2.5 The Periodic Table

The periodic table organizes all known elements by atomic number, revealing patterns in their properties that make predicting chemical behavior possible. Understanding how the table is structured and what trends it contains is foundational for nearly everything else in chemistry.

The Periodic Table

Organization of the Periodic Table

Elements are arranged in order of increasing atomic number, which is the number of protons in an atom's nucleus (H has 1, He has 2, Li has 3). Each element has a unique chemical symbol, usually one or two letters with the first letter capitalized (H, He, C, Fe).

The table is organized into periods (rows) and groups (columns):

• Periods correspond to the number of electron shells an atom has. An element in period 1 has 1 shell, period 2 has 2 shells, and so on.
• Groups tell you how many valence electrons an element has. Valence electrons are the electrons in the outermost shell, and they're the ones that determine an element's chemical properties and reactivity. For example, all Group 1 elements have 1 valence electron, which is why they're all highly reactive metals.

The table is also divided into four blocks based on which type of orbital the outermost electrons occupy:

• s-block: Groups 1 and 2 (plus helium), where the outermost electrons fill s-orbitals
• p-block: Groups 13 to 18, where the outermost electrons fill p-orbitals
• d-block: Groups 3 to 12 (the transition metals), where d-orbitals are being filled
• f-block: The lanthanides and actinides, shown below the main table, where f-orbitals are being filled

Predicting Properties from Position

One of the most useful things about the periodic table is that an element's position tells you a lot about its properties. Three major trends to know:

Atomic size (atomic radius)

• Decreases left to right across a period. As you move across, more protons in the nucleus pull the electrons in more tightly. That's why Na is larger than Cl, even though they're in the same period.
• Increases top to bottom within a group. Each new period adds another electron shell, pushing electrons farther from the nucleus. Cs is much larger than Li for this reason.

Ionization energy (the energy required to remove an electron)

• Increases left to right across a period. The stronger pull from the nucleus makes electrons harder to remove. Na is much easier to ionize than Cl.
• Decreases top to bottom within a group. Larger atoms have more electron shielding, so the outermost electron is held less tightly. Cs is easier to ionize than Li.

Electronegativity (an atom's ability to attract electrons in a chemical bond)

• Increases left to right across a period. Higher effective nuclear charge means a stronger pull on shared electrons. Cl is more electronegative than Na.
• Decreases top to bottom within a group. Larger atomic size weakens the pull on bonding electrons. F is the most electronegative element, far more so than I.

Metals vs. Nonmetals vs. Metalloids

The periodic table has a rough diagonal "staircase" line on the right side that separates three categories of elements:

• Metals sit to the left of the staircase (most of the table). They're shiny, malleable (can be hammered into shapes), ductile (can be drawn into wires), and conduct heat and electricity well. Examples: sodium (Na), iron (Fe), gold (Au).
• Nonmetals sit to the right of the staircase. They tend to be poor conductors, brittle in solid form, and dull in appearance. Examples: carbon (C), nitrogen (N), oxygen (O).
• Metalloids (semi-metals) sit right along the staircase line. They have properties in between metals and nonmetals, which makes some of them useful as semiconductors. Examples: boron (B), silicon (Si), germanium (Ge).
• Noble gases (Group 18) are nonmetals with a full set of valence electrons, which makes them extremely unreactive under normal conditions.

Note that hydrogen is a special case. Even though it sits in Group 1, it's a nonmetal, not a metal.

Additional Concepts

• Atomic mass is the weighted average mass of an element's naturally occurring isotopes, listed on the periodic table below the element's symbol.
• Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons. For example, carbon-12 and carbon-14 are both carbon, but carbon-14 has two extra neutrons.
• Electron configuration describes how electrons are arranged in an atom's orbitals. It's what determines where an element sits on the periodic table.
• The periodic table was originally developed by Dmitri Mendeleev in 1869. He organized elements by atomic mass and properties, even leaving gaps for elements that hadn't been discovered yet. The modern table arranges elements by atomic number instead, which resolved a few inconsistencies in Mendeleev's version.