How the Periodic Table Works

The Atomic Number as the Organizing Key

Every element on the periodic table is defined by its atomic number, the count of protons in its nucleus. Hydrogen has 1 proton, carbon has 6, iron has 26, and uranium has 92. When elements are lined up by atomic number, a pattern emerges: their chemical properties repeat at predictable intervals. This repetition is the periodic law, first identified by Dmitri Mendeleev in 1869 using atomic mass and later refined by Henry Moseley in 1913 using atomic number.

The atomic number determines how many electrons orbit the nucleus in a neutral atom, and the arrangement of those electrons, called the electron configuration, controls virtually every chemical property the element exhibits. Two elements with the same number of electrons in their outermost shell behave similarly, which is why they end up in the same column of the table.

Rows: Periods and Electron Shells

The seven horizontal rows are called periods. Each period corresponds to the filling of a new principal energy level, or electron shell. Period 1 has just two elements (hydrogen and helium) because the first shell can hold only 2 electrons. Period 2 has eight elements because the second shell accommodates 8 electrons across its s and p subshells. By period 4, the d subshell begins filling, adding 10 transition metals and stretching the row to 18 elements.

As you move across a period from left to right, each successive element has one more proton and one more electron than the last. The added electrons enter the same principal shell, but the growing nuclear charge pulls all the electrons closer to the nucleus. This progressive tightening explains why atoms get smaller across a period and why it takes more energy to remove an electron from elements on the right side of a period compared to the left.

Periods 6 and 7 each contain 32 elements because the f subshell begins filling, adding 14 lanthanides or actinides. These elements are placed in separate rows below the main table purely for visual convenience. Chemically, they belong between groups 2 and 3 of their respective periods.

Columns: Groups and Valence Electrons

The 18 vertical columns are called groups. Elements in the same group have the same number of valence electrons, the electrons in the outermost shell that participate in chemical bonding. Group 1 elements all have 1 valence electron. Group 2 elements have 2. Group 17 elements have 7. This shared valence electron count is the reason group members exhibit similar chemical behavior.

Sodium (group 1, period 3) and potassium (group 1, period 4) both have a single valence electron. Both are soft, silvery metals that react vigorously with water, form +1 ions, and produce similar compounds. Potassium is more reactive than sodium because its valence electron is in a higher shell, farther from the nucleus and easier to remove, but the fundamental chemistry is the same.

The periodic table groups page details all 18 groups and their characteristic properties.

The Four Blocks

The table divides into four blocks based on which type of orbital is being filled by the outermost electrons. The s-block (groups 1 and 2) contains elements whose last electron enters an s orbital. The p-block (groups 13 through 18) fills p orbitals. The d-block (groups 3 through 12) fills d orbitals, producing the transition metals. The f-block (lanthanides and actinides) fills f orbitals.

This block structure is a direct consequence of quantum mechanics. The shapes and energies of atomic orbitals determine the order in which electrons fill them, and the periodic table maps that filling order spatially. The s-block is two columns wide because each s subshell holds 2 electrons. The p-block is six columns wide (6 electrons per p subshell). The d-block is ten columns wide, and the f-block is fourteen columns wide.

Metals, Nonmetals, and Metalloids

A diagonal line running from boron (atomic number 5) to astatine (atomic number 85) roughly separates metals on the left from nonmetals on the right. Elements touching this line, such as silicon, germanium, arsenic, and tellurium, are metalloids with properties intermediate between metals and nonmetals.

Metals tend to lose electrons, forming positive ions. They conduct electricity, are malleable, and typically have high melting points. Nonmetals tend to gain or share electrons, forming negative ions or covalent bonds. They are generally poor conductors and more diverse in their physical states. The metallic character trends article explains how this property changes systematically across the table.

Predictive Power

The real value of the periodic table is prediction. If you know an element's position, you can estimate its atomic radius, ionization energy, electronegativity, common oxidation states, and the types of compounds it will form. You can predict that an element in group 1 will form a +1 ion, that an element in group 17 will form a -1 ion, and that noble gases in group 18 will resist forming compounds entirely.

Mendeleev famously used this predictive power to forecast the existence and properties of elements not yet discovered. He left gaps in his 1869 table and described what the missing elements should look like. When gallium, scandium, and germanium were later discovered and matched his predictions closely, it validated the entire framework.

Today, the periodic table's predictive capacity extends to the synthetic superheavy elements being created in particle accelerators. Researchers use periodic trends to estimate the properties of elements 119 and beyond, even before they are synthesized. The periodic trends guide covers all the major property patterns in detail.