Ch. 5: Periodic Table

1. Ch. 5: Periodic Table C. Goodman, Doral Academy Preparatory High School, 2011-2013

2. Essential Question: Section 5.1 • What is the history of the development of the Periodic Table? • What is the periodic law, and how can it be used to predict physical and chemical properties of elements? • What is the overall organization of the modern Periodic Table? • Three types of elements • Named groups • Other families

3. Section 5.1 Vocabulary • Mendeleev • Moseley • Periodic Law • Period • Group • Main group elements

4. Periodic Table - Definition Periodic Table-- an arrangement of the elements in order of their atomic numbers so that elements with the same chemical properties are in the same group (family). Examples: halogens, noble gases, alkali metals.

5. Why is it cool? • http://www.youtube.com/watch?v=u2ogMUDBaf4&playnext=1&list=PLAC3A0775D813045F&feature=results_main

6. History of the Periodic Table I • Mendeleev: 1869 • Atomic mass • Repeating (periodic) patterns of reactivity • In his favor: predicted the discovery of Gallium, which was isolated in his lifetime • Certain characteristic properties of elements can be foretold from their atomic weights • Problem: Iodine and Tellurium

7. History of the Periodic Table II • Moseley: 1914 • Atomic numbereach element has a unique atomic number; resequenced the table by electronic charge (=atomic #) rather than atomic weight. • Periodic Law • In his favor: solved the “Iodine and Tellurium” problem

8. Moseley’s Periodic Law - Definition • Periodic LawThe physical and chemical properties of the elements are periodic functions of their atomic numbers. http://www.youtube.com/watch?v=OduTDUGeAXEFind

10. How to use the periodic table… Atomic number: # of protons in the nucleus of an atom Symbol Basically the abbreviation for the element Average atomic mass # of Protons + # of Neutrons (amu) Weighted average mass of isotopes of the element Remember nuclear notation for isotopes? Notice that the atomic mass is a whole number – it’s not an average Also notice the different locations of the atomic mass and atomic number.

11. Groups (families) The Columns Elements in groups have similar chemical properties Periods The Rows Elements properties vary across periods The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period

12. Important terms • Main Group Elements s-block + p-block elements • Transition metalsd-block elements • Lanthanides and actinidesf-block elements • Metalloids, metals, non-metals (see below)

13. 2 Main Sections in Periodic Table Nonmetals Metals • - Majority of elements • Good Electrical & Heat Conductors • - Room temperature = most solids • Contain properties • Malleability • Ductility • High tensile strength • Poor Electrical Conductors • Poor Heat Conductors • Room temperature = most gases • One is a liquid at r.t. = Bromine • - Solid nonmetals generally brittle

14. Names of groups • Group 1a – Alkali metals • Group 2a – Alkali earth metals • Group 7a (17) – halogens • Group 8a (18) – Noble gases

15. Names of families– Transition metals

16. Names of families– Semiconductors

19. Section 2: Electron Configuration • What is the relationship between the location of atoms in a group, their electron configuration, and their chemical and physical properties? • What are the s-, p-, d-, and f-blocks, and how can their electron configurations of their elements be determined?

20. Section 5.2 vocabulary • Ion • Valence • Valence electrons • s-, p-, d- and f-block elements

21. Valence Electrons • Valence = outermost energy level in which contains electrons (in unexcited state). • Valence electrons are the electrons on the outermost energy level of the element. • The number of valence electrons determines the type of chemical reactions available to the element!

22. What does this have to do with groups? • All main group elements in a particular group have the same number of valence electrons. • Prove it? • Hehhehheh – that’s your job! • Write electron configurations, noble gas notation, of the s and p block elements in the first 5 rows. Write valence electrons (s/b only) in contrasting color • Elements hydrogen – xenon, for columns #1a, 2a, 3a, 4a, 5a, 6a, 7a, 8a • Huhhh? See next slide

23. Valence electrons &energy levels • Purpose: to determine the relationship between the group number and number of valence electrons. • Procedures • For each of the elements in the first 5 periods of the following groups… group (1a, 2a, 3a, 4a, 5a, 6a, 7a, and 8a) • Write the name of the element • Write the type of element • Write the electron configuration, using noble gas notation • Conclusion (Answer the following question)s: • 1. How does the group # relate to the number of valence electrons? • 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

24. Valence electrons & Energy Levels • Conclusion (Answer the following questions): • 1. How does the group # relate to the number of valence electrons? • 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?

