Chapter 5

1. Chapter 5 The Periodic Law

2. Objectives: • Describe the periodic tables of Moseley and Mendeleev. • Identify the various families of elements on the periodic table. • State the trends in atomic radius, ionization energy, electron affinity and ion size with a group or period on the periodic table. • Identify the relationship between these trends and the structure of the atom.

3. Chapter 5 • Section 1 • History of the Periodic Table

4. By 1860, more than 60 elements had been discovered. Chemists needed to learn the properties of these elements as well as those of the compounds they formed. At the time their was no method for accurately determining an element’s atomic mass. Things were disorganized and confusing among chemists.

5. Dmitri Mendeleev • Mendeleev hoped to organize the elements according to their properties. • He arranged the elements in order of increasing atomic mass and noticed certain similarities in their chemical properties would appear at regular intervals. • In 1869 Mendeleev created the first periodic table grouping the elements with similar properties.

7. Henry Moseley • In 1911 Moseley revised the periodic table of Mendeleev by arranging the elements in order of increasing atomic number (number of protons) and not atomic mass. • The elements were still grouped by similar physical and chemical properties.

8. Modern Periodic Table The periodic table has undergone extensive changes since Mendeleev’s time. More than 40 new elements have been discovered or synthesized.

9. 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.

10. Later additions to the periodic table include: Noble Gases (1868-1900) – group 18 elements that are characterized by their relative un-reactivity. The Lanthanides (early 1900’s) – the 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu)

11. The Actinides– the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr). • The lanthanides and actinides belong in periods 6 and 7 of the periodic table. • To save space and to group them together they are set off below the main portion of the periodic table.

12. Summary – Section 5.1 • Recognize the work of Mendeleev and Mosely. • How the modernperiodic table is arranged with respect to the elements. • Know the three sets of elements added to the periodic table after Mendeleev.

13. Chapter 5 • Section 2 • Electron Configuration • and the • Periodic Table

14. The Group 18 elements (Noble Gases) undergo few chemical reactions. • This stability results from the gases electron configuration. • Their highest occupied levels are completely filled with electrons (octet). • An atom’s electron configuration governs the atom’s chemical properties.

15. The elements arranged vertically in the periodic table share chemical properties. • The are also organized in horizontal rows, or periods. • The length of each period is determined by the number of electrons that can occupy the sublevels in that period.

17. In the first period, the 1s sublevel is filled. The 1s sub level can hold a total of two electrons. Therefore, the first period consists of two elements - hydrogen and helium. The second period can held 8 electrons and therefore consists of 8 elements. The third period is similar to the second period – 8 elements.

18. When you get to the fourth and fifth rows you have to include the transition metals. There are 10 transition metals in each of these rows. The total number of elements in the fourth and fifth rows are therefore 18 elements.

19. When you get to the sixth and seventh rows you have to include the lanthanides and actinides. • There are 14 elements in each of these groups. • The total number of elements in the six and seventh rows can therefore include 32 elements.

20. The periodic table is divided into 4 blocks

21. The period of an element can be determined from the element’s electron configuration. • For example, arsenic – As, has the electron configuration of: • [Ar]3d104s24p3 • The 4 in 4p3 indicates that arsenic is in the fourth period of the periodic table.

22. Without looking at the periodic table give the period number for the following electron configuration: • [Xe]6s2

23. The s-Block Elements The elements of the s block are chemically reactive metals. The Group 1 metals are more reactive than the Group 2 metals. The outermost energy level in an atom of the Group 1 elements contain a single electron. The ease with which the single electron is lost make the Group 1 metals extremely reactive.

24. Alkali Metals – The elements in Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium and francium). Because the alkali metals are so reactive they are not found in nature as free elements. They combine vigorously with most nonmetals. They react strongly with water to produce hydrogen gas.

25. Alkaline-earth Metals – The elements in Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium and radium). • Atoms of the alkaline-earth metals contain a pair of electrons in their outermost shell. • Slightly less reactive than the alkali metals. • Still too reactive to be found in nature as free elements.

