1. The order of filling sublevels as seen on the periodic table. LanthanoidsActinoids

2. Filled and half-filled sublevels are more stable than partially filled sublevels. Filled and half-filled sublevels are more stable than partially filled sublevels.     _ _ 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. 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.

3. Electron configuration of ions Ions gain or lose electrons to become more stable. Ions gain or lose electrons to become more stable. The electron configuration for calcium is The electron configuration for calcium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 The configuration for an ion of calcium would be The configuration for an ion of calcium would be 1s 2 2s 2 2p 6 3s 2 3p 6 Ca 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 Ca 2+ This is isoelectronic to (has the same electron configuration as) an Argon atom. This is isoelectronic to (has the same electron configuration as) an Argon atom.

4. Electron configuration in excited state. N 1s 2 2s 2 2p 3 ground state 1s 2 2s 2 2p 2 3s 1 excited state 1s 2 2s 2 2p 2 3s 1 excited state or 1s 2 2s 2 2p 1 3p 2

5. Paramagnetic- an atom having one or more unpaired electrons Paramagnetic- an atom having one or more unpaired electrons Diamagnetic- all electrons are paired Diamagnetic- all electrons are paired Which elements in period 2 are diamagnetic? Be and Ne

6. Coulomb’s Law Related to Ionization energy- energy needed to remove an electron charges Distance apart

7. Reactivity The most active metals are in the lower left corner, and the most active nonmetals are in the upper right corner. The most active metals are in the lower left corner, and the most active nonmetals are in the upper right corner. Periodic Properties have a repeating pattern. Ex.: density atomic radii atomic radii ionization energy ionization energy electronegativity electronegativity

8. Atomic Radii Increases down a group, since they have increasing energy levels. 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). Decreases across a period due to increasing nuclear charge (the force of attraction between nucleus and electrons).

9. Atomic radii What would this say about density trends?

11. 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. O 2-, F -, Ne, Na + Mg 2+

13. Coulomb’s Law Related to Ionization energy- energy needed to remove an electron charges Distance apart Electric force or potential energy Potential energy is negative for an attractive force

14. Effective nuclear charge Z eff = Z – S Z eff = Z – S #Protons#core e-

15. Ionization energies Energy needed to remove an electron from an atom (kJ/mol) Energy needed to remove an electron from an atom (kJ/mol) Metals have low ionization energies. Metals have low ionization energies. Nonmetals have high ionization energies (especially noble gases). 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 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. 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. Second ionization energy- the energy required to remove a second electron from an atom.

17. Ionization energies of aluminum (kJ/mol) element1st2nd3rd 4 th Al577.51810275011,580 1 st e - is from the 3p sublevel. 2 nd e - is one of the 3s pair. 3 rd e - is the other 3s electron. 4 th 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.

18. The first five ionization energies of an element are shown below. II2I2 I3I3 I4I4 I5I5 78615773232435616,091 kJ/mol The element is likely to form ionic compounds in which its charge is... A)1+ B)2+ C)3+ D)4+ E)5+ Answer: The first four electrons require relatively little energy to remove, resulting in a 4+ ion.

19. Electron affinity (ΔH ea ) -the attraction of an atom for an electron (kJ/mol). -the energy change when an electron is added to a gaseous atom. Same trend as ionization energy (metals- low, nonmetals- high).

21. Electronegativity - the power of an atom in a molecule to attract electrons to itself. - is related to... Ionization energy- measures how strongly an atom holds onto its electrons Ionization energy- measures how strongly an atom holds onto its electrons Electron affinity- measures how strongly an atom attracts additional electrons Electron affinity- measures how strongly an atom attracts additional electrons F has a high ionization energy and a very negative electron affinity. F has a high ionization energy and a very negative electron affinity.