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  7 periods (rows)  In general, each period represents an energy level, or an electron shell  Move down P. table: Principal quantum number (n) increases. Recall that in describing the highest energy electron of Cl, we found that electron in this orbital: 3p 5 n = 3 here. (Configuration is 1s 2 2s 2 2p 6 3s 2 3p 5.) Atomic radius increases if same group Metallic character increases if same group.

4. Mullis4 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

5. Mullis5 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.

6. Mullis6 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

7. Mullis7 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 - Gain Lose

8. Mullis8 Isoelectronic species  Definition: Has exactly the same number and configuration of electrons  Examples: Ne, Na +, O 2- Ne = 10 e - Na = 11 e - / Na + = 11-1 = 10 e - O = 8 e - / O 2- = 8 + 2 = 10 e -  Even though these have the same number of electrons, they retain all their protons. Therefore the sodium ion (Na + ) has the highest effective nuclear charge with 11 protons attracting its 10 electrons.

9. Mullis9 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?  In any isoelectronic series the species with the highest nuclear charge will have the smallest radius.

10. Mullis10 Metals  Metallic character increases down a group and from left to right across a period.  Metals are found to the left of the zig-zag line on the periodic table.  Metal properties: Lustrous (shiny) Malleable (can be shaped) Ductile (can be pulled into wire) Conduct electricity  Metals form cations (positive ions) This means they lose 1-4 electrons Therefore, they are usually found in IONIC compounds

11. Mullis11 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.

12. Mullis12 Nonmetals  Lower melting points than metals  Diatomic molecules are nonmetals.  The seven (7) diatomic molecules are: Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2  Two or more nonmetals form molecular compounds with COVALENT bonds.

13. Mullis13 Trends See your book for full explanation. Closer to F = more  ELECTRONEGATIVITY  ELECTRON AFFINITY  IONIZATION ENERGY Closer to Cs = more  METALLIC CHARACTER  ATOMIC RADIUS  REACTIVITY