a transition metal group, and in period 4, and is a metal.

Consider This 2.9

What element is the only nonmetal located on the left side of the period table?

Practice Problem 2.9

Identify the element that is described.

(a) the element that is in group VA (15) and in period 2, and is a nonmetal

(b) the element that is a noble gas and in period 3, and is a nonmetal

(d) the element that is in group IB (11), is a transition metal, and is in period 5

Further Practice: Questions 2.105 and 2.106 at the end of the chapter

The alkali metals (group IA) got their name because they react with water to form an alkaline (basic) solution. The word alkali comes from Arabic and means “ashes,” the material that was originally used to make alkaline solutions. The name for the alkaline earth metals (group IIA) is derived from their presence in metal oxides, which were once called earths.

Astatine (At), a halogen, is found in very small amounts in Earth’s crust. It is radioactive and very short lived. Since astatine has properties most similar to iodine, scientists believe astatine also exists as diatomic molecules.

Ions and the Periodic Table

In Section 2.3, we saw that ions have charges because they have a different number of electrons than there are protons in the nucleus. Megan noticed that some of the essential nutritional elements are ions that have the same charge. For example, sodium and potassium both exist as ions with a 1+ charge. Similarly, calcium and magnesium exist as 2+ ions. She found out that the pattern is no coincidence. The position of an element in the periodic table helps us to predict the charge on its ion. To understand this trend, let’s start by looking at the elements that do not form ions—the noble gases, group VIIIA (18).

The noble gases are the most stable—the least reactive—of all the elements. Their stability is associated with the number of electrons they contain. To achieve a similar stability, many atoms of the main-group elements gain or lose electrons to form ions with the same electron count as the nearest noble gas. The nonmetals Page 79usually gain electrons to form anions that have a noble gas electron count. Most main-group metals lose electrons to form cations that have a noble gas electron count. For example, a sodium atom, with 11 electrons, forms an Na⁺ ion by losing one electron. When it does, it has the same number of electrons as neon, 10. How many electrons does a magnesium atom lose when it forms an ion? When oxygen forms O²⁻ ions, each atom, with eight electrons, gains two electrons for a total of 10, like neon. How many electrons will a nitrogen atom gain to form an ion?

Because elements in the same group have an electron count that differs from that of the nearest noble gas by the same number, elements in the same group often form ions of the same charge. Like sodium, all group IA (1) elements lose one electron to form a cation with a charge of 1+. Group IIA (2) elements lose two electrons to form cations with a charge of 2+. Aluminum atoms of group IIIA (3) lose three electrons to form 3+ ions. Group VIIA (17) elements gain one electron to form anions with a charge of 1−. Group VIA (16) elements gain two electrons to form anions with a charge of 2−. Nitrogen and phosphorus of group VA (15) gain three electrons to form anions with a charge of 3−. The charges for ions that can be predicted from the periodic table are labeled on the abbreviated periodic table in Figure 2.23. Most of the other elements form two or more different ions of different charges. The charges of these ions cannot be predicted from the periodic table.

FIGURE 2.23 Many of the main-group elements form ions that have charges that can be predicted by their positions in the periodic table. The ions shown here are sometimes called common ions because these elements rarely form ions of other charges.

Although no compounds of noble gases exist naturally, a few stable compounds of krypton and of xenon have been prepared. A few compounds of argon are known, but they are stable only at low temperatures. This picture shows one synthetic compound of a noble gas, XeF₂, adhering to the interior of the reaction vessel.

©Gary J. Schrobilgen/McMaster University

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EXAMPLE 2.10Predicting Charges on Ions

Write the symbol for the ion that each of the following elements is predicted to form.

(a) magnesium (b) bromine (c) nitrogen

Solution:

These ions can be predicted by their positions in the periodic table.

(a) Magnesium is in group IIA (2), so it will lose two electrons to form Mg²⁺, giving it the same number of electrons as neon.

