1. Assignment 10/24 Read Chapter 6, section 1. You have 10 minutes. Answer the questions below: – 1. How is the modern periodic table arranged? Who organized elements with this arrangement? – 2. Who arranged the periodic table by increasing atomic mass? – 3. What is the name of Group 18 elements?

2. 11/8 Activity 1. Write the electron configuration of your element. 2. Name the group (if applicable) that your element falls in. List all properties of that group. Be detailed! 3. How many valence electrons does your element have in the outer shell. 4. Draw the Lewis dot structure for your element. When finished…transfer all the information about the elements and group to a piece of butcher paper. Show me your work and hang it on the wall somewhere. We will do a gallery walk of the groups at the end of the period today or beginning tomorrow.

3. 11/11 Today you will need… A clean sheet of paper, something to write with, and the paper from the side table. We will do a gallery walk activity using the information from Friday about the groups on the periodic table. You will color code a periodic table when the gallery walk is complete

4. Gallery Walk… When finished, answer the following questions on a clean piece of paper and turn in: – 1. Name Group 1, Group 2, Group 17, and Group 18 elements. – 2. Identify the properties of each of the groups listed in number 1. – 3. Determine the number of valence electrons in groups 1,2,13,14,15,16,17,18. – 4. Draw the electron configuration for ONE element from each group listed in number 1.

5. Chapter 5- The Periodic Law 5.1-History of the Periodic Table 5.2-Electron Configuration & the Periodic Table 5.3-Electron Configuration & Periodic Properties

6. 5.1-History of the Periodic Table Pages 123-127

7. Mendeleev Dmitri Mendeleev (1869, Russian) – Organized elements by increasing atomic mass. – Elements with similar properties were grouped together. – There were some discrepancies.

8. Mendeleev Dmitri Mendeleev (1869, Russian) – Predicted properties of undiscovered elements.

9. Moseley Henry Moseley (1913, British) – Organized elements by increasing atomic number. – Resolved discrepancies in Mendeleev’s arrangement. – Periodic Law-the physical and chemical properties of the elements are periodic functions of the atomic numbers.

10. Organization of the Elements Periodic table-an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

11. Additions to Mendeleev’s Periodic Table Noble gases – Group 18 – Argon discovered in 1894 – Took so long to discover because very unreactive Lanthanides – 14 elements with atomic numbers from 58-71 – Placed below the periodic table to conserve space Actinides – 14 elements with atomic numbers 90-103 – Also placed below periodic table

12. Warm-Up 11/11 Define Periodic Law. Where are the Lanthanides and Actinides located on the periodic table. Describe, specifically what contributions Mendeleev and Moseley made towards the creation of the periodic table. (i.e dates, arrangement, etc.)

13. 5.2-Electron Configuration & the Periodic Table Pages 128-139

14. Periods & Blocks of the Periodic Table Length of period (row) determined by how many electrons can occupy the sublevels being filled. – 1 st period-1s sublevel being filled with 2 electrons  2 elements, H & He – 3 rd period-3s & 3 p sublevels being filled with 2+6 electrons  8 elements Periodic table is divided into “blocks” based on the filling of sublevels with electrons.

15. Blocks of the Periodic Table

16. Determining Period from Configuration An element’s period can be determined by looking at its electron configuration The highest occupied energy level corresponds to the element’s period – As: [Ar]3d 10 4s 2 4p 3 4 in 4p 3 indicates that the highest energy level that electrons occupy is the 4 th. Therefore, As is located in the 4 th period of the periodic table.

