3.4 Periodic table and trends in atomic properties

The periodic table organizes elements by atomic number and electron configuration, revealing patterns in atomic properties that help predict chemical behavior. These patterns arise from the interplay between effective nuclear charge and electron shielding, and they govern everything from atomic size to reactivity and bonding tendencies.

Organization of the Periodic Table

Arrangement and Atomic Number

The periodic table is arranged in order of increasing atomic number, which equals the number of protons in an atom's nucleus. Elements fall into periods (rows) and groups (columns) based on their electron configurations and chemical properties.

Valence Electrons and Chemical Behavior

Elements in the same group share similar electron configurations in their outermost energy level. These outermost electrons are called valence electrons, and they're the reason elements in a group behave similarly in chemical reactions.

Moving left to right across a period, the number of valence electrons increases. This shift causes elements to become progressively more nonmetallic in character. Carbon, nitrogen, and oxygen illustrate this progression across Period 2.

Periodic Table Blocks

The table is divided into four blocks based on which subshell receives the last electron in the ground-state configuration:

• s-block: Groups 1 and 2 (alkali metals and alkaline earth metals), plus helium. These elements have their outermost electron in an s orbital.
• p-block: Groups 13–18 (main group elements on the right side). The last electron enters a p orbital.
• d-block: Groups 3–12 (transition metals). The last electron enters a d orbital.
• f-block: The two rows pulled out below the main table (lanthanides and actinides). The last electron enters an f orbital.

Periodic Trends in Atomic Properties

Atomic Radius

Atomic radius is the distance from the nucleus to the outermost electron orbital of an atom.

• Across a period (left → right): Atomic radius generally decreases. For example, lithium is much larger than neon. Each additional proton increases the effective nuclear charge, pulling the electron cloud inward, while electrons added to the same shell don't shield each other very well.
• Down a group (top → bottom): Atomic radius increases. Lithium is smaller than cesium because each new period adds another electron shell, pushing the outermost electrons farther from the nucleus.

Ionization Energy

Ionization energy (IE) is the minimum energy required to remove an electron from a neutral atom in the gaseous state.

• Across a period (left → right): IE generally increases. From sodium to argon, the growing effective nuclear charge holds electrons more tightly, making them harder to remove.
• Down a group (top → bottom): IE decreases. Fluorine has a much higher IE than iodine because the outermost electrons in larger atoms are farther from the nucleus and more shielded by inner electrons, so they're easier to remove.

Notice the inverse relationship with atomic radius: smaller atoms hold their electrons more tightly, so they have higher ionization energies.

Electron Affinity

Electron affinity (EA) is the energy change when an electron is added to a neutral atom in the gaseous state. A more negative value means the atom releases more energy upon gaining an electron, indicating a stronger "desire" for that extra electron.

• Across a period (left → right): EA generally becomes more negative. From beryllium to chlorine, higher effective nuclear charge creates a stronger pull on incoming electrons.
• Down a group (top → bottom): EA becomes less negative. Fluorine attracts an added electron more strongly than iodine does, because the larger atomic size in heavier elements weakens the attraction between the nucleus and an incoming electron.

Interplay of Effective Nuclear Charge and Shielding

All of these trends trace back to two competing factors:

• Effective nuclear charge ($Z_{\text{eff}}$): The net positive charge felt by valence electrons after accounting for shielding. It increases across a period because each added proton isn't fully offset by the added electron in the same shell.
• Shielding (screening) effect: Inner-shell electrons partially block the nuclear charge from reaching outer electrons. Shielding increases significantly down a group (more inner shells), but stays roughly constant across a period (electrons are added to the same shell).

Across a period, rising $Z_{\text{eff}}$ with nearly constant shielding explains why atoms shrink, IE rises, and EA becomes more negative. Down a group, the addition of whole new electron shells overwhelms the increase in nuclear charge, so atoms get larger and hold their outer electrons less tightly.

Predicting Element Properties

Similar Properties within Groups

Because group members share the same valence electron configuration, they display similar chemistry:

• Alkali metals (Group 1): One valence electron, highly reactive, readily form +1 cations. Reactivity increases down the group as IE drops (cesium is more reactive than lithium).
• Halogens (Group 17): Seven valence electrons, highly reactive, readily form −1 anions. Reactivity decreases down the group as electron affinity weakens (fluorine is more reactive than iodine).

Metallic and Nonmetallic Character

• Metallic character increases toward the bottom-left of the table. Metals are good conductors, malleable, and ductile. They tend to have low ionization energies and low electronegativities.
• Nonmetallic character increases toward the top-right. Nonmetals are poor conductors, often brittle as solids, and have high electronegativities.

This diagonal trend means elements near the boundary (like silicon and germanium) show intermediate, or metalloid, behavior.

Reactivity Trends

• Metals become more reactive going down a group. Lower ionization energies make it easier to lose electrons. Cesium reacts explosively with water; lithium reacts much more gently.
• Nonmetals become more reactive going up a group. Stronger electron affinities make it easier to gain electrons. Fluorine is the most reactive nonmetal; iodine is comparatively mild.

Gradual Changes within Periods

Within a single period, properties shift gradually as atomic number increases:

• Ionization energy trends upward (lithium to neon)
• Electronegativity trends upward (sodium to chlorine)
• Atomic radius trends downward

These aren't perfectly smooth progressions. Subshell effects create small irregularities (for example, boron has a slightly lower IE than beryllium because boron's outermost electron occupies a higher-energy 2p orbital rather than the filled 2s).

Atomic Properties and Reactivity

Influence on Chemical Bonding

The trends in atomic radius, ionization energy, and electron affinity directly shape how elements bond:

• Elements with smaller radii and higher ionization energies tend to share electrons, favoring covalent bonding. Carbon is a classic example, forming four strong covalent bonds.
• Elements with larger radii and lower ionization energies tend to transfer electrons, favoring ionic bonding. Sodium chloride (NaCl) forms when sodium loses an electron and chlorine gains one.

Reactivity of Metals and Nonmetals

• Alkali metals have low ionization energies and low electronegativities, so they lose electrons easily. Sodium reacts vigorously with water, producing hydrogen gas and sodium hydroxide.
• Halogens have high ionization energies and high electronegativities, so they gain electrons readily. Chlorine is a powerful oxidizing agent, stripping electrons from many metals and nonmetals.

Transition Metals

Transition metals sit in the d-block and have intermediate ionization energies and electronegativities. A defining feature is their ability to form multiple oxidation states. Iron, for instance, commonly exists as $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$, each with distinct chemical properties (rust is primarily $\text{Fe}^{3+}$ oxide, while $\text{Fe}^{2+}$ compounds tend to be pale green in solution). This versatility gives transition metals a wide range of reactivities.

Predicting Bonding and Stability

You can use electronegativity differences between two elements to predict bond type:

• Large electronegativity difference (typically > 1.7) → ionic bonding. Potassium fluoride (KF) is a stable ionic compound because potassium and fluorine differ greatly in electronegativity.
• Small electronegativity difference → covalent bonding. Water ($\text{H}_2\text{O}$) and carbon dioxide ($\text{CO}_2$) involve atoms with relatively similar electronegativities sharing electrons.

Stability also follows from these trends. Ionic compounds formed between elements at opposite ends of the electronegativity scale (like KF) tend to have very high lattice energies and are quite stable. Compounds between elements with similar electronegativities, such as hydrogen peroxide ($\text{H}_2\text{O}_2$), can be less stable and more reactive.