Electron Configurations & The Periodic Table

1. Electron Configurations & The Periodic Table

2. RECAP

3. Schrodinger’s Standing Waves • Louis De Broglie developed a theory that matter can have wave-like properties • Schrodinger extended this theory to electrons bound to a nucleus • Postulated that electrons resembled a standing wave • Certain orbitals exist at whole wavelengths of electron vibrations

4. Orbitals - Redefined • Orbital: region around the nucleus where there is a high probability of finding an electron • As per wave model of Schrodinger – because things are vibrating

5. Heisenberg Uncertainty Principle

6. Heisenberg Uncertainty Principle • Heisenberg studied statistics and developed matrix algebra • Developed a statistical approach to explaining how electrons works and realized… • IT IS IMPOSSIBLE TO KNOW THE EXACT POSITION AND SPEED OF ELECTRON AT A GIVEN TIME • At best, we can describe the probability of finding it at a specific place

7. Wave functions: the mathematical probability of finding an electron in a certain region of space • Wave functions give us: • Electron probability densities: the probability of finding an electron at a given location, derived from wave equations

10. Quantum Numbers • Quantum Numbers: numbers that describe the quantum mechanical properties (energies) of orbitals • From the solutions to Schrodinger’s equation • The most stable energy states is called the ground state

11. Summary of Quantum Numbers & Restrictions

12. Electron Configurations and Energy Level Diagrams • The four quantum numbers tell us about the energies of electrons in each atom • Unless otherwise stated were are talking about ground state energies

13. Hund’s rule – e- half-fill each orbital in a sublevel before pairing up Energy Level Diagrams 6s Pictorial representation of electron distribution in orbitals 5p 4d Aufbau principle – e- occupy the lowest energy orbital available p. 188 in text 5s 4p Pauli exclusion principle – max 2 e- per orbital (spin up and spin down) 3d E 4s 3p 3s 2p 2s n = 1 l = 0 ml = 0 ms = -½ 1s

15. Creating Energy-Level Diagrams • Pauli exclusion principle – • no two electrons in an atom may have the same four quantum numbers • no two electrons in the same orbital may have the same spin • only two electrons with opposite spins may occupy an orbital • aufbau principle – (German for “building up’) • each electron is added to the lowest available energy orbital • Hund’srule – • one electron is placed in each orbital at the same energy level before the second electron is placed

16. Orbital Diagrams O(z = 8) 1s 3s 3p 2s 2p P(z = 15) 1s 3s 3p 2s 2p Ar(z = 18) 1s 3s 3p 2s 2p

17. Energy Level Diagrams 6s 5p 4d 5s 4p 3d E 4s 3p 3s 2p 2s Anions - Add e- to lowest energy sublevel available. 1s

18. Energy Level Diagrams 6s 5p 4d 5s 4p Cations - Remove e- from sublevel with highest value of n. 3d E 4s 3p 3s 2p 2s 1s

19. Energy Level Diagrams 6s 5p 4d 5s 4p Cations - Remove e- from sublevel with highest value of n. 3d E 4s 3p 3s 2p 2s 1s

20. Electron Configuration & The Periodic Table

21. s-block f-block d-block p-block 2 groups 14 groups 10 groups 6 groups Explaining the Periodic Table Main Group metals are “S block”

22. 6p 6e- Period 6 10e- 5d 32e- 4f 14e- 2e- 6s 6e- 5p Period 5 4d 10e- 18e- 5s 2e- 6e- 4p Period 4 10e- 3d 18e- 2e- 4s 6e- 3p Period 3 8e- 3s 2e- 6e- 2p Period 2 8e- 2e- 2s Period 1 2e- 2e- 1s 1,2 13-18 3-12 La & Ac series Group

23. Electron configurations • Lists e- location from low to high energy in the following format iron atom: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

