1. Chapter 3: Electrons in atoms

2. Learning outcomes:  Energy levels and shapes of orbitals  Electronic configurations  Ionisation energy, trends across a period

3. The quantum mechanical model

4. How are electrons arranged?. Electrons are not evenly spread but exist in layers called shells. The arrangement of electrons in the shells is called the electron structure or electronic configuration. 3 rd shell n=3 2 nd shell n=2 1 st shell n=1 Max: 2 n 2

5. Each shell has a maximum number of electrons that it can hold. Electrons fill the shells nearest the nucleus first. 3 rd shell holds a maximum of 18 electrons 2 nd shell holds a maximum of 8 electrons 1 st shell holds a maximum of 2 electrons

7. Time for a break and practice Please make check-up 1 on page

8. Simplified electron configuration 2, 5 2, 8 2, 8, 7 2, 8 7N7N 2 nd shell holds a maximum of 8 electrons 1 st shell holds a maximum of 2 electrons 17 Cl 11 Na + 8 O 2- For “ion” the number of proton is NOT the same as electrons

9. Electronic configuration  This model assumes the electrons have the same location and energy  Until it was discovered that electrons have different locations and energy Li Li + + e - (1 st ionisation energy: E 1 ) Li + Li 2+ + e - (2 nd ionisation energy: E 2 ) Li 2+ Li 3+ + e - (3 rd ionisation energy: E 3 ) E 1 ≠ E 2 ≠ E 3

10. Ionisation energy Li Li + + e - (1 st ionisation energy: E 1 ) Li + Li 2+ + e - (2 nd ionisation energy: E 2 ) Li 2+ Li 3+ + e - (3 rd ionisation energy: E 3 ) Li 2,1 Which electron will be easiest to remove? E 1 < E 2 < E 3 ΔH i1 < ΔH i2 < ΔH i3

11. Table 3.2 in the book on p. 35  For every element, the successive ionisation energy increases; for every next electron it is more difficult to remove  We can in theory continue removing electrons until only the nucleus is left  We call this sequence the “successive ionisation energy”  Sometimes we find a big gap/jump in ionisation energy

12. Example: sodium  The first ionisation energy is quite low, it is likely quite far from the nucleus  The 2 nd to the 9 th ionisation energy are in a gradual successive increase indicating these electrons are in the same shell  The 10 th and 11 th electrons have high ionisation energies compared to the rest, they must be near the nucleus.  The jump between the 9 th and 10 th suggests a change in shell

13. Factors affecting the first ionization energy  Nuclear charge (number of protons) the bigger nuclear charge, the higher 1st ionization energy.  Atomic radius (distance effect) the bigger atomic radius, the lower 1st ionization energy.  Shielding effect (number of shells) the bigger Shielding effect, the lower 1st ionization energy

14. The first ionization energies of the first 20 elements in the periodic table is shown below:

15. Worked example

16. The model of the atom A model is what fits logic, experimental observations and mathematical calculations

17. 17 Cl2, 8, 73 2 nd shell, with a maximum of 8 electrons Symbol Simple electronic configurationNumber of shells (last number is Group) (=period) 1 st shell, with a maximum of 2 electrons 3 nd shell, with a maximum of 18 electrons

18. 6 C 10 Ne 11 Na Symbol Simple electronic configurationNumber of shells (last number is Group) (=period) 2, 42 2, 82 2, 8, 13 19 K4

19. Where in the atom is the electron? According to quantum mechanics it is most likely to find the electron for the of the H-atom at 0.0000000000529 meter (52.9 pm) from the nucleus

20. Shells  Principal quantum shells (n=1, n=2 etc.)  Remember for each the max number of electrons is 2n 2 (so for n=2, max 8 electrons)  We know from experiments and calculations these 8 electrons have different energies…. so we need a new model of the atom where we can distinguish between electron energy n = 1 n = 2 Quantum shellSubshells

