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Chapter 8 Periodic Properties of the Elements

Chemistry: A Molecular Approach , 2nd Ed. Nivaldo Tro. Chapter 8 Periodic Properties of the Elements. Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA. Nerve Transmission. Movement of ions across cell membranes is the basis for the transmission of nerve signals

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Chapter 8 Periodic Properties of the Elements

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  1. Chemistry: A Molecular Approach, 2nd Ed.Nivaldo Tro Chapter 8Periodic Properties of the Elements Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA

  2. Nerve Transmission • Movement of ions across cell membranes is the basis for the transmission of nerve signals • Na+ and K+ ions are pumped across membranes in opposite directions through ion channels • Na+ out and K+ in • The ion channels can differentiate Na+ from K+ by their difference in size • Ion size and other properties of atoms are periodic properties – properties whose values can be predicted based on the element’s position on the Periodic Table Tro: Chemistry: A Molecular Approach, 2/e

  3. Mendeleev (1834–1907) • 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, 2/e

  4. Periodic Pattern NM H2O a/b 1.0 H2 M Li2O b M BeO a/b NM B2O3 a NM CO2 a NM N2O5 a O2 NM NM 6.9 LiH 9.0 BeH2 10.8 BH3 12.0 CH4 14.0 NH3 16.0 H2O 19.0 HF Al2O3 a/b M Na2O b M MgO b M NM P4O10 a NM SO3 a NM Cl2O7 a SiO2 a M/NM 23.0 NaH 24.3 MgH2 AlH3 27.0 PH3 H2S HCl SiH4 28.1 31.0 32.1 35.5 M K2O b M CaO b 39.1 KH 40.1 CaH2 M = metal, NM = nonmetal, M/NM = metalloid a = acidic oxide, b = basic oxide, a/b = amphoteric oxide Tro: Chemistry: A Molecular Approach, 2/e

  5. Mendeleev's Predictions Tro: Chemistry: A Molecular Approach, 2/e

  6. 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 • and knowing Why allows us to predict What Tro: Chemistry: A Molecular Approach, 2/e

  7. Electron Configurations number of electrons in the orbital 1s1 principal energy level of orbital occupied by the electron sublevel of orbital occupied by the electron 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 Tro: Chemistry: A Molecular Approach, 2/e

  8. How Electrons Occupy Orbitals • Calculations with Schrödinger’s equation show hydrogen’s one electron occupies the lowest energy orbital in the atom • Schrödinger’s equation calculations for multielectron atoms cannot be exactly solved • due to additional terms added for electron-electron interactions • Approximate solutions show the orbitals to be hydrogen-like • Two additional concepts affect multielectron atoms: electron spin and energy splitting of sublevels Tro: Chemistry: A Molecular Approach, 2/e

  9. 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, 2/e

  10. Electron Spin Experiment Tro: Chemistry: A Molecular Approach, 2/e

  11. 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 Tro: Chemistry: A Molecular Approach, 2/e

  12. Spin Quantum Number, 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 Tro: Chemistry: A Molecular Approach, 2/e

  13. unoccupied orbital orbital with one electron orbital with two 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, 2/e

  14. 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 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, 2/e

  15. Q u a n t u m V a l u e s N u m b e r S i g n i f i c a n c e N u m b e r o f V a l u e s s i z e a n d P r i n c i p a l , n 1 , 2 , 3 , . . . - e n e r g y o f t h e o r b i t a l s h a p e o f A z i m u t h a l , l 0 , 1 , 2 , . . . , n - n 1 o r b i t a l o r i e n t a t i o n o f M a g n e t i c , - l , . . . , 0 , . . . + l 2 l + 1 o r b i t a l m l d i r e c t i o n o f S p i n , m - _ , + _ 2 s e l e c t r o n s p i n Allowed Quantum Numbers Tro: Chemistry: A Molecular Approach, 2/e

  16. 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 • Because they are in the same orbital, they must have opposite spins Tro: Chemistry: A Molecular Approach, 2/e 16

  17. Sublevel Splitting in Multielectron Atoms • The sublevels in each principal energy shell of Hydrogen all have the same energy • or other single electron systems • We call orbitals with the same energy degenerate • For multielectron atoms, the energies of the sublevels are split • caused by charge interaction, shielding and penetration • 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, 2/e

  18. Coulomb’s Law • Coulomb’s Law describes the attractions and repulsions between charged particles • For like charges, the potential energy (E) is positive and decreases as the particles get farther apart • as r increases • For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together • The strength of the interaction increases as the size of the charges increases • electrons are more strongly attracted to a nucleus with a 2+ charge than a nucleus with a 1+ charge Tro: Chemistry: A Molecular Approach, 2/e

  19. Shielding & Effective Nuclear Charge Each electron in a multielectron 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 Tro: Chemistry: A Molecular Approach, 2/e

  20. Penetration • 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 • The degree of penetration is related to the orbital’s radial distribution function • in particular, the distance the maxima of the function are from the nucleus Tro: Chemistry: A Molecular Approach, 2/e

  21. Shielding & Penetration Tro: Chemistry: A Molecular Approach, 2/e

  22. Penetration 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 Tro: Chemistry: A Molecular Approach, 2/e

  23. Effect of Penetration and Shielding • Penetration causes the energies of sublevels in the same principal level to not be degenerate • 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 Tro: Chemistry: A Molecular Approach, 2/e

  24. 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 (and can cause anomalous electron configurations for certain elements) Tro: Chemistry: A Molecular Approach, 2/e