(b) Bromine is in group VIIA (17), so it will gain one electron to form Br⁻, giving it the same number of electrons as krypton.

(c) Nitrogen is in group VA (15), so it will gain three electrons to form N³⁻, giving it the same number of electrons as neon.

Consider This 2.10

Based on its position on the periodic table, what charge would you predict for the hydrogen ion?

Practice Problem 2.10

Write the symbol for the ion that each of the following elements is predicted to form.

(a) lithium (b) sulfur (c) aluminum

Further Practice: Questions 2.105 and 2.106 at the end of the chapter

In our bodies, ions usually exist dissolved in body fluids. Ions can also exist as components of compounds. In Chapter 3 you will learn about these ionic compounds. You will learn about their properties and how to name them. You will also learn how to determine charges on ions that cannot be predicted from the periodic table.

CHAPTER REVIEW

KEY CONCEPTS

Our current understanding of the atom began with Dalton’s atomic theory two centuries ago. It was not until the early 1900s that there was experimental evidence that showed atoms consist of protons, neutrons, and electrons.

°Protons are the positively charged particles in the nucleus. The number of protons in an atom is unique to each element and is called the atomic number (Z). The atomic number is listed above the element symbol on the periodic table in this book.

°Neutrons, also in the nucleus, have no charge. Atoms of an element can differ in the number of neutrons (N), which results in isotopes. The symbol for an isotope includes the mass number (A), which is the sum of the numbers of protons and neutrons. The mass number is written as a left-hand superscript on the element symbol to distinguish the isotopes of an element.

°Electrons are negatively charged and exist in the relatively large space outside the nucleus. In a neutral atom, the number of electrons is equal to the number of protons.

Ions differ from atoms in that they have a positive or negative charge.

°The charge on an ion is written as a right-hand superscript on the element symbol.

°A positively charged ion, called a cation, contains fewer electrons than protons.

°A negatively charged ion, called an anion, contains more electrons than protons.

°When an ion is formed from a neutral atom, electrons are lost or gained, resulting in the net charge on the ion.

°The properties of ions differ from elements containing neutral atoms.

The mass of an atom is usually reported in atomic mass units (amu).

°The mass of an atom, usually known to many decimal places, is not the same as the mass number, although they are similar in value.

°The relative atomic mass is the weighted average of the masses of all the isotopes of an element, taking into account their relative abundances in a naturally occurring sample. This value is listed under the element symbol on the periodic table in this book.

Page 81•The periodic table is organized by elements with similar properties, which allows us to determine or predict certain properties of elements.

°Elements in the same vertical column are in the same group and display similar chemical properties. For example, elements in group IA, the alkali metals, all form ions with a 1+ charge.

°Elements can also be classified as metals, nonmetals, or metalloids, or as main-group elements, transition metals, lanthanides, or actinides.

KEY TERMS

alkali metal (2.5)

alkaline earth metal (2.5)

anion (2.3)

atomic mass unit (amu) (2.4)

atomic number (2.2)

cation (2.3)

Dalton’s atomic theory (2.1)

diatomic molecule (2.5)

inner-transition metal (2.5)

ion (2.3)

isotope (2.2)

isotope symbol (2.2)

law of conservation of mass (2.1)

law of definite proportions (2.1)

main-group element (2.5)

relative atomic mass (2.4)

subatomic particle (2.2)

transition metal (2.5)

QUESTIONS AND PROBLEMS

The following questions and problems, except those in Additional Questions and Concept Review Questions, are paired. Questions in a pair focus on the same concept. Answers to selected questions and problems are in Appendix D.

Matching Definitions with Key Terms

2.1Match the key terms with the descriptions provided.

