1. 1 Vanessa N. Prasad-Permaul Valencia Community College CHM 1045

2. 2 Electron Configuration of Atoms Electron Configuration of an atom: a particular distribution of electrons among available subshells. Li 3 electrons: 1s 2 2s 1 Orbital Diagram: a diagram that shows how the orbitals of a subshell are occupied by electrons. Li 3 electrons: 1s 2s

3. 3 Electron Configuration of Atoms Pauli Exclusion Principle: no two electrons in an atom can have the same four quantum Numbers. Rewritten: An orbital can hold at most two electrons, and then only if the electrons have opposite spins. SUBSHELLNUMBER OF ORBITALS MAXIMUM NUMBER OF ELECTRONS s (l = 0)12 p (l = 1)36 d (l = 2)510 f (l = 3)714

4. 4 Electron Configuration of Atoms EXAMPLE 8.1 Which one of the following orbital diagrams or electron configurations are possible and which are impossible, according to the Pauli Exclusion Principle? Explain: a. d. 1s 3 2s 1 1s 2s 2p b.e. 1s 2 2s 1 2p 7 1s 2s 2p c. f. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 8 4s 2 1s 2s 2p

5. 5 Electron Configuration of Atoms EXERCISE 8.1 Look at the following orbital diagrams and electron configurations, which are possible and which are not according to the Pauli Exclusion Principle? Explain: a. d. 1s 2 2s 2 2p 4 1s 2s 2p b.e. 1s 2 2s 4 2p 2 1s 2s 2p c. f. 1s 2 2s 2 2p 6 3s 2 3p 10 3d 10 1s 2s 2p

6. 6 Electron Configuration of Atoms The Building-Up Principle Ground State: The electron configuration associated with the lowest energy level of the atom. Na 1s 2 2s 2 2p 6 3s 1 Excited State: The electron configuration associated with an atom the energy levels other than the most stable (ground state). Na* 1s 2 2s 2 2p 6 3p 1 (emission of a yellow light at 589nm) Energy s < p < d < f

7. Electron Configuration of Atoms Rules of Aufbau Principle:  Lower n orbitals fill first.  Each orbital holds two electrons; each with different m s.  Half-fill degenerate (same energy level) orbitals before pairing electrons. (p, d, & f)    NOT   __ 3p x 3p y 3p z 7

8. Electron Configuration of Atoms 8 A mnemonic diagram of the Aufbau Principle Increasing EnergyIncreasing Energy

9. Electron Configuration of Atoms Element Diagram Configuration Li (Z = 3)   1s 2 2s 1 1s 2s Be (Z = 4)   1s 2 2s 2 1s 2s B (Z = 5)    __ __1s 2 2s 2 2p 1 1s 2s 2p x 2p y 2p z C (Z = 6)     __1s 2 2s 2 2p 2 1s 2s 2p x 2p y 2p z 9

10. Electron Configuration of Atoms Element Diagram Configuration O (Z = 8)      1s 2 2s 2 2p 4 1s 2s 2p x 2p y 2p z Ne (Z = 10)      1s 2 2s 2 2p 6 1s 2s 2p x 2p y 2p z S (Z = 16)          1s 2s 2p x 2p y 2p z 3s 3p x 3p y 3p z 1s 2 2s 2 2p 6 3s 2 3p 4 or [Ne] 3s 2 3p 4 abbreviations using the noble gases referred to as a pseudo-noble gas core. Valence Electrons: an electron in an atom outside the noble gas or pseudo-noble-gas core. 10

11. Electron Configuration of Atoms 11 Table of Electron Configuration using noble gas core

12. Electron Configuration of Atoms 12 Table of the Valence-shell configurations of the Elements

13. Electron Configuration of Atoms 13 The building-up order using the Periodic Table.

14. Electron Configuration of Atoms 14 EXAMPLE 8.2: Use the Aufbau Principle to obtain the complete electron configuration for the ground state of the Gallium atom (Z = 31). Abbreviate with the noble gas core and what is the valence shell configuration? Gallium (Ga) Z = 31 Full configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 Rearranged by shells: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1 Abbreviated configuration: [Ar] 3d 10 4s 2 4p 1 Valence-shell configuration: 4s 2 4p 1

15. Electron Configuration of Atoms 15 EXERCISE 8.2: Use the Aufbau Principle to obtain the complete electron configuration for the ground state of the Manganese atom (Z = 25). Abbreviate with the noble gas core and what is the valence shell configuration?

