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CH 8: Electron Configuration & periodicity

Vanessa N. Prasad-Permaul Valencia Community College CHM 1045. CH 8: Electron Configuration & periodicity. Electron Configuration of Atoms. Electron Configuration of an atom: a particular distribution of electrons among available subshells . Li 3 electrons: 1s 2 2s 1

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CH 8: Electron Configuration & periodicity

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  1. Vanessa N. Prasad-Permaul Valencia Community College CHM 1045 CH 8: Electron Configuration & periodicity

  2. Electron Configuration of Atoms Electron Configuration of an atom: a particular distribution of electrons among available subshells. Li 3 electrons: 1s2 2s1 Orbital Diagram: a diagram that shows how the orbitals of a subshell are occupied by electrons. Li 3 electrons: 1s 2s

  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.

  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. 1s32s1 • 1s 2s 2p • b. e. 1s22s12p7 • 1s 2s 2p • c. f. 1s22s22p63s23p63d84s2 • 1s 2s 2p

  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. 1s22s22p4 • 1s 2s 2p • b. e. 1s22s42p2 • 1s 2s 2p • c. f. 1s22s22p63s23p103d10 • 1s 2s 2p

  6. Electron Configuration of Atoms The Building-Up Principle Ground State: The electron configuration associated with the lowest energy level of the atom. Na 1s22s22p63s1 Excited State: The electron configuration associated with an atom the energy levels other than the most stable (ground state). Na* 1s22s22p63p1 (emission of a yellow light at 589nm) Energy s < p < d < f

  7. Electron Configuration of Atoms Rules of Aufbau Principle: • Lower norbitals fill first. • Each orbital holds two electrons; each with different ms. • Half-fill degenerate (same energy level) orbitals before pairingelectrons. (p, d, & f)    NOT   __ 3px 3py 3pz

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

  9. Electron Configuration of Atoms Element Diagram Configuration Li (Z = 3) 1s2 2s1 1s 2s Be (Z = 4) 1s2 2s2 1s 2s B (Z = 5)  __ __ 1s2 2s2 2p1 1s 2s 2px 2py 2pz C (Z = 6)   __ 1s2 2s2 2p2 1s 2s 2px 2py 2pz

  10. Electron Configuration of Atoms Element Diagram Configuration O (Z = 8)  1s2 2s2 2p4 1s 2s 2px 2py 2pz Ne (Z = 10)  1s2 2s2 2p6 1s 2s 2px 2py 2pz S (Z = 16)  1s 2s 2px 2py 2pz 3s 3px 3py 3pz 1s2 2s2 2p6 3s2 3p4 or [Ne] 3s2 3p4 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.

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

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

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

  14. Electron Configuration of Atoms 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: 1s22s22p63s23p64s23d104p1 Rearranged by shells: 1s22s22p63s23p63d104s24p1 Abbreviated configuration: [Ar]3d104s24p1 Valence-shell configuration: 4s24p1

  15. Electron Configuration of Atoms 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 • EXAMPLE 8.3: What are the configurations for the outer electrons of : • Tellurium Z = 52 • [Kr] 5s24d105p4 • [Kr] 4d105s25p4 • 5s25p4 • Nickel Z = 28 • [Ar]4s23d8 • [Ar] 3d84s2 • 3d84s2

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

  18. Electron Configuration of Atoms EXERCISE 8.4: The lead atom has a ground state configuration of [Xe]4f145d106s26p2. 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: 1s2 2s2 2p6 3s2 3p6 4s2 3d4   __ 4s 3d 3d3d3d3d actual configuration: 1s2 2s2 2p6 3s2 3p6 4s1 3d5     4s 3d 3d3d3d3d

  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 1s 2s 2p 1s 2s 2p

  21. Orbital Diagrams of Atoms; Hund’s Rule EXAMPLE 8.4: Write the orbital diagram for the ground state of the iron atom. Z = 26 Electron configuration: 1s22s22p63s23p63d64s2 Noble gas: [Ar] 3d64s2 Valence electron: 3d64s2 Orbital Diagram: 1s 2s 2p 3s 3p 4s 3d

  22. Orbital Diagrams of Atoms; Hund’s Rule 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:

  23. Magnetic Properties of Atoms 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

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

  26. Periodic Properties • 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. Periodic Properties 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

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

  29. Periodic Properties • Ionization Energy: the minimum energy needed to remove the highest-energy (the outermost) electron from the neutral atom in the gaseous state. • Li (1s22s1) Li+ (1s2) + 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

  30. Periodic Properties Trends of First Ionization Energy, Ei Increase Increase

  31. Higher Ionization Energy, Ei1234… • 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: 1s2 2s2 2p6 3s1 • Mg: 1s2 2s2 2p6 3s2 • Cl: 1s2 2s2 2p6 3s2 3p5

  32. Periodic Properties Ionization Energy, Ei • Some exceptions/irregularities to general trend • Ei Be > Ei B we would expect opposite • Be 4 e 1s2 2s2 • B 5 e 1s2 2s2 2p1 • 2s is closer to nucleus than 2p, Zeff for Be is stronger • 2s is held more tightly and is harder to remove

  33. Periodic Properties Ionization Energy, Ei • Ei N > Ei O we would expect opposite • N 7e 1s2 2s2 2p3 __ __ __ • O 8e 1s2 2s2 2p4 __ __ __ • 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

  34. Periodic Properties • 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

  35. Periodic Properties 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

  36. Ionic Radii or size • Atoms expand when converted to anions • III A ns2 np1 __ __ __ • IV A ns2 np2 __ __ __ • V A ns2 np3 __ __ __ • VI A ns2 np4 __ __ __ • VII A ns2 np5 __ __ __ 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 Zeff and the increase in the electron-electron repulsions

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

  38. Higher Ionization Energy, Ei1234… • Ionization is not limited to one electron M + Energy  M+ + e Ei1 M+ + Energy  M2+ + e Ei2 M2+ + Energy  M3+ + e Ei3 • Larger amounts of energy are needed for each successive ionization, harder to remove an electron from a positively charged cation

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

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

  41. Periodicity in the Main-Group Elements Alkali Metals • Group 1A (ns1) • 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

  42. Periodicity in the Main-Group Elements Alkaline Earth Metals • Group 2A (ns2) • 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

  43. Periodicity in the Main-Group Elements Group 3A (ns2np1) • All but Boron which is a metalloid • Silvery • Good conductor • Relatively soft • Less reactive than 1A & 2A • metals

  44. Periodicity in the Main-Group Elements Halogens • Group 7A (ns2np5) • Non-metals • Diatomic molecules • Tend to gain e- during redox reaction.

  45. Periodicity in the Main-Group Elements Noble Gases • Group 8A (ns2np6) • Colorless, odorless, unreactive gases • Stable because of the filled subshell • Makes it difficult to add electrons or remove electrons

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

  47. Example 2: Ionic Radii Which of the following in each pair has a larger atomic radius? • Carbon or Fluorine • Chlorine or Iodine • Sodium or Magnesium • O or O2- • Ca or Ca2+

  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!! • C • Na+ • S • N3-

  49. Example 4: Electron config. and NG Abb. • Cl- • F- • Ca2+ • Na+

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