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Chapter 8 Periodic Properties of the Elements

Chemistry: A Molecular Approach , 1 st Ed. Nivaldo Tro. Chapter 8 Periodic Properties of the Elements. Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA. 2007, Prentice Hall. Mendeleev. order elements by atomic mass saw a repeating pattern of properties

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Chapter 8 Periodic Properties of the Elements

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  1. Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro Chapter 8Periodic Properties of the Elements Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2007, Prentice Hall

  2. Mendeleev • order elements by atomic mass • saw a repeating pattern of properties • Periodic Law – When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically • put elements with similar properties in the same column • used pattern to predict properties of undiscovered elements • where atomic mass order did not fit other properties, he re-ordered by other properties • Te & I Tro, Chemistry: A Molecular Approach

  3. nm H2O a/b H 1 H2 m Li2O b 7 LiH m/nmBeO a/b 9 BeH2 nm B2O3 a 11 ( BH3)n nm CO2 a 12 CH4 nm N2O5 a 14 NH3 nm 19 HF Li Be B C N O F m Na2O b 23 NaH m MgO b 24MgH2 nm P4O10 a 31 PH3 nm SO3 a 32 H2S nm Cl2O7 a 35.5 HCl Na Mg Al Si P S Cl Periodic Pattern nm O2 16 H2O m Al2O3 a/b 27 (AlH3) nm/mSiO2 a 28 SiH4 m = metal, nm = nonmetal, m/nm = metalloid a = acidic oxide, b = basic oxide, a/b = amphoteric oxide Tro, Chemistry: A Molecular Approach

  4. Mendeleev's Predictions Tro, Chemistry: A Molecular Approach

  5. What vs. Why • Mendeleev’s Periodic Law allows us to predict what the properties of an element will be based on its position on the table • it doesn’t explain why the pattern exists • Quantum Mechanics is a theory that explains why the periodic trends in the properties exist Tro, Chemistry: A Molecular Approach

  6. 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 • as they spin, they generate a magnetic field • spinning charged particles generate a magnetic field • 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” Tro, Chemistry: A Molecular Approach

  7. Electron Spin Experiment Tro, Chemistry: A Molecular Approach

  8. Spin Quantum Number, ms • spin quantum number describes how the electron spins on its axis • clockwise or counterclockwise • spin up or spin down • spins must cancel in an orbital • paired • mscan have values of ±½ Tro, Chemistry: A Molecular Approach

  9. Pauli Exclusion Principle • no two electrons in an atom may have the same set of 4 quantum numbers • therefore no orbital may have more than 2 electrons, and they must have with opposite spins • knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel • s sublevel has 1 orbital, therefore it can hold 2 electrons • p sublevel has 3 orbitals, therefore it can hold 6 electrons • d sublevel has 5 orbitals, therefore it can hold 10 electrons • f sublevel has 7 orbitals, therefore it can hold 14 electrons Tro, Chemistry: A Molecular Approach

  10. Allowed Quantum Numbers Tro, Chemistry: A Molecular Approach

  11. Allowed Quantum Numbers n = 1, 2, 3, 4, 5, 6, 7; these represent the 7 shells l = s, p, d, f; or the shape of the orbital s = 0 p = 1 d = 2 f = 3 ml = ± l, 0; this is the direct along the x, y, z axis ms = ± ½; direction of spin, north or south Tro, Chemistry: A Molecular Approach

  12. Quantum Numbers of Helium’s Electrons • helium has two electrons • both electrons are in the first energy level • both electrons are in the s orbital of the first energy level • since they are in the same orbital, they must have opposite spins Tro, Chemistry: A Molecular Approach

  13. Quantum Numbers of Helium’s Electrons Tro, Chemistry: A Molecular Approach

  14. SPECTROSCOPIC NOTATION for H, atomic number = 1 1 no. of s 1 electrons value of l value of n Writing Atomic Electron Configurations Two ways of writing complete configurations. One is called the spectroscopic notation.