17. Metallic Character Metals Nonmetals Metalloids

18. Areas of the Periodic Table Main Group Elements Transition Metals Inner Transition Metals

19. s-Block Elements: Groups 1 & 2 Chemically reactive metals Include the alkali metals and the alkaline earth metals

20. Alkali metals Group 1 metals ns 1 Silvery appearance and very soft Not found pure naturally because so reactive Because of extreme reactivity with moisture, usually stored under kerosene Video: Disposal of Surplus SodiumDisposal of Surplus Sodium Video: Alkali Metals in WaterAlkali Metals in Water

21. Alkaline-Earth metals Group 2 metals ns 2 Harder, denser, & stronger than alkali metals Also too reactive to be found free in nature (but less reactive than Gp. 1) Video: Magnesium/silver nitrate mixture reacting with waterMagnesium/silver nitrate mixture reacting with water

22. d-Block Elements: Groups 3-12 Metals with typical metallic properties Called “transition elements” Typically less reactive than Gps. 1&2, & some are extremely unreactive d sublevels first appears at the 3 rd energy level Fills after 4s Variations from expected in d-block, so elements in the same group do not necessarily have the same outer e- configuration

23. p-Block Elements: Groups 13-18 p and s-block elements together called “main- group elements” Total number of electrons in highest energy level=group # - 10 – Group 17 elements have 17-10=7 outer “valence” electrons Properties of p-block elements vary greatly since metals, nonmetals, and metalloids are contained here

24. p-block Elements Halogens – Group 17 nonmetals – Most reactive nonmetals React with most metals to form salts Metalloids – Fall on both sides of a “stair-step” line separating metals and nonmetals – Semi-conductors

25. f-Block Elements: Lanthanides & Actinides Lanthanides – Top row of f-block – 14 elements – Shiny metals similar in reactivity to the alkaline-earth metals Actinides – Bottom row of f-block – 14 elements – All radioactive – 1 st 4 elements found naturally on Earth; remainder only lab-made elements

26. 5.3-Electron Configuration & Periodic Properties Pages 140-154

27. Remember the Periodic Law When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

28. Warm Up 11-12 Explain the joke: Yo momma so fat if she was on Mendeleev’s periodic table she would be at the end.

29.  ½ the distance between the nuclei of identical atoms that are bonded together Atomic Radius

30. Li Ar Ne K Na

31. Ionization Energy First Ionization Energy- the energy required to remove one electron from a neutral atom

32. Ionization Energy First Ionization Energy K Na Li Ar Ne He

33. Ionization Energy Why opposite of atomic radius? – In small atoms, e - are close to the nucleus where the attraction is stronger Why small jumps within each group? – Stable e - configurations don’t want to lose e -

34. Ionization Energy Successive Ionization Energies Mg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ Large jump in I.E. occurs when a CORE e - is removed.

35. Al1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ Ionization Energy Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed.

36. Electron Affinity Energy change that occurs when an electron is acquired by a neutral atom Tends to become less negative (less energy released) DOWN and to the LEFT

37. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons Most electronegative element is fluorine – Given arbitrary value of 4; all others relative

38. A C B D E F G Label the Groups:

39. Atomic Radius Q: Why larger going down? A: Higher energy levels have larger orbitals Q: Why smaller to the right? A: Increased nuclear charge without additional energy levels to shield the electrons pulls them in tighter.

40. Q: Why is Ionization Energy greatest in upper right corner? A: In small atoms, e - are close to the nucleus where the attraction is stronger. More protons = more positive charge. Ionization Energy

41. Ionic Radius: The distance from the nucleus to the outer orbital of an ion.

42. Ions Atoms with an electric charge Cations Anions + charge - charge Lose electrons Gain Electrons Smaller Bigger

43. Ionic Radius

44. Q: Why are cations smaller than the neutral atom? A: When cations lose their electrons they also lose an energy level. Q: Why are anions larger than the neutral atom? A: As they gain electrons there is more repulsion between them, which pushes them farther from the nucleus.

45. Examples Which atom has the larger radius? BeorBa CaorBr

46. Examples Which atom has the larger radius? BeorBa CaorBr Ba Ca

47. Examples Which atom has the higher 1st I.E.? NorBi BaorNe

48. Examples Which atom has the higher 1st I.E.? NorBi BaorNe N Ne

49. Examples Which has the greater electonegativity? KorLi AlorCl

50. Examples Which has the greater electonegativity? KorLi AlorCl Li Cl

51. Examples Which particle has the larger radius? SorS 2- AlorAl 3+

52. Examples Which particle has the larger radius? SorS 2- AlorAl 3+ S 2- Al