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Electron Configuration and the Periodic Table

Electron Configuration and the Periodic Table. Mallard Creek Chemistry - Rines. Electromagnetic Radiation. Wave Nature of Light. Property of Waves Frequency No. of waves per second Wave Length Distance between corresponding points in a wave Amplitude Size of the wave peak.

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Electron Configuration and the Periodic Table

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  1. Electron Configuration and the Periodic Table Mallard Creek Chemistry - Rines

  2. Electromagnetic Radiation • Wave Nature of Light • Property of Waves • Frequency • No. of waves per second • Wave Length • Distance between corresponding points in a wave • Amplitude • Size of the wave peak

  3. Electromagnetic Radiation • Mathematical Relations • C = speed of light = 3.0 x 108 m/s • λ (lamda) = wavelength (m) • f= frequency (Hz or s-1) • This is how we know what color light is emitted! C = λ f

  4. FrequencyisinverselyproportionaltoWavelength • If λincreasesfdecreases • Iffincreases λ decreases • Speed of the wave is always constant at 3.0 x 108m/s

  5. Bohr Model • Nucleus: Neutrons and Protons • Orbits: Electrons • We know both specific energy and location of each electron • Electrons orbit the nucleus in certain fixed energy levels (or shells) Energy Levels Nucleus

  6. Bohr Model • Bohr’s Atomic Model of Hydrogen • Bohr - electrons exist in energy levels AND defined orbits around the nucleus. • Each orbit corresponds to a different energy level. • The further out the orbit, the higher the energy level

  7. Bohr’s Model • The Photoelectric Effect • Light releases electrons • Not all colors work • Atomic Emission Spectra • Hydrogen gas emitted specific bands of light • Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 6 5 4 3 2 1

  8. Electromagnetic Radiation • Photoelectric Effect – There is a minimum frequency to eject the electron

  9. Electromagnetic Radiation • Photoelectric Effect • Only explained by “energy packets” of light called a quantum • Quantum- minimum amount of energy that can be gained or lost by an atom • Photonsare massless particles of light of a certainquantumof energy • Based on the frequency and wavelength of the photon

  10. Bohr’s Model • Excited electrons • Energy added to atom – electrons “jump” up energy levels • When the atom relaxes - electron “falls” to lower energy levels and emits photon • Bohr Model of hydrogen • Reference Sheets!!!!!

  11. Electromagnetic Radiation • Atomic Line Spectra • Electrons in an atom add energy to go to an “excited state”. • When they relax back to the ground state, they emit energy in specific energy quanta

  12. 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ hv hv Electromagnetic Radiation • These observations suggested that electrons must exist in defined energy levels First, the electron absorbs energy and jumps from the ground state to an excited state Next, the excited electron relaxes to a lower excited state or ground state

  13. Electromagnetic Radiation • Particle Nature of Light • Wave nature could not explain all observations (Plank & Einstein) Photoelectric Effect When light strikes a metal electrons are ejected Atomic Line Spectra • When elements are heated, they emit a unique set of frequencies of visible and non-visible light. E = hf

  14. Other Scientists Contributions • De Broglie • Heisenburg • Modeled electrons as waves • Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron • Electrons exist in orbital’sof probability • Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

  15. Other Scientists Contributions • Schrödinger • Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom • Quantum Mechanical Model of the atom – current model of the atom treating electrons as waves.

  16. Quantum Mechanical Model • Nucleus: Neutrons and protons • Orbitals: region in space surrounding the nucleus where there is a 95% probability of finding an electron. • We know either energy or location of each electron.

  17. Solutions to the Wave Equation • Quantum Numbers • Wave Equation generates 4 variable solutions • n - size • l - shape • m - orientation • s – spin • Address of an electron

  18. Quantum Numbers • n – Primary Quantum Number • Describes the size and energy of the orbital • n is any positive # • n = 1,2,3,4,…. • Found on the periodic table • Like the “state” you live in

  19. Quantum Numbers • l – Orbital Quantum Number • Sub-level of energy • Describes the shape of the orbital • l = 0,1,2,3,4,….(n-1) • “City” you live in n = 3 l = 0,1,2 n = 2 l = 0,1 n = 1 l = 0

  20. Quantum Numbers • l – Orbital Quantum Number • # level = # sublevels • 1st level – 1 sublevel • 2nd level – 2 sublevels • 4th level = 4 sublevels

  21. s p d f Quantum Numbers Sublevels are named for their shape • s • l = 0 • Spherical in shape • p • l = 1 • Dumbbell in shape • d • l = 2 • f • l = 3

  22. Quantum Numbers • m – Magnetic Quantum Number • Describes the orientation of the orbital in space • Also denotes how many orbital's are in each sublevel • For each sublevel there are2l +1 orbital's • “Street” you live on

  23. Quantum Numbers • Look at Orbital's as Quantum Numbers l = 1 m = -1, 0, +1 For each p sublevel there are 3 possible orientations, so three 3 orbital's l = 0 m = 0 Can only be one s orbital

  24. Orbital Designations

  25. Orbital Rules

  26. Reflection • How is the Bohr model different from the earlier models of the atom? • Who contributed to the modern model of the atom? How is it different from Bohr’s? • Why do atoms give unique atomic line spectra? • What are ground and excited states? • Is 2d possible? 4f ? 2s ? 6p? 1p? • How many total orbital's in the 2nd level? 4th level.

