Chapter Three: Periodic Properties and Trends in Chemistry

Periodic Table

Arrangement of elements by increasing atomic number.

Dobereiner's Triads

Groups of three elements with similar properties.

Law of Octaves

Element properties repeat every eighth element.

Mendeleev's Periodic Table

Elements arranged by increasing mass, showing periodicity.

Modern Periodic Table

Elements arranged by increasing atomic number.

Periods

Horizontal rows in the periodic table.

Groups

Vertical columns containing similar elements.

Electron Configuration

Distribution of electrons in an atom's orbitals.

Ground State

Lowest energy state of an electron.

Schrödinger's Equation

Describes electron behavior in atoms.

Electron Spin

Fundamental property indicating electron orientation.

Spin Quantum Number

Fourth quantum number indicating electron spin direction.

Pauli Exclusion Principle

No two electrons can have identical quantum numbers.

Orbital Diagrams

Visual representation of electron distribution in orbitals.

Degenerate Orbitals

Orbitals with the same energy level.

Effective Nuclear Charge

Net attraction an electron feels from the nucleus.

Shielding Effect

Reduction of nuclear charge felt by outer electrons.

Penetration

Ability of an electron to approach the nucleus.

Radial Distribution Function

Describes electron probability density around the nucleus.

Multi-electron Atoms

Atoms with more than one electron.

Quantum Numbers

Set of numbers describing electron states.

Charge Interaction

Forces between charged particles in an atom.

Energy Sublevels

Different energy levels within principal energy shells.

Aufbau Principle

Electrons fill orbitals from lowest to highest energy.

Pauli Exclusion Principle

No more than two electrons per orbital.

Hund's Rule

Electrons occupy unfilled orbitals before pairing.

Condensed Electronic Configuration

Uses noble gas notation for electron arrangement.

Valence Electrons

Electrons in the outermost energy level.

Core Electrons

Electrons not in the outermost energy level.

Anomalous Electron Configurations

Unexpected electron arrangements in certain elements.

Cation

Positively charged ion formed by losing electrons.

Anion

Negatively charged ion formed by gaining electrons.

Atomic Radius

Average distance between atomic nuclei in a molecule.

Effective Nuclear Charge (Zeff)

Charge experienced by an electron from the nucleus.

Paramagnetic

Atoms with unpaired electrons, attracted to magnets.

Diamagnetic

Atoms with all paired electrons, not attracted to magnets.

Periodic Table Blocks

Divided into s, p, d, and f sublevels.

Group Number

Indicates number of valence electrons in main-group elements.

Row Number

Indicates highest principal quantum number of elements.

Sodium Electron Configuration

1s2 2s2 2p6 3s1 or [Ne]3s1.

Ion Formation

Atoms gain or lose electrons to achieve stability.

Magnetic Properties

Determined by the presence of unpaired electrons.

Periodic Trend in Atomic Radius

Increases down a group, decreases across a period.

Electron Configuration of Ions

Ions mimic noble gas configurations for stability.

Electron Configuration and Reactivity

Valence electrons dictate chemical behavior of elements.

Zeff

Effective nuclear charge experienced by electrons.

s (shielding constant)

Represents electron shielding from nuclear charge.

Ionic radius

Size of an ion compared to its atom.

Cation radius

Smaller than corresponding neutral atom radius.

Anion radius

Larger than corresponding neutral atom radius.

Isoelectronic ions

Ions with the same number of electrons.

Ionization energy (IE)

Energy required to remove an electron from an atom.

First Ionization Energy

Energy to remove the first electron from neutral atom.

Second Ionization Energy

Energy to remove the second electron from an ion.

Trends in IE

IE decreases down a group, increases across a period.

Electron Affinity (EA)

Energy change when an electron is added to an atom.

Positive EA

Indicates energy is released when gaining an electron.

Negative EA

Indicates energy is absorbed when gaining an electron.

Trends in EA

No definite trend across periodic table groups.

Metallic Character

Tendency of an element to lose electrons.

Effective nuclear charge

Net positive charge experienced by valence electrons.

Successive Ionization Energies

Energy increases with each successive electron removal.

Periodic trend in ionic radii

Ion size increases down a group.

First IE exceptions

Notable deviations in ionization energy trends.

Electron pairing

Occurs when adding electrons to orbitals.

Atomic size arrangement

Order elements by decreasing atomic size.

Comparing ionic sizes

List ions by increasing ionic size.