(a)	an electrically neutral subatomic particle
(b)	a law stating that the mass of the substances produced in a reaction equals the mass of the substances that reacted
(c)	a positively charged subatomic particle
(d)	a member of one of the A groups of elements in the periodic table
(e)	the average mass of an atom of an element, taking into account the masses and abundances of all the naturally occurring isotopes
(f)	the sum of the protons and neutrons in an atom
(g)	an atom of an element with a specific number of neutrons
(h)	an ion with a positive charge
(i)	any particle found within an atom
(j)	any element in group IA (1)
(k)	a chart of all the known elements ordered by increasing atomic number and arranged in columns and rows to emphasize similar properties

2.2Match the key terms with the descriptions provided.

(a)	a member of one of the B groups of elements in the periodic table
(b)	a law stating that all samples of a pure compound always contain the same proportions of the component elements
(c)	a negatively charged subatomic particle
(d)	an ion with a negative charge
(e)	a neutral particle consisting of two bound atoms
(f)	any element in group VIIIA (18)
(g)	a set of elements in the same horizontal row in the periodic table
(h)	the number of protons in an atom
(i)	the unit of mass used to describe single atoms or small numbers of atomic particles, which is equal to one-twelfth the mass of a carbon-12 atom
(j)	a set of elements in the same vertical column in the periodic table
(k)	any element in group IIA (2)
(l)	the central core of the atom, which contains the protons and neutrons, and most of the mass of the atom

Dalton’s Atomic Theory

2.3Which laws did Dalton use to argue that matter consists of atoms?

2.4What modern technique allows us to “see” the surface of atoms?

2.5How does Dalton’s atomic theory describe atoms of different elements?

2.6How does Dalton’s atomic theory explain the law of conservation of mass?

2.7How does Dalton’s atomic theory explain the law of definite proportions?

2.8How does Dalton’s atomic theory explain the formation of compounds from elements?

Page 822.9Does the following diagram represent a chemical reaction that obeys the law of conservation of mass? If it does not, how should the diagram be modified to represent the reaction correctly?

2.10Does the following diagram represent a chemical reaction that obeys the law of conservation of mass? If it does not, how should the diagram be modified to represent the reaction correctly?

Structure of the Atom

2.11Which experiment showed that all atoms contain negatively charged particles with masses far less than the mass of a hydrogen atom?

2.12What information about the structure of the atom did Rutherford’s gold foil experiment provide?

2.13Which subatomic particles are negatively charged?

2.14Which subatomic particles are positively charged?

2.15Helium is used in balloons and blimps because of its lifting power. Draw a nuclear model of a helium atom with a mass number of 4 (helium-4). Show the locations of the nucleus, protons, neutrons, and electrons. Indicate the number of each type of subatomic particle.

2.16Hydrogen-3 (tritium) is used as a tracer in chemical, biological, and medical research. Draw a nuclear model of a hydrogen atom with a mass number of 3 (hydrogen-3). Show the locations of the nucleus, protons, neutrons, and electrons. Indicate the number of each type of subatomic particle.

2.17Which subatomic particle has approximately the same mass as a proton?

2.18Which of the following best describes about how many times greater in mass a proton is than an electron: 200, 2000, or 20,000?

2.19Explain why the mass of a carbon atom is about twice the mass of the protons in a carbon atom. What accounts for the difference in mass?

2.20Which subatomic particle was discovered last? Why was it difficult to detect?

2.21What is the atomic number for each of the following elements?

(a) hydrogen

(b) oxygen

2.22How many protons are in each atom of each of the following elements?

(a) helium

(b) argon

2.23The number of which subatomic particle determines the identity of an element?

2.24How do isotopes of an element differ?

2.25What information do you need to determine the atomic number of an atom?

2.26What information do you need to determine the mass number of an atom?

2.27Which of the following are always the same for atoms of the same element?

(a) mass number

(b) atomic number

(d) mass of an atom

2.28Which of the following are different for isotopes of an element?

(a) mass number

(b) atomic number

(d) mass of an atom

2.29Given the composition of the nuclei, what are the atomic number, neutron number, and mass number of each isotope of hydrogen shown?