21. The quantum mechanical model Simplified model Realistic model

22. Principal quantum shell Number of Sub-shells Name of the Sub-shell Max. number of electrons n = 11 1s 2 n = 22 2s 2p 2626 n = 33 3s 3p 3d 2 6 10

23. Subshells and their shapes Atomic orbital is a space around the nucleus holding 1 or 2 electrons

24. n = 11s Where in the atom are the electrons? energy 2 He2 2 Simple electronic configurationComplicated electronic configuration 1s 2 Principle quantum shell Sub-shell Number of electrons

25. n = 1 n = 2 1s (e<2) 2s (e<2) 2p (e<6) Where in the atom are the electrons? energy 8 O2, 6 2 2 4 Simple electronic configurationComplicated electronic configuration 1s 2 2s 2 2p 4

26. n = 1 n = 2 1s (e<2) 2s (e<2) 2p (e<6) Where in the atom are the electrons? energy 11 Na2, 8, 1 2 2 6 Simple electronic configurationComplicated electronic configuration 1 n = 3 3d (e<10) 3p (e<6) 3s (e<2)

27. n = 1 n = 2 1s (e<2) 2s (e<2) 2p (e<6) Where in the atom are the electrons? energy n = 3 3d (e<10) 3p (e<6) 3s (e<2) n = 4 4s (e<2) 4d (e<10) 4p (e<6) 4f

29. Subshells and atomic orbitals

30. From simple to complicated electron configuration to noble gas electronic configuration notation 2, 6 1s 2 2s 2 2p 4 [He] 2s 2 sp 4 2, 8, 7 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne] 3s 2 3p 5 8 O 17 Cl 19 K 35 Br Element: Simple: Complicated: Noble gas:

31. Note the following:  Potassium: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1  The 3d subshell: 3d<4p  Chromium and copper are exceptions:  Cr: [Ar] 4s 1 3d 5 rather than [Ar] 4s 2 3d 4 and  Cu: [Ar] 4s 1 3d 10 rather than [Ar] 4s 2 3d 9

32. The blocks of the periodic table  Elements in Group 1 and Group 2 are in the s-block and have their outer electrons in an s subshell.  Elements in Group 3 to 18 have outer electrons in a p subshell.  Elements that add electrons to the d subshells are called the d-block elements.

33. Use the electronic configuration to find the group …s 2 is in group: …p 1 is in group: …p 3 is in group: …p 6 is in group: …d 3 is in group: …d 7 is in group:

34. In which period, group and block of the following electron configuration? periodgroupblock 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 1s 2 2s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5

35. RULES FOR FILLING ENERGY LEVELS  Aufbau Principle “Electrons enter the lowest energy orbital first”  Pauli’s Exclusion “Sub-Orbitals can hold a max. of 2 electrons provided they have opposite spin”  Hund’s Rule “Orbitals of the same energy remain singly occupied before pairing up.

36. Examples N = 1s 2 2s 2 2p 3 O = 1s 2 2s 2 2p 4

37. From simple to complicated electron configuration 2, 5 becomes 1s 2 2s 2 2p 3 2, 6 becomes 1s 2 2s 2 2p 6 2, 8, 7 becomes 1s 2 2s 2 2p 6 3s 2 3p 5 2, 8, 8 becomes 1s 2 2s 2 2p 6 3s 2 3p 6 7 N 8 O 2- 17 Cl 19 K +

39. Ionisation: trend across a period  General increase across period  Rapid decrease between last element of a period and 1 st of a new period  Be and B because 2s and 2p  N and O

40. The first ionization energies of the first 20 elements in the periodic table is shown below:

41. Ionisation: Trend down a group  General trend decrease  further away from the nucleus  increased shielding  despite increased nuclear charge Li = 519 kJ/mol Na = 494 kJ/mol K = 418 kJ/mol Rb = 403 kJ/mol

42. Worked example