1. Periodic Relationships Among the Elements Chapter 5 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. When the Elements Were Discovered

3. Dimitri Mendeleev (1834 – 1907) Arranged the elements by increasing ATOMIC MASS, he saw a periodic repetition of properties Produced the first PERIODIC TABLE – 1871 The table placed elements with similar properties in the same column Kept “holes” for undiscovered elements, and predicted the properties in advance Development of the Periodic Table

4. Properties of elements predicted by Mendeleev Development of the Periodic Table

5. H. G. Moseley (1887-1915) Rearranged the elements by ATOMIC NUMBER This has become the MODERN PERIODIC TABLE Development of the Periodic Table

6. PERIODIC LAW When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic (repeating) pattern. patterns on the periodic table are called periodic trends

7. Valence Electrons valence electrons: electrons available to be lost, gained, or shared in the formation of chemical compounds ◦ electrons in the outermost energy level ◦ electrons that are responsible for reactions Elements in a group have similar properties because they have the same valence electron configuration

8. All electrons under the highest energy level Inner Core Electrons

9. Valence Electron Configuration Groupe- configurationValence electrons 1ns 1 2ns 2 13ns 2 np 1 14ns 2 np 2 15ns 2 np 3 16ns 2 np 4 17ns 2 np 5 18ns 2 np 6

10. +1+2+3 -2-3 Charges Of Representative Elements 8.2

11. Na + : [Ne]Al 3+ : [Ne]F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne] N 3- : 1s 2 2s 2 2p 6 or [Ne] What atoms are isoelectronic with Ne? an ion that has the same electron configuration Please note – metals will lose electrons to form cations, non-metals will gain electrons to form anions

12. Transition Metals Transition metals form ions by first losing electrons from the s-block. Consider Fe. ◦ Write the electron configuration for the neutral atom. ◦ [Ar]4s 2 3d 6 ◦ Now, write the electron configuration for the iron (II) ion. ◦ [Ar]3d 6

13. S-block elements The elements of Group 1, with the exception of hydrogen, are known as the alkali metals. Extremely reactive and will not be found in their pure form in nature. In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife. Form strong bases when placed in water. Electron configuration = ns 1 The elements of Group 2 are called the alkaline-earth metals. Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form. Will also form strong bases when placed in water. Electron configuration = ns 2

14. The d-block elements (Groups 3-12) electron configuration = (n - 1)d 1-10 ns 0-2 Transition metals: the d-block elements are metals with typical metallic properties ◦ Ductility ◦ Malleability ◦ Conductivity ◦ Having luster Groups and periods within the transition metals have varied properties and it is therefore hard to group metals by individual properties

15. p-block elements (group 13-18) electron configuration = ns 2 np 1-6 Main-group elements: the elements in the p-block of varying properties All non-metals are in the p-block (aside from hydrogen) All metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) are in the p-block

16. Halogens Elements in group 17 are the halogens They are highly reactive in similar ways ◦ All form salts with group 1 and 2 metals They are volatile ( gas forms easily) Electron configuration = ns 2 np 5

17. f-block elements Lanthanides and Actinides Lanthanides are shiny metals and have similar reactivity to group 2 metals Their position reflects the fact that they involve the filling of the 4f sublevel Actinides: ◦ Mainly synthetic and radioactive

18. Periodic Table Groups

19. Atomic Radii (the distance between the nucleus & the end of the cloud) One half the distance between the nuclei of identical atoms that are bonded together Atomic radii: 200 pm Bond length 400 pm

20. Atomic Radii DOWN a Group ↓ As you go down a group another energy level is added, the atom size gets larger ↓ The number of occupied orbitals between the nucleus and the outermost energy level increases Shielding Effect: reduction of attraction between positive nucleus and outer electrons, outer electrons are not held tight and can move away

21. Na Atomic Radius: down group P P P X X X P X P X P X P P P P P X X X X X X P

22. K P P P X X X P X P X P X P P P P P X X X X X P X X X X X X X X X

23. Atomic Radii DOWN a Group ↓ DOWN THE GROUP ATOMIC RADIUS INCREASES more energy levels, the larger the size of the atom

24. Atomic Radii ACROSS a Period → Each atom gains one proton and one electron in the same energy level → Each added electron is the same distance from the nucleus → The positive charge increases and exerts a greater force on the electrons thereby pulling it closer to the nucleus

