Periodic Properties of the Elements

Vocabulary flashcards covering periodic table basics, quantum numbers, electron configurations, and periodic trends as discussed in the video notes.

Periodic Table

Organizational chart for the elements based on repeating patterns in properties; arranged by increasing atomic number.

Period

A row in the periodic table; indicated by the period number.

Group

A column in the periodic table; contains families of elements with similar properties; indicated by the group number.

1A–8A (IA–VIIIA) and 1B–8B (IB–VIIIB)

Two common numbering conventions for groups in the periodic table.

Alkali Metals

Group 1A elements; highly reactive metals.

Alkaline Earth Metals

Group 2A elements; reactive metals.

Halogens

Group 7A elements; highly reactive nonmetals.

Noble Gases

Group 8A elements; inert, very unreactive gases.

Metals

Typically good conductors of heat and electricity; blue-colored in the notes’ scheme.

Nonmetals

Elements with varied properties, often poor conductors; yellow-colored in the notes’ scheme.

Metalloids (Semimetals)

Elements with properties between metals and nonmetals; green-colored in the notes’ scheme.

Transition Metals

Elements in the d-block; include many metals with variable oxidation states.

Inner Transition Metals

Lanthanides and Actinides; f-block elements.

Main Group Elements

Elements that include metals, nonmetals, and metalloids outside the transition and inner-transition blocks.

Lanthanides

Inner transition metals in the first row (f-block).

Actinides

Inner transition metals in the second row (f-block).

Recognize Elements Up to Krypton (Z=36)

Know names and symbols; example: Fe is iron; Na is sodium.

Electron Configuration

Arrangement of electrons in an atom’s orbitals; includes ground state and excited states.

Ground State

Lowest energy arrangement of electrons.

Principal Quantum Number (n)

Defines the energy level of an orbital.

Angular Momentum Quantum Number (l)

Defines the shape of an orbital (s, p, d, f).

Magnetic Quantum Number (ml)

Defines the orientation of an orbital in space.

Spin Quantum Number (ms)

Represents electron spin; two possible values: +1/2 and −1/2.

Aufbau Principle

Electrons fill the lowest-energy orbitals first, in order of increasing energy.

Pauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers; an orbital holds at most two electrons with opposite spins.

Hund’s Rule

For degenerate orbitals, electrons occupy separate orbitals with the same spin before pairing.

Anomalous Electron Configurations

Deviations like Cr and Cu due to extra stability of half-filled or full subshells.

Valence Electrons

Outermost electrons in the highest energy shell; most involved in bonding and determining properties.

Core Electrons

Inner-shell electrons that shield outer electrons from the nucleus.

Effective Nuclear Charge (Zeff)

net attractive force on a valence electron after shielding; Zeff = Zactual − shielding.

Shielding

Inner electrons reduce the attraction of the nucleus felt by outer electrons.

Atomic Radius (Covalent Radius)

Half the distance between nuclei of two identical bonded atoms; increases down a group, decreases across a period.

Ionic Radius

Radius of an ion; cations are smaller than their parent atoms, anions larger.

Isoelectronic Series

Ions that have the same electron configuration but different nuclear charges; size varies with Z.

Ionization Energy (IE1)

Energy required to remove one electron from a neutral atom in the gas phase.

Electron Affinity (EA)

Energy released when a neutral atom gains an electron; more negative EA means a greater tendency to accept an electron.

Zeff Trends

Zeff increases across a period and, generally, also increases with descending ionization energy trends; shielding influences these trends.

Block Designations (s, p, d, f blocks)

Period table blocks based on the last filled orbital: s-block, p-block, d-block, f-block.

s-Block

Block containing Group 1A–2A (and H/He in some views); last filled orbital is s.

p-Block

Block containing Groups 13–18; last filled orbital is p.

d-Block

Block of transition metals; last filled orbital is d.

f-Block

Block of lanthanides and actinides; last filled orbital is f.

Group Number and Valence Electrons (Main Group)

Group number indicates the number of valence electrons for main-group elements.

Period Number and Valence Shell (Main Group)

Period number corresponds to the principal energy level (n) of the valence electrons.

Noble Gas Configuration

Filled valence shell; atoms gain/lose electrons to achieve ns2 np6 configuration.

Cation Formation (Main Group Metals)

Metals lose valence electrons to form cations matching a noble gas configuration.

Anion Formation (Non-Metals)

Non-metals gain electrons to fill valence orbitals to ns2 np6, forming anions.

Transition Metal Cations

Transition metals typically lose their s electrons before d electrons when forming cations.

K+ Example

Potassium loses its 4s electron to form K+, achieving [Ar] 4s0 configuration.

Chromium Anomalous Configurations

Predicted vs actual: Cr often [Ar] 4s2 3d4 but actual is [Ar] 4s1 3d5.

Copper Anomalous Configurations

Predicted vs actual: Cu often [Ar] 4s2 3d9 but actual is [Ar] 4s1 3d10.

Electron Configuration Notation (Condensed)

Use nearest noble gas core in brackets, then the remaining valence electron configuration (e.g., Na: [Ne] 3s1).

Charge from Group Number (Main Group)

Group number minus 8 gives the charge of the resulting anion for main-group elements.

Examples of Ion Charges (Main Group)

Group 1 elements form +1 cations; Group 2 form +2 cations; Group 17 form −1 anions, etc.

Transition Metal Cation Electron Loss Order

Valence s electrons are lost before d electrons; higher n electrons are lost first.

Effective Nuclear Charge in Periods

Zeff generally increases when moving left to right across a period due to increasing nuclear charge and relatively less shielding.