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Many electron atoms and the Periodic Table

Many electron atoms and the Periodic Table. Objectives. Explain the scientific basis for the Periodic Table Apply the Aufbau principle, the Pauli Exclusion principle and Hund’s rule to electrons in an atom Explain the concept of energy levels in an atom and the order of filling these levels

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Many electron atoms and the Periodic Table

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  1. Many electron atoms and the Periodic Table

  2. Objectives • Explain the scientific basis for the Periodic Table • Apply the Aufbau principle, the Pauli Exclusion principle and Hund’s rule to electrons in an atom • Explain the concept of energy levels in an atom and the order of filling these levels • Write the electronic configuration of the first 20 elements • Draw and explain a block diagram of the Periodic Table • Explain the meaning and position of the transition elements • Explain the periodic variations of atomic size, ionisation energy, electron affinity and electronegativity

  3. The Periodic Law • The physical and chemical properties of the elements are a function of the electronic configuration of the their atoms which vary with increasing atomic number in a periodic manner • hence the Periodic Table • elements are grouped according to their electronic structure • provides information on chemical properties

  4. Many Electron Atoms • Describe in terms of hydrogen orbitals • Same quantum numbers and shapes • energies different • For Hydrogen: • Depend on n 3s, 3p, 3d all the same energy • For many e- atom: • Different subshells at different energies 2s < 2p depend on n and l

  5. Many e- Atoms Hydrogen 0 0 n=4 3d 4s 3p n=3 n=3 3s 3s 3p 3d Energy n=2 n=2 2p 2s 2s 2p n=1 n=1 1s 1s Orbital Energies Energy Need to consider Effect of increased nuclear charge Repulsions between electrons

  6. Effective Nuclear Charge • Net positive charge from nucleus attracting an electron • electron shielded by inner e- • effective nuclear chargeZeff = Z - SS = No. of e- between atom and nucleus • hence outer shell e- experience less + charge • effect of “screening” depends on e- distribution • need to consider orbital shape Assume 5 Z e- of interest 3e- found in sphere Eff. Nuclear charge = 5 - 3 = 2

  7. Orbital shape and Energy • s-orbital e- can be close to nucleus • p-orbital further away than an “s” e- • d-orbital further from nucleus than “p” e- • therefore, • “s” e- has least screening by other e- • so a larger effective nuclear charge and is more tightly bound(lower energy than “p” ie. more stable)

  8. Order of Energy Levels for many e- Atom In general, 1s < 2s < 2p < 3s < 3p…...

  9. An easy way to remember…... 1s Increasing Energy 2s 2p 3s 3p 3d 4s 4p 4d 4f 5p 5s 5d 5f 6s 6p 6d

  10. Energy

  11. Electron Spin • need to consider another property of electrons to determine how electrons populate orbitals • envisage electron as spinning on own axis • quantized • only 2 spin states • distinguished by the spin -magnetic quantum number ms 1 1 + - 2 2

  12. REMIND! Electron Spin • Stern-Gerlach experiment - when a beam of ground state H atoms (1s) is passed through a magnetic field, the beam splits into two beams

  13. Pauli and Hund Pauli Exclusion Principle • no two electrons in an atom can have the same 4 quantum numbers • there are only 2 values of ms • hence, an orbital can only hold 2 electrons and they must have opposite spins Hund’s Rule • If orbitals have the same energy, add electrons singly with spins parallel first

  14. The Aufbau Principle Building up • fill available orbitals with available electrons starting with lowest energy orbitals (most stable) • this gives ground state Note • don’t forget Pauli and Hund!!

  15. Building up atoms

  16. Aufbau Principle • The aufbau principle shows how orbitals are filled: in the order to the left. Two extra rules are needed.

  17. Aufbau Principle-2 • Hund’s rule states that when filling a set of orbitals at the same energy (sunshell), one electron is placed in each orbital before pairing occurs. • Pauli’s principle tells us that when placing a second electron in an orbital, its spin must be opposite to the electron already in the orbital. (Spin is usually represented by an arrow in an orbital box.)

  18. Aufbau Principle-3a • “Box” or “orbital diagrams for electron configurations. Start at 1s The 1s orbital is filled. Second electron is paired. (Pauli)

  19. Aufbau Principle-3b 1s2, 2s1 1s2, 2s2,2p2 Hund’s rule applies to p subshell

  20. Aufbau Principle-3c 1s2, 2s2,2p4 • In 2p subshell, Hund’s rule! Next electron follows Pauli principle

  21. Electrons and the Periodic Table • Electron fit logically into the periodic table. The s block elements (see next slide) start filling at level 1, the p block at level 2, and the d block at level 3.