25. REMEMBER! PROTONS are bigger and stronger! electrons are smaller and weaker! P ++ - - e

26. Atomic Radius: across period P P P X X X P X P X P X P P P P P X X X X X

27. Atomic Radii ACROSS a Period → ACROSS THE PERIOD ATOMIC RADIUS DECREASES stronger attraction of protons, easier to hold on to the electrons

28. Atomic Radii Decreases in size from left to right Increases in size from up to down Increases

30. Atomic Radii

31. Ionic Radii half the distance from center-center of 2 like ions

32. Ionic Radius ACROSS the Period Cation: positive ion formed from losing an electron → A cation is always smaller than the original atom → The more electrons lost the more protons available to attract a smaller number of electrons.

33. Na + Ionic Radius P P P X P X X P P P P P P P X X X X X X X X X P

34. Ionic Radius ACROSS the Period → ACROSS THE PERIOD IONIC RADIUS DECREASES greater pull on electrons, the shorter the radius

35. Ionic Radius ACROSS the Period Anion: negative ion formed from gaining an electron → A anion is always larger than the original atom → The more electrons gained, the less protons available to attract a larger number of electrons.

36. F Ionic Radius P P P X P X X P P P P P P P X X X X X X X X P -

37. Ionic Radius ACROSS the Period → ACROSS THE PERIOD IONIC RADIUS DECREASES more electrons added, more difficult to keep track of, increasing the size of the atom

38. Ionic Radius DOWN a Group ↓ As you go down a group another energy level is added, increasing the size of the atom. (just like the atomic radius)

39. Ionic Radius DOWN the Group ↓ DOWN THE GROUP IONIC RADIUS INCREASES more energy levels, increase in atom size

40. Ionic Radii Positive ion = cation – smaller than parent atom Negative ion = anion – larger than parent atom

41. Ionic Radii Decreases in size from left to right Increases in size from up to down Increases

42. amount of energy needed to remove an electron from an atom ↓As you go down a group atoms become larger, electrons are farther from the nucleus and more easily removed ↓The more electrons in an atom between the nucleus and valence shell, the greater the shielding effect Ionization Energy

43. Ionization Energy DOWN a Group ↓ DOWN THE GROUP IONIZATION ENERGY DECREASES greater distance from the nucleus, the easier to lose an electron (less energy needed)

44. Ionization Energy ACROSS a Period → As atomic radius decreases there is a greater attraction between protons and electrons. → The stronger the attraction, the more energy needed to remove an electron.

45. Ionization Energy ACROSS a Period → ACROSS THE PERIOD IONIZATION ENERGY INCREASES

46. Additional Ionization energies Removing a second electron is the 2 nd Ionization Energy, and it typically greater than the first ionization energy With each additional electron removed, the Ionization Energy will increase.

47. Ionization Energy ◦ Increases from left to right ◦ Decreases from up to down Increases

48. The change in energy associated with the addition of an electron to a neutral atom Dark green means no electron affinity Yellow means high electron affinity (trends by groups) Electron Affinity

49. Electron Affinity ACROSS the Period → As the atomic radius decreases and shielding is constant, it is easier to attract an electron. → The addition of protons in the nucleus creates a stronger attraction to an electron

50. Electron Affinity DOWN the Group ↓ The larger the atom, the more difficult to accept electrons

52. Electron Affinity ◦ Increases from left to right ◦ Decreases from up to down Increases

53. It is a “tug of war” between the two atoms of a bond Electronegativity H F : : : Which is the more electronegative element?... tendency for an atom to attract electrons

54. Electronegativity ACROSS the Period → ACROSS THE PERIOD ELECTRONEGATIVITY INCREASES stronger the attraction, the easier to add more electrons

55. Electronegativity DOWN the Group ↓ DOWN THE GROUP ELECTRONEGATIVITY DECREASES farther the distance from the nucleus, more difficult to attract electrons

56. Electronegativity ◦ Increases from left to right ◦ Decreases from up to down Increases

57. Other Trends Reactivity of Metals Video 1 Reactivity of Metals Video 2

58. Increasing reactivity METAL REACTIVITY

59. Increasing reactivity NONMETAL REACTIVITY