1 / 69

Chapter 8

Chapter 8. Electron Configuration and Chemical Periodicity. Electron Configuration and Chemical Periodicity. 8.1 Development of the Periodic Table. 8.2 Characteristics of Many-Electron Atoms. 8.3 The Quantum-Mechanical Model and the Periodic Table.

ita
Télécharger la présentation

Chapter 8

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 8 Electron Configuration and Chemical Periodicity

  2. Electron Configuration and Chemical Periodicity 8.1Development of the Periodic Table 8.2Characteristics of Many-Electron Atoms 8.3 The Quantum-Mechanical Model and the Periodic Table 8.4Trends in Some Key Periodic Atomic Properties 8.5The Connection Between Atomic Structure and Chemical Reactivity

  3. Table 8.1 Mendeleev’s Predicted Properties of Germanium (“eka Silicon”) and Its Actual Properties Predicted Properties of eka Silicon(E) Actual Properties of Germanium (Ge) Property atomic mass 72 amu 72.61 amu appearance gray metal gray metal density 5.5 g/cm3 5.32 g/cm3 molar volume 13 cm3/mol 13.65 cm3/mol specific heat capacity 0.31 J/g.K 0.32 J/g.K oxide formula EO2 GeO2 oxide density 4.7 g/cm3 4.23 g/cm3 sulfide formula and solubility ES2; insoluble in H2O; soluble in aqueous (NH4)2S GeS2; insoluble in H2O; soluble in aqueous (NH4)2S ECl4; (< 100 oC) chloride formula (boiling point) GeCl4; (84 oC) chloride density 1.9 g/cm3 1.844 g/cm3 element preparation reduction of K2EF6 with sodium reduction of K2GeF6 with sodium

  4. Observing the Effect of Electron Spin Figure 8.1

  5. principal n positive integers (1, 2, 3,…) orbital energy (size) angular momentum l integers from 0 to n-1 orbital shape (l values of 0, 1, 2 and 3 correspond to s, p, d and f orbitals, respectively.) magnetic ml integers from -l to 0 to +l orbital orientation spin ms +1/2 or -1/2 direction of e- spin Table 8.2 Summary of Quantum Numbers of Electrons in Atoms Name Symbol Allowed Values Property Each electron in an atom has its own unique set of four (4) quantum numbers.

  6. The Pauli Exclusion Principle No two electrons in the same atom can have the same four quantum numbers An atomic orbital can hold a maximum of two electrons and they must have opposite spins (paired spins)

  7. Spectral evidence of energy-level splitting in many-electron systems Figure 8.2 Many-electron atoms: have nucleus-electron and electron-electron interactions Leads to the splitting of energy levels into sublevels of differing energies: the energy of an orbital depends mostly on its n value (size) and somewhat on its l value (shape)

  8. Effect of nuclear charge (Zeffective) Effect of electron repulsions (shielding) Factors Affecting Atomic Orbital Energies Higher nuclear charge lowers orbital energy (stabilizes the system) by increasing nucleus-electron attractions. 1. Additional electron in the same orbital An additional electron raises the orbital energy through electron-electron repulsions. 2. Additional electrons in inner orbitals Inner electrons shield outer electrons more effectively than do electrons in the same sublevel.

  9. The effect of nuclear charge Greater nuclear charge lowers orbital energy Figure 8.3

  10. The effect of another electron in the same orbital Each electron shields the other from the full nuclear charge, thus raising orbital energy Figure 8.4

  11. The effect of other electrons in inner orbitals Inner electrons shield outer electrons very well and raise orbital energy greatly Figure 8.5

  12. The effect of orbital shape Figure 8.6 2s electron farther from nucleus than 2p electron, but penetrates near nucleus; increased attraction results in lower orbital energy

  13. General Rule for Predicting Relative Sublevel Energies For a given n value, the lower the l value, the lower the sublevel energy; thus…. s < p < d < f

  14. as s, p, d or f Figure 8.7 Order for filling energy sublevels with electrons Illustrating Orbital Occupancies A. The electron configuration # of electrons in the sublevel n l B. The orbital diagram (box or circle)

  15. A vertical orbital diagram for the Li ground state no color = empty light = half-filled Sublevel energy increases from bottom to top 1s22s1 dark = filled, spin-paired Figure 8.8

  16. Hund’s Rule When orbitals of equal energy are available, the electron configuration of lowest energy has the maximum number of unpaired electrons with parallel spins.

