Periodicity

Atomic Radii covalent radius metallic radius Atomic radius is one-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule (or one-half the distance between the nuclei in two neighboring atoms).

Ionic Radius radius of an atoms ion Trend compares ions to their neutral counterparts or cations with cations and anions with anions

Ions Variation in the number of electrons; results in an atom becoming charged Anion Cation gain of electrons overall charge is negative lose of electrons overall charge is positive 11 protons 12 neutrons 11electrons 10 electrons Sodium atom Sodium ion 1+ Na 1− F

Cations and Anions Of Representative Elements +1 +2 +3 -3 -2 -1

Cation vs Atom Increases Increases

Anion vs Atom Increases Increases

The Radii (in pm) of Ions of Familiar Elements

Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5 *Most transition metals can form more than one cation and frequently the cations are not isoelectronic with the preceding noble gas *The order of electron filling does not determine or predict the order of electron removal for transition metals.

Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

Ionization Energy Minimum energy required to remove an electron from a gaseous atom in its ground state The magnitude of IE is a measure of how “tightly” an atom’s nucleus holds onto its valence electron (columbic attraction)

IE1 + X (g) X+(g) + e- IE1 first ionization energy IE2 + X+(g) X2+(g) + e- IE2 second ionization energy IE3 + X2+(g) X3+(g) + e- IE3 third ionization energy IE1 < IE2 < IE3 *Chemical properties of any atom are determined by the configuration of the atom’s valence electrons. The stability of these outermost electrons is reflected directly in the atom’s ionization energies.

*Note: while valence electrons are relatively easy to remove from the atom, core electrons are much harder to remove. Thus, there is a large jump in ionization energy between the last valence electron and the first core electron.

General Trends in First Ionization Energies Increasing First Ionization Energy Increasing First Ionization Energy

Exceptions Between Group 2A and 3A (ex: Be and B) Group 3A is lower b/c of the single electron in the outermost p subshell which is shielded Between Group 5A and 6A (ex: N and O) Group 6A doubles up one electron, the proximity of two electrons in the same orbital results in greater electrostatic repulsion and lower energy

Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. *or think of it as the energy that must be supplied to remove an electron from the anion. X (g) + e− X− (g) F (g) + e− F− (g) DH = -328 kJ/mol EA = +328 kJ/mol *lg. pos. electron affinity means that the negative ion is very stable O (g) + e− O− (g) DH = -141 kJ/mol EA = +141 kJ/mol *The more positive the electron affinity is of an element, the greater the affinity of an atom of the element to accept an electron.

Electronegativity Chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself An atoms electronegativity is affected by both its atomic number and the distance at which its valance electrons reside from the charged nucleus (columbic attraction)