  25. Filling the Orbitals with Electrons • Energy levels and sublevels fill from lowest energy to high • s→p→d→f • Aufbau Principle • 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 Tro: Chemistry: A Molecular Approach, 2/e

  26. Electron Configuration of Atoms in their Ground State 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 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 Tro: Chemistry: A Molecular Approach, 2/e

  27. Order of Sublevel Fillingin 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 Tro: Chemistry: A Molecular Approach, 2/e

  28. Electron Configurations Tro: Chemistry: A Molecular Approach, 2/e

  29. Example: Write the full ground state orbital diagram and electron configuration of manganese s sublevel holds 2 e− p sublevel holds 6 e− d sublevel holds 10 e− 1s 2s 2p 3s 3p 4s f sublevel holds 14 e− 3 d Mn Z = 25, therefore 25 e− 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f                  2 e− +2 = 4e− +6 +2 = 12e−      +6 +2 = 20e− +10 = 30e− Therefore the electron configuration is 1s22s22p63s23p64s23d5 Based on the order of sublevel filling, we will need the first seven sublevels Tro: Chemistry: A Molecular Approach, 2/e

  30. Practice — write the full ground state orbital diagram and electron configuration of potassium. Tro: Chemistry: A Molecular Approach, 2/e

  31. Practice — write the full ground state orbital diagram and electron configuration of potassium, answer s sublevel holds 2 e− p sublevel holds 6 e− d sublevel holds 10 e− 1s 2s 2p 3s 3p 4s f sublevel holds 14 e− K Z = 19, therefore 19 e− 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f                 2 e− +2 = 4e− +6 +2 = 12e− Therefore the electron configuration is 1s22s22p63s23p64s1 +6 +2 = 20e− Based on the order of sublevel filling, we will need the first six sublevels Tro: Chemistry: A Molecular Approach, 2/e

  32. 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 number of valence electrons Tro: Chemistry: A Molecular Approach, 2/e

  33. 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) Tro: Chemistry: A Molecular Approach, 2/e

  34. 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, 2/e

  35. Tro: Chemistry: A Molecular Approach, 2/e

  36. s1 s2 p1 p2p3p4p5 p6 s2 1 2 3 4 5 6 7 d1 d2d3d4d5d6d7d8d9d10 f2f3f4f5f6f7f8f9 f10f11f12f13f14 f14d1 Tro: Chemistry: A Molecular Approach, 2/e

  37. 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 five valence electrons Tro: Chemistry: A Molecular Approach, 2/e

  38. 4s 3d 6s 4f Transition Elements 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 • 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 Tro: Chemistry: A Molecular Approach, 2/e

  39. 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 five valence electrons Tro: Chemistry: A Molecular Approach, 2/e

  40. Practice – Use the Periodic Table to write the short electron configuration and short orbital diagram for each of the following • Na (at. no. 11) • Te (at. no. 52) • Tc (at. no. 43) [Ne]3s1 3s [Kr]5s24d105p4 5s 5p 4d [Kr]5s24d5 5s 4d Tro: Chemistry: A Molecular Approach, 2/e

  41. Irregular Electron Configurations We know that because of sublevel splitting, the 4s sublevel is lower in energy than the 3d; and therefore the 4s fills before the 3d But the difference in energy is not large Some of the transition metals have irregular electron configurations in which the ns only partially fills before the (n−1)d or doesn’t fill at all Therefore, their electron configuration must be found experimentally Tro: Chemistry: A Molecular Approach, 2/e

  42. Irregular Electron Configurations • 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 Tro: Chemistry: A Molecular Approach, 2/e

  43. Properties & Electron Configuration • The properties of the elements follow a periodic pattern • elements in the same column have similar properties • the elements in a period show a pattern that repeats • The quantum-mechanical model explains this because the number of valence electrons and the types of orbitals they occupy are also periodic Tro: Chemistry: A Molecular Approach, 2/e

  44. The Noble Gas Electron 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 Tro: Chemistry: A Molecular Approach, 2/e

  45. Everyone Wants to Be Like a Noble Gas! The Alkali Metals The alkali metals have one more electron than the previous noble gas In their reactions, the alkali metals tend to lose one electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge Tro: Chemistry: A Molecular Approach, 2/e

  46. Everyone Wants to Be Like a Noble Gas!The Halogens The electron configurations of the halogens all have one fewer electron than the next noble gas In their reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas Forming an anion with charge 1− In their reactions with nonmetals, they tend to share electrons with the other nonmetal so that each attains the electron configuration of a noble gas Tro: Chemistry: A Molecular Approach, 2/e

  47. Eight Valence Electrons • Quantum mechanical calculations show that eight valence electrons should result in a very unreactive atom • an atom that is very stable • the noble gases have eight valence electrons and are all very stable and unreactive • He has two valence electrons, but that fills its valence shell • Conversely, elements that have either one more or one less electron should be very reactive • the halogen atoms have seven valence electrons and are the most reactive nonmetals • the alkali metals have one more electron than a noble gas atom and are the most reactive metals • as a group Tro: Chemistry: A Molecular Approach, 2/e

  48. 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, 2/e

  49. Tro: Chemistry: A Molecular Approach, 2/e

  50. Electron Configuration of Anions in Their Ground State • Anions are formed when nonmetal atoms gain enough electrons to have eight valence electrons • filling the s and p sublevels of the valence shell • The sulfur atom has six valence electrons S atom = 1s22s22p63s23p4 • To have eight valence electrons, sulfur must gain two more S2− anion = 1s22s22p63s23p6 Tro: Chemistry: A Molecular Approach, 2/e

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