1. The Periodic Table GPS 7

2. History of the Periodic Table Mendeleev –credited for creating the first periodic table –arranged elements in order of increasing atomic mass Video

3. History of the Periodic Table Moseley –arranged elements in order of increasing atomic number –results in interesting periodic trends

4. Moseley proposed the modern periodic law. Modern periodic law: –the chemical properties of elements are periodic functions of their atomic numbers History of the Periodic Table

5. Groups and Periods Group or Family –elements within the same column on the periodic table Period –elements within the same row on the periodic table

6. Groups on the Periodic Table

7. Periods on the Periodic Table Lanthanides Actinides

8. Metals, Nonmetals, Metalloids

9. Metalloids Metalloids are commonly used in semiconductors Semiconductors are used in solar cells, computer chips, and LEDs (among other applications)

10. Physical Properties of Metals and Nonmetals MetalsNonmetals Aluminum Gold Copper wire Sulfur powder Hydrogen gas bubbling through a solution Liquid Bromine

11. Physical Properties of Metals and Nonmetals PropertyMetalsNonmetals DensityHigher densitiesLower densities ConductivityVery highVery low Melting pointHigher temps.Lower temps. Boiling pointHigher temps.Lower temps. MalleabilityVery goodVery poor DuctilityVery goodVery poor

12. Chemical Properties of Groups Octet Rule –an atom that has a full outer-most energy level is unreactive (usually it is full with 8 electrons, but there are some exceptions) all atoms want to satisfy the octet rule atoms react with each other in order to satisfy the octet rule

13. Octet Rule: Forming Bonds Atoms will gain, lose, or share electrons in order to fulfill the octet rule When electrons are transferred, ions are formed. Ions are attracted to each other and form ionic bonds When electrons are shared, covalent bonds are formed

14. Chemical Properties of Groups Oxidation Numbers

15. Chemical Properties of Groups Example: For the following elements, show the ion that it most commonly forms (using proper notation): Chlorine Sodium Nitrogen Sulfur Neon Boron Magnesium Lead Cl - Na + N 3- S 2- No ion forms B 3+ Mg 2+ Pb 4+, Pb 4-

16. Oxidation Numbers and Ions Ions –atoms with charges (i.e. not neutral) Oxidation numbers –indicate the ion that an atom most always forms during chemical reactions Anion –a negatively charged ionEx: O 2- Cation –a positively charged ionEx: Na 1+

17. Periodic Trends: Ionic Radius Ionic Radius CationsAnions

18. Explaining the Trends: Ionic Radius Ionic radii of cations are smaller than their atomic radii because electrons were lost. –Smaller if all electrons from outer-most energy level were lost –If outer-most energy level remains, then the ion is smaller because there is now less electron-electron repulsion Ionic radii of anions are larger than their atomic radii because electrons were gained. –Larger if electrons added create a new, higher energy level –If a new outer-most energy level is not created, then the ion is larger because there is now more electron-electron repulsion

19. Ionic Radius vs Atomic Radius Is the ionic radius of the calcium ion bigger or smaller than the atomic radius of the calcium atom? Justify your answer.

20. Ionic Radius vs Atomic Radius Is the ionic radius of the nitrogen ion bigger or smaller than the atomic radius of the nitrogen atom? Justify your answer.

21. Periodic Trends: Which of the particle-view diagrams below represents Na and which represents Na + ? Justify your answer.

22. Periodic Trends: Which of the particle-view diagrams below represents Cl and which represents Cl - ? Justify your answer.

23. Periodic Trends: Atomic Radius Atomic Radius ½ the distance between the nuclei of two atoms when the atoms are bonded

24. Explaining the Trends: Atomic Radius Explaining a trend across a period ↔: –The effective nuclear charge (Z eff ) increases as the number of protons increases (from left to right across the table →) –Increased Z eff results in a smaller radius Explaining a trend up or down a group ↕: –The number of energy levels in a certain atom increases as you go down a group, making the atomic radius larger down a group

25. Effective Nuclear Charge (Zeff) The effective nuclear charge (Z eff ) is the net force (pull) that an electron experiences from the protons in the nucleus.

