Electron Configurations and Periodic Trends

Orbital Shapes Quantum Numbers Atom Electron Configurations Electron Configuration – complete description of orbitals occupied by all the electrons in an atom • for atoms in ground state – electrons occupy energy shells, subshells and orbitals that give the lowest energy for the atom • start with 1s orbital and work up 1 Electron Configurations Electron Configurations for orbitals with same energy (degenerate) such as the three 2p orbitals • half-fill each orbital with first 3 electrons • go back and pair electrons with 4, 5 and 6 Hund’s Rule – electrons pair only after each orbital in a subshell is occupied by a single electron C: 1s 2 2s 2 2p2 1s 2s 2p or: [He] 2s 2 2p 2 Electron Configurations and the Periodic Table 2 Valence Electrons Chemically similar behavior occurs among elements within a group in the periodic table. Valence electrons – electrons held in outer shell Core electrons – electrons held in filled inner shells Electron Configurations of Transition Elements • • d subshell being filled valence electrons - s and p electrons in outermost shell plus electrons in incompletely filled (n - 1)d subshell (n - 1)d orbitals are filled after ns orbitals and before filling np orbitals • Co: [Ar] 3d  Cr: 4s half-filled d subshells are favored if possible (Hund’s Rule) [Ar] 3d 4s Example 5 What family of elements is characterized by having an ns2np3 valence-electron configuration? Group 5A 3 Ion Electron Configurations Metal atoms lose electrons to form cations with a positive charge equal to the group number. Nonmetal atoms gain electrons to form anions with a negative charge equal to the A group number minus eight. Ca (Group 2)  Ca2+ O (Group 6A)  O2- Transition Metal Ions The ns electrons are at a higher energy than the (n-1)d electrons so they are always removed before the (n-1)d electrons when TMs form cations. Fe  Fe2+ + 2eFe0 [Ar] Fe2+ [Ar] 3d 4s 3d 4s Magnetic Properties of Compounds • Diamagnetic – atoms and ions with all electrons paired (will be repelled by magnetic field) • Paramagnetic – atoms and ions with unpaired electrons (will be attracted to magnetic field) 4 Periodic Trends: Atomic Radii • move down periodic table – increasing n • as n gets larger that shell’s radius increases • atomic size (radius) increases as you move down the periodic table • atomic size (radius) decreases as you move right in the periodic table Periodic Trends: Atomic Radii Periodic Trends: Atomic Radii 5 Example 6 Which will be closer to the nucleus, the n = 3 electron shell in Ar or the n = 3 electron shell in Kr? Kr because # protons is higher Bond Length Example 7 Predict which will be greater, the P―Br bond length in PBr3 or the As―Br bond length in AsBr3? As―Br atomic radius of As > P 6 Example 8 Arrange the following atoms in order of increasing atomic radius: K, Mg, Ca Mg < Ca < K Periodic Trends: Ionic Radii • radii of ions of elements in same group increase in size going down the group • radius of cation always smaller than that of neutral atom from which derived • radius of anion always larger than that of neutral atom from which derived Periodic Trends: Ionic Radii 7 Periodic Trends: Ionization Energies • ionization energy – energy needed to remove one electron from atom • s- and p-block elements – first ionization energies decrease down a group and across a period Periodic Trends: Ionization Energies Periodic Trends: Ionization Energies 8 Example 9 Predict which of the following atoms – C, Si, N or P – has the lowest first ionization energy. Si Periodic Trends: Electron Affinities • electron affinity (EA) – energy change when an electron is added to a gaseous atom to form 1- ion • measure of attraction an atom has for an additional electron • large negative value indicates a ready acceptance of an electron F(g) + e-  F-(g) Periodic Trends: Electron Affinities • electron affinity (EA) – energy change when an electron is added to a gaseous atom to form 1- ion • measure of attraction an atom has for an additional electron • large negative value indicates a ready acceptance of an electron F(g) + e-  F-(g) Ne + e-  Ne- ΔE = EA = -328 kJ/mol EA > 0 kJ/mol 9 Periodic Trends: Electron Affinities 10