1. Periodic Trends

2. Group Trends Group one, Alkali metals Group two, Alkaline earth metals Groups 3-12, Transition metals Group 17, Halogens Group 18, Noble gases At the bottom: Lanthanides and Actinides, also known as Rare earth metals

3. The Octet Rule All atoms want their outer (valence) shell to have eight electrons. The first shell (closest to nucleus) always holds two electrons. As you move away from the nucleus, the second shell will hold eight, the third shell will hold eight, the next two shells will hold 18. As long as the outer shell has eight, the atom is stable.

4. Orbitals or Shells Bohr's model of the atom showed electrons orbiting the nucleus. The quantum mechanical model instead shows energy levels called “orbitals” or “shells.” Each energy level corresponds to a period on the periodic table. Level 1, the ground state, is the same as period 1 on the table and so on...

5. Organizing Orbitals Orbitals are organized on the periodic table according to the shape, energy and number of electrons. This organization is called an “electron configuration.” There are four sections in the configuration; s, p, f and d. Compare your periodic table to the next slide.

7. Ions Ions are formed when an element either gives up electrons from their outer (valence) shell or gains electrons in their valence shell so their outer shell has eight. This transfer of electrons creates an element that is no longer neutral (same number of protons and electrons) Cations are ions with a positive charge Anions are ions with a negative charge

8. Cations (The “t” in “cation” looks like a plus sign) All metals form positive cations when they donate electrons from their valence shell. Alkali metals, for example, have one extra electron in their valence shell and would like to have a full outer shell. It is easier to give away one than gain seven! This leaves them with one more proton than electron, so they have a positive charge of +1.

9. Anions (“A-negative-ion”) Anions are formed by non-metals when they take electrons from metals to fill their outer shell with a total of eight. More electrons than protons means they are no longer neutral. They are now more negative. A halogen, for example, that has one more electron is more negative with a -1 charge.

10. Your Turn! What charge will a potassium ion have? – K +1 What charge will a fluorine ion have? – F -1 What charge will a magnesium ion have? – Mg +2 What about noble gases? – Already stable!

11. Atomic Radii What does the atomic number for a neutral atom represent? What is the atomic radius? What is the trend toward size of the atomic radii as you move across the periodic table?

12. Atomic Radii Decrease as You Move Across a Period The atomic number increases as you move left- to-right across the periodic table. The atomic radius decreases as you move left-to-right, because the greater positive charge in the nucleus pulls the electrons closer.. Draw arrows on your periodic table and label the trend.

13. Atomic Radii Increase as you Move Down a Group The energy levels increase as you move down the periodic table (from period 1 to 7) Each energy shell of electrons enlarges the overall radius.

15. Ionization Energy The IE is the minimum amount of energy required to remove an electron from the outer shell of a neutral atom. When several e- are removed (example, group 2), the energy required to remove the first e- is called the “first ionization energy”. As more e- are removed, the energy is called the second IE, and so on.

16. Ionization Energy What is this ion? 4 protons, 3 electrons…

17. Ionization Energy Trend IE decreases as you move down the periodic table because the electrons are farther away from the nucleus and therefore are easier to remove. IE increases as you move across from left to right because the more electrons there are in the valence shell, the more energy it takes to remove one. Draw arrows on your periodic table and label this trend.

19. Electron Affinity The measure of the change in energy when you add an electron to an atom. EA increases as you move from left to right on the periodic table because you produce more stable anions as you move to the right. The elements to the right have a larger positive charge in the nucleus, so attract the electrons more easily.

20. Electron Affinity Trend Stops at Noble Gases Noble gas valence shells are full. Noble gases are very stable. They do not want additional electrons. Their EA values are therefore positive.

22. Electronegativity Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. This differs from electron affinity in that EA relates to the attraction of electrons by an unbonded atom becoming an anion.

23. Electronegativity Trends Electronegativity generally increases as you move left to right across the periodic table. Electronegativity generally increase as you move up a group.

24. Electronegativity Trends

25. Fluorine, F Fluorine, in group 17 and period 2, is the element with the highest electronegativity.

26. Noble Gases, Lanthanides, Actinides Noble gases have a full valence shell, so do not normally attract e-. Lanthanides and actinides posses a more complex chemistry and therefore do not follow electronegativity trends.