Properties That Can Be Predicted from the Periodic Table

Figure 5.14 shows the relative sizes of the atoms of the representative elements. Notice that atom size increases from top to bottom in a column and from right to left across a row. This trend is related to electron configuration. As we look at the elements in column 1, for example, we see that the single valence electron for each successive element is in a higher principal energy level than the last, and the electron is thus farther away from the positively charged nucleus; hence, the atomic radius increases going from top to bottom. This same regular increase in size can be observed in each column of the periodic table.

Atoms decrease in size going across a period from left to right. For elements within a period, electrons are being added one by one to the same principal energy level. At the same time, protons are also being added one by one to the nucleus, increasing its positive charge. This increasing positive charge increases the attraction of the nucleus for all electrons and pulls them all closer to the nucleus, decreasing the atom's radius. Thus, atomic size is a periodic property that increases from top to bottom within a column and from right to left across a period.

 FIGURE 5.14 The relative sizes of the atoms of the representative elements.

B. Ionization Energy
The ionization energy of an element is the minimum energy required to remove an electron from a gaseous atom of that element, leaving a positive ion. An equation expressing the ionization of sodium would be:

Electrons are held in the atom by the attractive force of the positively charged nucleus. The farther the outermost electrons are from the nucleus, the less tightly they are held. Thus, the ionization energy within a group of elements decreases as the elements increase in atomic number. Among the atoms of naturally occurring alkali metals, the single valence electron of cesium is farthest from the nucleus (in the sixth principal energy level), and we can correctly predict that the ionization energy of cesium is the lowest of all the alkali metals. (Recall that francium is not naturally occurring.)

From left to right across a period, the ionization energy of the elements tends to increase. The number of protons in the nucleus (the nuclear charge) increases, yet the valence electrons of the elements are in the same energy level. It becomes increasingly more difficult to remove an electron from the atom. The ionization energy of chlorine is much greater than that of sodium, an element in the same period.

The ionization energies of elements 1 through 36 are plotted versus their atomic numbers in Figure 5.15. The peaks of the graph are the high ionization energies of the noble gases. The height of the peaks decreases as the number of the highest occupied energy level increases. The low points of the graph are the ionization energies of the alkali metals, which have only one electron in their valence shell. These points, too, decrease slightly as the number of the highest occupied energy level increases. The graph shows that ionization energy is periodically related to atomic number. Even within a row of the periodic table, the variations in ionization energy are closely related to electron configuration.

The ionization energy of an element is a measure of its metallic nature. From Figure 5.15, we see that each alkali metal has the lowest ionization energy of the elements in its period. Therefore, alkali metals are the most metallic elements. From bottom to top in the periodic table and from left to right across it, the metallic nature of the elements decreases.

 FIGURE 5.15 The ionization energies of elements 1 through 36 plotted versus their atomic numbers.

Nonmetals, located in the upper-right section of the periodic table, have high ionization energies. Except for the noble gases, fluorine has the highest ionization energy. Therefore, excluding the noble gases, fluorine is the least metallic (or most nonmetallic) element. From top to bottom in a column or to the left of fluorine, elements become more metallic. In summary, ionization energy increases from bottom to top of a column and from left to right across a period.

Electron affinity is closely related to but the opposite of ionization energy. Electron affinity is the energy change that occurs when an electron is added to a neutral atom. For a nonmetal this change is usually a release of energy. The equation showing this reaction for chlorine is:

Electron affinities are fairly difficult to measure. Accurate values have been determined for only a few elements. In general the values become increasingly negative from left to right across a period. Consequently, a halogen will have the most-negative electron affinity of all the representative elements in its period (remember that the noble gases are not representative elements). Electron affinity does not change with the same regularity as does atomic radius or ionization energy, and relative values cannot be predicted as easily.

Figure 5.16 summarizes how atomic radii, ionization energies, and metallic properties change within the periodic table.