2.30Identify the element from the atom shown. What are the atomic number, neutron number, and mass number of this isotope?

2.31What are the atomic number, neutron number, and mass number of the following isotopes?

(a)

(b)

(c)

2.32Identify the element from the atom shown. What are the atomic number, neutron number, and mass number of this isotope?

2.33How many protons, neutrons, and electrons are in each of the following?

(a) an oxygen atom with a mass number of 15

(b) a silver atom with a mass number of 109

2.34How many protons, neutrons, and electrons are in each of the following?

(a) a hydrogen atom with a mass number of 1

(b) a magnesium atom with a mass number of 26

2.35What is the isotope symbol for atoms that contain the following numbers of subatomic particles?

(a) 1 proton, 1 electron, and 2 neutrons

(b) 4 protons, 4 electrons, and 5 neutrons

2.36What is the isotope symbol for atoms that contain the following numbers of subatomic particles?

(a) 2 protons and 1 neutron

(b) 47 protons and 62 neutrons

2.37How many protons and neutrons are in an atom represented by the following?

(a)

(b) ³⁹K

2.38How many protons and neutrons are in an atom represented by the following?

(a)

(b) ⁶⁸Zn

2.39A study of blood flow in heart muscle uses nitrogen-13 incorporated into a chemical agent. How many protons and neutrons are found in the N-13 nucleus?

2.40Phosphorus-32 can be used to treat eye tumors. How many protons and neutrons are found in the P-32 nucleus?

2.41Complete the following table for the designated atoms.

Isotope Symbol	Number of Protons	Number of Neutrons	Number of Electrons

	25	31
		10	8

2.42Complete the following table for the designated atoms.

Isotope Symbol	Number of Protons	Number of Neutrons	Number of Electrons

	12	11
		14	16

Ions

2.43How do an atom and an ion of the same element differ?

2.44What changes when an ion forms from an atom, or when an atom forms from an ion?

2.45What forms when a neutral atom undergoes each of the following changes?

(a) It gains one electron.

(b) It loses two electrons.

2.46What forms when an ion with a 1+ charge undergoes each of the following changes?

(a) It gains one electron.

(b) It loses two electrons.

2.47Write the symbol for the ion that is formed after each of the following changes. Identify each as a cation or an anion.

(a) A zinc atom loses two electrons.

(b) A phosphorus atom gains three electrons.

2.48Write the symbol for the ion that is formed after each of the following changes. Identify each ion as a cation or an anion.

(a) A selenium atom gains two electrons.

(b) A mercury atom loses two electrons.

2.49How many protons and electrons are found in each of the following?

(a) Zn²⁺

(b) F⁻

2.50How many protons and electrons are found in each of the following?

(a) P³⁻

(b) Al³⁺

Page 842.51Complete the following table for the designated ions.

Isotope Symbol	Number of Protons	Number of Neutrons	Number of Electrons

	12	13	10
	7	6	10

2.52Complete the following table for the designated ions.

Isotope Symbol	Number of Protons	Number of Neutrons	Number of Electrons

	38	50	36
	1	1	2

2.53What element has 18 electrons when it forms a cation with a 1+ charge?

2.54What element has 18 electrons when it forms an anion with a 2− charge?

2.55What element has 27 electrons when it forms a cation with a 2+ charge?

2.56What element has 46 electrons when it forms a cation with a 1+ charge?

2.57How do ⁷Li⁺ and ⁶Li each differ from a neutral lithium-7 atom? Which differs from lithium-7 by the greatest mass?

2.58How do and differ in their numbers of subatomic particles? Which should have the greater mass?

2.59Potassium citrate is a compound found in some sparkling beverages. The potassium in this compound is a cation with a 1+ charge. How many protons are found in the nucleus of a potassium ion? How many electrons surround this nucleus?

