ELECTRIC NATURE OF ATOMS - Atomic Structure and the Periodic Table of the Elements - REVIEW OF MAJOR TOPICS - SAT Subject Test Chemistry

SAT Subject Test Chemistry




Atomic Structure and the Periodic Table of the Elements

These skills are usually tested on the SAT Subject Test in Chemistry. You should be able to . . .

• Describe the history of the development of atomic theory.

• Explain the structure of atoms, their main energy levels, sublevels, orbital configuration, and the rules that govern how they are filled.

• Place atoms in groups and periods based on their atomic structure.

• Write formulas and names of compounds.

• Explain how chemical and physical properties are related to positions in the Periodic Table, including atomic size, ionic size, electronegativity, acid-forming properties, and base-forming properties.

• Explain the nature of radioactivity, the types and characteristics of each, and the inherent dangers.

• Identify the changes that occur in a decay series.

• Do the mathematical calculations to determine the age of a substance using its half-life.

This chapter will review and strengthen these skills. Be sure to do the Practice Exercises at the end of the chapter.

The idea of small, invisible particles being the building blocks of matter can be traced back more than 2,000 years to the Greek philosophers Democritus and Leucippus. These particles, considered to be so small and indestructible that they could not be divided into smaller particles, were calledatoms, the Greek word for indivisible. The English word atom comes from this Greek word. This early concept of atoms was not based upon experimental evidence but was simply a result of thinking and reasoning on the part of the philosophers. It was not until the eighteenth century that experimental evidence in favor of the atomic hypothesis began to accumulate. Finally, around 1805, John Dalton proposed some basic assumptions about atoms based on what was known through scientific experimentation and observation at that time. These assumptions are very closely related to what scientists presently know about atoms. For this reason, Dalton is often referred to as the father of modern atomic theory. Some of these basic ideas were:


Know Dalton”s five basic ideas about atoms.

1. All matter is made up of very small, discrete particles called atoms.

2. All atoms of an element are alike in weight, and this weight is different from that of any other kind of atom.

3. Atoms cannot be subdivided, created, or destroyed.

4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged.

By the second half of the 1800s, many scientists believed that all the major discoveries related to the elements had been made. The only thing left for young scientists to do was to refine what was already known. This came to a suprising halt when J. J. Thomson discovered the electron beam in a cathode ray tube in 1897. Soon afterward, Henri Becquerel announced his work with radioactivity, and Marie Curie and her husband, Pierre, set about trying to isolate the source of radioactivity in their laboratory in France.


Know Niels Bohr”s model based on a planetary model as opposed to quantum theory based on a probability model.

During the late nineteenth and early twentieth centuries, more and more physicists turned their attention to the structure of the atom. In 1913 the Danish physicist Niels Bohr published a theory explaining the line spectrum of hydrogen. He proposed a planetary model that quantized the energy of electrons to specific orbits. The work of Louis de Broglie and others in the 1920s and 1930s showed that quantum theory described a more probabilistic model of where the electrons could be found that resulted in the theory of orbitals


From around the beginning of the twentieth century, scientists have been gathering evidence about the structure of atoms and fitting the information into a model of the atomic structure.

Basic Electric Charges

The discovery of the electron as the first subatomic particle is credited to J. J. Thomson (England, 1897). He used an evacuated tube connected to a spark coil as shown in Figure 3. As the voltage across the tube was increased, a beam became visible. This was referred to as a cathode ray. Thomson found that the beam was deflected by both electrical and magnetic fields. Therefore, he concluded that cathode rays are made up of very small, negatively charged particles, which became known as electrons.

Figure 3. Cathode Ray Tube

Figure 3. Cathode Ray Tube

Further experimentation led Thomson to find the ratio of the electrical charge of the electron to its mass. This was a major step toward understanding the nature of the particle. He was awarded a Nobel Prize in 1906 for his accomplishment.

It was an American scientist, Robert Millikan, who in 1909 was able to measure the charge on an electron using the apparatus pictured in Figure 4.


Millikan”s experiment determined the mass of an electron.

FIGURE 4. Millikan Oil Drop Experiment

FIGURE 4. Millikan Oil Drop Experiment

Oil droplets were sprayed into the chamber and, in the process, became randomly charged by gaining or losing electrons. The electric field was adjusted so that a negatively charged drop would move slowly upward in front of the grid in the telescope. Knowing the rate at which the drop was rising, the strength of the field, and the mass of the drop, Millikan was able to calculate the charge on the drop. Combining the information with the results of Thomson, he could calculate a value for the mass of a single electron. Eventually, this number was found to be 9.11 × 10−28 gram.

Ernest Rutherford (England, 1911) performed a gold foil experiment (Figure 5) that had tremendous implications for atomic structure.

FIGURE 5. Rutherford”s Experiment


Rutherford”s experiment using alpha particles confirmed that there was mostly empty space between the nucleus and electrons.