16. Electron Configuration of Atoms 16 EXAMPLE 8.3: What are the configurations for the outer electrons of : a.Tellurium Z = 52 [Kr] 5s 2 4d 10 5p 4 [Kr] 4d 10 5s 2 5p 4 5s 2 5p 4 a. Nickel Z = 28 [Ar]4s 2 3d 8 [Ar] 3d 8 4s 2 3d 8 4s 2

17. Electron Configuration of Atoms 17 EXERCISE 8.3: What are the configurations for the noble gas and the outer electrons of : a.Arsenic b.Bromine c.Silver d.Calcium

18. Electron Configuration of Atoms 18 EXERCISE 8.4: The lead atom has a ground state configuration of [Xe]4f 14 5d 10 6s 2 6p 2. find the period and group for this element. From it’s position in the periodic table, classify it as main-group element, a transition element or an inner transition element.

19. Anomalous Electron Configurations  19 of the predicted configurations from the periodic table are wrong  Largely due to unusual stability of both half-filled and fully filled subshells Cr (Z=24) expected configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4      __ 4s 3d 3d 3d 3d 3d actual configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5    4s 3d 3d 3d 3d 3d 19

20. Orbital Diagrams of Atoms; Hund’s Rule Hund’s Rule: 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 the electrons. * 1s 2s 2p 20

21. 21 EXAMPLE 8.4: Write the orbital diagram for the ground state of the iron atom. Z = 26 Electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 Noble gas: [Ar] 3d 6 4s 2 Valence electron: 3d 6 4s 2 Orbital Diagram: 1s 2s 2p 3s 3p 4s 3d Orbital Diagrams of Atoms; Hund’s Rule

22. 22 EXERCISE 8.5: Write the orbital diagram for the ground state of the phosphorus atom. Z = 15 Electron configuration: Noble gas: Valence electron: Orbital Diagram: Orbital Diagrams of Atoms; Hund’s Rule

23. 23 Paramagnetic Substance: a substance that is weakly attracted by a magnetic field this attraction is generally the result of unpaired electrons Diamagnetic Substance: a substance tht is not attracted by a magnetic field or is very slightly repelled by such a field. This property generally means that the substance has only paired electrons Magnetic Properties of Atoms

24. Periodic Properties 24 The Periodic Law: When the elements are arranged by atomic number, their physical and chemical properties vary periodically. Atomic Radius Ionization Energy Electron Affinity (important in discussions of chemical bonding)

25. Periodic Properties 25 Representation of Atomic Radii of the Main-Group Elements

26. Periodic Properties 26 Two Factors that primarily determine the size of the outermost orbital: Principle quantum number (n) of the orbital; the larger the n of the orbital, the larger the size of the orbital. The effective nuclear charge acting on an electron in the orbital; as the effective nuclear charge increases, the size of the orbital decreases by pulling the electrons inward. Effective nuclear charge: the positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution.

27. 27 EXAMPLE 8.5: Refer to the periodic table use the trends noted for size of atomic radii to arrange the following in order of increasing atomic radius: Al, C, Si C is above Si in Group IVA  the radius of C is smaller than that of Si. Al and Si are in the same period, going to the right of the table  the radius of Si is smaller than that of Al C, Si, Al In order of increasing radius Periodic Properties

28. 28 EXERCISE 8.6: Using the periodic table, arrange the following in order of increasing atomic radius: Na, Be, Mg. Periodic Properties

29. 29 Ionization Energy: the minimum energy needed to remove the highest-energy (the outermost) electron from the neutral atom in the gaseous state. Li (1s 2 2s 1 ) Li + (1s 2 ) + e - Within a period, values tend to increase with atomic number  the lowest values are found in Group 1A. Elements with the lower ionization energy lose electrons easily Noble gases have high ionization energy Generally, as atomic numbers increase, ionization energy increases Periodic Properties

30. Trends of First Ionization Energy, E i 30 Increase Periodic Properties

31. Higher Ionization Energy, E i1234…  Easy to remove an electron from a partially filled valence shell  Difficult to remove an electron from a filled valence shell  Large amount of stability associated with filled s & p subshells  Na: 1s 2 2s 2 2p 6 3s 1  Mg: 1s 2 2s 2 2p 6 3s 2  Cl:1s 2 2s 2 2p 6 3s 2 3p 5 31

32. Ionization Energy, E i  Some exceptions/irregularities to general trend  E i Be > E i B we would expect opposite  Be 4 e 1s 2 2s 2  B 5 e 1s 2 2s 2 2p 1  2s is closer to nucleus than 2p, Z eff for Be is stronger  2s is held more tightly and is harder to remove 32 Periodic Properties

