The How and Why

1. Chapter 6 The Periodic Table The How and Why

2. History Elements known to ancients: C, Cu, Au, Fe, Pb, Hg, Ag, S, Sn Added before 1700: As, Sb, Bi, P, Zn Dobereiner, Johann (1780-1849): arranged elements in triads (Ca, Sr, Ba; Cl, Br, I) John Newlands (1837-1898): arranged elements in group of eight Properties repeat every 8th elements: Li, Be, B, C, N, O, F, Na Na, Mg, Al, Si, P, S, Cl, K

3. History Dmitri Mendeleev (1834-1907):used the masses of elements as most of the masses were determined in XIX century (1869) Arranged elements in order of increasing atomic masses Found a pattern of repeating properties

4. Mendeleev’s Table Grouped elements in columns by similar properties in order of increasing atomic mass. Found some inconsistencies - felt that the properties were more important than the mass, so switched order ( Te, I). Found gaps in the trends- maybe undiscovered elements. Predicted their properties before they were found ( eka boron – Sc; eka aluminum- Ga; eka silicon: Ge).

5. The Modern Table Elements are still grouped by properties. Similar properties are in the same column. Order is in increasing atomic number (Moseley, 1914). Added a column of elements Mendeleev didn’t know about (Noble Gases). The noble gases weren’t found because they didn’t react with anything.

6. Periodic Law • Mendeleev (1869): Properties of elements are a function of the atomic masses of the elements. • Modern periodic Law (Mosley, 1914) properties of elements are a periodic function of their atomic numbers ( # of protons in the nucleus). Explains Mendeelev’s irregularities.

7. Horizontal rows are called periods • There are 7 periods 1 2 3 4 5 6 7

8. Vertical columns are called groups or families • Elements are placed in columns by similar properties.

9. 8A0 1A The elements in the A groups are called the representative elements (also numbered 1-18) 2A 3A 4A 5A 6A 7A

10. The group B are called the transition elements Inner Transition elements

11. Group 1A(1) are the alkali metals Group 2A(2) are the alkaline earth metals

12. Group 7A (17) is called the Halogens Group 8A (18) are the noble or inert gases

13. Why the similarities in Properties? The part of the atom another atom sees is the electron cloud. These are the outside or valence orbitals. The orbitals fill up in a regular pattern. The outside orbital electron configuration repeats. The properties of atoms repeat.

14. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

15. He 2 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

16. S- block s1 Alkali metals all end in s1 Alkaline earth metals all end in s2 Have to include He but it fits better later. He has the properties of the noble gases. s2

17. Transition Metals -d block s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

18. The P-block p1 p2 p6 p3 p4 p5

19. f6 f13 f1 f2 f3 f4 f5 f7 f8 f10 f12 f14 f11 f9 F - block • inner transition elements

20. 1 2 3 4 5 6 7 Each row (or period) is the energylevel fors and porbitals.

21. d orbitals fill up after previous energy level … first d is 3d even though it’s in row 4. 1 2 3 4 5 6 7 3d

22. 1 2 3 4 5 6 7 f orbitals start filling at 4f 4f 5f

23. Writing Electron configurations the easy way Review Notes

24. Electron Configurations Repeat • The shape of the periodic table is a representation of this repetition of electron configurations. • When we get to the end of the column the outermost energy level is full. • This is the basis for our shorthand.

25. Trends and Properties of Elements in the Periodic Table • The following properties will be examined: • Radius of the atom • Ionization energy • Electron affinity • Electron negativity

26. Effective Nuclear Charge • Effective nuclear charge is experienced by an outer electron at the outer edge of an atom • Zeff = Z – S Z is the atomic number S is the number of core electrons. See blackboard for examples.

27. Atomic Size (Radius) First problem where do you start measuring. The electron cloud doesn’t have a definite edge. Determined by measuring more than 1 atom at a time.

28. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

29. Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level is further away from nucleus Charge on nucleus ( Zeffective) More charge pulls electrons in closer.

30. Group Trends H As we go down a group Each atom has another energy level So the atoms get bigger. Li Na K Rb

31. Periodic Trends As you go across a period the radius gets smaller. Filling the same energy level. ***More nuclear charge (higher Zeff). Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

32. Atomic Radii in the PT

33. Rb Overall K Na Li Atomic Radius (nm) Kr Ar Ne H Atomic Number 10

34. Ionization Energy (IE) The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion. The energy required to remove the first e- is called the first ionization energy. A(g) + IE → A(g) +1

35. Ionization Energy The 2nd ionization energy is the energy required to remove the second electron. 2nd IE is always greater than 1st IE. The 3rd IE is the energy required to remove a third electron. 3rd IE > 2nd IE > 1st IE

36. Symbol First Second Third 810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 13122731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

37. Symbol First Second Third 810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

38. What Determines IE 1. The greater the nuclear charge the greater IE. 2. Distance form nucleus influences IE. 3. Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. 4. Shielding effect

39. Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus. It is shielded from the nucleus by all the inner electrons A second electron in the same energy level has the same shielding.

40. Group Trends • As you go down a group first IE decreases because The electron is further away More shielding by inner electrons as there are more energy levels.

41. Periodic trends All the atoms in the same period have the same energy level. Same shielding. Increasing nuclear charge increases the force of attraction between the nucleus and the electrons. IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals.

42. He He >IE than H same shielding greater nuclear charge H First Ionization energy Atomic number

43. He Li < IE than H more shielding further away (>n) outweighs greater nuclear charge H First Ionization energy Li Atomic number

44. He Be > IE than Li same shielding greater nuclear charge H First Ionization energy Be Li Atomic number

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

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

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

48. He Breaks the pattern because removing an electron gets to 1/2 filled p orbital N O C H First Ionization energy Be B Li Atomic number

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

50. Ne He F Ne < IE than He Both are full, Ne has more shielding Greater distance (>n) N O C H First Ionization energy Be B Li Atomic number