1. C HAPTER 11 “T HE P ERIODIC T ABLE & P ERIODIC T RENDS ”

2. O RGANIZING THE E LEMENTS OBJECTIVES: Explain how elements are organized in a periodic table.

3. O RGANIZING THE E LEMENTS OBJECTIVES: Compare early and modern periodic tables.

4. O RGANIZING THE E LEMENTS OBJECTIVES: Identify three broad classes of elements.

5. O RGANIZING THE E LEMENTS A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about 13 had been identified by the year 1700. As more were discovered, chemists realized they needed a way to organize the elements.

6. O RGANIZING THE E LEMENTS Chemists used the properties of elements to sort them into groups. In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties One element in each triad had properties intermediate of the other two elements

7. M ENDELEEV ’ S P ERIODIC T ABLE By the mid-1800s, about 70 elements were known to exist Dmitri Mendeleev – a Russian chemist and teacher Arranged elements in order of increasing atomic mass Thus, the first “Periodic Table”

8. M ENDELEEV He left blanks for yet undiscovered elements When they were discovered, he had made good predictions But, there were problems: Such as Co and Ni; Ar and K; Te and I

9. A BETTER ARRANGEMENT In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number The arrangement used today The symbol, atomic number & mass are basic items included

10. P ERIODIC L AW a periodic pattern of an elements properties when arranged in order of increasing atomic # properties are related to the elements electron configuration

11. COMPONENTS OF THE PERIODIC TABLE 1. Periods go across represent energy levels, “n” “n” = 1-7 properties of elements change as “Z” increases

12. COMPONENTS OF THE PERIODIC TABLE 2. Groups/Families go down represent sublevels: s, p, d, f have similar properties designated by letters A/B

13. P ERIODIC T RENDS OBJECTIVES: Describe trends among the elements for atomic size.

14. P ERIODIC T RENDS OBJECTIVES: Explain how ions form.

15. P ERIODIC T RENDS OBJECTIVES: Describe periodic trends for first ionization energy, ionic size, and electronegativity.

16. P ERIODIC T RENDS 1. Atomic # & Atomic Properties 2. Electron Configuration (s, p, d, f)

17. P ERIODIC T RENDS 3. Atomic Size & Radius Period: generally size decreases as “Z” increases (go across)

18. N UCLEAR C HARGE E FFECT  holds the electrons in the atom.  The positive nucleus pulls in on the negative electrons

19. P ERIODIC T RENDS 3. Atomic Size & Radius Group: generally size increases as “n” increases (go down)

20. Atomic Number Atomic Radius (pm) H Li Ne Ar 10 Na K Kr Rb 3 Period 2

21. P ERIODIC T RENDS 4. Ionic Size the gain/loss of e- produces an ion that carries a charge (+/-) Metals lose e-, are smaller than a neutral atom because they “lost” an energy level & nuclear charge increases Nonmetals gain e-, are larger than a neutral atom because the nuclear charge decreases with more e-

22. P ERIODIC T RENDS 4. Ionic Size (cont) Group: generally size increases as you go down Period: generally cations & anions decrease in size as you go across

24. P ERIODIC T RENDS 5. Ionization Energy energy needed to remove an e- from an atom Group generally IE decreases as you go down Period generally IE increases as you go across

25. I ONIZATION E NERGY The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE.

26. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

27. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276 Why did these values increase so much ?

28. W HAT FACTORS DETERMINE IE A. Nuclear Charge (across) more protons means a greater attraction for e- & need more energy to remove B. Shielding Effect (down) more e- means less attraction for e-, easier to remove C. Radius (down) more nrg levels means less attraction for e-, easier to remove D. Sublevels (across) full/half-full sublevels are more stable & need more energy to remove

29. S HIELDING The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Second electron has same shielding, if it is in the same period

30. I ONIZATION E NERGY - G ROUP TRENDS As you go down a group, the first IE decreases because... The electron is further away from the attraction of the nucleus, and There is more shielding.

31. I ONIZATION E NERGY - P ERIOD TRENDS All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals.

32. First Ionization energy Atomic number He He has a greater IE than H. Both elements have the same shielding since electrons are only in the first level But He has a greater nuclear charge H

33. First Ionization energy Atomic number H He l Li has lower IE than H l more shielding l further away l These outweigh the greater nuclear charge Li

34. First Ionization energy Atomic number H He l Be has higher IE than Li l same shielding l greater nuclear charge Li Be

35. First Ionization energy Atomic number H He l B has lower IE than Be l same shielding l greater nuclear charge l By removing an electron we make s orbital half-filled Li Be B

36. First Ionization energy Atomic number H He Li Be B C

37. First Ionization energy Atomic number H He Li Be B C N

38. First Ionization energy Atomic number H He Li Be B C N O Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital

39. First Ionization energy Atomic number H He Li Be B C N O F

40. First Ionization energy Atomic number H He Li Be B C N O F Ne Ne has a lower IE than He Both are full, Ne has more shielding Greater distance

41. First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

42. I ONIZATION E NERGY

43. D RIVING F ORCES Full Energy Levels require lots of energy to remove their electrons. Noble Gases have full orbitals. Atoms behave in ways to try and achieve a noble gas configuration.

44. 2 ND I ONIZATION E NERGY For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected. True for s 2 Alkaline earth metals form 2+ ions.

45. 3 RD IE Using the same logic s 2 p 1 atoms have an low 3rd IE. Atoms in the aluminum family form 3+ ions. 2nd IE and 3rd IE are always higher than 1st IE!!!

46. P ERIODIC T RENDS 6. Electron Affinity attraction of an atom for an e- higher # (pos. #) means stronger attraction neg.#’s: don’t want any e-; must put energy in to attract an e- Group: generally affinity decreases as “n” increases Period: generally affinity increases as “Z” increases Metals - low;Nonmetals - high affinities

48. METALS conduct electricity/heat malleable lustrous ductile tenacious lose electrons (+ #’s) high melting point left of line low IE low e- affinity low electronegativity larger size

49. NONMETALS insulators nonmalleable dull nonductile brittle low melting point smaller size gain electrons (- #’s) right of line high IE high e- affinity high electronegativit y

50. E LECTRONEGATIVITY tendency for an atom to attract e- to itself when chemically combining with another element Group: generally electronegativity decreases as “n” increases (larger size, the less it can attract e- to itself) Period: generally electronegativity increases as “Z” increases (size stays same, nuclear charge increases, attraction increases) except transition metals