Chapter 4: Periodic Trends and Properties of Elements

Flashcards covering periodic trends (atomic radius, ionization energy, electron affinity), element classifications, and ionic electron configurations.

What extra credit opportunity is available for attending the chemistry department cookout on Friday the 25th?

Students can earn $+5$ points on a quiz by writing about the food, a professor, or a classmate they met at the cookout and submitting it on Canvas.

What is atomic radius?

Atomic radius is the distance between the nucleus and the last electron or the last shell of electrons of an atom.

What are the periodic trends for atomic radius?

Atomic radius increases from top to bottom down a group (due to additional electron shells) and increases from right to left across a period (due to decreasing effective nuclear charge, $Z_{\text{eff}}$).

What is the order of phosphorus, sulfur, and oxygen in order of increasing atomic radius?

Oxygen < Sulfur < Phosphorus

What is ionization energy?

Ionization energy is the minimum energy required to remove an electron from an atom in its gas state.

How do anions and cations differ in charge and formation?

Anions are negatively charged ions formed by gaining electrons, while cations are positively charged ions formed by losing electrons.

What is the periodic trend for first ionization energy?

First ionization energy increases from left to right across a period and increases from bottom to top up a group.

How does subshell type ($s$, $p$, $d$, $f$) affect the ease of electron removal?

Electrons in $s$ subshells are held most tightly and are hardest to remove, whereas electrons in $f$ orbitals are furthest from the nucleus and easiest to remove.

Why is removing a second electron from a sodium ($Na$) atom significantly harder than removing the first?

The first electron removed is a valence electron from the $3s$ subshell, whereas the second electron must be removed from a full core subshell ($2p$), which experiences a much higher effective nuclear charge.

What is electron affinity?

Electron affinity is the energy released when an atom in the gas phase accepts an electron.

What is the general periodic trend for electron affinity across a period?

Electron affinity increases from left to right across a period as the effective nuclear charge increases.

Why do subsequent electron affinities require energy input compared to the first electron affinity?

Adding an electron to an ion that is already negatively charged causes electrostatic repulsion, making each subsequent electron addition harder.

Why does phosphorus ($P$) require more energy to add an electron compared to silicon ($Si$)?

Phosphorus has a half-filled $3p$ subshell ($3p^3$) where each orbital has one electron; adding an extra electron requires doubling up in an orbital, which is higher in energy according to Hund's rule.

What are the primary properties of metals?

Metals are shiny/lustrous, malleable, ductile, good conductors of heat and electricity, have low ionization energies, and tend to form cations.

What are the primary properties of nonmetals?

Nonmetals are dull, brittle, poor conductors of heat and electricity, have high electron affinities, and tend to form anions.

What is the electron configuration of a sodium cation ($Na^+$)?

$1s^2 2s^2 2p^6$ (which is equivalent to the noble gas configuration $[Ne]$).

What does the term 'isoelectronic' mean?

Isoelectronic describes chemical species (atoms or ions) that possess the exact same electron configuration.

Which ions are isoelectronic with neon ($Ne$)?

Examples include $Na^+$ (sodium cation), $F^-$ (fluoride anion), and $O^{2-}$ (oxide anion).