26. Periodic Trends: Ionization Energy energy + Li → Li + + e - Ionization Energy –the energy required to remove an electron from an atom (thus, creating a cation) –also known as First Ionization Energy when only removing one electron from the outer energy level Example:

27. Periodic Trends: Ionization Energy Trend down a Group: Shielding effect –inner electrons block the attraction of the nucleus for outer electrons, thus lowering first ionization energy –The greater the shielding effect, the lower the ionization energy (the opposite is also true). Ionization energy decreases as you go down a group

28. Explaining the Trends: Ionization Energy Explaining a trend across a period ↔: –The effective nuclear charge (Z eff ) increases as the number of protons increases (from left to right across the table →) –Increased Z eff results in higher ionization energy required to remove a valence electron from an atom Explaining a trend up or down a group ↕: –The number of energy levels in a certain atom increases as you go down a group –The further the electron from the nucleus, the more it is shielded from the positive pull of the nucleus, the less ionization energy required to remove a valence electron from an atom

29. Periodic Trends: Electronegativity Electronegativity –the tendency of an atom to attract electrons to itself when it is bonded to another atom –Imagine a “tug-of-war” over valence electrons involved in making the bond Example: A and B have the same electronegativity and are sharing valence electrons equally B is slightly more electronegative than A, so the electrons are more highly attracted to B B is much more electronegative than A, so the electrons are so highly attracted to B that they leave A and transfer to B

30. Explaining the Trends: Electronegativity Explaining a trend across a period ↔: –The effective nuclear charge (Z eff ) increases as the number of protons increases (from left to right across the table →) –Increased Z eff results in an atom having a greater electronegativity Explaining a trend up or down a group ↕: –Increased number of energy levels decreased an atom’s electronegativity due to the increased distance between bonding electrons and the nucleus

31. Periodic Trends Example: Put the following elements in order of: a. increasing ionization energy b. increasing electronegativity c. increasing atomic radius zinc, rubidium, fluorine, calcium, silver, oxygen, potassium, sulfur

32. Periodic Trends: Example: a.Explain why the first ionization energy for fluorine is greater than the first ionization energy for lithium. b. Explain why the atomic radius of lithium is smaller than the atomic radius of cesium.

33. Periodic Trends: Use principles of atomic structure to explain the following: a.Fluorine is more electronegative than Iodine. b. The atomic radius of lithium is 0.546 nanometers whereas the atomic radius of carbon is 0.442 nanometers.

34. Periodic Trends: Which of the particle-view diagrams below represents Mg, Ca, and Mg 2+ ? Justify your answer.

35. Periodic Trends: Reactivity of Nonmetals Reactivity –The tendency to change into something else when mixed with another substance

36. Periodic Trends: Reactivity of Metals

37. Periodic Trends: Reactivity Which two groups are the most reactive? Why? Which is the most reactive nonmetal? Why? Which is the most reactive metal? Why?

38. Periodic Trends: Diatomic Molecules Naturally occurring diatomic molecules: Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 - These occur as gases at room temperature * except Br 2 (liquid at room temp) * except I 2 (solid at room temp) - Naming: Ex: O 2 oxygen gas N 2 nitrogen gas Br 2 liquid bromine I 2 solid iodine

39. Resources http://images.google.com/url?q=http://spiff.rit.ed u/classes/phys314/lectures/period/period.htmlhttp://images.google.com/url?q=http://spiff.rit.ed u/classes/phys314/lectures/period/period.html http://www.chem.wisc.edu/areas/organic/content -chem.htmhttp://www.chem.wisc.edu/areas/organic/content -chem.htm http://serc.carleton.edu/usingdata/nasaimages/in dex4.htmlhttp://serc.carleton.edu/usingdata/nasaimages/in dex4.html http://www.crystalmaker.com/support/tutorials/cr ystalmaker/resources/VFI_Atomic_Radii.jpghttp://www.crystalmaker.com/support/tutorials/cr ystalmaker/resources/VFI_Atomic_Radii.jpg http://intro.chem.okstate.edu/1314F00/Lecture/C hapter7/CommonCations.htmlhttp://intro.chem.okstate.edu/1314F00/Lecture/C hapter7/CommonCations.html

40. SOL covered during lesson CH 2 d, e, f, g, h