 FIGURE 5.16 Trends of various atomic properties as related to position in the periodic table.

C. The Formation of Ions
Atoms are electrically neutral. The number of positively charged protons in the nucleus of an atom equals the number of negatively charged electrons outside the nucleus. If electrons are added or lost as an atom reacts, the atom acquires a charge and becomes an ion.

1. The octet rule
We have already observed that the noble gases are very unreactive (Section 5.6A). This lack of reactivity is attributable to a stable electron configuration. Looking back to Section 5.5C (Table 5.3), you can see that all the noble gases but helium have eight electrons (two s and six p) in the highest occupied energy level. When atoms of the other representative elements react, they lose, gain, or share enough electrons to attain the noble-gas electron structure - a complete octet, eight electrons, in their outer shell. This tendency is expressed by the octet rule: An atom generally reacts in ways that give it an octet of electrons in its outer shell. Hydrogen and lithium are exceptions; they react in ways that give them the same electron configuration as helium, with two outer-shell electrons.

An atom with one, two, or three valence electrons usually reacts by losing these electrons to acquire the electron configuration of the noble gas next below it in atomic number. An atom with six or seven valence electrons will usually react by adding enough electrons to acquire the electron configuration of the noble gas next above it in atomic number. Other atoms may attain a complete octet by sharing electrons with a neighboring atom (discussed in Section 7.1.)

2. Positive ions, or cations
When a neutral atom loses an electron, it forms a positively charged ion, called a cation (pronounced "cát-i-on"). In general, metals lose electrons to form cations. The atom thereby attains the electron configuration of the noble gas next below it in atomic number.

For example, an alkali metal loses one electron to form a cation with a single positive charge. Sodium loses its single 3s valence electron to form the ion Na+, which has the electron configuration of neon:

An alkaline earth metal loses two electrons to form a cation with a charge of +2. In forming the magnesium ion, Mg2+, a magnesium atom loses its two valence electrons:

Aluminum loses its three valence electrons to form a cation with a charge of +3:

The names of these cations are the same as the metals from which they are formed (see Table 5.7).

TABLE 5.7 Cations of metals
Alkali metal cations
Symbol Name
Li+ lithium ion
Na+ sodium ion
K+ potassium ion
Rb+ rubidium ion
Cs+ cesium ion
Alkaline earth metal cations
Symbol Name
Mg2+ magnesium ion
Ca2+ calcium ion
Sr2+ strontium ion
Ba2+ barium ion
Other metal cations
Al3+ aluminum ion

Transition elements and the metals to their right do not always follow the octet rule; frequently they form more than one cation. For example, iron forms Fe2+ and Fe3+; cobalt forms Co2+ and Co3+. The names of these ions must indicate the charge they carry. The preferred system of nomenclature (naming) is that recommended by the International Union of Pure and Applied Chemistry (IUPAC). In this system, the name of the metal is followed by a Roman numeral (in parentheses) showing the charge on the ion. No extra space is left between the name and the number. Thus, Fe2+ is iron(II) (pronounced "iron two"), and Fe3+ is iron(III). In the older system, the name of the cation of lower charge ends in ous, and the name of the cation of higher charge ends in ic. Examples of both systems of naming are given in Table 5.8.

TABLE 5.8 Naming of cations
Symbol IUPAC name Older name
Co2+ cobalt(II) cobaltous
Co3+ cobalt(III) cobaltic
Cu+ copper(I) cuprous
Cu2+ copper(II) cupric
Cr2+ chromium(II) chromous
Cr3+ chromium(III) chromic
Fe2+ iron(II) ferrous
Fe3+ iron(III) ferric

Notice that none of the cations discussed here have a charge greater than +3. When ions are formed, electrons are pulled off one by one from the atom. Thus, the first electron is removed from a neutral atom, the second electron from an ion of charge +1, the third electron from an ion of charge +2, and so on. The amount of energy necessary to remove an electron increases dramatically as the positive charge of the ion increases. To remove a fourth electron and form an ion of charge +4 is energetically unlikely.