2.60Calcium citrate is a compound found in some calcium supplement medications. The calcium in this compound consists of ions containing 18 electrons. What is the charge of the calcium ions? How many protons are found in the nucleus of a calcium ion?

Atomic Mass

2.61What is the basis for the atomic mass unit (amu) scale?

2.62What is the mass of a carbon-12 atom on the amu scale? Is this value exact or approximate?

2.63What is the approximate mass, in atomic mass units, of the following isotopes?

(a)

(b)

2.64Estimate the combined mass (in atomic mass units) of 10 cobalt-59 atoms.

2.65Approximately how much greater in mass is a D₂ molecule than an H₂ molecule?

2.66Approximately how much greater in mass is a D₂O molecule than an H₂O molecule? (Assume the oxygen atoms have a mass number of 16.)

2.67Approximately how much greater in mass is a krypton-80 atom than an argon-40 atom?

2.68Approximately how much greater in mass is a magnesium-24 atom than a carbon-12 atom?

2.69Why do we use the amu mass scale instead of the gram mass scale when discussing masses of atoms?

2.70What is the relationship between grams and atomic mass units? How many atomic mass units are in 1 g?

2.71What is the difference between the mass of an atom and the mass number of an atom?

2.72Which isotope has exactly the same mass (in atomic mass units) as its mass number?

2.73How is the mass of an individual atom determined? How do you determine the mass number of an atom?

2.74Which of the following is most likely the mass of a nickel-62 atom: 62.0000000 amu, 62.495654 amu, 61.5871338 amu, or 61.9283461 amu. Why?

2.75Naturally occurring calcium is composed of two isotopes: calcium-40 and calcium-44. Which isotope is most abundant?

2.76Naturally occurring silicon is composed mostly of one isotope. Write its symbol.

2.77An unknown element (X) discovered on a planet in another galaxy was found to exist as two isotopes. Their atomic masses and percent abundances are listed in the following table. What is the relative atomic mass of the element?

Isotope	Mass (amu)	Natural Abundance (%)
²²X	21.995	75.00
²⁰X	19.996	25.00

2.78Suppose that the isotope abundance of magnesium on another planet has been determined to be different from that on Earth, as shown in the following table. Calculate the relative atomic mass of magnesium on this planet.

Isotope	Mass (amu)	Natural Abundance (%)
²⁴Mg	23.985	20.00
²⁵Mg	24.985	20.00
²⁶Mg	25.983	60.00

Page 852.79The mass spectrum of nickel is shown.

(a)	What is the most abundant isotope of nickel?
(b)	What is the least abundant isotope of nickel?
(c)	Will the relative atomic mass of nickel be closer to a value of 58 or of 60?
(d)	List the number of protons, electrons, and neutrons in each of the ions shown in the mass spectrum.

2.80The mass spectrum of magnesium is shown.

(a)	What is the most abundant isotope of magnesium?
(b)	What is the least abundant isotope of magnesium?
(c)	Will the relative atomic mass of magnesium be closer to a value of 24 or 25 or 26?
(d)	List the number of protons, electrons, and neutrons in each of the ions shown in the mass spectrum.

2.81What is the mass in amu of 1000 boron atoms?

2.82What is the mass in amu of 1000 mercury atoms?

2.83Which contains more atoms, 2500 amu of boron atoms or 2500 amu of mercury atoms?

2.84A sample of pure silver and a sample of pure gold have the same mass. Which contains the greatest number of atoms?

The Periodic Table

2.85Identify which of the elements Br, K, Mg, Al, Mn, and Ar can be classified in each of the following ways.

(a)	alkali metal
(b)	halogen
(c)	transition metal
(d)	alkaline earth metal
(e)	noble gas
(f)	main-group element

2.86Identify which of the elements He, Zn, Pb, I, Ca, and Na can be classified in each of the following ways.