  15. ORBITAL BOX NOTATION for He, atomic number = 2 Arrows 2 depict s 1 electron spin 1s Writing Atomic Electron Configurations Two ways of writing complete configurations. Other is called the orbital box notation. One electron has n = 1, l = 0, ml = 0, ms = + 1/2 Other electron has n = 1, l = 0, ml = 0, ms = - 1/2

  16. Electron Configurations • the ground state of the electron is the lowest energy orbital it can occupy • the distribution of electrons into the various orbitals in an atom in its ground state is called its electron configuration • the number designates the principal energy level • the letter designates the sublevel and type of orbital • the superscript designates the number of electrons in that sublevel • He = 1s2 Tro, Chemistry: A Molecular Approach

  17. unoccupied orbital orbital with 1 electron orbital with 2 electrons Orbital Diagrams • we often represent an orbital as a square and the electrons in that orbital as arrows • the direction of the arrow represents the spin of the electron Tro, Chemistry: A Molecular Approach

  18. Sublevel Splitting in Multielectron Atoms • the sublevels in each principal energy level of Hydrogen all have the same energy – we call orbitals with the same energy degenerate • or other single electron systems • for multielectron atoms, the energies of the sublevels are split • caused by electron-electron repulsion • the lower the value of the l quantum number, the less energy the sublevel has • s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) Tro, Chemistry: A Molecular Approach

  19. Penetrating and Shielding • the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p • the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus • the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively • the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p Tro, Chemistry: A Molecular Approach

  20. Penetration & Shielding Tro, Chemistry: A Molecular Approach

  21. Shielding Effect animation Tro, Chemistry: A Molecular Approach

  22. 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Energy • Notice the following: • because of penetration, sublevels within an energy level are not degenerate • penetration of the 4th and higher energy levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level • the energy difference between levels becomes smaller for higher energy levels

  23. Order of Subshell Fillingin Ground State Electron Configurations start by drawing a diagram putting each energy shell on a row and listing the subshells, (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 Tro, Chemistry: A Molecular Approach

  24. Filling the Orbitals with Electrons • energy shells fill from lowest energy to high • subshells fill from lowest energy to high • s → p → d → f • Aufbau Principle • orbitals that are in the same subshell have the same energy • no more than 2 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 Tro, Chemistry: A Molecular Approach

  25. Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium. • Determine the atomic number of the element from the Periodic Table • This gives the number of protons and electrons in the atom Mg Z = 12, so Mg has 12 protons and 12 electrons Tro, Chemistry: A Molecular Approach

  26. 1s 2s 2p 3s 3p Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium. • Draw 9 boxes to represent the first 3 energy levels sandp orbitals • since there are only 12 electrons, 9 should be plenty Tro, Chemistry: A Molecular Approach

  27. Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium. • Add one electron to each box in a set, then pair the electrons before going to the next set until you use all the electrons • When pair, put in opposite arrows          1s 2s 2p 3s 3p Tro, Chemistry: A Molecular Approach

  28.  1s  2s  2p    3s 3p Example 8.1 – Write the Ground State Electron Configuration and Orbital Diagram and of Magnesium. Write the 4 quantum numbers • Use the diagram to write the electron configuration • Write the number of electrons in each set as a superscript next to the name of the orbital set 1s22s22p63s2 = [Ne]3s2 Tro, Chemistry: A Molecular Approach

  29. Valence Electrons • the electrons in all the subshells 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 number of valence electrons Tro, Chemistry: A Molecular Approach

  30. Electron Configuration of Atoms in their Ground State • Kr = 36 electrons 1s22s22p63s23p64s23d104p6 • there are 28 core electrons and 8 valence electrons • Rb = 37 electrons 1s22s22p63s23p64s23d104p65s1 [Kr]5s1 • for the 5s1 electron in Rb the set of quantum numbers is n = 5, l = 0, ml = 0, ms = +½ • for an electron in the 2p sublevel, the set of quantum numbers is n = 2, l = 1, ml = -1 or (0,+1), and ms = - ½ or (+½)

  31. Electrons and Quantum Numbers • Write the 4 quantum numbers for: • Se • Ag • Cf-247 Tro, Chemistry: A Molecular Approach

  32. Electron Configurations Tro, Chemistry: A Molecular Approach

  33. 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 Tro, Chemistry: A Molecular Approach

  34. Tro, Chemistry: A Molecular Approach

  35. s1 s2 p1 p2 p3 p4 p5 p6 s2 1 2 3 4 5 6 7 d1 d2d3 d4 d5 d6 d7 d8 d9 d10 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1 Tro, Chemistry: A Molecular Approach

  36. Electron Configuration fromthe Periodic Table 8A 1A 1 2 3 4 5 6 7 3A 4A 5A 6A 7A 2A Ne P 3s2 3p3 P = [Ne]3s23p3 P has 5 valence electrons Tro, Chemistry: A Molecular Approach