24. Cl(z = 17) 1s 3s 3p 2s 2p Shorthand e-configurations becomes . . . Cl: 1s2 2s2 2p6 3s2 3p5 • e-config. written only for the outer shell electrons • Inner shell e- expressed as a noble gas. So . . . Cl: 1s2 2s2 2p6 3s2 3p5 becomes . . . Cl: [Ne]3s2 3p5 Shorthand: [noble gas]

25. Anomalous E Configurations • Completely filled sub-levels are more stable than half filledorbitals which are more stable than partially filledorbitals. • > >

26. n=4 n=3 s p d Anomalies in Configurations

27. The configuration of Copper What do you think the Actual e- configuration is?

28. e- e- e- e- Electrostatic repulsion North Magnetic field South Magnetic field Explaining the multivalent ions The configuration of stable ions can also be predicted by electron configuration • As orbitals are being filled there are varying levels of stability due to the interaction of forces or electrostatic repulsion between electrons and force associates with the magnetic field due to the electron’s spin. • Filled orbitals are most stable because the electrostatic repulsion is balanced against the magnetic attraction. Magnetic attraction

29. Explaining the multivalent ions • Orbitals that are completely filled, like Noble gases, have the most stable structure due to the balanced forces between electrostatic repulsion and magnetic attraction. • Transitional metals often have partially filled d-orbitals and complete s-orbitals with similar energy levels. Electrons are lost to achieve the best combination of stability. • Best stability – completely filled orbitals (2 electrons/orbital) Electrostatic repulsion and Magnetic attraction • Next best stability – half filled orbitals (1 electron/orbital) Electrostatic repulsion with minimal crowding • Co: [Ar] 4s2 3d7 Co: [Ar] - Filled 4s and half filled 3d Co2+: [Ar] - Half filled 4s & 3d less stable than Co2+ so less common ion Co3+: [Ar]

31. Some materials exhibit strong magnetic properties naturally these are referred to as being ferromagnetic. • As moving charges have magnetic fields so too do spinning electrons. Those spinning in one direction would have one magnetic polarity as those spinning in the opposite direction the opposing magnetic pole. • For example – clockwise – south pole and counter clockwise –north pole • Those atoms that have a number of similarly spinning electrons would have similar magnetic fields. • If these “magnetic atoms” were free to align themselves with neighbouring atoms of similar characteristics they would create regions of magnetism in the material called domains. • The arrangement of these domains in a material results in a magnet.

32. Iron, nickel and cobalt are such naturally occurring atoms. Iron, nickel and cobalt are small enough atoms that they can realign themselves due to the magnetic properties of their surrounding and thereby create domains. There are other such atoms that have similar electron configurations but limited ability to migrate. Hence, they are reduced in their magnetic properties.

33. Ferromagnetic – materials with strong magnetic properties. Their presence increases a magnetic field substantially. Paramagnetic – materials with weak magnetic properties. Their presence only slightly strengthens a magnetic field. Diamagnetic – materials that have reduced magnetic properties. Their presence weakens a magnetic field.

34. Shape of orbitals • The diagram we used to represent oxygen is; - - - 8 Protons - - - - - 16 O 8

35. Shape of orbitals • The diagram we might currently use to represent oxygen is;

36. Try These… • p. 191 #3alt,4 • p. 197 #5 • p. 194 #6-11 • p. 197 #1,2,6-10,13

37. Confidence building questions Write out the shorthand notation for the electron configuration of B. Write out the shorthand notation for the electron configuration of Cl. Write out the shorthand notation for the electron configuration of F. Write out the shorthand notation for the electron configuration of Ca. Write out the shorthand notation for the electron configuration of Kr. Write out the shorthand notation for the electron configuration of O2-. Notice that this is an anion! Write out the shorthand notation for the electron configuration of Na+. Notice that this is a cation! Why are Groups 1 and 2 referred to as the s-block of the periodic table? Why are Groups 3 through 12 referred to as the d-block of the periodic table? Using what you now know about electron configurations explain the notion that elements in the same column in the periodic table have similar chemical and physical properties.