  17. PROBLEM: Write a set of quantum numbers for the third electron and a set for the eighth electron of the fluorine (F) atom. 9F 1s 2s 2p n = n = l = l = ml = ml = ms= ms= Determining Quantum Numbers from Orbital Diagrams Sample Problem 8.1 PLAN: Use the orbital diagram to find the third and eighth electrons. Up arrow = +1/2 Down arrow = -1/2 SOLUTION: The third electron is in the 2s orbital. Its quantum numbers are: 2 0 0 +1/2 The eighth electron is in a 2p orbital. Its quantum numbers are: 2 1 -1 -1/2

  18. Orbital occupancy for the first 10 elements, H through Ne Figure 8.9 He and Ne have filled outer shells: confers chemical inertness

  19. Condensed ground-state electron configurations in the first three periods Figure 8.10 Similar outer electron configurations correlate with similar chemical behavior.

  20. Similar Reactivities within A Group Orbitals are filled in order of increasing energy, which leads to outer electron configurations that recur periodically, which leads to chemical properties that recur periodically. Figure 8.11

  21. Cr and Cu: Half-filled and filled sublevels are unexpectedly stable!

  22. A periodic table of partial ground-state electron configurations Figure 8.12

  23. The relation between orbital filling and the Periodic Table Figure 8.13

  24. Common tool used to predict the filling order Of sublevels n values are constant horizontally l values are constant vertically combined values of n+1 are constant diagonally p. 302

  25. Categories of Electrons Inner (core) electrons: fill all the lower energy levels of an atom Outer electrons: those electrons in the highest energy level (highest n value) of an atom Valence electrons: those involved in forming compounds; the bonding electrons; among the main-group elements, the valence electrons are the outer electrons

  26. General Observations about the Periodic Table A. The group number equals the number of outer electrons (those with the highest value of n) (main-group elements only) B. The period number is the n value of the highest energy level. C. The n value squared (n2) gives the total number of orbitals in that energy level; 2n2 gives the maximum number of electrons in the energy level.

  27. PROBLEM: Using the periodic table, give the full and condensed electron configurations, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements: full configuration: condensed configuration: partial orbital diagram: 4s1 SAMPLE PROBLEM 8.2 Determining Electron Configuration (a) potassium (K: Z = 19) (b) molybdenum (Mo: Z = 42) (c) lead (Pb: Z = 82) PLAN: Use the atomic number for the number of electrons and the periodic table for the order of filling of the electron orbitals. Condensed configurations consist of the preceding noble gas plus the outer electrons. SOLUTION: (a) for K (Z = 19) 1s22s22p63s23p64s1 [Ar] 4s1 K has 18 inner electrons.

  28. full configuration: condensed configuration: partial orbital diagram: 5s1 4d5 full configuration: condensed configuration: partial orbital diagram: 6s2 6p2 SAMPLE PROBLEM 8.2: (continued) (b) for Mo (Z = 42) 1s22s22p63s23p64s23d104p65s14d5 [Kr] 5s14d5 Mo has 36 inner electrons and 6 valence electrons. (c) for Pb (Z = 82) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2 [Xe] 6s24f145d106p2 Pb has 78 inner electrons and 4 valence electrons.

  29. KEY PRINCIPLE All physical and chemical properties of the elements are based on the electronic configurations of their atoms.

