18.1 Periodicity

The periodic table organizes elements by atomic number and electron configuration, grouping them into periods and groups so that elements with similar properties line up together. This arrangement reveals predictable trends in atomic size, ionization energy, electronegativity, and reactivity. Understanding these patterns is how you move from memorizing individual elements to actually predicting how they'll behave in chemical reactions.

Periodic Table Organization and Trends

Classification of Elements in the Periodic Table

Every element has a spot on the periodic table determined by its atomic number (the number of protons in its nucleus). Atomic number increases from left to right and from top to bottom.

An element's electron configuration follows a specific orbital filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Where an element falls in this filling sequence determines its position on the table. The electrons in the outermost shell, called valence electrons, are the ones that drive chemical properties and reactivity.

The table is divided into periods (horizontal rows) and groups (vertical columns):

• The period number tells you the highest principal energy level (n) occupied by that element's electrons. For example, elements in Period 3 have valence electrons in the n = 3 shell.
• The group number relates to the number of valence electrons. Groups 1 and 2 have 1 and 2 valence electrons, respectively. Groups 13–18 have 3–8 valence electrons.

Elements in the same group share similar electron configurations, which is why they behave similarly in reactions:

• Main group elements (Groups 1, 2, 13–18) have valence electrons in s and p orbitals (configurations like $ns^1$, $ns^2$, $ns^2np^{1-6}$)
• Transition metals (Groups 3–12) have partially filled d-orbitals ($ns^2(n-1)d^{1-10}$)
• Noble gases (Group 18) have completely filled outer electron shells, satisfying the octet rule, which makes them extremely unreactive

Trends in Atomic Properties

Atomic Size (Atomic Radius)

Atomic radius follows two clear trends:

• Across a period (left to right): Atomic size decreases. As you add protons, the increasing effective nuclear charge pulls electrons closer to the nucleus, even though you're adding electrons to the same shell. For example, lithium is larger than carbon: Li > Be > B > C.
• Down a group (top to bottom): Atomic size increases. Each new period adds another electron shell, pushing valence electrons farther from the nucleus. So Li < Na < K < Rb.

Ionization Energy

First ionization energy is the energy needed to remove the outermost electron from a neutral, gaseous atom. It trends opposite to atomic size:

• Across a period: Ionization energy increases because the stronger nuclear charge and smaller atomic size make electrons harder to remove (Li < Be < B < C).
• Down a group: Ionization energy decreases because valence electrons are farther from the nucleus and inner electrons shield them from the nuclear charge (Li > Na > K > Rb).

Electronegativity

Electronegativity measures how strongly an atom attracts electrons in a chemical bond. Fluorine is the most electronegative element.

• Across a period: Electronegativity increases due to greater effective nuclear charge.
• Down a group: Electronegativity decreases due to larger atomic size and more electron shielding.

Reactivity

Reactivity trends differ for metals and nonmetals because they achieve stability in opposite ways:

• Metals (left side of the table) react by losing electrons to form cations. Reactivity increases going down a group because ionization energy drops, making it easier to lose electrons. Rubidium reacts more violently with water than sodium does (Li < Na < K < Rb in reactivity).
• Nonmetals (right side of the table) react by gaining electrons to form anions. Reactivity generally increases going up a group and to the right across a period, because smaller atoms with higher electronegativity attract electrons more effectively. Fluorine is the most reactive nonmetal (N < O < F).

Metals vs. Nonmetals vs. Metalloids

Metals make up the majority of the periodic table and share a set of characteristic properties:

• Good conductors of heat and electricity because of delocalized electrons that move freely through the metallic lattice (Cu, Ag, Au)
• Malleable (can be hammered into sheets) and ductile (can be drawn into wires) because metallic bonding allows layers of atoms to slide past each other without breaking (Al, Fe)
• Generally high melting and boiling points due to strong metallic bonds; tungsten (W) has the highest melting point of any metal at 3,422°C
• Form basic oxides and hydroxides in water (e.g., $Na_2O$, KOH)
• Tend to lose electrons in reactions: $Na \rightarrow Na^+ + e^-$, $Mg \rightarrow Mg^{2+} + 2e^-$

Nonmetals have contrasting properties:

• Poor conductors of heat and electricity (insulators) because they lack delocalized electrons (S, P)
• Brittle in solid form due to covalent or molecular bonding rather than metallic bonding
• Lower melting and boiling points because intermolecular forces between their molecules are weaker than metallic bonds ($O_2$, $Cl_2$, $Br_2$)
• Form acidic oxides in water (e.g., $CO_2$ dissolved in water forms carbonic acid; $SO_3$ forms sulfuric acid)
• Tend to gain electrons in reactions: $Cl + e^- \rightarrow Cl^-$, $O + 2e^- \rightarrow O^{2-}$

Metalloids (also called semi-metals) sit along the "staircase" line dividing metals from nonmetals (B, Si, Ge, As, Sb, Te). They have intermediate properties. Most notably, several metalloids are semiconductors, meaning their electrical conductivity falls between that of metals and insulators and can be tuned by adding impurities. Silicon is the backbone of computer chips, and germanium is used in transistors, precisely because of this tunable conductivity.

Special Groups of Elements

The lanthanides and actinides are the two rows that sit below the main body of the periodic table:

• Lanthanides: 15 elements from lanthanum (La, Z = 57) to lutetium (Lu, Z = 71). These elements are characterized by the progressive filling of 4f orbitals. They're sometimes called "rare earth elements," though most aren't actually rare. They're used in magnets, lasers, and phone screens.
• Actinides: 15 elements from actinium (Ac, Z = 89) to lawrencium (Lr, Z = 103). These fill 5f orbitals. Many actinides are radioactive, and elements beyond uranium (Z = 92) are synthetic, meaning they don't occur naturally.

Within each series, the elements share very similar chemical properties because the f-electrons being added are buried deep inside the atom and don't significantly change how the outermost electrons interact with other atoms.