2. ELECTRON CONFIGURATION By Hilary Scurlock Edited by Mrs. Rosenfield

3. Electron Configurations Describe the location of electrons Orbital = region of space (same as electron cloud) Principle quantum # (n) = energy level –n = 1, 2, 3, 4, 5 … with increasing distance from the nucleus Energy Sublevels = s, p, d, f (see p. 133-134) –s sphere1 orbitalmax 2 electrons –pdumbbell3 (x, y, z)6 –dclover leaf510 –f714 Max 2 electrons in each orbital

4. Atomic Orbitals http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html This site shows 3-D versions of the atomic orbitals – I just thought you would be interested

6. Some important info… There are 4 quantum numbers that are associated with electron configurations: –Principal quantum number (n) = energy level –Angular momentum quantum number ( l ) = sublevel, which is the type of orbitals (s, p, d, or f) –Azmuthal quantum number (m l ) = number of orbitals related to the sublevel –Electron spin quantum number (m s ) = tells if electrons are spinning clockwise or counter clockwise

7. We can use the four quantum numbers to label each unique electron in any orbital in any atom, thus giving us electron configuration. The quantum numbers are just like your address! No two electrons will have the same four quantum numbers

8. Principal quantum number (n) n can be any integer from 1 to infinity It tells you the energy level The larger the number the higher the energy level

9. Angular momentum quantum number ( l ) l is any integer from 0 to n – 1 It tells you the type of orbital l value Orbital Type 0 sSsS 1 pPpP 2 dDdD 3 fFfF

10. Azmuthal quantum number (m l ) Tells you the number of orbitals related to the sublevel m l is any integer from – l to + l

11. Electron spin quantum number (m s ) This tells if electrons are spinning clockwise or counter clockwise m s is either + ½ or – ½ For all electron configurations you will not know if the electrons are clockwise or counter clockwise but with orbital notation you will.

12. For example the electron configuration for Hydrogen is: 1s 1 Principal quantum # (n) Angular momentum quantum # ( l ) # of electrons in orbital or subshell NOTICE: For electron configurations m l and m s are not used

13. Orbital Diagrams These show us the number of orbitals and the spin of the electrons in those orbitals. –Here m l and m s are used H 1s 1 This box represents the orbital number (m l ) It can also be given as lines or circles The arrow denotes one of the two possible spinning motions (m s )

14. Knowing this, now we can draw out a configuration and an orbital notation Helium –2 electrons Beryllium –4 electrons Lithium –3 electrons Helium 1s 2 ___ 1s 2 Beryllium 1s 2 2s 2 ___ ___ 1s 2 2s 2 Lithium1s 2 2s 1 ___ ___ 1s 2 2s 1

15. An Electron configuration is how the electrons are distributed among the various atomic orbitals. Some Rules to Configure By: If you want to get these electron configurations/orbital diagrams right there are some rules that we must follow.

16. PAULI EXCLUSION PRINCIPLE This principle states that two electrons in the same orbital must have opposite spins. An example of the Pauli exclusion principle: He (1s 2 ): ____ ____ 1s 2 1s 2

17. HUND’S RULE Electrons entering a subshell containing more than one orbital (Ex: p, d or f - orbitals) will be the most stable if the electrons are arranged by themselves in separate orbitals alone before being paired up. An example of Hund’s rule: Nitrogen (1s 2 2s 2 2p 3 ): ___ ___ ___ ___ ___ 1s 2 2s 2 2p 3 What would fluorine look like?

18. AUFBAU PRINCIPLE Electrons start at the lowest energy orbitals first and then continue to fill orbitals of increasing energy. How do we know the order????

19. The electrons follow a general pattern when filling energy levels and orbitals and guess what object we can thank in helping us figure this out???? THE PERIODIC TABLE

20. Those wacky d orbitals! When dealing with transition metals, which have d orbitals, always fill the s-orbital in the next energy level before filling the d- orbital. For example, the configurations of elements such as Sc are different than we might expect, because, in the cases of this element, the 4s orbital is filled before the 3d orbital. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1

21. Another way to do it: If the periodic table is too confusing for you, or you just like to memorize things, try: Energy level diagram - basically a picture of the aufbau principle –It shows you the order in which the orbitals are filled.

22. Order of filling orbitals (p.135)

23. Review of Writing Electron Configurations

24. Ask yourself these questions every time you have to write an electron configuration or orbital notation: 1.Where is the element on the periodic table? 2.What is the atomic number? 3.How many electrons? 4.What is the row number? 5.How many energy levels? 6.What subshell(s) does the element have? 7.What is the electron configuration? ONLY CONTINUE IF YOU HAVE TO WRITE ORBITAL NOTATIONS: 8.How many orbitals in each subshell? 9.What is the orbital notation?

25. Some Examples: Give the electron configuration for the following: HeBe ONa ArTe Give the orbital diagram for the following: LiC VK

26. Quick Quiz! What is wrong with this configuration? Al: 1s 2 2s 2 2p 4 3s 2 3p 3

27. Where do the Lanthanide’s and Actinide’s fit in??? Lanthanides (f-block) Ce-Lu Actinides (f-block) Th-Lr Ce 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 1 5d 1 Cm 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 7s 2 6d 1 5f 7 Hg 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10

28. A SHORTCUT You do not have to write out all the info Noble gas core Li = [He] 2s 1 P = [Ne] 3s 2 3p 3 TRY Oxygen Zinc Cesium

29. Exceptions to the Rules Just as with any good rule there are exceptions!!! An atom with almost half or completely - filled d and f orbitals are more stable when you can fill them from an s-orbital Ex: Chromium = [Ar]4s 1 3d 5 OR Copper = [Ar]4s 1 3d 10 You try: –Molybdenum = Mo –Gold = Au –Europium = Eu

30. Valence Electrons These electrons determine the chemical properties of an element They are generally the electrons in the atom’s highest energy levels. –Determine the number of valence electrons using electron configurations: –S [Ne] 3s 2 3p 4 = 6 valence e - –Cs [Xe] 6s 1 = 1 valence e - –Fe[Ar] 4s 2 3d 6 = 2 valence e - –Br[Ar] 4s 2 3d 10 4p 5 = 7 valence e - On the periodic table elements in the same column tend to have the same number of valence electrons.

31. Electron-dot Structures Visual way to show valence electrons of an element for bonding. Has element’s symbol in the middle and is surrounded by dots representing the atom’s valence electrons. –Dots are placed one at a time on the four sides of the symbol and then paired up until they are all used.

33. Examples of Electron-dot Structures S Cs Fe Br See pg 140 – 141 for more examples

34. DIAMAGNETISM and PARAMAGNETISM Paramagnetic substances are those that are attracted by a magnet. Because the electrons have parallel spins, the magnetic fields reinforce each other. Diamagnetic substances always repel magnets. Because the electrons have antiparallel spins, the two magnetic fields cancel each other out. NSNS SNSN NSNS NSNS NSNS DIAMAGNETICPARAMAGNETIC