Electron Configurations & The Periodic Table

1. Electron Configurations & The Periodic Table

2. Quantum Atomic Theory Schrödinger (1887-1961) Heisenberg (1901-1976) Plank (1858-1947) Pauli (1900-1958) Einstein (1879-1955)

3. Bohr-Rutherford vs Quantum Similarities • Electrons have discrete amounts of energy • Positive nucleus in the centre Differences: Heisenburg Uncertainty Principle– the position and magnitude of an electron cannot both be known. If one is measure, the other is altered.

4. Orbitals - Redefined • The Quantum model of the atom describes electrons in different orbitals (or energy levels) around the nucleus. The traditional orbits of the Bohr model are subdivided. • Orbital/Sub-orbital: region around the nucleus where there is a high probability of finding an electron • The Period on the periodic table tells you the energy level, the blocks contained within that period lets you know the kinds of orbitals the atom has or where the last electrons of that element are likely to be at any given moment.

6. The maximum number of electrons in each type of suborbital: s = 2 electrons maximum p = 6 electrons maximum d = 10 electrons maximum f = 14 electrons maximum

7. Electron Configuration & The Periodic Table

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

9. So what does carbon look like? • Carbon has 6 e-, so… • Therefore, the electron configuration of carbon is: 1s2 2s2 2p2 The electron configuration of potassium is: 1s2 2s2 2p6 3s2 3p6 4s1

10. Shorthand e-configurations Cl: 1s2 2s2 2p6 3s2 3p5 becomes Cl: [Ne]3s2 3p5 • e-config. Is written only for the outer shell electrons • The noble gas indicates all inner shells are full Shorthand: [noble gas]

11. 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

12. 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

13. 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

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

15. 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

16. 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

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

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

19. 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.