(a)	alkali metal
(b)	halogen
(c)	transition metal
(d)	alkaline earth metal
(e)	noble gas
(f)	main-group element

2.87Name the element that is a halogen in period 3.

2.88Name the element that is an alkaline earth metal in period 5.

2.89What element is in group IVB (4) and in period 4?

2.90What element is in group IVA (14) and in period 2?

2.91Identify each of the following elements as a metal, nonmetal, or metalloid.

(a)	calcium
(b)	carbon
(c)	potassium
(d)	silicon

2.92Identify each of the following elements as a metal, nonmetal, or metalloid.

(a)	phosphorus
(b)	chromium
(c)	arsenic
(d)	sodium

2.93Identify the following elements as a main-group element, transition metal, lanthanide, or actinide.

(a)	oxygen
(b)	magnesium
(c)	tin
(d)	uranium
(e)	chromium

2.94Identify the following elements as a main-group element, transition metal, lanthanide, or actinide.

(a)	sulfur
(b)	iron
(c)	plutonium
(d)	calcium
(e)	xenon

2.95In which group of the periodic table do all the elements occur as diatomic molecules?

2.96Seven elements occur as diatomic molecules. Write molecular formulas for each.

2.97Which of the following elements does not occur as a diatomic molecule: nitrogen, fluorine, or neon?

2.98Which of the following elements does not occur as a diatomic molecule: sulfur, hydrogen, or oxygen?

2.99In which group of the periodic table do all the elements exist naturally as gases of uncombined atoms?

2.100The noble gases are sometimes called “inert gases.” Why? Why do they not form ions?

2.101The ions of many of the main-group elements have the same number of ________ as the noble gas nearest to them in the periodic table.

2.102In which group of the periodic table do all the elements not form ions?

2.103Identify the groups of the periodic table in which all the elements form ions of the charge indicated.

(a)	1+ cations
(b)	2+ cations
(c)	1− anions
(d)	2− anions

2.104For each group listed, identify the charge on the ions that they form.

(a)	group VIA (16)
(b)	halogens
(c)	alkali metals
(d)	group IIA (2)

2.105For each element listed, write the symbol for the ion it forms.

(a)	sodium
(b)	oxygen
(c)	sulfur
(d)	chlorine
(e)	bromine

2.106For each element listed, write the symbol for the ion it forms.

(a)	nitrogen
(b)	phosphorus
(c)	magnesium
(d)	potassium
(e)	aluminum

2.107Sodium reacts vigorously with water to form hydrogen gas and a compound containing sodium ions. Which other elements are expected to react with water in a similar way?

2.108Chlorine gas reacts with potassium to form a compound with properties similar to table salt. Which other elements are expected to react with potassium in a similar way?

Additional Questions

2.109When iron rusts, a compound forms with the formula Fe₂O₃. When iron nails rust, should their mass increase, decrease, or remain the same? Explain your answer.

2.110Is the law of conservation of mass obeyed in the reaction shown? Explain.

2.111A 100-g sample of zinc sulfide contains 67.1 g zinc and 32.9 g sulfur. If a 1.34-g sample of zinc is heated with excess sulfur, 2.00 g of zinc sulfide form. Show how these data are in agreement with the law of definite proportions.

2.112A 100-g sample of HgO contains 92.6 g mercury and 7.4 g oxygen. What mass of mercury atoms and what mass of oxygen atoms are contained in any 100-g sample of HgO?

2.113What properties of electrons contributed to why they were the first subatomic particles to be discovered?

2.114What experimental evidence led to each of the following advances in our understanding of atomic structure?

(a)	discovery of the electron and its charge-to-mass ratio
(b)	determination of the mass and charge of the electron
(c)	the nuclear model of the atom

2.115Which isotope has a mass number of 60 and an atomic number of 28?

2.116Why is it impossible for an element to have a mass number less than its atomic number?

2.117How many protons and neutrons are in a potassium-39 atom?

