Chapter 8: Electron Configuration and Chemical Periodicity

Merged flashcards from Chapter 8 of Chemistry: The Molecular Nature of Matter and Change.

Pauli exclusion principle

States that no two electrons in the same atom can have the same four quantum numbers

• Within a specific sublevel and orientation, spin numbers must oppose each other, limiting the number of electrons in each orbital to 2

Nuclear charge (Z)

As this number increases, denoting more protons in the nucleus, sublevels are stabilized due to increased nucleus-electron attractions

• Results in more energy being needed to remove a specific electron

Shielding

The ability of other electrons to lessen the nuclear attraction for an outer electron through repulsion, creating a lower effective nuclear charge (Z_eff) and enabling easier transfers

• Primarily performed by electrons in the same or inner energy levels

Penetration

The phenomenon in which an outer electron moves through the region occupied by the core electrons to spend part of its time closer to the nucleus

• This increases effective nuclear charge and decreases the effects of shielding by other electrons

Ground-state configuration

The lowest energy distribution of electrons in the levels and sublevels of atoms

Aufbau principle

An approach for building up atoms by adding one proton and one electron at a time to obtain the ground-state configurations of elements

Electron configuration

A way to notate the distribution of electrons in sublevels using the form nl^# (like 2s²)

Orbital diagram

A way to visualize electron distributions via boxes filled with up to 2 arrows representing each energy level and electron spin

• Filled according to the lowest and least repulsive electron orbital available

Hund’s rule

States that lowest energy electron configurations maximize unpaired electrons with parallel spins

Partial orbital diagram

An electron diagram with only the filled sublevels present

Condensed electron configuration

A way to write the electron configuration using the noble gas from the previous row to represent the configuration prior to the filled electrons

• Cl can be written as [Ne] 3s²3p⁵

Transition elements

An element that occupies the d or f block of the periodic table with those orbitals filled last

• Fills after the next energy level’s s sublevel due to penetration causing lowered energies

  • Half-filled and filled sublevels, though, may fill first due to greater stability (Cr has [Ar] 4s¹3d⁵ and not 4s²3d⁴ — this also applies to Cu filling 3d early)

Inner electrons (core electrons)

Electrons an atom has in common with the previous noble gas and any completed series

Outer electrons

Electrons with the highest energy level (n) that are farthest from the nucleus

• These correlate with similar chemical behavior due to electron interactions

Valence electrons

Electrons involved in forming compounds

• These are the outer electrons for main-group elements, and include the previous block’s d sublevel as well for transition elements

Atomic size (atomic radius)

A measure of how closely one atom lies next to another, determined from the distances between nuclei of atoms

• Knowing the value of this for an element in a bond can help in finding this for another atom

• Increases with n and decreases with Z_eff due to electrons being present in more distant energy levels or pulled closer to the nucleus

Metallic radius

The atomic size as defined for metals, equal to half of the shortest distance between atoms in a crystal of the element

Covalent radius

The atomic size as defined for molecules and nonmetals, equal to half of the shortest distance between nuclei of bonded atoms

Ionization energy (IE)

The required energy input for the removal of 1 mol of electrons from 1 mol of gaseous atoms or ions

• Can be described as part of a series, numbered with subscripts to refer to removal order

• Can be used to determine reactivity, as low values create cations while high values create anions given greater nuclear attraction

  • Decreases down a group (lower Z_eff, except for group 13)

  • Increases across a period (greater Z_eff, except for group 13 and 16 for periods 2 and 3 due to greater stability)

Core electrons

These are the most difficult to remove due to needing a large ionization energy after removing the valence electrons, thus not being involved in reactions

Electron affinity (EA)

The energy change after the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions

• First is typically negative due to attraction, as energy is released

• Second is always positive, as energy must be absorbed to add another electron

• While sublevel energies and repulsions affect trend regularity, some can be stated:

  • Very exothermic for reactive nonmetals due to negative attraction

  • Slightly exothermic for reactive metals, due to positive ion affinity instead

  • Slightly endothermic for noble gases, due to little electron attraction

Metallic behavior

The properties of being shiny, conductive, ductile, malleable, and generally positive; these decrease going across a period and increase going down a group

Electron loss

A property of metals relating to charge that:

• Increases down a group due to lower Z_eff and IE

• Decreases across a period due to higher Z_eff and IE

• Creates strong reducing agents in lower groups, and strong oxidizing agents in higher groups

Redox behavior

A property of metals in aqueous solutions that:

• In lower groups (metals): Involves electron transfer to form ionic oxides, which create bases by generating OH^- in water (like MgO + H₂O → Mg²⁺ + 2OH^-)

• In higher groups (nonmetals): Involves electron sharing to form covalent oxides, which create acids by releasing H⁺ in water (like SO₃ + H₂O → 2H⁺ + SO₄^2-)

Isoelectronic

Having the same number and configuration of electrons as another species

• Elements at either end of a period typically target a similar configuration to noble gases, forming anions or cations for full sublevels

  • Transition metals lose ns^# sublevels first

• Electrons lost or gained must be denoted as e^-

Paramagnetic

A species with at lest one unpaired electron that is attracted by an external magnetic field

• Includes many transition metals and their compounds due to unpaired electrons

Dimagnetic

A species with all paired electrons that are not attracted by a magnetic field

Ionic radius

The size of an ion, measured through nuclear distance between ions in a crystal

• Cations become smaller due to weaker pull of electrons

• Anions become larger due to greater pull of electrons

• Can be ordered by first accounting for n, then realizing that increasing Z_eff reduces size after adding or removing electrons (important when isoelectronic)