  27. Aufbau Principle • Aufbau Principal • Lowest energy orbital available fills first • “Lazy Tenant Rule”

  28. Pauli’s Exclusion Principle • No two electrons have the same quantum #’s • Maximum electrons in any orbital is two () • Pauli Exclusion Principle

  29. Hund’s Rule • When filling degenerateorbital's, electrons will fill an empty orbital before pairing up with another electron. • Empty room rule • Hund’s Rule RIGHT WRONG

  30. Periodic Table & Electron Configuration

  31. Periodic Table & Electron Configuration Using the periodic table for the filling order of orbitals, by going in atomic number sequence until you use all the needed electrons in the element

  32. Level (n) Increasing Energy Orbitals (m) Orbital Energy Diagram Sub-level (l) d ______ ______ ______ ______ ______ p ______ ______ ______ 3 s ______ p ______ ______ ______ 2 s ______ 1 s ______ An energy diagram for the first 3 main energy levels

  33. 2px2 2py2 2pz2 Electron Spin Increasing Energy Orbital Energy Diagram and Electron Configuration p ______ ______ ______ 3 s ______ p ______ ______ ______ 2 s ______ 1 s ______ 1s2 2s2 1s2 2s2 2p6 Electron Configuration Notation An energy diagram for Neon

  34. Orbital Notation • Orbital Notation shows each orbital • O(atomic number 8) ____ ____ ____ ____ ____ ____ 1s2s2px2py2pz3s • 1s22s22p4electron configuration!

  35. Orbital Notation • Orbital Notation shows each orbital • O(atomic number 8) ____ ____ ____ ____ ____ ____ 1s2s2px2py2pz3s • !

  36. Orbital Notation • Write the orbital notation for S • S(atomic number 16) ___ __ __ __ __ __ __ __ __ 1s2s2p3s 3p • 1s22s22p63s23p4 • How many unpaired electrons does sulfur have? 2 unpaired electrons!

  37. Valence Electrons • Valence Electrons • As (atomic number 33) • 1s22s22p63s23p64s23d104p3 • The electrons in the outermost energy level. • s and p electrons in last shell • 5 valence electrons

  38. Core Electrons Valence Electrons Valence Electrons • Longhand Configuration 2p6 S 16e- 2s2 1s2 3s2 3p4 • Shorthand Configuration S 16e- [Ne]3s2 3p4

  39. Noble Gas Configuration • Example - Germanium X X X X X X X X X X X X X [Ar] 4s2 3d10 4p2

  40. Electron Configuration Let’s Practice • P (atomic number 15) • 1s22s22p63s23p3 • Ca (atomic number 20) • 1s22s22p63s23p64s2 • As (atomic number 33) • 1s22s22p63s23p64s23d104p3 • W (atomic number 74) • 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 Noble Gas Configuration [Ne]3s23p3 [Ar]4s2 [Ar]4s23d104p3 [Xe]6s24f145d4

  41. Electron Configuration Your Turn • N (atomic number 7) • 1s22s22p3 • Na (atomic number 11) • 1s22s22p63s1 • Sb(atomic number 51) • 1s22s22p63s23p64s23d104p65s24d105p3 • Cr (atomic number 24) • 1s22s22p63s23p64s23d4 Noble Gas Configuration [He]2s22p3 [Ne]3s1 [Kr]5s24d105p3 [Ar]4s23d4

  42. Stability • Full energy level • Full sublevel • Half full sublevel

  43. Exceptions • Copper • Expect: [Ar] 4s2 3d9 • Actual: [Ar] 4s1 3d10 • Silver • Expect:[Kr] 5s2 4d9 • Actual: [Kr] 5s1 4d10 • Chromium • Expect:[Ar] 4s2 3d4 • Actual: [Ar] 4s1 3d5 • Molybdenum • Expect: [Kr] 5s2 4d4 • Actual: [Kr] 5s1 4d5 Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel

  44. Stability • Atoms create stability by losing, gaining or sharing electrons to obtain a full octet • Isoelectronic with noble gases 0 +1 +2 +4 -2 +3 -3 -1 Atoms take electron configuration of the closest noble gas

  45. Stability • Na (atomic number 11) • 1s22s22p63s1 • 1s22s22p6= [Ne] 1 Valence electron Metal = Loses Ne Na

  46. Try Some • P-3(atomic number 15) • 1s22s22p63s23p6 • Ca+2(atomic number 20) • 1s22s22p63s23p6 • Zn+2(atomic number 30) • 1s22s22p63s23p63d10 • Lost valence electrons (s and p) Full Octet

  47. Lewis Structures • Shows valence electrons only! • s & p electrons • Write noble gas configuration for the element • Place valence electrons around element symbol in order 6 3 X s electrons 4 1 p electrons 2 7 5 8

  48. Try Some • Write the Lewis structures for: • Oxygen (O) • [He] 2s2 2p4 • Iron (Fe) • [Ar] 4s23d6 • Bromine (Br) • [Ar] 4s2 3d104p5 • • • • O • • Valence electrons • Fe • • • Br • • • • •

  49. What Do I Need to Know? • How the periodic table is arranged • Be able to identify subcategories of the periodic table • How the elements within a group are similar • How the elements within a period are similar • Be able to compare and contrast the electronegativities, ionization energies, and radii of metals and non-metals

  50. Periodic Table Dmitri Mendeleev – Father of the Periodic Table Some Problems He left blank spaces for what he said were undiscovered elements (he was right!) He broke the pattern of increasing atomic weight to keep similar reacting elements together What He Did • Put elements in rows by increasing atomic weight • Put elements in columns by similar properties

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