Periodic Properties of the Elements

Classification and Properties of Elements

The periodic table organizes elements based on recurring chemical properties. Understanding the classification of elements is essential for predicting their behavior in chemical reactions.

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

• Nonmetals: Poor conductors, brittle in solid form, and often found as gases or dull solids. Located on the right side of the periodic table.

• Metalloids (Semi-metals): Exhibit properties intermediate between metals and nonmetals. Found along the staircase line (e.g., B, Si, Ge, As, Sb, Te).

• Families/Groups: Vertical columns with similar chemical properties. Examples include alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), and noble gases (Group 18).

Example: Sodium (Na) is an alkali metal, while chlorine (Cl) is a halogen and a nonmetal.

Periodic Table Blocks and Electron Configuration

Elements are grouped into blocks based on their valence electron configurations:

• s-block: Groups 1 and 2, plus helium.

• p-block: Groups 13-18.

• d-block: Transition metals (Groups 3-12).

• f-block: Lanthanides and actinides (inner transition metals).

Subshells within a shell are arranged in order of increasing energy: .

• Maximum electrons per orbital: 2 (due to the Pauli Exclusion Principle).

• Electron configuration notation: Shorthand uses noble gas core (e.g., [Ne] 3s2 3p5 for Cl).

Electron Configuration Principles

• Aufbau Principle: Electrons fill orbitals of lowest energy first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Example: The electron configuration of oxygen (O):

Coulomb’s Law, Shielding, and Penetration

• Coulomb’s Law: The force between two charged particles increases when the charges are larger and decreases when they are farther apart.

  • (this is why IE inc. L->R and Atomic Radius dec. L->R)

• Shielding: Core electrons reduce the effective nuclear charge felt by valence electrons.

• Penetration: The ability of an electron to get close to the nucleus; s orbitals penetrate more than p, d, or f.

Effective Nuclear Charge (Zeff)

• Definition: The net positive charge experienced by valence electrons.

• Calculation: , where Z is the atomic number and S is the number of core electrons.

Isoelectronic Species

• Definition: Atoms and ions with the same number of electrons.

• Example: , , , and are isoelectronic (10 electrons each).

Magnetism

• Paramagnetism: Atoms with unpaired electrons; attracted to magnetic fields.

• Diamagnetism: Atoms with all electrons paired; weakly repelled by magnetic fields.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy (IE): Energy required to remove an electron; decreases down a group, increases across a period.

• Electron Affinity (EA): Energy change when an electron is added; generally becomes more negative across a period.

• Metallic Character: Increases down a group, decreases across a period.

Ion Size Comparisons

• Cations are smaller than their parent atoms; anions are larger.

• Within an isoelectronic series, higher nuclear charge means smaller radius.

Ionization Energy Equations

• First Ionization:

• Second Ionization:

Large jumps in ionization energy indicate removal of core electrons.

Molecules and Compounds

Elemental Forms and Diatomic Molecules

• Certain elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

• Phosphorus exists as P4; sulfur as S8.

Chemical Formulas and Interpretation

• Chemical formula: Indicates the types and numbers of atoms in a compound (e.g., H2O: 2 H, 1 O).

• To write a formula, use the number of each atom present.

Ion Charges and Polyatomic Ions

• Group Charges: IA: +1, IIA: +2, IIIA: +3, VA: -3, VIA: -2, VIIA: -1.

• Scandium: +3, Aluminum: +3, Zinc: +2, Cadmium: +2, Silver: +1.

• Polyatomic ions: Memorize common ones (e.g., nitrate, sulfate, ammonium).

Naming Compounds

• Ionic Compounds: Name cation first, then anion. Use Roman numerals for transition metals with variable charge (e.g., iron(III) chloride).

• Molecular Compounds: Use Greek prefixes (mono-, di-, tri-, etc.) to indicate number of atoms (e.g., CO2: carbon dioxide).

• Hydrates: Compounds with water molecules; named with prefixes (e.g., CuSO4·5H2O: copper(II) sulfate pentahydrate).

• Common Names: Some compounds have accepted common names (e.g., H2O: water, NH3: ammonia).

Percent Composition and Molar Mass

• Percent composition:

• Molar mass: Sum of atomic masses of all atoms in the formula (g/mol).

• Avogadro’s number: particles/mol.

Empirical and Molecular Formulas

• Empirical formula: Simplest whole-number ratio of elements in a compound.

• Molecular formula: Actual number of atoms of each element; a multiple of the empirical formula.

• To determine molecular formula: , where

Combustion Analysis

• Used to determine empirical/molecular formulas of compounds containing C, H, N, O by measuring masses of CO2, H2O, etc. produced.

Chemical Bonding

Valence Electrons and Bond Types

• Valence electrons: Electrons in the outermost shell; determine chemical reactivity.

• Bond types: Single (2 electrons), double (4 electrons), triple (6 electrons).

• Bond length: single > double > triple; bond strength: triple > double > single.

Lewis Dot Structures

• Show valence electrons as dots around element symbols.

• For molecules and polyatomic ions, obey the octet rule (8 electrons around each atom, except H: 2 electrons).

• Draw resonance structures if multiple valid arrangements of electrons exist.

Bonding Definitions

• Ionic bond: Electrostatic attraction between cations and anions.

• Covalent bond: Sharing of electron pairs between atoms.

• Polar covalent bond: Unequal sharing of electrons due to difference in electronegativity.

• Nonpolar covalent bond: Equal sharing of electrons.

• Electronegativity: Tendency of an atom to attract electrons in a bond; increases across a period, decreases down a group.

Bond Polarity and Strength

• Bond polarity increases with greater difference in electronegativity.

• Bond length decreases and bond strength increases with more shared electron pairs.

Formal Charge

• Used to determine the most stable Lewis structure.

• Formula:

Octet Rule Exceptions

• Odd electron species: Molecules with an odd number of electrons (e.g., NO).

• Incomplete octets: Atoms with fewer than 8 electrons (e.g., Be, B).

• Expanded octets: Atoms in period 3 or higher can have more than 8 electrons (e.g., SF6).

VSEPR Theory and Molecular Geometry

• VSEPR (Valence Shell Electron Pair Repulsion) Theory: Predicts molecular shapes based on repulsion between electron pairs.

• Electronic geometry: Arrangement of all electron groups (bonding and lone pairs).

• Molecular geometry: Arrangement of atoms (ignoring lone pairs).

• Bond angles: Determined by geometry and presence of lone pairs (e.g., tetrahedral: 109.5°, trigonal planar: 120°, linear: 180°).

Molecular Polarity

• Determined by both bond polarity and molecular geometry.

• Show bond polarity with arrows; molecule is polar if dipoles do not cancel.