  37. 4s 3d 6s 4f Transition Elements • for the d block metals, the principal energy level is one less than valence shell • one less than the Period number • sometimes s electron “promoted” to d sublevel Zn Z = 30, Period 4, Group 2B [Ar]4s23d10 • for the f block metals, the principal energy level is two less than valence shell • two less than the Period number they really belong to • sometimes d electron in configuration Eu Z = 63, Period 6 [Xe]6s24f 7

  38. Electron Configuration fromthe Periodic Table 8A 1A 1 2 3 4 5 6 7 3A 4A 5A 6A 7A 2A 3d10 Ar As 4s2 4p3 As = [Ar]4s23d104p3 As has 5 valence electrons Tro, Chemistry: A Molecular Approach

  39. Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each of the following • Na (at. no. 11) • Te (at. no. 52) • Tc (at. no. 43) Tro, Chemistry: A Molecular Approach

  40. Practice – Use the Periodic Table to write the short electron configuration and orbital diagram for each of the following • Na (at. no. 11) [Ne]3s1 • Te (at. no. 52) [Kr]5s24d105p4 • Tc (at. no. 43) [Kr]5s24d5 3s 5s 5p 4d 5s 4d Tro, Chemistry: A Molecular Approach

  41. Properties & Electron Configuration • elements in the same column have similar chemical and physical properties because they have the same number of valence electrons in the same kinds of orbitals Tro, Chemistry: A Molecular Approach

  42. Electron Configuration & Element Properties • the number of valence electrons largely determines the behavior of an element • chemical and some physical • since the number of valence electrons follows a Periodic pattern, the properties of the elements should also be periodic • quantum mechanical calculations show that 8 valence electrons should result in a very unreactive atom, an atom that is very stable – and the noble gases, that have 8 valence electrons are all very stable and unreactive • conversely, elements that have either one more or one less electron should be very reactive – and the halogens are the most reactive nonmetals and alkali metals the most reactive metals • as a group Tro, Chemistry: A Molecular Approach

  43. Electron Configuration &Ion Charge • we have seen that many metals and nonmetals form one ion, and that the charge on that ion is predictable based on its position on the Periodic Table • Group 1A = +1, Group 2A = +2, Group 7A = -1, Group 6A = -2, etc. • these atoms form ions that will result in an electron configuration that is the same as the nearest noble gas Tro, Chemistry: A Molecular Approach

  44. Tro, Chemistry: A Molecular Approach

  45. Electron Configuration of Anions in their Ground State • anions are formed when atoms gain enough electrons to have 8 valence electrons • filling the s and p sublevels of the valence shell • the sulfur atom has 6 valence electrons S atom = 1s22s22p63s23p4 • in order to have 8 valence electrons, it must gain 2 more S2- anion = 1s22s22p63s23p6 Tro, Chemistry: A Molecular Approach

  46. Electron Configuration of Cations in their Ground State • cations are formed when an atom loses all its valence electrons • resulting in a new lower energy level valence shell • however the process is always endothermic • the magnesium atom has 2 valence electrons Mg atom = 1s22s22p63s2 • when it forms a cation, it loses its valence electrons Mg2+ cation = 1s22s22p6 Tro, Chemistry: A Molecular Approach

  47. Trend in Atomic Radius – Main Group • Different methods for measuring the radius of an atom, and they give slightly different trends • van der Waals radius = nonbonding • covalent radius = bonding radius • atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds • Atomic Radius Increases down group • valence shell farther from nucleus • effective nuclear charge fairly close • Atomic Radius Decreases across period (left to right) • adding electrons to same valence shell • effective nuclear charge increases • valence shell held closer Tro, Chemistry: A Molecular Approach

  48. Effective Nuclear Charge • in a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other • outer electrons are shieldedfrom full strength of nucleus • screening effect • effective nuclear charge is net positive charge that is attracting a particular electron • Z is nuclear charge, S is electrons in lower energy levels • electrons in same energy level contribute to screening, but very little • effective nuclear charge on sublevels trend,s > p > d > f Zeffective= Z - S Tro, Chemistry: A Molecular Approach

  49. Screening & Effective Nuclear Charge

  50. Trends in Atomic RadiusTransition Metals • increase in size down the Group • atomic radii of transition metals roughly the same size across the d block • must less difference than across main group elements • valence shell ns2, not the d electrons • effective nuclear charge on the ns2 electrons approximately the same Tro, Chemistry: A Molecular Approach

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