Chapter 8 Periodic Properties & Electron Configurations Bushra Javed

1. Chapter 8 Periodic Properties & Electron Configurations BushraJaved

2. Contents • Electron Spin • Electron Configurations of elements • The development of Periodic Table • Periodic Trends

3. Electron Configurations • Quantum-mechanical theory describes the behavior of electrons in atoms • The electrons in atoms exist in orbitals • A description of the orbitals occupied by electrons is called an electron configuration number of electrons in the orbital 1s1 principal energy level of orbital occupied by the electron sublevel of orbital occupied by the electron

4. How Electrons Occupy Orbitals • Calculations with Schrödinger’s equation show hydrogen’s one electron occupies the lowestenergy orbital in the atom • Schrödinger’s equation calculations for multi-electron atoms cannot be exactly solved due to electron-electron interactions in multi-electron atoms

5. How Electrons Occupy Orbitals • The interactions that occur in multi-electron atoms are due to : • electron spin • & energy splitting of sublevels

6. The Property of Electron Spin • Spin is a fundamental property of all electrons • All electrons have the same amount of spin • The orientation of the electron spin is quantized, it can only be in one direction or its opposite • spin up or spin down • The electron’s spin adds a fourth quantum number to the description of electrons in an atom, called the Spin Quantum Number, ms • not in the Schrödinger equation

7. The Property of Electron Spin • Experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field. • The experiment reveals that the electrons spin on their axis • spinning charged particles generate a magnetic field

8. Electron Spin • If there is an even number of electrons, about half the atoms will have a net magnetic field pointing “north” and the other half will have a net magnetic field pointing “south”

9. Electron Spin

10. The two possible spin orientations of an electron and the conventions for msare illustrated here.

11. ms, and Orbital Diagrams • mscan have values of +½ or −½ • Orbital Diagrams use a square to represent each orbital and a half-arrow to represent each electron in the orbital • By convention, a half-arrow pointing up is used to represent an electron in an orbital with spin up • Spins must cancel in an orbital • paired

12. unoccupied orbital orbital with one electron orbital with two electrons Orbital Diagrams • We often represent an orbital as a square or a circle and the electrons in that orbital as arrows • the direction of the arrow represents the spin of the electron

13. Pauli Exclusion Principle • No two electrons in an atom may have the same set of four quantum numbers • Therefore no orbital may have more than two electrons, and they must have the opposite spins • Knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel

14. Pauli Exclusion Principle ssublevel has 1 orbital, therefore it can hold 2 electrons psublevel has 3 orbitals, therefore it can hold 6electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons

15. Sublevel Splitting in Multi-electron Atoms • The sublevels in each principal energy shell of Hydrogen all have the same energy • We call orbitals with the same energy degenerate • For multi-electron atoms, the energies of the sublevels are split • caused by charge interaction, shielding and penetration • The lower the value of thelquantum number, the less energy the sublevel has • s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)

16. Shielding & Effective Nuclear Charge • Each electron in a multi-electron atom experiences both the attraction to the nucleus and repulsion by other electrons in the atom • These repulsions cause the electron to have a net reduced attraction to the nucleus – it is shielded from the nucleus • The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron

17. Shielding & Penetration

18. penetration& Effective Nuclear Charge • The closer an electron is to the nucleus, the more attraction it experiences • The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus • Penetration causes the energies of sublevels in the same principal level to not be degenerate

19. Effect of Penetration and Shielding • In the fourth and fifth principal levels, the effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level • The energy separations between one set of orbitals and the next become smaller beyond the 4s • the ordering can therefore vary among elements • causing variations in the electron configurations of the transition metals and their ions

20. Filling the Orbitals with Electrons • Energy levels and sublevels fill from lowest energy to high • s→p→d→f • AufbauPrinciple • Orbitals that are in the same sublevel have the same energy • No more than two electrons per orbital • Pauli Exclusion Principle • When filling orbitals that have the same energy, place one electron in each before completing pairs • Hund’s Rule

21. Filling the Orbitals with Electrons The lowest-energy configuration of an atom is called its ground state. Any other allowed configuration represents an excited state.

22. Electron Configuration of Atoms The electron configuration is a listing of the sublevels in order of filling with the number of electrons in that sublevel written as a superscript. Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

23. ‘Building up’ order of Filling in sublevels (Ground State Electron Configurations) Start by drawing a diagram putting each energy shell on a row and listing the sublevels, (s, p, d, f), for that shell in order of energy (left-to-right) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s Next, draw arrows through the diagonals, looping back to the next diagonal each time

24. Electron Configuration & the Periodic Table • The Group number corresponds to the number of valence electrons • The length of each “block” is the maximum number of electrons the sublevel can hold • The Period number corresponds to the principal energy level of the valence electrons

25. s1 s2 p1 p2p3p4p5 p6 s2 1 2 3 4 5 6 7 d1 d2d3d4d5d6d7d8d9d10 f2f3f4f5f6f7f8f9 f10f11f12f13f14 f14d1

26. Blocks of Elements

27. Blocks of Elements For main-group (representative) elements, ansor ap subshell is being filled. Ford-block transition elements, a d subshell is being filled. Forf-block transition elements, anf subshell is being filled.

