PROPERTIES RELATED TO THE PERIODIC TABLE - 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


Metals are found on the left of the chart (see Table 4 ) with the most active metal in the lower left corner. Nonmetals are found on the right side with the most active nonmetal in the upper right-hand corner. The noble or inert gases are on the far right. Since the most active metals react with water to form bases, the Group 1 metals are called alkali metals. As you proceed to the right, the base-forming property decreases and the acid-forming properties increase. The metals in the first two groups are the light metals, and those toward the center are heavy metals.

The elements found along the dark line in the Periodic Table (Table 4) are called metalloids. These elements have certain characteristics of metals and other characteristics of nonmetals. Some examples of metalloids are boron, silicon, arsenic, and tellurium.

Here are some important general summary statements about the Periodic Table:

• Acid-forming properties increase from left to right on the table.

• Base-forming properties are high on the left side and decrease to the right.

• The atomic radii of elements decrease from left to right across a period.

• First ionization energies increase from left to right across a period.

• Metallic properties are greatest on the left side of the table and decrease to the right.

• Nonmetallic properties are greatest on the right side of the table and decrease to the left.


These are important trends to remember.

Study Table 4 carefully because it summarizes many of these properties. For a more detailed description of metals, alloys, and metalloids see Chapter 13.

Radii of Atoms

The size of an atom is difficult to describe because atoms have no definite outer boundary. Unlike a volleyball, an atom does not have a definite circumference.

To overcome this problem, the size of an atom is estimated by describing its radius. In metals, this is done by measuring the distance between two nuclei in the solid state and dividing this distance by 2. Such measurements can be made with xray diffraction. For a nonmetallic element that exists in pure form as a molecule, such as chlorine, measurements can be made of the distance between nuclei for two atoms covalently bonded together. Half of this distance is referred to as the covalent radius. The method for finding the covalent radius of the chlorine atom is illustrated in the following diagram.

Figure 10 shows the relative atomic and ionic radii for some elements. As you review this chart, you should note two trends:

1. Atomic radii decrease from left to right across a period in the Periodic Table (until the noble gases).

2. Atomic radii increase from top to bottom in a group or family.

The reason for these trends will become clear in the following discussions.


Know these trends in atomic radii.

Atomic Radii in Periods

Since the number of electrons in the outer principal energy level increases as you go from left to right in each period, the corresponding increase in the nuclear charge because of the additional protons pulls the electrons more tightly around the nucleus. This attraction more than balances the repulsion between the added electrons and the other electrons, and the radius is generally reduced. The inert gas at the end of the period has a slight increase in radius because of the electron repulsion in the filled outer principal energy level. For example, lithium”s atomic radius in Figure 10 is 0.152 nm at the one end of period 2 whereas fluorine has a radius of only 0.064 nm at the far end of the period. This trend can be seen in Figure 10 across every period.


This is the explanation of these trends in periods.

Atomic Radii in Groups

For a group of elements, the atoms of each successive member have another outer principal energy level in the electron configuration and the electrons there are held less tightly by the nucleus. This is so because of their increased distance from the nuclear positive charge and the shielding of this positive charge by all the core electrons. Therefore the atomic radius increases down a group. For example, oxygen”s atomic radius in Figure 10 is 0.066 nm at the top of group 16, whereas polonium has a radius of 0.167 nm at the bottom of the same group. This trend can be seen in Figure 10 down every group.


. . . and in groups.

Ionic Radius Compared with Atomic Radius

Metals tend to lose electrons in forming positive ions. With this loss of negative charge, the positive nuclear charge pulls in the remaining electrons closer and thus reduces the ionic radius below that of the atomic radius.

Nonmetals tend to gain electrons in forming negative ions. With this added negative charge, which increases the inner electron repulsion, the ionic radius is increased beyond the atomic radius. See Figure 10 for relative atomic and ionic radii values.


Know this relationship between the atomic radius and the ionic radius.

FIGURE 10. Radii of Some Atoms and Ions (in nanometers)

FIGURE 10. Radii of Some Atoms and Ions (in nanometers)

Notes: The atomic radius is usually given for metal atoms, which are shown in gray, and the covalent radius is usually given for atoms of nonmetals, which are shown in black.


The electronegativity of an element is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond. This electronegativity number is based on an arbitrary scale going from 0 to 4. In general, a value of less than 2 indicates a metal.

Notice in Table 5 that the electronegativity decreases down a group and increases across a period. The inert gases can be ignored. The lower the electronegativity number, the more electropositive an element is said to be. The most electronegative element is in the upper right corner—F, fluorine. The most electropositive is in the lower left corner of the chart—Fr, francium.

Table 5. First Ionization Energies and Electronegativities


The most electronegative element is F, in the upper right corner.


The least electronegative element is Fr, in the lower left corner.

Ionization Energy

Atoms hold their valence electrons, then, with different amounts of energy. If enough energy is supplied to one outer electron to remove it from its atom, this amount of energy is called the first ionization energy. With the first electron gone, the removal of succeeding electrons becomes more difficult because of the imbalance between the positive nuclear charge and the remaining electrons. The lowest ionization energy is found with the least electronegative atom.


Know this definition of the first ionization energy.

Ionization energies can be plotted against atomic numbers, as shown in the graph below. Follow this discussion on the graph to help you understand the peaks and valleys. Not surprisingly, the highest peaks on the graph occur for the ionization energy needed to remove the first electron from the outer energy level of the noble gases, He, Ne, Ar, Kr, Xe, and Rn, because of the stability of the filled p orbitals in the outer energy level. Notice that, even among these elements, the energy needed gradually declines. This can be explained by considering the distance of the involved energy level from the positively charged nucleus. With each succeeding noble gas, a more distant p orbital is involved, therefore making it easier to remove an electron from the positive attraction of the nucleus. Besides this consideration, as more energy levels are added to the atomic structure as the atomic number increases, the additional negative fields associated with the additional electrons screen out some of the positive attraction of the nucleus. Within a period such as that from Li to Ne, the ionization energy generally increases. The lowest occurs when a lone electron occupies the outer s orbital, as in Li. As the s orbital fills with two electrons at atomic number 4, Be, the added stability of a filled 2s orbital explains the small peak at 4. At atomic number 5, B, a lone electron occupies the 2p orbital. This electron can be removed with less energy, and therefore a dip occurs in the graph. With the 2p orbitals filling according to Hund”s Rule (refer to Table 2), with only one electron in each orbital before pairing occurs, again a slightly more stable situation and, therefore, another small peak occur at atomic number 7. After this peak, a dip and continual increases occur until the 2p orbitals are completely filled with paired electrons at the noble gas Ne. As you continue to associate the atomic number with the line in the chart, you find peaks occurring in the same general pattern. These peaks are always related to the state of filling of the orbitals involved and the distance of these orbitals from the nucleus.


Know how this trend relates to the chart.
Can you explain the peaks?


Know the reason for the peaks and valleys.