33. Ionization Energy, E i  E i N > E i O we would expect opposite  N 7e 1s 2 2s 2 2p 3 __ __ __  O 8e 1s 2 2s 2 2p 4 __ __ __  Only difference is that an electron is being removed from a half-filled orbital (N) and one from a filled orbital (O)  Electrons repel each other and tend to stay as far apart as possible, electrons that are forced together in a filled orbital are slightly higher in energy so it is easier to remove one  O < N 33 Periodic Properties

34. 34 EXAMPLE 8.6: Using the periodic table, arrange the following in order of increasing ionization energy: Ar, Se, S. Se is below S I Group VIA  ionization energy of Se should be lower than S S and Ar are in the same period with Z increasing from S to Ar  the ionization energy of S should be lower than that of Ar. Se > S> Ar Periodic Properties

35. 35 EXERCISE 8.7: The first ionization energy of the chlorine atom is 1251 kJ/mol. State which of the following values would be the more likely ionization energy for the iodine atom. Explain. a. 1000kJ/mol or b. 1400kJ/mol Periodic Properties

36. Ionic Radii or size  Atoms expand when converted to anions  III Ans 2 np 1 __ __ __  IV Ans 2 np 2 __ __ __  V Ans 2 np 3 __ __ __  VI Ans 2 np 4 __ __ __  VII Ans 2 np 5 __ __ __ Adding one electron to each of these will not add another shell it will just fill an already occupied p subshell  Therefore the expansion is due to the decrease in Z eff and the increase in the electron-electron repulsions 36

37. Ionic Radii or size  Atoms contract when an electron is removed to form a cation.  Dec. # of shells  Inc. Z eff : Less electrons, less shielding, outer electrons more attracted to nucleus, therefore smaller more compact 37

38. Higher Ionization Energy, E i1234…  Ionization is not limited to one electron M + Energy  M + + eE i1 M + + Energy  M 2+ + eE i2 M 2+ + Energy  M 3+ + eE i3  Larger amounts of energy are needed for each successive ionization, harder to remove an electron from a positively charged cation 38

39. Electron Affinity, E ea  Energy change that occurs when an electron is added to an isolated atom in the gaseous state.  The more negative the E ea, the greater the tendency of the atom to accept an electron  Group 7A (halogens) have the most negative E ea, high Z eff and room in valence shell  Group 2A and 8A have near zero or slightly positive E ea 39 Periodic Properties

40. 40 EXERCISE 8.8: Using the general comments that were discussed in this section, decide which has the larger negative electron affinity: C or F. Periodic Properties

41. Alkali Metals  Group 1A (ns 1 )  Metallic  Soft  Good Conductors  Low melting point  Lose 1 electron in redox reactions; powerful reducing agent  Very reactive  Not found in elemental state in nature 41 Periodicity in the Main-Group Elements

42. Alkaline Earth Metals  Group 2A (ns 2 )  Harder, but still relatively soft  Silvery  High melting point than group 1A  Less reactive than group 1A  Loses 2e - in redox reaction; powerful reducing agent  Not found in elemental form in nature 42 Periodicity in the Main-Group Elements

43. Group 3A (ns 2 np 1 )  All but Boron which is a metalloid  Silvery  Good conductor  Relatively soft  Less reactive than 1A & 2A  metals 43 Periodicity in the Main-Group Elements

44. Halogens  Group 7A (ns 2 np 5 )  Non-metals  Diatomic molecules  Tend to gain e - during redox reaction. 44 Periodicity in the Main-Group Elements

45. Noble Gases  Group 8A (ns 2 np 6 )  Colorless, odorless, unreactive gases  Stable because of the filled subshell  Makes it difficult to add electrons or remove electrons 45 Periodicity in the Main-Group Elements

46. Example 1: Electron Config. And NG Abb. 1. Sodium 2. Titanium 3. Argon 46

47. Example 2: Ionic Radii Which of the following in each pair has a larger atomic radius? 1. Carbon or Fluorine 2. Chlorine or Iodine 3. Sodium or Magnesium 4. O or O 2- 5. Ca or Ca 2+ 47

48. Example 3: Quantum Numbers and Electron Configuration What are the 4 quantum numbers for the following? Remember you are only interested in the last electron!! 1. C 2. Na + 3. S 4. N 3- 48

49. Example 4: Electron config. and NG Abb. 1. Cl - 2. F - 3. Ca 2+ 4. Na + 49