Chapter 5 Periodic Table

1. Chapter 5 Periodic Table

2. Mendeleev and Chemical Periodicity • Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. • Repeating patterns are referred to as periodic. • Mendeleev created a table in which elements with similar properties were grouped together—a periodic table of the elements.

3. Mendeleev’s Periodic Table Dmitri Mendeleev Refer to pg 124 Text – For Elements out of Place. Dmitri Mendeleev

4. Properties of Some Elements Predicted By Mendeleev

5. Moseley and the Periodic Law • In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic weight. • The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. • What does this mean?

6. Periodicity of Atomic Numbers

7. Noble Gases – Group 18 or VIII – Unreactive gases with eight valence electrons. Lanthanides– Atomic # 58-71 added to the periodic table in 1900. Actinides – Atomic # 90 - 103

8. The Modern Periodic Table The Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. Period or series - A horizontal row across the periodic table (7 total) Group or family – A vertical column on the periodic table (18 total)

9. A Spiral Periodic Table

10. Period The Periodic Table Group or Family Group or family Period

11. The Properties of Group I: the Alkali Metals • Never found pure in nature • Easily lose valence electron • React violently with water • Intheir pure states,they have a • soft silvery appearance and can • be cut with a knife • React with halogens to form salts • Group configuration = ns1 • n arrangement of the elements in

12. Group 2 of the periodic table • alkaline-earth metals. • beryllium, magnesium, calcium, strontium, barium, and radium • Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form. • Group configuration ns2 Radium was widely used in self-luminous clock and watch hands, until too many watch factory workers had died of it. This antique watch is still quite radioactive and will remain for thousands of years.

13. Halogens – Group 17 or VII • React vigorously with most metals toform salts. • Fluorine and chlorine are gases at room temp., bromine is a reddish liquid, and iodine is a dark purple solid. • Astantine is a synthetic element prepared in very small quantities. • Group configuration ns2np5

14. Transition Metals“Group B Metals”Many exceptions to Aufbau exist in this area. Why? Why was the mad hatter mad?

15. Properties of Metals • Metals are good conductors of heat and electricity • Metals are malleable • Metals are ductile • Metals have high tensile strength • Metals have luster • Located to the left of • the zigzag line on the • periodic table

16. Examples of Metals Potassium, K reacts with water and must be stored in kerosene Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Zinc, Zn, is more stable than potassium Mercury, Hg, is the only metal that exists as a liquid at room temperature

17. Propertiesof Nonmetals Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element. • Nonmetals are poor conductors of heat and electricity • Nonmetals tend to be brittle • Many nonmetals are gases at room temperature • Located to the right of the zigzag line on the periodic table

18. Examples of Nonmetals Microspheres of phosphorus, P, a reactive nonmetal Sulfur, S, was once known as “brimstone” Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

19. Properties of Metalloids Metalloids straddle the border between metals and nonmetals on the periodic table. • They have properties of both metals and nonmetals. • Metalloids are more brittle than metals, less brittle than most nonmetallic solids • Metalloids are semiconductors of electricity • Some metalloids possess metallic luster

20. Silicon, Si – A Metalloid • Silicon has metallic luster • Silicon is brittle like a nonmetal • Silicon is a semiconductor of electricity Other metalloids include: • Boron, B • Germanium, Ge • Arsenic, As • Antimony, Sb • Tellurium, Te

21. S P Blocks of the Periodic Table dbd lock s block f f block The s and p blocks = the main group elements. The d block = the transition metals. The f block = the inner transition metals.

22. Periods and Blocks of the Periodic Table Sample Problem A 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. b. 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)?

23. Solution for Sample Problem A • It is in the sixth period, as indicated by the highest principal quantum number in its configuration, 6. • The element is in the s block. The element is in Group 2, as indicated by the group configuration of ns2.

24. Sample Problem B An element has the electron configuration [Kr]4d55s1. 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.

25. Sample Problem B Solution • The number of the highest occupied energy level is 5, so the element is in the fifth period. • There are five electrons in the d sublevel, which means that it is incompletely filled. The d sublevel can hold 10 electrons. Therefore, the element is in the d block. • For d-block elements, the number of electrons in the ns sublevel (1) plus the number of electrons in the (n 1)d sublevel (5) equals the group number, 6. This Group 6 element is molybdenum.

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

27. Sample Problem C Solution • The group number is higher than 12, so the element is in the p block. • The total number of electrons in the highest occupied s and p sublevels is therefore equal to the group number minus 10 (14  10 = 4). • Two electrons are in the s sublevel, so two electrons must also be present in the 2p sublevel. • The outer electron configuration is 2s22p2. • The element is carbon, C, which is a nonmetal.

28. Sample Problem D Name the block and group in which each of the following elements is located in the periodic table. Then, 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. a. [Xe]4f145d96s1 c. [Ne]3s23p6 b. [Ne]3s23p5 d. [Xe]4f66s2

29. Sample Problem D Solution a.The4f 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.

30. Sample Problem D Solution 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.

31. Periodicity (RepeatingTrends)

32. Summation of Periodic Trends

33. Determination of Atomic Radius Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii” Periodic Trends in Atomic Radius • Radius decreases across a period • Increased effective nuclear charge due • to decreased shielding • Radius increases down a group • Addition of principal quantum levels

34. Table of Atomic Radii

35. Lithium configuration • Lithium ion configuration Compare two cations

36. Oxygen configuration • Oxygen ion configuration Compare two anions

37. Ionization Energy - the energy required to remove an electron from an atom • Increases for successive electrons taken from the same atom • Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove • Tends to decrease down a group Outer electrons are farther from the nucleus

38. Table of 1st Ionization Energies

39. Electron Affinity - The energy change that occurs when an electron is acquired by a neutral atom. • Most atoms release energy when they acquire an electron

40. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across a period • Electronegativities tend to decrease down a group or remain the same

41. Periodic Table of Electronegativities

42. Ionic Radii • Positively charged ions formed when an atom of a metal loses one or more electrons Cations • Smaller than the corresponding atom • Negatively charged ions formed when nonmetallic atoms gain one or more electrons Anions • Larger than the corresponding atom

43. Table of Ion Sizes