28. Electron Configurations Example 1 Write the ground state electron configuration of the chlorine atom, Cl, For chlorine, Cl, Z = 17. 1s2 2s2 2p6 3s2 3p5

29. Electron Configurations Example 2 Write the ground state electron configuration of the manganese atom, Mn

30. Electron Configurations Example 3: What is the ground-state electron configuration of tantalum (Ta)? a. 1s22s22p63s23p64s23d104p65s24d105p6 6s25d104f7 b. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f145d3 c. 1s22s22p63s23p64s23d104p65s24d105p6 6s2 5d3 d. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f14

31. Electron Configurations Example 4 Which of the following electron configurations corresponds to the ground state electron configuration of an atom of a transition element? a. 1s22s22p2 b. 1s22s22p63s23p5 c. 1s22s22p63s23p64s2 d. 1s22s22p63s23p63d54s2

32. Exceptions to the ‘Building up’ order of Filling • About 21 of the predicted configurations are inconsistent with the actual configurations observed. One Possible Reason: • Half filled and fully filled subshells are highly Stable • Cr, Cu ,Ag and U are some of the exceptions.

33. Exceptions to the ‘Building up’ order of Filling Chromium (Z=24) and copper (Z=29) have been found by experiment to have the following ground-state electron configurations: Cr: 1s2 2s2 2p6 3s2 3p6 3d54s1 Cu: 1s2 2s2 2p6 3s2 3p6 3d104s1 In each case, the difference is in the 3d and 4ssubshells.

34. Exceptions to the ‘Building up’ order of Filling • Expected • Cr = [Ar]4s23d4 • Cu = [Ar]4s23d9 • Mo = [Kr]5s24d4 • Ru = [Kr]5s24d6 • Pd = [Kr]5s24d8 • Found Experimentally • Cr = [Ar]4s13d5 • Cu = [Ar]4s13d10 • Mo = [Kr]5s14d5 • Ru = [Kr]5s14d7 • Pd = [Kr]5s04d10

35. Electron Configurations Example 5 Which of the following electron configurations represents an allowed excited state of the indicated atom? a. He: 1s2 b. Ne: 1s2 2s2 2p6 c. Na: 1s2 2s2 2p6 3s2 3p2 4s1 d. P: 1s2 2s2 2p6 3s2 3p2 4s1

36. The Noble Gas coreElectron Configuration The noble gases have eight valence electrons. except for He, which has only two electrons We know the noble gases are especially non-reactive He and Ne are practically inert The reason the noble gases are so non-reactive is that the electron configuration of the noble gases is especially stable

37. Noble Gas Core Electron Configuration • A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [] to represent all the inner electrons, then just write the last set. Rb= 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1

38. Noble Gas Core Electron Configuration Example 6 Which ground-state electron configuration is incorrect? a. Fe: [Ar] 3d5 b. Ca: [Ar] 4s2 c. Mg: [Ne] 3s2 d. Zn: [Ar] 3d10 4s2

39. Writing Electron Configurations The pseudo-noble-gas core includes the noble-gas subshells and the filled inner, (n – 1), d subshell. For bromine, the pseudo-noble-gas core is [Ar]3d10

40. Hund’s rule Hund’s rule states: that the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons

41. Electron Configurations

42. 1s 2s 2p Electron Configurations Example 7 Draw an orbital diagram for nitrogen.

43. 1s 2s Electron Configurations Example 8 Which of the following electron configurations or orbital diagrams are allowed and which are not allowed by the Pauli exclusion principle? If they are not allowed, explain why? • 1s22s12p3 • 1s22s12p8 • 1s22s22p63s23p63d8 • 1s22s22p63s23p63d11

44. Electron Configurations Example 9 write the complete ground state orbital diagram and electron configuration of potassium

45. Valence Electrons • The electrons in all the sublevels with the highest principal energy shell are called the valence electrons • Electrons in lower energy shells are called core electrons • Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the numberof valence electrons

48. Valence- shell configuration For main-group elements, the valence configuration is in the form nsAnpB The sum of A and B is equal to the group number. So, for an element in Group VA of the third period, the valence configuration is 3s23p3

49. Valence-shell configuration Example 10 What are the valence-shell configuration for arsenic and zinc? Arsenic is in period 4, Group VA. Its valence configuration is 4s24p3. Zinc, Z = 30, is a transition metal in the first transition series. Its noble-gas core is Ar, Z = 18. Its valence configuration is 4s23d10.

50. Valence-shell configuration Example: 11 What is the valence -shell configuration of Tc, Z = 43