26. Classwork Review the practice problems 1 and 2a-d on page 133.

27. The d-Block Elements The elements of the d block are known as the transition metals. They are good conductors of electricity and have a high luster. They are typically less reactive than the alkali and alkaline-earth metals. Palladium, platinum and gold are among the least reactive of all elements.

28. Classwork Review the practice problems 1 and 2a-b on page 136.

29. The p-Block Elements The elements of the p-block and s-block are known as the main-group elements. They consist of the elements in Group 13-18 (except helium). The properties of the p-block vary greatly. The p-block contains metals (Al), metalloids (Si) and nonmetals (Br).

30. The electron configuration of bromine is: • [Ar]3d104s24p5

31. Halogens – the elements of Group 17 (fluorine, chlorine, bromine, iodine and astatine). The halogens are the most reactive nonmetals. They react vigorously with most metals to form compounds known as salts (NaCl). Fluorine and chlorine are gases. Bromine is a liquid and iodine is a solid.

32. The metalloids divide the metals from the nonmetals. • They are mostly brittle solids with some properties of metals and some of nonmetals. • The electrical conductivity falls between the metals and nonmetals. • The metals of the p block are reactive enough to be found in nature only as compounds and not free elements (except bismuth).

33. Classwork Review the practice problems 1a-b and 2a-b on page 138.

34. Homework Section Review – page 139 Questions 1, 2, 4 and 5 End of chapter problems – page 156-157 Questions 27, 28 and 29 Due:

35. Chapter 5 • Section 3 • Electron Configuration • and the • Periodic Properties

36. So far you have learned that the elements are arranged in the periodic table according to their atomic number. There is also a correlation between the arrangement of the elements and their electronic configuration. We will look at the relationship between the electron configurations and the periodic trends of the elements.

37. Atomic Radii • Ideally, the size of an atom is defined by the edge of its orbital. • However, this boundary varies under different conditions. • One way to express the atomic radius is to measure the distance between the nuclei of two identical atoms that are bonded together, then divide this distance by two.

38. Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

39. Period Trends There is a gradual decrease across a row.

40. There is a gradual decrease across a row. • The trend to smaller atoms across a period is caused by increasing positive charge of the nucleus. Adding of protons or increasing atomic number. • The electrons are pulled closer to the nucleus. • The increase pull results in a smaller atomic radius.

41. Group Trends There is an increase down a group.

42. There is an increase down a group. • As electrons are added to sublevels in higher energy levels located further from the nucleus, the size of the atoms increase. • An exception is between aluminum (radius – 143 pm) and gallium (radius – 135 pm). • This is due to the gallium being proceeded by the 10 d-block elements. • The nuclear charge is considerably higher.

43. Problem • Of the elements magnesium-Mg, chlorine-Cl, sodium-Na, and phosphorus-P, which has the largest atomic radius and why? • Of the elements calcium-Ca, beryllium-Be, barium-Ba and strontium-Sr, which as the largest atomic radius and why?

44. Ionization Energy • Ionization energy (IE) – the energy required to remove one electron from a neutral atom of an element. • A + energy A+ + e- • An ion is an atom or group of bonded atoms that has a positive or negative charge.

47. Period Trends In general, ionization energies of the elements increase across a period. Group 1 elements – have the lowest ionization energies. Therefore they lose electrons most easily. Very reactive. Group 18 elements - have the highest ionization energies . They do not lose electrons easily. Very low reactivity.

48. The increase is due to increasing nuclear charge (more protons going across a period). • A higher positive charge more strongly attracts electrons in the same energy level. • Therefore, it is tougher to remove an electron from an atom. • Increasing nuclear charge is responsible for both an increasing ionization energy and decreasing atomic radius across a period.

49. Group Trends Ionization energies generally decrease down a group. Electrons removed from atoms of the elements down a group are farther from the nucleus. Also, the electrons from the lower energy levels shield the outer electrons. Therefore, they are removed more easily.

50. Classwork Review practice problems 1and 2 on page 142. Review practice problems 1 (a-c) on page 146.