  30. Defining metallic and covalent radii A. Metallic radius: 1/2 the distance between adjacent nuclei in a crystal B. Covalent radius: 1/2 the distance between bonded nuclei in a molecule Figure 8.14

  31. Atomic radii of main-group and transition elements Opposing forces: Changes in n and changes in Zeff Overall Trends (A) n dominates within a group; atomic radius generally increases in a group from top to bottom (B) Zeff dominates within a period; atomic radius generally decreases in a period from left to right Figure 8.15

  32. Periodicity of atomic radius Large size shifts when moving from one period to the next Figure 8.16

  33. PROBLEM: Using only the periodic table, rank each set of main group elements in order of decreasing atomic size. SAMPLE PROBLEM 8.3 Ranking Elements by Atomic Size (a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb PLAN: Size increases down a group; size decreases across a period. SOLUTION: (a) Sr > Ca > Mg These elements are in Group 2A. (b) K > Ca > Ga These elements are in Period 4. (c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr. (d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

  34. Ionization Energy The amount of energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions; an energy- requiring process; value is positive in sign IE1 = first ionization energy: removes an outermost electron from the gaseous atom: atom(g) ion+(g) + e- ∆E = IE1 > 0 IE2 = second ionization energy: removes a second electron from the gaseous ion: ion+(g) ion+2(g) + e- ∆E = IE2 > IE1 Atoms with a low IE1 tend to form cations during reactions, whereas those with a high IE1 (except noble gases) often form anions.

  35. Ionization Energies: Correlations with Atomic Size 1. As size decreases, it take more energy to remove an electron. 2. Ionization energy generally decreases down a group. 3. Ionization generally increases across a period.

  36. Periodicity of first ionization energy (IE1) Figure 8.17 Lowest values for alkali metals; highest values for noble gases

  37. First ionization energies of the main-group elements Increase within a period and decrease within a group Figure 8.18

  38. PROBLEM: Using the periodic table, rank the elements in each of the following sets in order of decreasing IE1: SAMPLE PROBLEM 8.4 Ranking Elements by First Ionization Energy (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs PLAN: IE decreases down in a group; IE increases across a period. SOLUTION: (a) He > Ar > Kr Group 8A elements- IE decreases down a group. (b) Te > Sb > Sn Period 5 elements - IE increases across a period. (c) Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I is to the left of Xe; Cs is further to the left and down one period.

  39. The first three ionization energies of beryllium (in MJ/mol) Successive IEs increase, but a large increase is observed to remove the first core electron (for Be, IE3) Figure 8.19

  40. PROBLEM: Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration: IE1 IE2 IE3 IE4 IE5 IE6 1012 1903 2910 4956 6278 22,230 SAMPLE PROBLEM 8.5 Identifying an Element from Successive Ionization Energies PLAN: Look for a large increase in energy that indicates that all of the valence electrons have been removed. SOLUTION: The largest increase occurs at IE6, that is, after the 5th valence electron has been removed. The element must have five valence electrons with a valence configuration of 3s23p3, The element must be phosphorus. P (Z = 15). The complete electronic configuration is: 1s22s22p63s23p3.

  41. Electron Affinity (EA) The energy change accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions. atom(g) + e- ion-(g) ∆E = EA1 (usually negative) EA2 is always positive (adding negative charge to negatively charged ion).

  42. Electron affinities of the main-group elements Negative values = energy is released when the ion forms Positive values = energy is absorbed to form the anion Figure 8.20

  43. General Trends Involving IEs and EAs Reactive non-metals: Groups 6A and 7A; in their ionic compounds they form negative ions (have high IEs and very (-) EAs) Reactive metals: Group 1A; in their ionic compounds, they form positive ions (have low IEs and slightly (-) EAs) Noble gases: Group 8A; they do not lose or gain electrons (have very high IEs and slightly (+) EAs)

  44. Trends in three atomic properties Figure 8.21

  45. Metallic Behavior Metals: shiny solids; tend to lose electrons in reactions with non-metals (left and lower 3/4 of periodic table) Non-metals: tend to gain electrons in reactions with metals; upper right-hand quarter of periodic table Metalloids: have intermediate properties; located between the metals and non-metals in the periodic table Metallic behavior decreases left to right and increases top to bottom in the periodic table

  46. Trendsin metallic behavior Figure 8.22

  47. Moving down a GROUP: elements at the top tend to form anions and those at the bottom tend to form cations Moving across a PERIOD: elements at the left tend to form cations and those at the right lend to form anions

More Related