2.118What number of each type of subatomic particle is found in most H⁺ ions?

2.119Explain why the relative atomic mass of cobalt is greater than that of nickel, although the atomic number of nickel is one greater than that of cobalt.

2.120Which element has an isotope that has a mass approximately three times greater than its most abundant isotope?

2.121Some tables list relative atomic mass to as many significant figures as can be reported accurately. The relative atomic mass of argon, 39.948 amu, is known to five significant figures, while that of fluorine, 18.9984032 amu, is known to nine significant figures. Natural samples of argon contain three isotopes, but natural samples of fluorine contain only one. Explain why their relative atomic masses differ in the number of significant figures reported.

2.122How many atoms are in each of the following samples of carbon-12?

(a)	120 amu
(b)	12,000 amu
(c)	7.22 × 10²⁴ amu

2.123Naturally occurring iodine is composed of only one isotope. What is its mass number?

2.124How many atoms are in each of the following samples of naturally occurring iodine? (Iodine has only one naturally occurring isotope.)

(a) 127 amu

(b) 12,700 amu

Page 872.125A sample of pure carbon and a sample of pure iodine have the same mass. Which has the greatest number of atoms?

2.126Two equal-volume balloons contain the same number of atoms. One contains helium and one contains argon. Comment on the relative densities of the gases in these balloons.

2.127Naturally occurring boron comprises two isotopes, boron-10 and boron-11. The atomic mass of boron-10 is 10.013 amu. The atomic mass of boron-11 is 11.009 amu. Which of the following is the best estimate of the percent abundance of each isotope of boron? Why?

50.0% boron-10 and 50.0% boron-11

20.0% boron-10 and 80.0% boron-11

80.0% boron-10 and 20.0% boron-11

95.0% boron-10 and 5.0% boron-11

5.0% boron-10 and 95.0% boron-11

2.128Sodium metal reacts with water to form the compound sodium hydroxide and hydrogen gas. What is the formula for the hydrogen gas? Would you expect a similar reaction when potassium is added to water? How about copper or silver? Explain why.

2.129Bromine is a reddish-brown liquid at room temperature. Write the formula that represents liquid bromine.

2.130Nitrogen and oxygen are the main components of the air we breathe. Write formulas to represent these gases.

2.131Which element is in group IA (1) but is not an alkali metal?

2.132Which element is the only nonmetal in group IVA (14)?

2.133Why couldn’t we form a cation of an element by adding a proton to its atom?

2.134Compare and contrast Mendeleev’s periodic table with the modern periodic table.

2.135Complete the following table for the designated subatomic particles.

	Particle	Mass (g)	Relative Charge
(a)		1.6749 × 10⁻²⁴	0
(b)		9.1094 × 10⁻²⁸
(c)		1.6726 × 10⁻²⁴

2.136Atoms of the element shown in this photo contain 28 protons. What is the element?

2.137Positron emission tomography (PET) is used in medicine to produce an image of functions inside the body and can identify tumors. To make a tumor visible in the scan, a compound containing an isotope of a certain element is injected into the patient. The most common isotope used for PET scans is fluorine-18. How many protons, neutrons, and electrons are present in an atom of this isotope?

2.138When silver tarnishes, a chemical reaction occurs in which a compound of silver and sulfur forms on the surface of the silver metal. The silver in the compound with sulfur is in the form of an ion with a 1+ charge. In the process of going from a neutral Ag atom to an Ag⁺ ion, what happens to the number of protons and electrons?

2.139The iron found in the human body is in the form of Fe²⁺ or Fe³⁺ ions. How do neutral atoms of iron and the Fe²⁺ and Fe³⁺ ions differ?

2.140Which has the greater density, an ice cube containing the hydrogen-1 isotope or an ice cube containing the hydrogen-2 isotope (see Figure 2.13)? Why?

2.141Iodine is an essential trace element required by the body and is associated with the function of the thyroid gland. To which group and period on the periodic table does iodine belong?

