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The first Periodic Table? The order of filling sublevels as seen on the periodic table. Lanthanoids Actinoids Filled and half-filled sublevels are more stable than partially filled sublevels. E E EEE E EEE h h h h _ 1s 2s 2p 3s 3p 3d _ 4s Thus Cr takes an electron from 4s to put one electron in each of its 3d orbitals and Cu takes a 4s electron to fill each of its 3d orbitals. Electron configuration of ions Ions gain or lose electrons to become more stable. The electron configuration for calcium is 1s2 2s2 2p6 3s2 3p6 4s2 The configuration for an ion of calcium would be 1s2 2s2 2p6 3s2 3p6 Ca2+ This is isoelectronic to (has the same electron configuration as) an Argon atom. Electron configuration in excited state. N 1s2 2s2 2p3 ground state 1s2 2s2 2p2 3s1 or 1s2 2s2 2p1 3p2 excited state Reactivity The most active metals are in the lower left corner, and the most active nonmetals are in the upper right corner. Periodic Properties Ex.: density atomic radii ionization energy electronegativity have a repeating pattern. Atomic Radii Increases down a group, since they have increasing energy levels. Decreases across a period due to increasing nuclear charge (the force of attraction between nucleus and electrons). Atomic radii What would this say about density trends? Ionic radii Metals lose electrons when they ionize, so their ionic radii are smaller than their atomic radii. Nonmetals gain electrons, so are larger than their atomic radii. “Isoelectronic”- ions with the same number of electrons. Ex. O2-, F- , Ne, Na+ Mg2+ Ionization energies Energy needed to remove an electron from an atom (kJ/mol) Metals have low ionization energies. Nonmetals have high ionization energies (especially noble gases). Going down a group, the ionization energy decreases due to increased atomic radius and the shielding effect. Going across a period it increases due to increasing nuclear charge. Second ionization energy- the energy required to remove a second electron from an atom. Ionization energies of aluminum (kJ/mol) element 1st 2nd 3rd 4th Al 577.5 1810 2750 11,580 1st e- is from the 3p sublevel. 2nd e- is one of the 3s pair. 3rd e- is the other 3s electron. 4th e- would be from a full 2p sublevel. Looking at ionization energies can help us predict oxidation numbers. Al usually has a 3+ oxidation number. Electron affinity The attraction of an atom for an electron (kJ/mol). Same trend as ionization energy (metals- low, nonmetals- high). Electronegativity The power of an atom in a molecule to attract electrons to itself. Electronegativity is related to. . . Ionization energy- measures how strongly an atom holds onto its electrons Electron affinity- measures how strongly an atom attracts additional electrons F has a high ionization energy and a very negative electron affinity.