1. 1 Electron Shells  Move down P. table: Principal quantum number (n) increases.  Distribution of electrons in an atom is represented with a radial electron density graph. Radial electron density is probability of finding an electron at a particular distance from the nucleus. Electron shells are diffuse and overlap a great deal.

2. 2 Examples of Electron shells  He: 1s 2  Radial plot shows 1 maximum  Ne: 1s 2 2s 2 2p 6  Radial plot shows 2 maxima ( 1 each for the 1 st and 2 nd energy levels )  Ar: 1s 2 2s 2 2p 6 3s 2 3p 6  Radial plot shows 3 maxima ( 1 each for the 1 st,2 nd and 3 rd energy levels )

3. Effective Nuclear Charge 3  Charge actually “experienced” by the valence electrons  Two factors: 1. distance from the nucleus 2. shielding by the core electrons  Attraction nucleus declines as a function of 1/distance 2.  If an electron is shielded, it is less attracted to the nucleus and is more easily removed.

4. Effective Nuclear Charge (Z eff ) Z eff = Z - S  Z = the atomic number of the element  S is equal to the number of electrons that lie in shells with n values smaller than the n for the electron you're interested in. Ex: Calculate the shielding for a 3p electron: All electrons with n values of 1 or 2 contribute to the shielding: 2 e- (n=1) + 8 e- (n=2) = 10 total electrons so S = 10 4

5. Z eff = Z - S  Consider only the electrons in lower shells to contribute to shielding. The small shielding contribution by other electrons in the same shell is ignored.  s and p and d electrons each contribute differently to the shielding of an outer electron, but this simple method does not take that into account. 5 Effective Nuclear Charge (Z eff )

6. 6 ProtonsSpeciesElectron Configuration Protons - Core e- Zeff 12Mg1s 2 2s 2 2p 6 3s 2 12 - 10+2 12Mg 2+ 1s 2 2s 2 2p 6 12 - 2+10 7N1s 2 2s 2 2p 3 7 - 2+5 7N 3- 1s 2 2s 2 2p 6 7-2+5 13Al1s 2 2s 2 2p 6 3s 2 3p 1 13 - 10+3 13Al 3+ 1s 2 2s 2 2p 6 13 - 2+11

7. 7 ProtonsSpeciesElectron Configuration Protons - Core e- Zeff 4Be1s 2 2s 2 4-2+2 5B1s 2 2s 2 3s 1 5-2+3 7N1s 2 2s 2 2p 3 7 - 2+5 7N 2s↑↓ 2p ↑ ↑ ↑. 8O1s 2 2s 2 2p 4 8-2+6 8O 2s↑↓ 2p ↑↓ ↑ ↑.

8. Z eff Example: Iron (Fe)  Each of the 3p electrons: Z eff = 26 - 10 = 16  Each of the 3d electrons: Z eff = 26 - 10 = 16 (so even though these electrons go in later, they are still considered the n=3 shell, so they have the same shielding as the 3s and 3p e-.)  Each of the 4s electrons: Z eff = 26 - 24 = 2 (Add the 3d electrons into the shielding S, since they are in the n=3 shell and will therefore shield the 4s electrons.) 8

9. 9 Atomic Sizes  As we move down a group, atoms become larger. Larger n = more shells = larger radius  As we move across a period, atoms become smaller. More protons = more effective nuclear charge, Zeff More positive charge increases the attraction of nucleus to the electrons in the outermost shell, so the electrons are pulled in more “tightly,” resulting in smaller radius

10. Atomic Radius 10

11. 11 Ion size Isoelectronic = Having exactly the same number and configuration of electrons)  The oxide ion is isoelectronic with neon, and yet O 2– is bigger than Ne. O 2– : 8 protons (+)Ne: 10 protons (+) 10 electrons(-)  This is Coulomb's law at work. In any isoelectronic series the species with the highest nuclear charge will have the smallest radius.

12. 12 Ionization energy  Ionization energy of an ion or atom is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.  The first ionization energy, I 1 is the energy required to remove one electron from an atom. Na(g)  Na + (g) + e -  The 2 nd ionization energy, I 2, is the energy required to remove an electron from an ion. Na + (g)  Na 2+ (g) + e -  Larger ionization energy, harder to remove electron.

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14. 14 Periodic Trends in Ionization Energy  Highest = Fluorine  Ionization energy decreases down a group. Easier to remove electrons that are farther from the nucleus.  Ionization energy increases across a period. Zeff increases, so it’s harder to remove an electron. Exceptions: Removing the 1 st and 4 th p electrons

15. 15 Electron Affinity  Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion.  Electron affinity: Cl(g) + e -  Cl - (g)  Ionization energy: Cl(g)  Cl + (g) + e -  Affinity for reaction above is exothermic: ∆ E = -349 kJ/mol  If adding the electron makes the species more stable, it will be exothermic. Gain Lose

16. 16 Metals  Metallic character increases down a group and from left to right across a period.  Metal properties: Lustrous (shiny) Malleable (can be shaped) Ductile (can be pulled into wire) Conduct electricity  Metal oxides form basic ionic solids: Metal oxide + water  metal hydroxide  Metal oxides react with acids to form salt and water

17. Metallic Character  A metal tends to lose its outer electron(s)  The more likely it is to lose the electrons, the more metallic it is  Low ionization energy is associated with high metallic character 17

18. 18 Metals  Metal oxides form basic ionic solids: Metal oxide + water  metal hydroxide MgO(s) + H 2 O(l)  Mg(OH) 2 (s)  Metal oxides react with acids to form salt and water MgO(s) + 2HCl(aq)  MgCl 2 (aq) + H 2 O(l)  Most neutral metals are oxidized rather than reduced.  Metals have low ionization energies.

19. 19 Metal reactivity  Which of the alkali metals would you expect to react most violently with water? Li, Na, K, Rb  Of these four, rubidium has the lowest ionization energy, making it the most reactive. Rubidium reacts explosively with water.

20. 20 Nonmetals  Lower melting points than metals  Diatomic molecules are nonmetals.  Most nonmetal oxides are acidic: Nonmetal oxide + water  acid P 4 O 10 (s) + 6H 2 O(l)  4H 3 PO 4 (aq)  Nonmetal oxides react with bases to form salt and water: CO 2 (g) + 2NaOH(aq)  Na 2 CO 3 (aq) + H 2 O(l)

21. 21 Nonmetallic oxides  Which nonmetallic oxide would you expect to be the strongest acid? NO 2, N 2 O, N 2 O 4, N 2 O 5 N 2 O 5 : Nitrogen has an oxidation state of +5 in this compound. In general, the higher the oxidation state of the nonmetal, the more acidic the nonmetal oxide.

22. 22 General Trend Summary Electronegativity, Ionization Energy, Electron Affinity Electronegativity, Ionization Energy, Electron Affinity F Atomic Radius, Metallic Character Atomic Radius, Metallic Character Fr

23. Mullis23

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25. A. 1s 2 2s 2 2p 6 3s 2 C. 1s 2 2s 2 2p 6 3s 2 3p 4 B. 1s 2 2s 2 2p 6 3s 2 3p 2 D. 1s 2 2s 2 2p 6 3s 2 3p 6 E. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Use the choices above to answer each item below. 1. Represents an element in the oxygen family. 2. Represents an electron configuration for a chloride ion. 3. Represents an electron configuration of a common ion of an alkali metal. 4. Represents an atom with four valence electrons. 25