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Every element in a given horizontal row has its outermost electrons in the same principle energy region or level. Within each row, the number of electrons increases with the atomic number from left to right. The number of elements in each row varies, and reflects the sequence in which electrons are added to various quantum levels as the atoms are formed. For example, row 1 has only two elements because the first quantum level can contain only two 1s electrons. The two elements are hydrogen (1s₁) and helium (1s₂). Row 2 contains eight elements that reflect the progressive addition of 2s, then 2p electrons during the formation of lithium (helium + 2s₁) through neon (helium + 2s₂, 2p₆). Row 3 contains eight elements that reflect the filling of the 3s and 3p quantum regions respectively during the addition of electrons in sodium (neon + 3s₁) through argon (neon + 3s₂, 3p₆) as indicated in Table 2.3. Row 4 contains 18 elements which reflects the addition of 10 3d‐subshell electrons after the 4s electrons and before the 4p elections. Row 5 contains 32 elements because of the addition of 14 3f electrons after the two 5s electrons prior to the addition 10 4d subshell electrons and the six 5p electrons (Figure 2.5). The process continues through rows 6 and 7 ending with uranium. The rows that contain up to 14 f‐subshell electrons are too long to fit conveniently in a table, so the elements (lanthanides and actinides) to which f‐subshell electrons have been added are shown separately at the bottom of the table. In summary, elements are grouped into rows on the periodic table according to the highest ground state principle quantum level (1–7) occupied by their electrons. Their position within each row depends on the distribution and numbers of electrons within the principle quantum levels.

Table 2.3 Periodic table of the naturally‐occurring elements displaying atomic symbols, atomic number (Z), average mass, ground state electron configuration, common valence states and electronegativity of each element.

Simplified periodic table of the elements showing symbols atomic numbers and mass numbers

Simplified periodic table of the elements showing symbols and electron configurations

Simplified periodic table of the elements electronegativities common and valence states

      2.2.2 Ionization

The periodic table not only organizes the elements into rows based on their electron properties, but also into vertical columns based upon their tendency to gain or lose electrons to become more stable (Table 2.3). Ideal atoms are electrically neutral because they contain the same numbers of positively charged protons and negatively charged electrons (p⁺ = e⁻). Many atoms are not electrically neutral; instead they are electrically charged particles called ions. The process by which they acquire their charge is called ionization (Box 2.1). In order for an ion to form, the number of positively charged protons and negatively charged electrons must become unequal. Cations are positively charged ions because they have more positively charged protons than negatively charged electrons (p⁺ > e⁻). Their charge is equal to the number of excess protons (p⁺ − e⁻). Cations form when electrons are lost from the electron cloud. Ions that have more negatively charged electrons than positively charged protons, such as the ion chlorine (Cl⁻ ), will have a negative charge and are called anions. The charge of an anion is equal to the number of excess electrons (e⁻ − p⁺). Anions form when electrons are added to the electron cloud during ionization.

Box 2.1 Ionization energy

Ionization energy (IE) is the amount of energy required to remove an electron from its electron cloud. Ionization energies are periodic as illustrated for 20 elements in Table B2.1.

      The first ionization energy is the amount of energy required to remove one electron from the electron cloud; the second ionization energy is the amount required to remove a second electron and so forth. Ionization energies are lowest for electrons that are weakly held by the nucleus and higher for electrons that are strongly held by the nucleus or are in stable configurations. Ionization energies decrease down the periodic table because the most weakly held outer electrons are shielded from the positively charged nucleus by a progressively larger number of intervening electrons. Elements with relatively low first ionization energies are called electropositive elements because they tend to lose one or more electrons and become positively charged cations. Most elements with high first ionization energies are electronegative elements because they tend to add electrons to their electron clouds and become negatively charged anions. Since opposite charges tend to attract, you can imagine the potential such ions have for combining to produce other Earth materials. The arrangement of elements into vertical columns or groups within the periodic table helps us to comprehend the tendency of specific atoms to lose, gain or share electrons. For example, on the periodic table (see Table 2.3), column 2 (IIA) elements commonly exist as divalent (+2) cations because the first and second ionization energies are fairly similar and much lower than the third and higher ionization energies. This permits two electrons to be removed fairly easily from the electron cloud, but makes the removal of additional electrons much more difficult. Column 13 (IIIA) elements commonly exist as trivalent (+3) cations (see Table 2.3). These elements have somewhat similar first, second and third ionization energies, which are much smaller than the fourth and higher ionization energies. The transfer of electrons is fundamentally important in the understanding of chemical bonds and the development of mineral crystals.

Table B2.1 Ionization energies for hydrogen through calcium (units in kJ/mole).

Element	Ionization energy
	First		Second		Third		Fourth		Fifth		Sixth		Seventh		Eighth
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