1.3 Periodic Table and Elemental Properties

The periodic table organizes elements based on atomic structure and properties. It's divided into blocks, periods, and groups, each reflecting specific electron configurations and trends. Understanding these patterns helps predict chemical behavior and reactivity.

Elements are classified into main groups, transition metals, and inner transition series. This organization reveals similarities in properties, valence electrons, and chemical reactions. Recognizing these groupings is the foundation for understanding periodic law and elemental trends.

Periodic Table Organization

Structure and Blocks of the Periodic Table

The periodic table is arranged in periods (horizontal rows) and groups (vertical columns). Periods correspond to the principal quantum number $n$ of the outermost electron shell, so moving across a period means filling orbitals within the same shell. Groups share the same valence electron configuration, which is why elements in the same group tend to behave similarly in chemical reactions.

The table is further divided into four blocks based on which subshell is being filled:

• s-block (Groups 1–2 plus He): outermost electrons occupy s orbitals. These elements have configurations ending in $ns^1$ or $ns^2$.
• p-block (Groups 13–18): outermost electrons fill p orbitals, with configurations ending in $np^1$ through $np^6$.
• d-block (Groups 3–12): the transition metals, where electrons fill $(n-1)d$ orbitals. Note that the d orbitals being filled belong to the shell one below the valence shell.
• f-block (lanthanides and actinides): displayed as two rows below the main table, where $(n-2)f$ orbitals are being filled.

Properties and Trends Within the Periodic Table

Periodic trends arise from the interplay between effective nuclear charge ($Z_{\text{eff}}$) and electron shielding. As you move across a period, protons are added to the nucleus while electrons enter the same shell, so $Z_{\text{eff}}$ increases and the electron cloud is pulled inward. Moving down a group, new shells are added, increasing the distance from the nucleus and the amount of shielding from inner electrons.

These two factors drive the major trends:

• Atomic radius decreases across a period (higher $Z_{\text{eff}}$ pulls electrons closer) and increases down a group (additional shells expand the atom).
• Ionization energy increases across a period (electrons are held more tightly) and decreases down a group (outer electrons are farther from the nucleus and easier to remove).
• Electronegativity increases across a period and decreases down a group, following the same logic as ionization energy.
• Electron affinity generally becomes more exothermic across a period and less exothermic down a group, though this trend has notable exceptions (e.g., nitrogen's half-filled $2p$ subshell gives it a lower electron affinity than its neighbors).
• Metallic character decreases across a period and increases down a group, since metals tend to lose electrons easily, which is favored by low $Z_{\text{eff}}$.

A common pitfall: don't treat these trends as perfectly smooth. Irregularities show up at half-filled and fully filled subshells. For instance, the first ionization energy of oxygen is lower than that of nitrogen because removing one electron from oxygen relieves electron-electron repulsion in the doubly occupied $2p$ orbital.

Element Classifications

Main Group Elements and Noble Gases

Main group elements include the s-block and p-block elements. Their chemistry is largely governed by the number of valence electrons, which matches their group number (using the 1–8 convention for main groups).

• Alkali metals (Group 1) have a single $ns^1$ electron and lose it readily, forming $+1$ ions. They react vigorously with water to produce hydroxides and hydrogen gas. Reactivity increases down the group: cesium reacts explosively, while lithium reacts relatively gently.
• Alkaline earth metals (Group 2) have $ns^2$ configurations and form $+2$ ions. Magnesium and calcium react with water (calcium more readily), producing hydroxides and $H_2$ gas.
• Chalcogens (Group 16) have $ns^2np^4$ configurations. Oxygen and sulfur are essential to biological systems, with oxygen being the most electronegative element in this group.
• Halogens (Group 17) exist as diatomic molecules ($F_2$, $Cl_2$, $Br_2$, $I_2$) and are highly reactive nonmetals. They need just one electron to complete their valence shell, so they readily form $-1$ ions. Fluorine is the most electronegative element on the entire table.
• Noble gases (Group 18) have completely filled valence shells ($ns^2np^6$, except for helium at $1s^2$), giving them exceptional stability. While they were once thought to be entirely inert, heavier noble gases like xenon do form compounds (e.g., $XeF_2$, $XeF_4$).

Transition Metals, Lanthanides, and Actinides

Transition metals occupy the d-block and are defined by having partially filled d orbitals in at least one common oxidation state. Their chemistry differs from main group elements in several important ways:

• Variable oxidation states. Because the $(n-1)d$ and $ns$ electrons are close in energy, transition metals can lose different numbers of electrons. Iron, for example, commonly exists as $Fe^{2+}$ and $Fe^{3+}$.
• Colored compounds. Partially filled d orbitals allow d-d electronic transitions that absorb visible light, producing characteristic colors (e.g., $Cu^{2+}$ solutions appear blue).
• Catalytic activity. Many transition metals and their compounds serve as catalysts because they can adopt multiple oxidation states and provide surfaces for reactant adsorption. Platinum and palladium are widely used in catalytic converters and hydrogenation reactions.

The lanthanides ($4f$ series) and actinides ($5f$ series) make up the f-block:

• Lanthanides share very similar chemical properties because the $4f$ electrons being added are buried deep inside the atom and don't strongly influence bonding. They most commonly form $+3$ ions. The lanthanide contraction, a steady decrease in ionic and atomic radii across the series, occurs because $4f$ electrons shield each other poorly from the increasing nuclear charge. This contraction has a downstream effect: it makes the third-row transition metals (Hf through Hg) nearly the same size as their second-row counterparts.
• Actinides include several radioactive elements. Uranium and plutonium are the most well-known, used in nuclear energy and weapons. Early actinides show more variable oxidation states than lanthanides, but the heavier actinides tend to settle into the $+3$ state.

Metalloids and Their Unique Properties

Metalloids sit along the diagonal staircase separating metals from nonmetals on the periodic table. They display properties intermediate between the two classes.

• Their most distinctive feature is semiconductivity: metalloids like silicon and germanium conduct electricity better than nonmetals but worse than metals, and their conductivity increases with temperature (opposite to metals). This property makes them the backbone of the electronics industry.
• Metalloids can form alloys with metals and covalent bonds with nonmetals, giving them versatile bonding behavior.
• Boron is notable for its electron-deficient bonding. It forms complex hydrides (boranes) with unusual three-center two-electron bonds, and it's used in neutron-absorbing materials for nuclear reactors.
• Silicon forms an extensive network of $Si-O$ bonds in silicate minerals, making it the second most abundant element in Earth's crust.

The exact list of metalloids varies slightly depending on the source, but boron, silicon, germanium, arsenic, antimony, and tellurium are the most commonly cited.