1. Chemistry Topic 2- Atomic Structure Electric configuration Made by: Mr. Hennigan’s class (Class 2014 CHEM HL crew )

2. Electric Configuration 12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of main energy levels and sub-levels in atoms. 12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom. 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 the p x, p y and p z 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.

3. Orbitals There are different orbitals in an atom: s orbitals p orbitals d orbitals f orbitals These have different shapes – and energy levels

4. S - orbitals

5. s-orbitals The s-orbital can hold up to 2 electrons (which pair up. Every electron shells has a single s-orbital with 2 electrons in it.

6. P - Orbitals

7. PEE-orbitals The p orbitals are the orbitals that fill up after the s orbitals are filled. The p orbitals can have up to 6 electrons (that is 3 pair of electrons) The n=1 shell doesn't have a p orbital. n=2 onwards have s and p orbitals

8. D - Orbitals

9. d orbitals d orbitals can hold up to 10 electrons which is 10/2 = 5 pair of electrons You are not required to know the shape of d orbitals but you should know it exists and the way it is filled with electrons will be explained later.

10. F - Orbitals

11. YOU ARE NOT REQUIRED TO KNOW THE SHAPE OF F ORBITALS!

12. Aufbau Principle (this is German for “building up”) It states that electronic configuration follows an order, in which the lower levels fill first. The first electron goes into the lowest energy orbital available (the 1s orbital) the next electron pairs up with it in the same orbital and the third electron (that of lithium) fits into the next orbital up, the 2s orbital.

13. Aufbau Principle

14. Pauli's exclusion principle Says that there can only be 2 electrons in each orbital (with opposite spins). We draw these as arrows pointing in different directions. (just like the last slide!)

15. Hund's rule Says that each orbital should be half filled before any is completely filled (since there is less repulsion if all electrons have the same spin). Electrons will therefore fill the lowest energy levels (ie 1 then 2 and so on) with two going in each orbital, but only doubling up when all orbitals in the level are filled.

16. Electron orbital arrangement

17. Orbitals do not always follow the logical order. Some orbitals have less energy than others, even if the energy level is greater. Ex: 4s comes before 3d. When ionization happens, 4s empties before 3d.

18. A way to memorize the order

19. Beware of exceptions: There are two exception in the pattern 1. Chromium 2. Copper Atoms are more stable having a full sub-orbital or by having half filled sub orbital (with the same spin direction) You would expect: Cr= (Ar) 4s 2 3d 4 but for stability =4s 1 3d 5 Try for Copper

20. HOWEVER When removing electrons, 4s electrons are removed first to for cations

21. Now lets look at the evidence for what we’ve just studied!

22. Definition of first ionization energy The first ionization energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

23. Successive ionization means consecutive ionizations 1st IE (ionization energy) = energy required to remove the outermost electron 2nd IE= energy required to remove the outermost electron from the cation (+ve ion) formed by 1st IE

24. An example of Ionization energy graph

25. Why is there a slight drop between Mg and Al?

26. Shouldn't the line be straight? According to the Bohr model electrons from level 3 are all at the same distance to the nucleus. Therefore the increasing number of protons should increase the attraction for the outermost electron, and 1 st IE should keep increasing across a period.

27. Why is there a slight drop between Mg and Al? As we can see Al, has a slight drop, this is because the electronic configuration is (Ne)3s 2 3p 1, this gives evidence that p orbitals are further away from nucleus, hence the electrostatic attraction is lower. This causes the slight drop, as less energy is required to remove the outermost electron What causes the drop between P and S? This is where electrons start to pair up in the same p=orbital, and we get increased repulsion

29. Slide six represents how successive ionization energy and the energy level relate. The closer an electron is to the nucleus the harder it is to remove that electron. Every change in energy level there is a big jump in the energy needed to remove an electron. As we remove electrons in the same energy level ionization energy increases as the electron number is less than the proton number in the nucleus.

31. Guess which element?

32. The element represented in the chart has 11 ionization energies, meaning it has 11 electrons. There are two electrons with the highest ionization energy. Those are located in level 1. The next 8 electrons are close in ionization energy. These are located in level 2 The last electron, with the lowest ionization energy is located in level 3. The element is in period 3 and group 1. The element is Na

34. Here we can see a pattern between periods of the period table Ionization energies generally increase across the period. The peaks are where there are changes in period.

35. So dude to conclude our message in the prelude..... What is the what is the answer to life the universe and everything?? Answer: 42 (If you got it correct you are currently the top of you class and you get a PC!!!)