25. Group 14

26. Valence Electrons • Main group elements have characteristic numbers of valence electrons. • Group 1 – 1 valence electron • Group 2 – 2 valence electrons • Groups 13-18 • # valence electrons = Group # - 10 • Example: Group 13 elements have 13-10 = 3 valence electrons

27. Valence Electrons • Main group elements have characteristic numbers of valence electrons. • S block • Group 1 – 1 valence electron • Group 2 – 2 valence electrons • P block • Groups 13-18 • Group # - 10 • Example: Group 13 elements have 13-10 = 3 valence electrons

28. Relationship Between Periodicity and Electron Configurations

29. Sample problem A • Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located. • Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?

30. Sample problem B An element has the electron configuration [Kr] 5s24d5. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.

31. Sample problem C Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, use your periodic table to name the element, and identify it as a metal, nonmetal, or metalloid.

32. In book, p. 135 1: Identify period, block, group, element [Kr]5s2 2. write configuration of… a. Group 2 elements b. the group 2 element in the fourth period. c. the element in the 3rd period, group 15

33. Sample Problem C Solution • p-block (group # >12 • 14-10 = 4 electrons in s, p • 2 e- in s, 2e- in p • The outer electron configuration is 2s22p2. • The element is carbon, C, which is a nonmetal.

34. More practice problems! Name the block and group in which each of the following elements is located in the periodic table. Use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity. • [Xe]6s24f145d8 • [Ne]3s23p2 • [Ne]3s23p5 • [Xe]4f66s1

35. How do I know which groups are more reactive than others? For main group elements, look at the number of valence (s/p) electrons 8 valence electrons (noble gases)= not reactive Less reactive 4<3<2<1 More reactive Less reactive 4<5<6<7 More reactive

36. Sample Problem D Solution The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nine electrons. Therefore, this element is in the d block. The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity. b. The incompletely filled p sublevel shows that this element is in the p block. A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens. The element is chlorine, Cl, and is highly reactive. Section2 Electron Configuration and the Periodic Table Chapter 5 Periods and Blocks of the Periodic Table, continued

37. Sample Problem D Solution, continued c. This element has a noble-gas configuration and thus is in Group 18 in the p block. The element is argon, Ar, which is an unreactive nonmetal and a noble gas. d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide. Group numbers are not assigned to the f block. The element is samarium, Sm. All of the lanthanides are reactive metals. Section2 Electron Configuration and the Periodic Table Chapter 5 Periods and Blocks of the Periodic Table, continued

38. Section 3: Periodic TrendsEssential Questions • Compare the periodic trends of atomic radii, ionization energy, electronegativity, and state the reasons for these variations. • What are valence electrons, and how many are present in atoms of each main-group element? • Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.

40. Section 5.3 Vocabulary • Atomic radius • Ion • Ionization energy • Cation • Anion • Electron affinity • Electronegativity

41. Atomic Radii:½ the distance between the nuclei of identical atoms bonded

42. Periodic Trends: Atomic Radii DECREASES across periods Because of increasing positive charge of the nucleus Holds electrons more tightly INCREASES down groups Higher principle quantum number Valence electrons in higher main energy levels Located farther from the nucleus

43. Watch it kid, I’ve got my ion you. Ions

44. Ion – Definition An Ion is… An atom or group of bonded atoms, which has a positive (+) or negative (-) charge.

45. There are two kinds of ions… Cation • CRUNCH (subtract) an electron • This results in a positive charge • When an electron is removed, the atom loses bulk (like a muscle which shrinks when it atrophies) • So, the radius of a cation is smaller than the atomic radius

46. Anions • ADD an electron • This results in a NEGATIVE charge • When an electron is added, the atom gains bulk (like a muscle which grows when you work out) • So, the radius of a anion is larger than the atomic radius

47. Two sodium atoms bumped into each other. One said: "Why do you look so sad?“ The other responded: "I lost an electron.“ The first one asked "Are you sure?“ The other replied "I'm positive."

48. Which elements form which type of ion? • Metal on left tend to form cations (+) • Nonmetals at the upper right tend to form anions (-) • Hydrogen is a non-metal, but it forms a cation (+)

49. Periodic Trends: Ionic Radii – very similar to trends for atomic radii

50. Ionization Energy (IE) Definition Ionization energy is… The energy required to remove one electron from a neutral atom of an element, forming a cation. Note: does not apply to formation of anions! Unit of measure: kJ/mol