What you need today

1. What you need today • You need your Atomic Structure Packet • Periodic table • Paper for notes • Writing Utensil

2. Tonight’s Homework • Page 6 in Atomic Structure Packet #1-6 orbital diagram and electron configuration

3. Videos Watch videos 1-8 and take notes on the handout • Click here to show videos

4. Orbital Diagrams

5. Quantum Mechanical Model Max Planck and Werner Heisenberg expanded upon Bohr’s model There are energy levels, but they are not circular orbits Orbitals are regions in which you may find electrons in the electron cloud Bohr Quantum Image from the Higher Education Academy Physical Sciences Centre

6. Electrons are organized in four ways 1st = Energy levels or shells 1,2,3,4,etc 4th = Spin ‘up’ or ‘down’ 2nd = Sublevels or subshells each energy level contain sublevels called s,p,d,f 3rd = Orbitalseach sublevel contain orbitals (different # for different sublevel) Definition: an orbital is the space occupied by two (a pair of) electrons

7. Electron Sublevels

8. Practice – Rally Coach • On the next slide there are a list of questions over what you just learned. • A will tell B the answer to the first question then B will tell A the answer to the next questions and so forth

9. A:How many electrons can a d sublevel hold? B:Name the sublevels. A:What energy level does sublevel f start on? B:How many electrons can the second energy level hold? A: How many orbitals are in a p sublevel? B:How many electrons can an s sublevel hold? A:How are energy levels labeled? B:What energy level does sublevel p start on?

10. Since sublevels and orbitals are too complicated to draw all the time, we simplify with orbital diagrams. Each orbital is represented by a box or a line. or Each electron is represented by an up or down arrow.  (means 2 electrons) _______

11. s sublevel s sublevel has one orbital, so we draw one box. p sublevel p sublevel has 3 orbitals, so we draw 3 boxes.

12. d sublevel d sublevel has 5 orbitals, so we draw 5 boxes. f sublevel f sublevel has 7 orbitals, so we draw 7 boxes.

13. Orbital and sublevel information is like a map, telling you where an electron can be found in an atom. There are three rules that govern why an electron will be in one sublevel rather than another: • Pauli Exclusion Principle • Hund’s Rule • Aufbau Principle

14. Pauli Exclusion Principle: No two electrons can be organized in exactly the same way (E.L, sublevel, orbital, spin). What that means is that 2 electrons may occupy one orbital, but they must have opposite spin direction. This is why we draw electrons as arrows facing opposite directions when they share a box: WRONG!   RIGHT! 

15. Hund’s Rule of Maximum Multiplicity:Electrons occupy vacant orbitals before pairing   Example: 4 electrons in a d sublevel WRONG -This is not as stable.Electrons repel each other.     RIGHT -This is stable. Hund’s Rule paraphrased – spread them out before you pair them up!

16. Aufbau Principle: Each electron must occupy the lowest energy orbital available. Not all sublevels and orbitals have the same energy! s  p  d  f energy increases 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p This chart shows the order that electrons fill sublevels. Although the d and f sublevels are on lower energy levels, they have high energy and do not fill until after the s and p for higher energy levels.

17. The number of columns in each block corresponds to the number of electrons that fit in that sublevel

18. You can use the periodic table like a game board to see the order in which the sublevels fill. Each square can mean the position of one electron The period number tells you the energy level (for s and p) The block tells you the sublevels

19. An element’s location within the block tells you how many electrons it has in that sublevel Example: Nitrogen is in the 3rd column of the p block. It has 3 electrons in the 2p sublevel

20. Rally Coach A: To which sublevel do we go after filling 4s? B: To which sublevel do we go after filling the 6s? A: To which sublevel do we go after filling the 4d? B: To which sublevel do we go after filling the 2s?

21. Example: Manganese has 25 electrons. Draw its orbital diagram. 1s 2s 2p 3s 3p 4s 3d Use Aufbau Principle to decide which order to fill sublevels, Pauli Exclusion Principle to place arrows in opposite directions in each orbital, Hund’s Rule to spread out before pairing up!

22. Example: Manganese has 25 electrons. Draw its orbital diagram.          1s 2s 2p 3s 3p       4s 3d Use Aufbau Principle to decide which order to fill sublevels, Pauli Exclusion Principle to place arrows in opposite directions in each orbital, Hund’s Rule to spread out before pairing up!

23. Practice • Individually: write the orbital diagram for • Chlorine • Switch papers with your shoulder partner and check for accuracy (correct/celebrate) • Individually: write the orbital diagram for • Iodine • Switch papers with your face partner and check for accuracy

24. Electron Configuration

25. 1s 4s 2s 2p 3s 3p Electron Configurations • Draw the orbital diagram for calcium.

26. Shorthand Electron Configuration • In Aufbau order of filling electrons for Ca 1s2 2s2 2p6 3s2 3p6 4s2 • Superscripts are number of electrons in that sublevel • How many electrons does this neutral atom have?1s2 2s2 2p6 3s2 3p6 4s1 2+2+6+2+6+1= 19 electrons • What element is it? potassium

27. Writing Electron Configuration • Always fill electron orbitals in Aufbau Order • When you write your final electron configuration, put them in the same order electrons fill: 1s2 2s2 2p6 3s2 3p6 4s23d10 4p6 5s24d10 5p6

28. Now you practice! • Write the electron configuration for cadmium. • Write the electron configuration for tantalum. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d3

29. Write the electron configuration for tantalum. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d3

30. Your Turn: Nickel has 28 electrons. Draw its orbital diagram. Then, write its shorthand electron configuration.

31. Ions and Exceptions

32. What are ions? • Ions have lost or gained electrons • Gained electrons = anions, negatively charged • Lost electrons = cations, positively charged

33. If an atom gains one additional electron, just follow the aufbau rules to put the electron in its proper location For example: Cl-1 Write the electron configuration for the neutral atom: Cl: 1s22s22p63s23p5 2. Add the addition electrons into the vacant orbitals (put the extra in the 3p) Cl -1: 1s22s22p63s23p6 Anions

34. What if it gains 2 electrons? For example: O2- (Neutral) O: 1s22s22p4 (Anion) O2-: 1s22s22p6 What do you notice about anions? They always have electron configurations that look like noble gases! Anions

35. If an atom loses one electron, you must take the electron from the highest energy sublevel! For example: Na+1 Write the electron configuration for the neutral atom: Na: 1s22s22p63s1 2. The highest energy sublevel is the 3s, so remove the electron from there Na+1: 1s22s22p6 Cations

36. Remember, you must take the electrons from the highest energy sublevel! For example: Ti2+ Write the electron configuration for the neutral atom: Ti: 1s22s22p63s23p64s23d2 2. The highest energy sublevel is the 4s, so remove the electrons from there Ti2+: 1s22s22p63s23p63d2 Cations

37. Valence Electrons • Electrons in the highest ENERGY LEVEL • Let’s practice with the examples so far: • Cl-1 = 8 • O2- = 8 • Na+1 = 8 • Ti2+ =10

38. exit quiz • Write the electron configuration of gold(I) and identify the number of valence electrons [gold(I) = Au+] • Circle the number of valence electrons so it is easy to find • Draw the orbital diagram for fluoride