1. 2.3 Electron Arrangement 2.3.1 Describe the electromagnetic spectrum 2.3.2 Distinguish between a continuous spectrum and a line spectrum 2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels 2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20

2. Bohr’s Model Why don’t the electrons fall into the nucleus?Why don’t the electrons fall into the nucleus? Move like planets around the sun.Move like planets around the sun. In circular orbits at different levels.In circular orbits at different levels. Amounts of energy separate one level from another.Amounts of energy separate one level from another.

3. Bohr postulated that: Fixed energy related to the orbitFixed energy related to the orbit Electrons cannot exist between orbitsElectrons cannot exist between orbits The higher the energy level, the further it is away from the nucleusThe higher the energy level, the further it is away from the nucleus An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive) Think of Noble gasesThink of Noble gases

4. Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low energy High energy Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light

5. Wavelength and frequency

6. How did he develop his theory? He used mathematics to explain the visible spectrum of hydrogen gasHe used mathematics to explain the visible spectrum of hydrogen gas Lines are associated with the fall of an excited electron back down to its ground state energy level.Lines are associated with the fall of an excited electron back down to its ground state energy level. http://www.mhhe.com/physsci/chemistry/essentialchemis try/flash/linesp16.swfhttp://www.mhhe.com/physsci/chemistry/essentialchemis try/flash/linesp16.swfhttp://www.mhhe.com/physsci/chemistry/essentialchemis try/flash/linesp16.swfhttp://www.mhhe.com/physsci/chemistry/essentialchemis try/flash/linesp16.swf

7. The line spectrum electricity passed through a gaseous element emits light at a certain wavelength Can be seen when passed through a prism Every gas has a unique pattern (color)

8. Line spectrum Continuous line spectrum

9. Those who are not shocked when they first come across quantum theory cannot possibly have understood it. (Niels Bohr on Quantum Physics)

10. Wavelengths and energy Understand that different wavelengths of electromagnetic radiation have different energies.Understand that different wavelengths of electromagnetic radiation have different energies. c=vλc=vλ –c=velocity of wave (2.998 x 10 8 m/s) –v=(nu) frequency of wave –λ=(lambda) wavelength

11. Bohr also postulated that an atom would not emit radiation while it was in one of its stable states but rather only when it made a transition between states.Bohr also postulated that an atom would not emit radiation while it was in one of its stable states but rather only when it made a transition between states. The frequency of the radiation emitted would be equal to the difference in energy between those states divided by Planck's constant.The frequency of the radiation emitted would be equal to the difference in energy between those states divided by Planck's constant.

12. E high -E low = hv = hc/ λ h=3.983 x 10 -13 Jsmol -1 = Plank’s constant E= energy of the emitted light (photon) v = frequency of the photon of light λ = is usually stated in nm, but for calculations use m. This results in a unique emission spectra for each element, like a fingerprint.This results in a unique emission spectra for each element, like a fingerprint. electron could "jump" from one allowed energy state to another by absorbing/emitting photons of radiant energy of certain specific frequencies.electron could "jump" from one allowed energy state to another by absorbing/emitting photons of radiant energy of certain specific frequencies.

13. Energy must then be absorbed in order to "jump" to another energy state, and similarly, energy must be emitted to "jump" to a lower state.Energy must then be absorbed in order to "jump" to another energy state, and similarly, energy must be emitted to "jump" to a lower state. The frequency, v, of this radiant energy corresponds exactly to the energy difference between the two states.The frequency, v, of this radiant energy corresponds exactly to the energy difference between the two states. In order for the emitted energy to be seen as light the wavelength of the energy must be in between 380 nm to 750 nmIn order for the emitted energy to be seen as light the wavelength of the energy must be in between 380 nm to 750 nm

14. For Hydrogen only! E n = -R/n 2, where R is -1312 kJ/mol and n is principle quantum number (energy level)E n = -R/n 2, where R is -1312 kJ/mol and n is principle quantum number (energy level) Example: Calculate the energy required to ionize a mole of electrons from the 4 th to the 2 nd energy level in a hydrogen atom?Example: Calculate the energy required to ionize a mole of electrons from the 4 th to the 2 nd energy level in a hydrogen atom? E 4 = -1312 / 4 2 = - 82 kJ E 4 = -1312 / 4 2 = - 82 kJ E 2 = -1312 / 2 2 = - 328 kJ E 2 = -1312 / 2 2 = - 328 kJ E 4 – E 2 = - 82 kJ – (- 328 kJ)= 246 kJ

15. What is the wavelength of light emitted when electrons go from n=4 to n=2 ? Is it visible to our eyes?What is the wavelength of light emitted when electrons go from n=4 to n=2 ? Is it visible to our eyes? E = hc/ λ, therefore λ = hc/E λ = [(3.983 x 10 -13 kJsmol -1 )(2.998 x 10 8 ms -1 )]/(246 kJmol -1 ) = 4.85 x 10 -7 m Convert to nm and see if its visible! (1 nm = 1 x 10 -9 m) (4.85 x 10 -7 m)( 1nm) = 485 nm (Its probably the green line) 1 x 10 -9 m

