Periodic Trends and Properties of Elements: Study Guide (Chapters 3, 4, 5, 6)

Periodic Trends and Properties of Elements

Classification of Elements

The periodic table organizes elements based on their properties. Elements can be classified as metals, nonmetals, or metalloids, each with distinct physical and chemical characteristics.

• Metals: Good conductors of heat and electricity, malleable, ductile, and typically have a shiny appearance. Found on the left and center of the periodic table.

• Nonmetals: Poor conductors, not malleable or ductile, often brittle in solid form, and found on the right side of the periodic table.

• Metalloids: Have properties intermediate between metals and nonmetals. Located along the zigzag line (staircase) on the periodic table.

Example: Silicon (Si) is a metalloid, while sodium (Na) is a metal and chlorine (Cl) is a nonmetal.

Element Families and Their Properties

Certain columns (groups) in the periodic table are known as families, with members sharing similar properties.

• Alkali Metals (Group 1): Li, Na, K, Rb, Cs, Fr. Highly reactive, form +1 ions.

• Alkaline Earth Metals (Group 2): Be, Mg, Ca, Sr, Ba, Ra. Reactive, form +2 ions.

• Halogens (Group 17): F, Cl, Br, I, At. Very reactive nonmetals, form -1 ions.

• Noble Gases (Group 18): He, Ne, Ar, Kr, Xe, Rn. Inert gases, rarely form ions.

Example: Sodium (Na) is an alkali metal; chlorine (Cl) is a halogen.

Predictable Ion Charges

Many elements form ions with predictable charges based on their group number.

• Group 1: +1

• Group 2: +2

• Group 13: +3

• Group 15: -3

• Group 16: -2

• Group 17: -1

Example: Magnesium (Mg) forms Mg2+; oxygen (O) forms O2−.

Trends in Atomic Size (Atomic Radius)

Atomic size varies systematically across periods and down groups in the periodic table.

• Across a Period (Left to Right): Atomic radius decreases due to increasing nuclear charge pulling electrons closer.

• Down a Group (Top to Bottom): Atomic radius increases as additional electron shells are added.

Example: Atomic radius: Na > Mg > Al (across period); Li < Na < K (down group).

Arranging Atoms by Size

Given a list of elements, you can arrange them by increasing or decreasing atomic radius using periodic trends.

• Order increases down a group and decreases across a period.

Example: For Li, Na, K: K > Na > Li (increasing size). >greater than < lesser than

Trends in Ionic Radii

Ionic radii refer to the size of ions. Trends are similar to atomic radii but also depend on the charge of the ion.

• Cations (positive ions): Smaller than their parent atoms due to loss of electrons.

• Anions (negative ions): Larger than their parent atoms due to gain of electrons and increased electron-electron repulsion.

• Across a period: Ionic radius decreases for cations and then for anions.

• Down a group: Ionic radius increases.

Example: Na+ < K+; Cl− < Br−.

Isoelectronic Species

Isoelectronic species are atoms and ions that have the same number of electrons.

• Definition: Species with identical electron configurations.

• Examples: Ne, Na+, F− (all have 10 electrons).

Arranging Isoelectronic Species by Size: For isoelectronic species, the more positive the charge, the smaller the radius.

• Example: O2− > F− > Na+ > Mg2+ (all have 10 electrons). >greater than < less than

Trends in Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• First Ionization Energy (IE1): Energy to remove the first electron.

• Second Ionization Energy (IE2): Energy to remove the second electron.

• Across a Period: Ionization energy increases due to higher nuclear charge.

• Down a Group: Ionization energy decreases as electrons are farther from the nucleus.

Equation:

$\text{X}(g) \rightarrow \text{X}^+(g) + e^- \quad (\text{IE}_1)$

$\text{X}^+(g) \rightarrow \text{X}^{2+}(g) + e^- \quad (\text{IE}_2)$

Example: IE1 of Na < IE1 of Mg.

Exceptions to Ionization Energy Trends

Some elements do not follow the general trend due to electron configurations (e.g., full or half-filled subshells).

• Example: Boron (B) has a lower IE1 than beryllium (Be) due to the start of p-orbital filling.

• Example: Oxygen (O) has a lower IE1 than nitrogen (N) due to electron pairing in the p-orbital.

Successive Ionization Energies

Each successive ionization energy is higher than the previous one because it is harder to remove an electron from a positively charged ion.

• Reason: Increased effective nuclear charge after each electron is removed.

Example: For magnesium: $\text{IE}_1 < \text{IE}_2 < \text{IE}_3$

Summary Table: Periodic Trends

Property	Across a Period (→)	Down a Group (↓)
Atomic Radius	Decreases	Increases
Ionic Radius	Decreases (for cations/anions separately)	Increases
Ionization Energy	Increases	Decreases

Additional info: These notes synthesize and expand on the provided study guide points, adding definitions, examples, and a summary table for clarity and completeness.