  22. Excited State Atoms • Occur when energy has been supplied to raise e- energy Ne 1s22s22p6 Ground state 1s22s22p5….5s1 High energy excited state 1s22s22p5….3p1 Low energy excited state

  23. Transition metals Consider the elements in the 4th period(after Ar 1s22s22p63s23p6) • after 3p natural sequence would be 3d • but 4s has (slightly) lower energy than3d • according to Aufbau must fill 4s before 3d • ground state for K is [Ar]4s1 and Ca is [Ar]4s2 • as charge increases the energy of 3d decreases • in Sc 3d < 4s • Sc+ [Ar]3d14s1 Sc2+[Ar]3d1

  24. Transition metals • energy drop in 3d continues through to Zn • consequences • for elements Cu  Zn oxidation state 1+ or 2+ • beyond Zn 3d electrons have no chemical role • elements from Sc to Zn called d-block elements • filling up the d orbital

  25. Transition metals • anomalies • Cr [Ar]3d54s1 expect [Ar]3d44s2 • Cu [Ar]3d104s1 expect [Ar]3d94s2 • similar occur in fifth period • also Ru [Kr]4d75s1 expect [Kr]4d65s2 • sixth period filling is erratic • energies of 4f, 5d & 6s comparable • seventh period all are radioactive

  26. The Periodic Table Organisation of the elements • electronic configurations related to position of element • elements grouped according to type of orbital the outer shell electrons are in BLOCK: Named for last subshell occupied GROUP: the columns • all elements have same outer orbital e- configuration • similar chemical properties PERIODS: rows • all elements same shell

  27. The Periodic Table • Subshell orbitals with same energy eg. 2p • Shellorbitals with similar energy eg. 2s, 2p • Valence Electronsoccupy outermost shell • Core Electronsoccupy filled inner shells Cl 1s22s22p6 3s23p5 Ne core valence • Closed Shell Atomsfull outer shell - very stable - noble gases

  28. Periodic Properties Predicted by considering e- configurations • Sizes of atoms and ions • Ionisation energies • Electron affinities • Electronegativities • Polarising powers and polarisabilities

  29. Sizes Of Atoms and Ions Atoms do not have sharply defined boundaries • Hence, need to define atomic size • Atomic size depends on chemical environment • ie. Bonding etc

  30. This shielding means that each valence electron in effect only “feels” a +1 charge form the nucleus; this occurs for an highly excited valence electron. Otherwise the shielding makes the “seen” charge is higher than +1

  31. 2r Defining Atomic and Ionic Size estimating size atomic radius = half thedistance between nearest atomsin element (in condensed phases) for ions, base ionic radii on interatomicdistance in ionic crystals. (depends on charge...) Cu Cu Cu+ Cu+2 atom covalent bonding 1.28 1.17 0.96 0.69 Ao

  32. Decrease Increase Sizes of Atoms and Ions Why? Consider: 1. Principle Quantum number (shell) 2. Effective nuclear charge

  33. Across a Period……. Increase nuclear charge but no. of core electrons stay the same so effective nuclear charge increases while shell remains the same hence electrons drawn closer to nucleus hence decrease atomic size eg. Na 1s22s22p63s1 1.91Ao Mg 1s22s22p63s2 1.60 Ao

  34. Down A Group…. • More distant electron shell occupied while effective nuclear charge the same • hence atomic size increases eg. Li 1s22s1 1.57 A0 Na 1s22s22p63s1 1.91 Ao

  35. Atomic radii

  36. Radius of Ions Cation < Atom eg. Na+ < Na0.96 Ao 1.91 Ao 1s22s22p6 1s22s22p63s1  lost an e- core electrons exposed more tightly bound Decreases across a period eg. Na+ > Mg2+

  37. Radius of Ions Atom < Anion eg. Cl < Cl0.99 Ao 1.81 Ao [Ne]3s23p5 [Ne]3s23p6  gained an e- electron cloud greater decreases nuclear pull by each electron Decreases across a period e.g. S2- > Cl-

  38. Ionisation Energy The energy required to remove an electron from a ground state atom X(g) X+(g)+ e-E = IE1 Measure of stability of outer shell electron configuration • Depends on • size of the atom • effective nuclear charge • screening effect ofinner electrons • type of electron

  39. Increase Decrease Ionisation Energy Why? Consider 1. Effective nuclear charge 2. Distance of e- from nucleus

  40. Across a Period…. Increase in effective nuclear charge Decrease in radius hence increase attraction between e- and nucleus hence increase IE Exceptions: “p” less stable than “s” (B < Be) orbitals “singly occupied” more stable than “doubly occupied” (O < N)

  41. Down a Group…... Increase radius while effective nuclear charge the same hence Decrease attraction between e- and nucleus hence decrease IE

  42. Ionization energy

  43. Electron Affinity The energy released when an e- added to atom to form anion eg. F(g) + e- F-(g) EA = 328 kJ/mol • a small EA means e- must be forced to stick • measure of ability of atom to accept e-

  44. Increase Decrease Low for Noble gases Electron Affinities Same as IE Why? Consider 1. Size 2. Effective Nuclear charge

  45. Electron Affinities 1. F + e- F- EA = 328.0 kJ mol-1 1s2 2s2 2p6 - stable closed shell = Ne 2. Ne + e- Ne- EA = negative 1s2 2s2 2p6 3s1 - new shell, further from nucleus - almost totally screened from nuclear charge - so unstable

  46. Electron Affinities Be 1s22s2 Low EA Filled s subshell Next e- higher energy level so need energy to add e- N 1s22s22p3 Low EA Half filled “p” Adding another e- will cause e- repulsion hence unfavourable

  47. Electronegativity The ability of an atom to draw e- to itself in a chemical bond useful for predicting extent of chargetransfer between atoms eg. “Covalent”  “Ionic” H—H C—H N—H NaCl

  48. Cs and F IE1 EA Cs low small F high large (Cs gives up e- easily, while F accepts e- easily.) Electron donor Electropositive Electron acceptor Electronegative Electronegativity Related to EA and IE

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