3. Negative ions, or anions
When a neutral atom gains an electron, it forms a negatively charged ion, called an anion (pronounced "án-i-on"). Typically, nonmetals form anions, gaining enough electrons to acquire the electron configuration of the noble gas of next higher atomic number. Elements of group 6, with six valence electrons, form anions by gaining two electrons; the halogens, with seven valence electrons, form anions by gaining one electron. The names of these anions include the root name of the element and the ending ide. Table 5.9 lists several common anions and their names; in each case, the root of the name is italicized.

Symbol Name Symbol Name F- fluoride ion O2- oxide ion Cl- chloride ion S2- sulfide ion Br- bromide ion I- iodide ion

4. Polyatomic ions
The ions described in the preceding paragraphs are monatomic ions; that is, each contains only one atom. Many polyatomic ions are also known. Polyatomic ions are groups of atoms bonded together that carry a charge due to an excess or deficiency of electrons. Table 5.10 lists the formulas and names of several common polyatomic ions. The symbols in the formula show which elements are present. The subscripts ("1" is understood) tell how many atoms of each element are present in the ion.

D. Metals and Nonmetals; Acids and Bases
So far we have shown that metals usually have one, two, or three valence electrons. They have low ionization energies and are found to the left in the periodic table. Nonmetals have four, five, six, or seven valence electrons, have high ionization energies, and are in the upper-right section of the periodic table. All these properties are closely related to electron configurations, and we have used electron configurations in discussing them. However, before electron configurations were known, people did know the differences between metals and nonmetals.

Early chemists observed how the physical properties of a metal (malleability, luster, and conductivity), described in Section Section 3.3, contrasted with those of nonmetals. These chemists also identified differences in chemical properties. The compounds formed when oxides of metals react with water are very much alike and very different from those formed when oxides of nonmetals react with water.

When a metallic oxide reacts with water, a hydroxide is formed:

Metal oxide Equation Hydroxide
sodium oxide Na2O + H2O 2 NaOH sodium hydroxide
magnesium oxide MgO + H2O Mg(OH)2 magnesium hydroxide
aluminum oxide Al2O3 + 3 H2O 2 Al(OH)3 aluminum hydroxide

When a nonmetallic oxide reacts with water, an acid is formed:

Nonmetal oxide Equation Acid
carbon dioxide CO2 + H2O H2CO3 carbonic acid
sulfur trioxide SO3 + H2O H2SO4 sulphuric acid
oxide of phosphorus P4O10 + 6 H2O 4H3PO4 phosphoric acid

Table 5.11 lists several common hydroxides and acids.

 Common hydroxides Common acids sodium hydroxide NaOH hydrochloric acid HCl potassium hydroxide KOH acetic acid HC2H3O2 calcium hydroxide Ca(OH)2 nitric acid HNO3 aluminum hydroxide Al(OH)3 sulfuric acid H2SO4 ammonium hydroxide NH4OH carbonic acid H2CO3 phosphoric acid H3PO4

Hydroxides are a subset of a larger group called bases, although not all bases are hydroxides. In all but ammonium hydroxide, the cation of a hydroxide is a metallic ion. A hydroxide dissolves in water to yield hydroxide ion. The solution of a hydroxide feels slippery because of the action of these ions on the skin. (You may have noticed this property in household ammonia, a dilute solution of ammonium hydroxide). The solution of a hydroxide gives a class of compounds called indicators characteristic colors (see Table 5.12).

 Acids Hydroxides In aqueous solutions release H+ release OH- Indicators litmus red blue phenolphthalein colorless red methyl orange red yellow

Most common acids contain hydrogen, a nonmetal, and frequently oxygen. Acids dissolve in water to yield hydrogen ions. Indicators in acid solutions show different colors than they do in solutions of hydroxides.