Chapter 3 Periodic Properties

Periodic Law & Periodic Table

Periodic Law → Physical and chemical properties repeat predictably when elements are arranged by increasing atomic number.

Mendeleev arranged elements by → Increasing atomic mass and similar properties.

Why did Mendeleev leave gaps? → He predicted undiscovered elements would fill them.

What did Mendeleev successfully predict? → Elements such as gallium (Ga) and germanium (Ge).

Modern periodic table is arranged by → Increasing atomic number.

Electron Configurations

Pauli Exclusion Principle → No two electrons in the same atom can have identical sets of quantum numbers.

Electron configuration notation: 1s² — what does 1 mean? → Energy level.

Electron configuration notation: 1s² — what does s mean? → Type of orbital.

Electron configuration notation: 1s² — what does ² mean? → Number of electrons in those orbitals.

Orbital diagram box → Represents one orbital.

Orbital diagram arrow → Represents an electron.

Arrow direction → Represents relative electron spin.

Hund’s Rule → Equal-energy orbitals fill singly first with parallel spins.

Which fills first: 4s or 3d? → 4s.

Why does 4s fill before 3d? → 4s is lower in energy.

Valence electrons → Outermost electrons; important in chemical bonding.

Core electrons → Electrons in complete inner energy levels.

Shielding & Effective Nuclear Charge

Shielding/screening → Repulsion of an electron by other electrons from the full nuclear charge.

Effective nuclear charge formula → Zeff = Z − S

Z in Zeff = Z − S → Atomic number.

S in Zeff = Z − S → Screening constant, usually close to the number of inner electrons.

Greater penetration toward nucleus means → Electron experiences more nuclear attraction.

Why are there different energy sublevels? → Shielding and penetration affect electron energy.

Periodic Table & Electron Configuration

Periodic table blocks correspond to → Different orbital types.

s-block → Left side of periodic table.

p-block → Right side.

d-block → Transition metals.

f-block → Inner transition metals.

Condensed electron configuration uses → Previous noble gas in brackets.

Se condensed configuration → [Ar] 4s² 3d¹⁰ 4p⁴

Bi condensed configuration → [Xe] 6s² 5d¹⁰ 4f¹⁴ 6p³

Cu electron configuration exception → [Ar] 4s¹ 3d¹⁰

Why does Cu have an unusual configuration? → 4s and 3d orbitals are very close in energy.

Element Types

Metals are located → Lower left and middle of periodic table.

Metal properties → Shiny, malleable, ductile, good conductors.

Metals tend to → Lose electrons.

Nonmetals are located → Upper right.

Nonmetal properties → Often gases or dull/brittle substances; poor conductors.

Nonmetals tend to → Gain electrons.

Metalloids are located → Along the stair-step line.

Metalloids have → Properties of both metals and nonmetals.

Many metalloids are → Semiconductors.

Important Families

Group 1A → Alkali metals.

Group 2A → Alkaline earth metals.

Group 7A → Halogens.

Group 8A → Noble gases.

Alkali metal valence configuration → ns¹

Alkali metals usually form → +1 ions.

Alkaline earth metals have → ns²

Alkaline earth metals usually form → +2 ions.

Halogens need how many electrons for a full shell? → 1.

Halogens usually form → −1 ions.

Noble gases have how many valence electrons? → 8, except He has 2.

Why are noble gases unreactive? → Their electron configurations have full, especially stable shells.

What largely determines an element's chemical properties? → Number of valence electrons.

Atomic Radius — VERY IMPORTANT

Bonding atomic radius → Half the distance between two covalently bonded nuclei.

Atomic radius across a period → → DECREASES

Atomic radius down a group ↓ → INCREASES

Why does radius decrease left → right? → Increasing Zeff pulls electrons closer.

Why does radius increase going down? → Increasing principal energy level n.

Where are the largest atoms? → Bottom-left of periodic table.

Cation compared with parent atom → Smaller.

Why is a cation smaller? → It loses valence electron(s).

Anion compared with parent atom → Larger.

Why is an anion larger? → It gains electron(s).

Ionic size down a group → Increases.

Ionization Energy — VERY IMPORTANT

Ionization energy (IE) → Energy required to remove an electron from a gaseous atom or ion.

First ionization energy (IE₁) → Energy required to remove the first electron.

Successive ionization energies → Get progressively larger.

Huge jump in IE means → All valence electrons have been removed and you're starting to remove core electrons.

Alkali metals have a huge jump at → IE₂

Alkaline earth metals have a huge jump at → IE₃

Ionization energy across a period → → Generally INCREASES

Ionization energy down a group ↓ → DECREASES

Why does IE increase left → right? → Zeff increases.

Why does IE decrease going down? → Valence electrons are farther from the nucleus.

Ionization Energy Exceptions

IE exception: Group IIA → IIIA → Electron begins being removed from a p orbital, which is farther from the nucleus.

IE exception: Group VA → VIA → Electron is removed from a doubly occupied orbital, where electron-electron repulsion makes removal easier.

Electron Affinity

Electron affinity (EA) → Energy change when an electron is added to a gaseous atom.

Example of electron affinity → Cl(g) + e⁻ → Cl⁻(g)

Electron affinity left → right → Generally becomes more exothermic.

EA exception: Group 1A → 2A → Added electron must enter a p orbital instead of an s orbital.

EA exception: Group 4A → 5A → Added electron must enter an already occupied orbital, causing repulsion.

⭐ The Trend Card to Memorize

Atomic Radius:
← INCREASES
↓ INCREASES

Ionization Energy:
→ INCREASES
↑ INCREASES

Metallic Character:
← INCREASES
↓ INCREASES

Electron Affinity: → generally becomes more exothermic

🧠 Easy Memory Trick

Think:

BOTTOM LEFT = BIG + METALLIC

TOP RIGHT = TIGHT + HARD TO REMOVE ELECTRONS