EMR and the atom: Part Deux

1. 0 Electron Configurations EMR and the atom: Part Deux http://imagers.gsfc.nasa.gov/ems/waves3.html

2. First, a review… • What do scientists mean when they say light is quantized? • Compare a ground state and an excited state electron. • Explain the origin of the atomic emission spectrum of an element. • What do we know so far about the arrangement of electrons in atoms?

3. 0 What you’ve seen so far…. Model of a Nitrogen (z=7) atom

4. 0 Which is really not true- why? • Because orbitals- the “electron cloud” are • 3-D, not flat • are not round in most cases • e-s are spread out as much as possible • (e-s are moving very rapidly)

5. 0 Orbitals • The electrons are spread out in orbitals that have varying • Shapes • Energy (distance from nucleus) • The orbitals are described in regards to their quantum numbers • Descriptions that are descriptive and hierarchical • There are 4 numbers that describe an orbital • Written as follows: (#, #, #, ±#)

6. 0 Principal quantum number (n) The first number (1, #, #,±#) • Describe the • distance from the nucleus of the orbital • The energy of the orbital • Values for n are integers • The smallest possible value is 1 • As the distance from the nucleus (and therefore energy) increases, the number increases

7. 0 Quantum numbers

8. 0 There periodic table and n • The 7 periods on the periodic table correspond to n values • Each period has a unique n value • For the 1st period, n=1 • For the 2nd period, n=2 • And so on….

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12. 0 The shape of things • Is the shape of the orbital • There are 4 shapes (although we only deal with the first three) • s = sphere • p= peanut • d= double peanut • f = flower

13. The s orbital http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif

14. p orbitals

15. d orbitals

16. d orbitals

17. f orbitals

18. General tutorials for electron configuration stuff • some slides in this PowerPoint are from this site already • http://www.wwnorton.com/college/chemistry/gilbert/tutorials/ch3.htm • See key equations and concepts (select from menu on the left), as well as the looking through the overview where to the tutorials are listed (links for just those are on the left, too)

20. 0 Orbitals • Denote the orbital sublevel that is filled • It is the third number in the description (#,#,1, ±#) • s orbitals have one sublevel; a sphere has one orientation in space • p orbitals have three sublevels; 3 orientations in space • d orbitals have five sublevels; 5 orientations in space • f orbitals have seven sublevels; 7 orientations in space

21. “Flavors” of ml 0 • s orbitals have one sublevel; a sphere has one orientation in space

22. “Flavors” of ml 0 • p orbitals have three sublevels; 3 orientations in space

23. “Flavors” of ml 0 • d orbitals have five sublevels; 5 orientations in space

24. “Flavors” of ml 0 • f orbitals have seven sublevels; 7 orientations in space

25. 0 Spin • It is the last number in the description (#,#,#,±½) • Spin is +½ or -½ • Up or down

26. 0 How we use this…. • There is a specific order to how the e- fill the orbitals; it is not random • Although there are exceptions to the rules (last thing we do)

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28. The principles of e- configuration • The Aufbau (next) Principle: • That e- fill the lowest energy sublevel before going to the next sublevel • The Pauli Exclusion Principle: • That e-s are paired according to opposite spins • Hund’s Rule: • e-s spread out in equal energy sublevels before placing electrons

29. The first level to fill is the 1s level • It is the lowest energy sublevel • It holds two electrons • They are oppositely paired (up and down- ↑↓) • Each sublevel (each __) holds 2 electrons

30. Next… • The second sublevel is the 2s sublevel • It also holds 2 electrons (because s holds 2, not because of the number), • also oppositely paired ↑↓

32. 1s2, 2s2,then comes 2p6 • So, as it states above • 1s fills, 2s fills ,then comes 2p • It holds up to six electrons • Because p orbitals hold 6 electrons

33. 0 Next… • From 2p, • 3s fills with 2e-, then onto • 3p, with 6e- then • 4s with 2e- followed by • 3d with10e- (because d holds 10e-) • Then 4p with 6e- • Notice, you follow the arrows • Remember, the number of electrons comes from the letter (the orbital’s momentum,ml)

34. 0 • The sublevels of the orbitals are first filled, then you continue onto the next level (Aufbau) • Also be sure to place one electron in each sublevel prior to filling the level (↑↑↑ and not ↑↓↑ _) (Hund) • e-s must be paired with e-s of opposite spin (↑↓, not ↑↑ or ↓↓) (Pauli)

35. 0 Putting it all together… • Carbon (neutral, so 6 electrons) • What this would look like: ↑↓↑↓↑↑ _ 1s 2s 2p (notice there are 6 arrows for 6 electrons) • This can also be written as 1s2 2s2 2p2 • Notice the superscripts add up to 6

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37. 0 There are some exceptions… • This is because some energy levels are very close together • electrons are able to move between close orbitals in order to minimize repulsion • Example: the 4s and 3d orbitals are very close in energy • So exceptions for some period 4 d block elements occur • Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4 • Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5 • Because it takes less energy to split the electrons between the 5 sublevels than it does to put them together in the 4s and 3d

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