16. Bohr’s Triumph His theory helped to explain periodic law (the trends from the periodic table)His theory helped to explain periodic law (the trends from the periodic table) Halogens (gp.17) are so reactive because it has one e- less than a full outer orbitalHalogens (gp.17) are so reactive because it has one e- less than a full outer orbital Alkali metals (gp. 1) are also reactive because they have only one e- in outer orbitalAlkali metals (gp. 1) are also reactive because they have only one e- in outer orbital

17. Drawback Bohr’s theory did not explain or show the shape or the path traveled by the electrons. His theory could only explain hydrogen and not the more complex atoms

18. The Quantum Mechanical Model Energy is quantized. It comes in chunks.Energy is quantized. It comes in chunks. A quanta is the amount of energy needed to move from one energy level to another.A quanta is the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy.Since the energy of an atom is never “in between” there must be a quantum leap in energy. Schrödinger derived an equation that described the energy and position of the electrons in an atomSchrödinger derived an equation that described the energy and position of the electrons in an atom

19. Energy level populations Electrons found per energy level of the atom.Electrons found per energy level of the atom. The first energy level holds 2 electronsThe first energy level holds 2 electrons The second energy level holds 8 electrons (2 in s and 6 in p)The second energy level holds 8 electrons (2 in s and 6 in p) The third energy level holds 18 electrons (2 in s, 6 in p and 10 in d) There is overlapping here, so when we do the populations there will be some changes.The third energy level holds 18 electrons (2 in s, 6 in p and 10 in d) There is overlapping here, so when we do the populations there will be some changes. That is as far as this course requires us to go!

20. Examples for group 1 Li 2.1Li 2.1 Na 2.8.1Na 2.8.1 K 2.8.8.1K 2.8.8.1

21. A good site: http://www.chemguide.co.uk/basicorg/bo nding/orbitals.html http://www.chemguide.co.uk/basicorg/bo nding/orbitals.html http://www.chemguide.co.uk/basicorg/bo nding/orbitals.html

22. Electron Configuration HL only 12.1.3 State the relative energies of s, p, d, and f orbitals in a single energy level 12.1.4 State the maximum number of orbitals in a given energy level. 12.1.5 Draw the shape of an s orbital and the shapes of px, py and pz orbitals 12.1.6 Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=54.

23. S orbitals 1 s orbital for1 s orbital for every energy level 1s 2s 3s 1s 2s 3s Spherical shapedSpherical shaped Each s orbital can hold 2 electronsEach s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitalsCalled the 1s, 2s, 3s, etc.. orbitals

24. P orbitals Start at the second energy levelStart at the second energy level 3 different directions3 different directions 3 different shapes3 different shapes Each orbital can hold 2 electronsEach orbital can hold 2 electrons

25. The D sublevel contains 5 D orbitals The D sublevel starts in the 3 rd energy levelThe D sublevel starts in the 3 rd energy level 5 different shapes (orbitals)5 different shapes (orbitals) Each orbital can hold 2 electronsEach orbital can hold 2 electrons

26. The F sublevel has 7 F orbitals The F sublevel starts in the fourth energy levelThe F sublevel starts in the fourth energy level The F sublevel has seven different shapes (orbitals)The F sublevel has seven different shapes (orbitals) 2 electrons per orbital2 electrons per orbital

27. Summary Starts at energy level

28. Electron Configurations The way electrons are arranged in atoms.The way electrons are arranged in atoms. Aufbau principle- electrons enter the lowest energy first.Aufbau principle- electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies.This causes difficulties because of the overlap of orbitals of different energies. Pauli Exclusion Principle- at most 2 electrons per orbital - different spinsPauli Exclusion Principle- at most 2 electrons per orbital - different spins Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.

29. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

30. Phosphorous, 15 e- to place The first to electrons go into the 1s orbital Notice the opposite spins only 13 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

31. The next electrons go into the 2s orbital only 11 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

32. The next electrons go into the 2p orbital only 5 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

33. The next electrons go into the 3s orbital only 3 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

34. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes 3 unpaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3

35. Orbitals fill in order Lowest energy to higher energy.Lowest energy to higher energy. Adding electrons can change the energy of the orbital.Adding electrons can change the energy of the orbital. Half filled orbitals have a lower energy.Half filled orbitals have a lower energy. Makes them more stable.Makes them more stable. Changes the filling orderChanges the filling order

36. Write these electron configurations Titanium - 22 electronsTitanium - 22 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 21s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Chromium - 24 electronsChromium - 24 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected But this is wrong!!But this is wrong!!

37. Chromium is actually 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 51s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why?Why? This gives us two half filled orbitals.This gives us two half filled orbitals. Slightly lower in energy.Slightly lower in energy. The same principal applies to copper.The same principal applies to copper.

38. Copper’s electron configuration Copper has 29 electrons so we expectCopper has 29 electrons so we expect 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 91s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration isBut the actual configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 101s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This gives one filled orbital and one half filled orbital.This gives one filled orbital and one half filled orbital. Remember these exceptionsRemember these exceptions

39. Great site to practice and instantly see results for electron configuration. electron configuration electron configuration