2.142After he conducted cathode ray tube experiments proving the existence of negatively charged particles we now call electrons, Thomson proposed a model of the atom called the plum pudding model. Describe how Rutherford’s gold foil experiment disproved Thomson’s model.

2.143Analysis of carbon-14 content is used to date fossils up to about 60,000 years old. There are two other naturally occurring isotopes of carbon, ¹²C and ¹³C. Given the relative atomic mass of carbon, 12.01 amu, which is the most likely abundance of carbon-14 in nature: approximately 99%, approximately 33%, less than 0.1%?

2.144A naturally occurring isotope of potassium has a mass number of 40. One of argon’s naturally occurring isotopes also has a mass number of 40. How does an atom of ⁴⁰K differ from an atom of ⁴⁰Ar?

Concept Review Questions

2.145Which of the following has a positive charge?

A.	a sodium atom
B.	a sodium ion
C.	a proton
D.	an electron
E.	both a sodium ion and a proton

For those that do not have a positive charge, identify whether they have a negative charge or no charge.

2.146A monatomic ion with a 4+ charge has the same number of electrons as an argon atom. What is the formula for this ion?

A.	C⁴⁺
B.	Ti⁴⁺
C.	Si⁴⁺
D.	Ar⁴⁺
E.	Sc⁴⁺

For those ions that do not have the same number of electrons as argon, determine the number of electrons that they do have.

Page 882.147Which of the following has the greatest number of electrons?

A.	Ca²⁺
B.	K⁺
C.	Cl⁻
D.	S²⁻
E.	These all have the same number of electrons.

Which has the greater number of protons?

2.148Chlorine gas is a yellow-green gas that is highly toxic. What is the formula for chlorine gas?

Cl₂
Cl^–₂
Cl^2–₂
Cl^–
Cl

Which of these is the correct formula for chloride ion?

2.149Atoms that are isotopes of the same element differ in

A.	number of neutrons
B.	number of protons
C.	number of electrons
D.	charge
E.	atomic number

For each incorrect answer, identify a pair of substances that would differ in that way.

2.150Which of the following statements about atoms and subatomic particles is true?

A.	The nucleus of an atom has no mass.
B.	Neutrons are negatively charged.
C.	Electrons exist in the nucleus.
D.	The nucleus makes up most of the mass of an atom.
E.	A cation has more electrons than protons.

Modify each incorrect statement to make it correct.

2.151Elements in which group always form monatomic ions with a negative 1 (1−) charge?

A.	alkali metals
B.	alkaline earth metals
C.	halogens
D.	noble gases
E.	transition metals

For each incorrect group, identify the charge those elements would have when they form ions. Note if any groups do not form ions that are not predictable or if they do not form ions at all.

2.152Identify the number of protons and neutrons in an atom of the isotope ⁶⁴Zn.

A.	64 protons, 34 neutrons
B.	64 protons, 30 neutrons
C.	30 protons, 30 neutrons
D.	30 protons, 34 neutrons
E.	More information is needed.

For zinc-64, to which of the following do the values 30, 34, and 64 correspond: mass number, atomic number, neutron number?

2.153Naturally occurring gallium (Ga) is made of two isotopes: gallium-69 and gallium-71. Which of the following statements is true?

A.	Gallium’s relative atomic mass is 70.00 amu.
B.	Both isotopes have the same mass: 69.72 amu.
C.	The isotopes are present in the same percentages (50% and 50%).
D.	Gallium-71 is present in the largest percent abundance.
E.	Gallium-69 is present in the largest percent abundance.

Explain why the incorrect answers are false.

2.154Which of the following elements is classified as a nonmetal in period 3?

A.	potassium
B.	phosphorus
C.	bromine
D.	silicon
E.	krypton

For each incorrect answer, identify the period the element is in and whether the element is a metal, nonmetal, or metalloid.

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