Chemistry Chapter on Noble Gases & Electron Config

Noble Gases and Electron Configuration

• Noble gases have completely filled p subshells, with helium having a fully filled 1s subshell.
• Transition metals are located in Groups 1B and 3B through 8B and have incompletely filled d subshells or can form cations with such characteristics.
• Group 2B includes zinc (Zn), cadmium (Cd), and mercury (Hg), which are not classified as representative elements or transition metals.
• Lanthanides and actinides are characterized by incompletely filled f subshells.

Electron Configuration and Chemical Reactivity

• Valence electrons are the outermost electrons that determine chemical reactivity, while core electrons are nonvalence electrons.
• Elements in the same group (column) of the periodic table share the same number and type of valence electrons, suggesting similar chemical behaviors.
• Alkali metals (Group 1) have a valence electron configuration of ns1 and typically lose one electron to form +1 cations.
• Alkaline earth metals (Group 2) possess a configuration of ns2 and tend to lose two electrons, forming +2 cations.
• Transition metals have a valence configuration of (n-1)dx ny, where the values x and y vary.
• Halogens (Group 17) have the configuration ns2 np5, while noble gases (Group 18) have ns2 np6, except helium (1s2), signifying stability and low reactivity.

Physical Properties and Classification

• Periodic Table elements can be categorized as metals, nonmetals, or metalloids.
• Metals are typically good heat and electricity conductors, malleable, and ductile, with most being solid except mercury, which is liquid.
• Nonmetals are usually poor conductors, with gases, liquids (like bromine), and solids (like carbon) represented among them.
• Metalloids possess intermediate properties, behaving like nonmetals but capable of conducting electricity under specific conditions.

Periodicity in Properties

• Effective nuclear charge (Z_eff) represents the nuclear charge experienced by an electron, factoring both the actual charge (Z) and the shielding effects from other electrons.
• Within a period, Z_eff increases, leading to a greater pull on valence electrons.
• In groups, the effective nuclear charge remains constant, but the increase in principal quantum number decreases the attraction of the nucleus to valence electrons.

Atomic and Ionic Radius

• The atomic radius decreases from left to right across a period due to increasing Z_eff.
• Conversely, the atomic radius increases down a group because of an increase in the principal quantum number (n).
• Ionic radii increase with added negative charges and decrease with increased positive charges; ionic radii comparisons can be made using isoelectronic series and charge considerations.

Ionization Energy

• Ionization energy is the minimum energy required to remove an electron from a gaseous atom; energy levels are defined for successive ionization processes.
• Ionization energy generally increases across a period due to greater Z_eff and decreases down a group due to increased distance between the nucleus and valence electrons.
• Alkali metals exhibit the lowest first ionization energies among groups due to their single valence electron being effectively shielded.