Alpha particles (helium nuclei) passed through the foil with few deflections. However, some deflections (1 per 8,000) were almost directly back toward the source. This was unexpected and suggested an atomic model with mostly empty space between a nucleus, in which most of the mass of the atom was located and which was positively charged, and the electrons that defined the volume of the atom. After two years of studying the results, Rutherford finally came up with an explanation. He reasoned that the rebounded alpha particles must have experienced some powerful force within the atom. And he assumed this force must occupy a very small amount of space, because so few alpha particles had been deflected. He concluded that the force must be a densely packed bundle of matter with a positive charge. He called this positive bundle the nucleus. He further discovered that the volume of a nucleus was very small compared with the total volume of an atom. If the nucleus were the size of a marble, then the atom would be about the size of a football field. The electrons, he suggested, surrounded the positively charged nucleus like planets around the sun, even though he could not explain their motion.

Further experiments showed that the nucleus was made up of still smaller particles called protons. Rutherford realized, however, that protons, by themselves, could not account for the entire mass of the nucleus. He predicted the existence of a new nuclear particle that would be neutral and would account for the missing mass. In 1932, James Chadwick (England) discovered this particle, the neutron.

Today the number of subatomic particles identified and named as discrete units has risen to well over 90.

Bohr Model of the Atom

In 1913, Niels Bohr (Denmark) proposed his model of the atom. This pictured the atom as having a dense, positively charged nucleus and negatively charged electrons in specific spherical orbits, also called energy levels or shells, around this nucleus. These energy levels are arranged concentrically around the nucleus, and each level is designated by a number: 1, 2, 3, . . . The closer to the nucleus, the less energy an electron needs in one of these levels, but it has to gain energy to go from one level to another that is farther away from the nucleus.

Because of its simplicity and general ability to explain chemical change, the Bohr model still has some usefulness today.


Bohr”s electron distribution to principal energy levels has the formula 2n2.

Components of Atomic Structure

The chart below lists the basic particles of the atom and important information about them.

(There are now some 30 or more named particles or units of atomic structure, but the above are the most commonly used.)

When these components are used in the model, the protons and neutrons are shown in the nucleus. These particles are known as nucleons. The electrons are shown outside the nucleus.

The number of protons in the nucleus of an atom determines the atomic number. All atoms of the same element have the same number of protons and therefore the same atomic number; atoms of different elements have different atomic numbers. Thus, the atomic number identifies the element. An English scientist, Henry Moseley, first determined the atomic numbers of the elements through the use of x rays.

The sum of the number of protons and the number of neutrons in the nucleus is called the mass number.

Table 1 summarizes the relationships just discussed. Notice that the outermost energy level can contain no more than eight electrons. The explanation of this is given in the next section.

In some cases, different types of atoms of the same element have different masses. For example, three types of hydrogen atoms are known. The most common type of hydrogen, sometimes called protium, accounts for 99.985% of the hydrogen atoms found on Earth. The nucleus of a protium atom contains one proton only, and it has one electron moving about it. The second form of hydrogen, known as deuterium, accounts for 0.015% of Earth”s hydrogen atoms. Each deuterium atom has a nucleus containing one proton and one neutron. The third form of hydrogen, tritium, is radioactive. It exists in very small amounts in nature, but it can be prepared artificially. Each tritium atom contains one proton, two neutrons, and one electron.

Protium, deuterium, and tritium are isotopes of hydrogen. Isotopes are atoms of the same element that have different masses. The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. In all three isotopes of hydrogen, the positive charge of the single proton is balanced by the negative charge of the electron. Most elements consist of mixtures of isotopes. Tin, for example, has ten stable isotopes, the most of any element.


Isotopes have the same atomic number but a different atomic mass. This means they differ in the number of neutrons, not protons.

The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. The percentage at which each of an element”s isotopes occurs in nature is taken into account when calculating the element”s average atomic mass.Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.

* A complete list of the names and symbols of the known elements can be found in the Tables for Reference section.
** PEL is used to represent the principal energy levels.

Calculating Average Atomic Mass

The average atomic mass of an element depends on both the mass and the relative abundance of each of the element”s isotopes. For example, naturally occurring copper consists of 69.17% copper-63, which has an atomic mass of 62.919 598 amu, and 30.83% copper-65, which has an atomic mass of 64.927 793 amu. The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

0.6917 × 62.919 598 amu + 0.3083 × 64.927 793 amu = 63.55 amu

Therefore, the calculated average atomic mass of naturally occurring copper is 63.55 amu. Average atomic masses of the elements listed in the Periodic Table, rounded to one decimal place for use in calculations and also in full to four decimal places, are given in the Chemical Elements table in the Tables for Reference section.

Valence Electrons

Each atom attempts to have its outer energy level complete and accomplishes this by borrowing, lending, or sharing its electrons. The electrons found in the outermost energy level are called valence electrons. The remainder of the electrons are called core electrons. The absolute number of electrons gained, lost, or borrowed is referred to as the valence of the atom.

This picture can be simplified to , showing only the valence electrons as dots in an electron dot notation. This is called the Lewis structure of the atom. To complete its outer orbit to eight electrons, chlorine must borrow an electron from another atom. Its valence number then is 1. As stated above, when electrons are gained, we assign a – sign to this number, so the oxidation number of chlorine is –1.


A Lewis structure shows the atomic symbol to represent the nucleus and inner shell electrons. It shows dots to represent the valence electrons.

Since sodium tends to lose this electron, its oxidation number is +1.