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# Electron arrangements

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### Electron arrangements

1. 1. – Electrons andPeriodic Behavior Cartoon courtesy of NearingZero.net
2. 2. Wave-Particle DualityJJ Thomson won the Nobel prize for describing theelectron as a particle.His son, George Thomson won the Nobel prize fordescribing the wave-like nature of the electron. The electron is The a particle! electron is an energy wave!
3. 3. The Wave-like Electron The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.Louis deBroglie
4. 4. Electromagnetic radiation propagatesthrough space as a wave moving at thespeed of light. c = νλC = speed of light, a constant (3.00 x 108 m/s) ν = frequency, in units of hertz (hz, sec-1) λ = wavelength, in meters
5. 5. Types of electromagnetic radiation:
6. 6. Max Planck• figured out that when a solid substance is heated, it gives off energy in "chunks“• later called quantums of energy • quantum means fixed amount• noticed that different substances released different "chunks" of energy
7. 7. The energy (E ) of electromagneticradiation is directly proportional to thefrequency (ν ) of the radiation. E = hν E = Energy, in units of Joules (kg·m2/s2) h = Planck’s constant (6.626 x 10-34 J·s) ν = frequency, in units of hertz (hz, sec-1)
8. 8. Long Wavelength = Wavelength TableLow Frequency = Low ENERGY Short Wavelength =High Frequency = High ENERGY
9. 9. Spectroscopic analysis of the visible spectrum……produces all of the colors in a continuous spectrum
10. 10. How does matter produce light?• When an electron in the ground state (lowest energy level that is natural) is promoted to an excited (higher) state it is temporary!!! The electron will fall back down to the ground state releasing light! (when viewed through a spectroscope- line emission spectra- gives a fingerprint of the atom)
11. 11. Spectroscopic analysis of the hydrogen spectrum……produces a “bright line” spectrum
12. 12. Electron transitionsinvolve jumps ofdefinite amounts ofenergy.This produces bandsof light with definitewavelengths.
13. 13. How does light behave as a particle? • Photoelectric effect- when light of a particular frequency hits the surface of a metal an electron is ejected off the surface!
14. 14. The Bohr Model of the Atom I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! WRONG!!! They’re moreNeils Bohr like bees around a hive.
15. 15. Things you must accept to do orbital notation (energy diagrams)• Energy builds further away from the nucleus• Each line represents an orbital• Each orbital can hold only two electrons• We as a group will decide to place positive spin arrows in first..this is arbitrary NOT A RULE Just so all our QNS are the same• Each electron is represented by an arrow• In an orbital the two electrons must point in different directions• Remember from the diagonal rule 4s fills before 3d breaking Aufbau’s Rule
16. 16. Rules for electron filling:• Aufbaus Rule- must fill the lowest energy level available first!• Hunds Rule -1 electron in each orbital of a sublevel before pairing begins
17. 17. Pauli Exclusion Principle Two electrons occupying the same orbital must have opposite spins- this eliminates 2 electrons within the same atom having theWolfgang same QNS Pauli
18. 18. Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not Werner where it is!Heisenberg
19. 19. How do electrons fill in an atom? The Diagonal Rule•
20. 20. Energy Levels, Sublevels, ElectronsEnergy Sublevels in Number of Number of Number ofLevel main energy orbitals per Electrons electrons per (n) level sublevel per sublevel main energy (n sublevels) level (2n2) 1 s 1 2 2 2 s 1 2 8 p 3 6 3 s 1 2 18 p 3 6 d 5 10 4 s 1 2 32 p 3 6 d 5 10 f 7 14
21. 21. Quantum Mechanical Model of the AtomMathematical laws can identify the regionsoutside of the nucleus where electrons are mostlikely to be found.These laws are beyond the scope ofthis class…VERY SIMPLY PUT… eachelectron has four numbers todescribe the probability of findingthe electron within the atom…
22. 22. Schrodinger Wave Equation h d ψ +V ψ 2 2 − = Eψ 8 π m dx 2 2 Equation for probability of a single electron being found along a single axis (x-axis)Erwin Schrodinger
23. 23. 4 Quantum Numbersn= Principle Quantum Number distance from the nucleus (Denotes Size) values 1-7 (note 7 periods on the P.T.)L = Sublevel shape of the cloud values 03 (0 = s, 1 = p, 2 = d, 3 = f)m = magnetic orientation about the axis values –3 3S= spinDirection of movement within orbital + ½ or – ½
24. 24. Electron Energy Level (Shell)Generally symbolized by n, it denotes theprobable distance of the electron from thenucleus.Number of electronsthat can fit in a shell: 2n 2
25. 25. An orbital is a region within an energy level wherethere is a probability of finding an electron. This is aprobability diagram for the s orbital in the firstenergy level…Orbital shapes are defined as the surface thatcontains 90% of the total electron probability.
26. 26. Sizes of s orbitalsOrbitals of the same shape (s, for instance) growlarger as n increases… Nodes are regions of low probability within an orbital.
27. 27. The s orbital has a spherical shape centered aroundthe origin of the three axes in space. s orbital shape
28. 28. P orbital shapeThere are three dumbbell-shaped p orbitals ineach energy level above n = 1, each assigned toits own axis (x, y and z) in space.
29. 29. Things get a bit more complicated with the five dd orbital shapes orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells” …and a “dumbell with a donut”!
30. 30. Shape of f orbitals
31. 31. Electron SpinElectron spin describes the behavior (direction ofspin) of an electron within a magnetic field. Possibilities for electron spin: 1 1 + − 2 2
32. 32. Electron Configuration Simpler way to write whichenergy levels and sublevels are filled within the atom
33. 33. Electron configuration of theelements of the first three series
34. 34. Element Configuration Orbital notation Noble gas notation notationLithium 1s22s1 [He]2s1 ____ ____ ____ ____ ____ 1s 2s 2pBeryllium 1s22s2 [He]2s2 ____ ____ ____ ____ ____ 1s 2s 2pBoron 1s22s2p1 [He]2s2p1 ____ ____ ____ ____ ____ 1s 2s 2pCarbon 1s22s2p2 [He]2s2p2 ____ ____ ____ ____ ____ 1s 2s 2pNitrogen 1s22s2p3 [He]2s2p3 ____ ____ ____ ____ ____ 1s 2s 2pOxygen 1s22s2p4 [He]2s2p4 ____ ____ ____ ____ ____ 1s 2s 2pFluorine 1s22s2p5 [He]2s2p5 ____ ____ ____ ____ ____ 1s 2s 2pNeon 1s22s2p6 [He]2s2p6 ____ ____ ____ ____ ____ 1s 2s 2p
35. 35. Orbital filling table
36. 36. Blocks of Elements
37. 37. Irregular conformations of Cr and Cu THE DEVIANT DsChromium steals a 4s electron tohalffill its 3d sublevel (more stable) Copper steals a 4s electron to FILL its 3d sublevel
38. 38. PeriodGroup or Family The Periodic Table Group or family Period
39. 39. Noble Gas Electron Configuration• Even shorter way to • Neon 1s22s22p6 write how electrons are • Aluminum 1s22s22p63s23p3 distributed in atom• Shows all electrons in shells common to a Noble • As you can see Al has the Gas as the symbol in a same distribution of inner bracket [symbol] then electrons as neon therefore it list leftover electrons in can be written as: their energy levels and shells • [Ne]3s23p3