The periodic table was first developed by Mendeleev and Meyer on the basis of the similarity in chemical and physical properties exhibited by certain elements. Moseley established that each element has a unique atomic number, which added more order to the periodic table.

We now know that elements in the same column of the periodic table have the same number of electrons in their valence orbitals. This similarity in valence electronic structure leads to the similarities among elements in the same group. Differences among elements in the same group arise because their valence orbitals are in different shells.

Many properties of atoms depend on the effective nuclear charge, which is the portion of the nuclear charge that an outer electron experiences after accounting for repulsions by other electrons in the atom. The core electrons are very effective in screening the outer electrons from the full charge of the nucleus, whereas electrons in the same shell do not screen each other very effectively. Because the actual nuclear charge increases as we progress through a period, the effective nuclear charge experienced by valence electrons increases as we move left to right across a period.

Sizes of Atoms and Ions

The size of an atom can be gauged by its bonding atomic radius, which is based on measurements of the distances separating atoms in their chemical compounds. In general, atomic radii increase down a column in the periodic table andand decrease left to right across a row.

Cations are smaller than their parent atoms, whereas anions are larger than their parent atoms. For ions of the same charge, size increases going down a column of the periodic table. We can write electron configurations for ions by first writing the electron configuration of the neutral atom and then removing or adding the appropriate number of electrons. For cations, electrons are removed first from the orbitals of the neutral atom with the largest value of n. If there are two valence orbitals with the same value of n (such as 4s and 4p), then the electrons are lost first from the orbital with a higher value of l (in this case, 4p). For anions, electrons are added to orbitals in the reverse order.

An isoelectronic series is a series of ions that has the same number of electrons. For such a series, size decreases with increasing atomic number as the electrons are attracted more strongly to the nucleus as its positive charge increases.

Ionization Energy and Electron Affinity

The first ionization energy of an atom is the minimum energy needed to remove an electron from the atom in the gas phase, forming a cation. The second ionization energy is the energy needed to remove a second electron from the cation, and so forth. Ionization energies show a sharp increase after all the valence electrons have been removed because of the much higher effective nuclear charge experienced by the core electronsand decrease left to right across a row.

Cations are smaller than their parent atoms, whereas anions are larger than their parent atoms. For ions of the same charge, size increases going down a column of the periodic table. We can write electron configurations for ions by first writing the electron configuration of the neutral atom and then removing or adding the appropriate number of electrons. For cations, electrons are removed first from the orbitals of the neutral atom with the largest value of n. If there are two valence orbitals with the same value of n (such as 4s and 4p), then the electrons are lost first from the orbital with a higher value of l (in this case, 4p). For anions, electrons are added to orbitals in the reverse order.

An isoelectronic series is a series of ions that has the same number of electrons. For such a series, size decreases with increasing atomic number as the electrons are attracted more strongly to the nucleus as its positive charge increases.

Ionization Energy and Electron Affinity

The first ionization energy of an atom is the minimum energy needed to remove an electron from the atom in the gas phase, forming a cation. The second ionization energy is the energy needed to remove a second electron from the cation, and so forth. Ionization energies show a sharp increase after all the valence electrons have been removed because of the much higher effective nuclear charge experienced by the core electrons. The first ionization energies of the elements show periodic trends that are opposite those seen for atomic radii, with smaller atoms having higher first ionization energies. Thus, first ionization energies decrease as we go down a column and increase as we proceed left to right across a row.

The electron affinity of an element is the energy change upon adding an electron to an atom in the gas phase, forming an anion. A negative electron affinity means that energy is released when the electron is added; hence, when the electron affinity is negative, the anion is stable. By contrast, a positive electron affinity means that the anion is not stable relative to the separated atom and electron. In general, electron affinities become more negative as we proceed from left to right across the periodic table. The halogens have the most negative electron affinities. The electron affinities of the noble gases are positive because the added electron would have to occupy a new, higher-energy subshell.

Metals, Nonmetals, and Metalloids

The elements can be categorized as metals, nonmetals, and metalloids. Most elements are metals; they occupy the left side and the middle of the periodic table. Nonmetals appear in the upper-right section of the table. Metalloids occupy a narrow band between the metals and nonmetals. The tendency of an element to exhibit the properties of metals, called the metallic character, increases as we proceed down a columnand decreases as we proceed from left to right across a row.

Metals have a characteristic luster, and they are good conductors of heat and electricity. When metals react with nonmetals, the metal atoms are oxidized to cations and ionic substances are generally formed. Most metal oxides are basic; they react with acids to form salts and water.

Nonmetals lack metallic luster and are generally poor conductors of heat and electricity. Several are gases at room temperature. Compounds composed entirely of nonmetals are generally molecular. Nonmetals usually form anions in their reactions with metals. Nonmetal oxides are acidic; they react with bases to form salts and water. Metalloids have properties that are intermediate between those of metals and nonmetals.

Trends for Group 1A and 2A Metals

The periodic properties of the elements can help us understand the properties of groups of the representative elements. The alkali metals (group 1A) are soft metals with low densities and low melting points. They have the lowest ionization energies of the elements. As a result, they are very reactive toward nonmetals, easily losing their outer s electron to form 1+ ions.

The alkaline earth metals (group 2A) are harder and denser, and have higher melting points than the alkali metals. They are also very reactive toward nonmetals, although not as reactive as the alkali metals. The alkaline earth metals readily lose their two outer s electrons to form 2+ ions. Both alkali and alkaline earth metals react with hydrogen to form ionic substances that contain the hydride ion, H−

Trends for Selected Nonmetals

Hydrogen is a nonmetal with properties that are distinct from any of the groups of the periodic table. It forms molecular compounds with other nonmetals, such as oxygen and the halogens.

Oxygen and sulfur are the most important elements in group 6A, the chalcogens. Oxygen is usually found as a diatomic molecule, O2. Ozone, O3, is an important allotrope of oxygen. Oxygen has a strong tendency to gain electrons from other elements, thus oxidizing them. In combination with metals, oxygen is usually found as the oxide ion, O2−, although salts of the peroxide ion, O22−, and superoxide ion, O2−, are sometimes formed. Elemental sulfur is most commonly found as S8 molecules. In combination with metals, it is most often found as the sulfide ion, S2−.

The halogens (group 7A) exist as diatomic molecules. The halogens have the most negative electron affinities of the elements. Thus, their chemistry is dominated by a tendency to form 1− ions, especially in reactions with metals.

The noble gases (group 8A) exist as monatomic gases. They are very unreactive because they have completely filled s and p subshells. Only the heaviest noble gases are known to form compounds, and they do so only with very active nonmetals, such as fluorine.