2. Periodic Trends Elemental Properties and Patterns

3. The Periodic Law Dmitri Mendeleev - first to publish an organized periodic table of known elements.

4. The Periodic Law Mendeleev even predicted the properties of undiscovered elements. Very accurate in his predictions, led the world to accept his ideas about periodicity and a logical periodic table. https://www.youtube.com/watch?v=fPnwBIT SmgU

5. The Periodic Law Mendeleev understood the ‘Periodic Law’ which states: “When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.”

6. The Periodic Law Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

7. Valence Electrons For Groups 1, 2, 13, 14, 15, 16, 17, 18: The digit farthest to the right is the number of valence electrons. Example: Groups 3-12? Many have 1-2, but figuring it out is more complicated.

8. A Different Type of Grouping Besides the 4 blocks of the table, there is another way of classifying element: Metals Nonmetals Metalloids or Semi-metals

9. Metals, Nonmetals, Metalloids

10. Most elements that border the stair case are metalloids. Have properties of both metals and nonmetals!

11. Metals Lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. Mostly hard – not brittle They are mostly solids at room temp.

12. Nonmetals The opposite! They are dull, brittle, nonconductors (insulators). Some are solid, but many are gases, and a few are liquid.

13. Metalloids Have properties both metals & nonmetals! Shiny but brittle. Semiconductors.

14. Periodic Trends There are several important atomic characteristics that show predictable trends that you should know. The first and most important is atomic radius.

15. Atomic Radius Is the distance from the center of the nucleus to the “edge” of the electron cloud. Since that is difficult to define, scientists use covalent radius, half the distance between the nuclei of 2 bonded atoms. Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10 -10 m.

16. Covalent Radius Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å. 2.86 Å 1.43 Å

17. Atomic Radius The trend for atomic radius in a GROUP is to go from smaller at the top to larger at the bottom. Why? With each step down the family, we add energy levels to the electron cloud, making the atoms larger.

18. Atomic Radius As you move ACROSS a PERIOD, atoms are smaller because of an increased attraction between nucleus and electron cloud. The increased attraction pulls the cloud in, making atoms smaller as we move from left to right.

19. Effective Nuclear Charge What keeps electrons from simply flying off into space? Effective nuclear charge is the pull that an electron “feels” from the nucleus. The closer an electron is to the nucleus, the more pull it feels. As effective nuclear charge increases, the electron cloud is pulled in tighter.

20. Ionization If an electron is given enough energy to overcome its attraction to the nucleus, it can leave the atom completely. The atom has been “ionized” or charged. Number of protons ≠ number of electrons.

21. Ionization Energy Energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ) The larger the atom is, the easier its electrons are to remove. Ionization energy and atomic radius are inversely proportional. As you move down, it decreases! As you move across, it increases!

22. Electronegativity Electronegativity is a measure of an atom’s attraction for another atom’s electrons. Metals are usually electron givers and have low electronegativity. Nonmetals are electron takers and have high electronegativity. What about the noble gases? Make a prediction to your table partners.

23. The Octet Rule The “goal” of most atoms (except H, Li and Be) is to have an octet (group of 8 electrons) in their valence energy level. Metals generally give electrons, nonmetals take them from other atoms. Remember… Atoms that have gained or lost electrons are called ions.

24. Ions When an atom gains an electron, it becomes negatively charged and is called an (anion.) When an atom loses an electron, it becomes positively charged (cation). Think of the “t” like a “+” sign.

25. Ionic Radius Cations are always smaller than the original atom. Conversely, anions are always larger than the original atom.

26. Cation Formation 11p+ Na atom 1 valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na +

27. Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.

28. Review Video for Atomic Trends http://www.youtube.com/watch?v=G3qboo Mh6Fc&list=TL6TJ_jx1X7SKnBP20kAt90 eSdTpgQOR8H http://www.youtube.com/watch?v=G3qboo Mh6Fc&list=TL6TJ_jx1X7SKnBP20kAt90 eSdTpgQOR8H http://www.youtube.com/watch?v=ViIDHZ b9iE8&list=TL6TJ_jx1X7SKnBP20kAt90e SdTpgQOR8H http://www.youtube.com/watch?v=ViIDHZ b9iE8&list=TL6TJ_jx1X7SKnBP20kAt90e SdTpgQOR8H

29. Link to bonding animations http://bcs.whfreeman.com/thelifewire/conte nt/chp02/02020.html http://bcs.whfreeman.com/thelifewire/conte nt/chp02/02020.html http://www.youtube.com/watch?v=cZy8tGF V8QE&list=TL6TJ_jx1X7SKnBP20kAt90e SdTpgQOR8H http://www.youtube.com/watch?v=cZy8tGF V8QE&list=TL6TJ_jx1X7SKnBP20kAt90e SdTpgQOR8H