CHEM 1311H - Electron configurations and periodic trends

atomic radius

½ the distance between the nuclei of two identical atoms joined by a single bond

• increases from right to left across the period

  • due to Z_eff increasing as protons inc and draws the valance electrons closer together 

• increases down a group 

  • valence electrons in larger atomic orbitals result in larger atoms

ionic radius

the distance between the nuclei of ions joined by an ionic bond

• increases from right to left across the period

• increases down a group

trends in ionic sizes

• cations: generally smaller compared to neutral parent atom

• anions: generally larger compared to neutral parent atom; extra electrons cause repulsion to be larger

isoelectronic series

• anion > neutral atom > cation

• the higher number of protons in isoelectronic series, the smaller the ionic radii

ionization energy

the amount of energy required to remove an electron from a neutral atom or an ion in the gaseous phase

• X (g) → X⁺ (g) + 1e^-

• always positive or endothermic

• increases as electrons are being removed (due to electron being pulled from an inc positive ion)

  • IE₁ > IE₂ > IE₃

first ionization energy trend

• higher Z_eff → atomic radius dec → stronger nuclear attraction → removing an electron from a shell is hard

• increases across a period

• decreases down a group

• exceptions:

  • group 2 is higher than expected

  • group 13 is lower than expected

  • group 16 elements is lower than group 15

  • group 18 is very high

electron attachment enthalpy (ΔH_ea)

the energy released when an electron is added to a neutral atom in the gaseous phase to form a gaseous phase anion

• X (g) + e^- —→ X^- (g) ΔH_ea

• can be negative due to the attraction of the electron to the nucleus → exothermic

• the more energy release (the more negative), the easier to add an electron

• more favorable from left to right across the period

  • Z_eff inc across a period

• less favorable going down a group

  • n inc and e^- are less attracted to nucleus

• exceptions:

  • period 2 is less favorable than period 3 because p2 has small electron clouds and higher density, and electron repulsion is greater in 2 than 3

  • group 2 is less favorable than expected because of full subshells

  • group 15 is less favorable because of half-full subshells

  • novle gases is less favorable because of full subshells

polarizability

the ease with which the electron cloud of an atom, ion, or molecule can be distorted by an electric field or another atom

• the larger the atomic size, the greater the polarizability

• dec from left to right across a period

• inc from top to bottom within a group

electronegativity

the ability of an atom to attract electrons to itself in a chemical bond

• inc across the period

• dec down a group

F, O, N, Cl 

make polar bonds (if bonded with a non-metal) or ionic bonds (if bonded to a metal)

lattice energy

the energy released when ions in a gaseous phase combine to give one mole of crystalline solid

• M⁺ (g) + X^- (g) → MX (s) ΔH_lattice = (-)

• generally exothermic

• gives an estimate of the strength of the attractive forces between two oppositely charged ions in an ionic compound

• inc as charges on the ions become larger

• inc as the ionic sizes become smaller

• ionic sizes are more important than ionic size

effective nuclear charge Z_eff

• actual attraction that an electron feels from a nucleus

• shielding causes lower Z_eff = higher energy (due to the less attraction to the nucleus)

• increases left to right across a period

Cr electron configuration

[Ar]4s¹3d⁵

Mo electron configuration

[Kr]5s¹4d⁵

Cu electron configuration

[Ar]4s¹3d¹⁰

Ag electron configuration

[Kr]5s¹4d¹⁰

Au electron configuration

[Xe]6s¹4f¹⁴5d¹⁰

shielding

the ability of electrons closer to the nucleus to reduce the effectiveness of the nucleus in attracting more-distant electrons, reducing the nuclear charge felt by the more-distant electrons

penetration

the ability of inner electrons to get closer to the nucleus versus outer electrons

• inner e have greater attraction resulting to a lower energy

• an electron in an orbital with good penetration is better at shielding than one with low penetration

  • s > p > d > f