Lecture 17

The Periodic Law of Elemental Properties

• Recurring patterns in physical and chemical properties of elements.

• Aid scientists in predicting undiscovered elements and understanding elemental behavior.

• Periodic Property: A property that is predictable based on an element's position on the periodic table.

• Mendeleev (1869): Summarized observations in the periodic law:

  • When elements are arranged in order of increasing mass, certain sets of properties recur periodically.

The Modern Periodic Table

• Group Numbering System:

  • A: Main group elements

  • B: Transition elements

Valence Electrons Determine Chemistry

• Number of valence electrons is constant down a given group of the periodic table; this is by design.

Determining Valence Electrons in Nitrogen

• Without writing an electron configuration, determine the number of valence electrons in nitrogen:

  • (a) 3

  • (b) 4

  • (c) 5

  • (d) 6

Elemental Groups (Families)

• Group 1: Alkali metals

• Group 2: Alkaline earth metals

• Group 16: Chalcogens

• Group 17: Halogens

• Group 18: Noble gases

Group 8: Noble Gases

• Have eight valence electrons (except He).

• Nonreactive due to stable electron configuration:

  • All outermost orbitals are filled.

• Octet Rule: Atoms often gain, lose, or share electrons to achieve the same number of electrons as the closest noble gas.

Group 1: Alkali Metals

• Characteristics of Alkali Metals:

  • Soft, silver metals.

  • Low melting points.

  • Highly reactive; often stored in mineral oil.

  • Not naturally found in elemental forms.

  • One valence electron.

  • Predictably form cations with +1 charge when reacting.

Group 2: Alkaline Earth Metals

• Characteristics of Alkaline Earth Metals:

  • Two valence electrons.

  • Predictably form cations with +2 charge when reacting.

  • Shiny, silver-white metals.

  • Higher melting points than alkali metals.

  • Reactive (but less than alkali metals).

  • Not naturally found in elemental form.

Group 7: Halogens

• Characteristics of Halogens:

  • Nonmetals.

  • Have one fewer electron than the next noble gas.

  • Strong desire for one more electron; commonly form anions with -1 charge.

Transition Metals and Inner Transition Metals

• Differences from Main-Group Elements: Exhibit trends differing from main-group elements (s block and p block).

• The 4s sublevel is lower in energy than 3d:

  • The 4s orbital fills before the 3d orbital.

• Irregular Electron Configurations:

  • Some have irregular configurations in which the ns orbital doesn’t fill or only partially fills before the (n−1)d.

  • Examples: Chromium (Cr), Molybdenum (Mo), Copper (Cu), Silver (Ag).

  • These anomalies happen when an s electron jumps to a d orbital to create a half-filled or completely filled sublevel.

• Experimental Electron Configurations:

  • Transition metals configurations must be found experimentally.

  • Example Configurations:

  • $[Ar] 3d^4 4s^2$ for Cr.

  • $[Ar] 3d^{10} 4s^1$ for Cu.

Actinide Series

• Characteristics of Actinide Series:

  • Unstable & radioactive.

  • Silvery or silvery-white luster in metallic form.

  • Can form stable complexes with ligands such as chloride or sulfate.

  • Many occur in nature as sea water or minerals.

Rare Earth Metals

• Known as "the seeds of technology" or "technology metals."

• Unique magnetic, phosphorescent, and catalytic properties.

• Abundant in Earth’s crust.

Metals, Nonmetals, and Metalloids

• Metals:

  • Malleable and ductile.

  • Shiny, lustrous, reflect light.

  • Conduct heat and electricity.

  • Form cations in solution.

  • Lose electrons in reactions (get oxidized).

• Nonmetals:

  • Brittle in solid state.

  • Dull, nonreflective, solid surface.

  • Electrical and thermal insulators.

  • Form anions and polyatomic anions.

  • Gain electrons in reactions (get reduced).

• Metalloids:

  • Shiny & brittle.

  • Exhibit metallic or nonmetallic reactivity.

  • Semiconductors (can either lose electrons to form cations or gain electrons to form anions).

Periodic Trend: Metallic Character

• Metallic Character: How closely an element’s properties match the ideal properties of a metal.

  • More malleable and ductile.

  • Better conductor.

  • Easier to ionize (lose electrons or undergo oxidation).

• Trends:

  • Metallic character decreases from left to right across a period.

  • Metallic character increases down a column.

Example of Metallic Character

• On the basis of periodic trends, choose the more metallic element from each pair:

  • a. Sn or Te: Sn is more metallic than Te.

  • b. P or Sb: Sb is more metallic than P.

  • c. Ge or In: In is more metallic than Ge.

  • d. S or Br: Cannot determine which is more metallic due to counteracting trends.

Periodic Trends: Atomic Radii and Effective Nuclear Charge

• Atomic radius:

  • Increases down a group.

  • Decreases across a period (left to right).

• Why? - To be explored further.

Coulomb’s Law and Charged Particles

• Coulomb’s Law: Describes potential energy ($E$) between charged particles:

  • For identical charges, $E$ is positive (> 0) and decreases as the particles get farther apart (as $r$ increases).

  • For opposite charges, $E$ is negative (< 0) and becomes more negative as the particles get closer.

  • The strength of interaction increases with the size of the charges.

  • Example: Electrons are more strongly attracted to a nucleus with a $2+$ charge than to a nucleus with a $1+$ charge.

Shielding and Effective Nuclear Charge

• In a multi-electron atom, an electron experiences:

  • Attraction to the protons in the nucleus and repulsion by other electrons in the atom.

• Shielding: A net reduction in the attraction to the nucleus felt by an electron due to repulsions from other electrons.

• Effective Nuclear Charge ($Z_{\text{eff}}$): The total attraction that an electron feels for the nucleus’s protons.

  • A shielded electron does not experience the full attraction by protons because other electrons interfere with/block attractive forces.