1. Mullis1 The Periodic Table  Elements are arranged in a way that shows a repeating, or periodic, pattern.  Dmitri Mendeleev created the first periodic table of the elements in 1869.  He ordered the ~70 known elements by their atomic masses and their chemical properties.  He found that some elements could not be put into groups with similar properties and at the same time stay in order.

2. Mullis2 Modern Periodic Table  Later, Henry Moseley carried on the work. Moseley put the elements in order of increasing atomic NUMBER. He found that the position of the element corresponded to its properties.  The modern periodic table shows the position of the element is related to : Atomic number AND Arrangement of electrons in its energy levels

3. Mullis3 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.

4. Mullis4 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 )

5. Mullis5 Atomic Sizes: Single atoms  Colliding argon atoms ricochet apart because electron clouds cannot penetrate one another to a significant extent.  The apparent radii are determined by the closest distances separating the nuclei during such collisions.  This radius is called the nonbonding radius.

6. Mullis6 Atomic Sizes: Bonded atoms  The distance between two nuclei is called the bond distance.  If the two atoms making up the molecule are the same, then ½ the bond distance is called the bonding atomic radius of the atom.  This radius is shorter than the nonbonding radius.

7. Mullis7 Atomic Sizes using Periodic Table  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

8. Mullis8 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.

9. Mullis9 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

10. Mullis10 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

11. Mullis11 Coulomb’s law  Which law can best be used to explain why addition of an electron to the O 2– ion is an endothermic process?  Coulomb’s law: The energy required for the process is necessary to overcome the electrostatic repulsion between the electron and the already negatively charged O 2– ion.

12. Mullis12 Ion size  The oxide ion is isoelectronic (has exactly the same number and configuration of electrons) with neon, and yet O 2– is bigger than Ne. Why?  This is Coulomb's law at work. In any isoelectronic series the species with the highest nuclear charge will have the smallest radius.

13. Mullis13 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

14. Mullis14 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.

15. Mullis15 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.

16. Mullis16 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)

17. Mullis17 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.