1. 1 Energy is Quantized! Max Planck first hypothesized that energy produced by atoms can only have certain values and is therefore quantized. That’s the reason why only distinct lines are seen in element line spectras. Energy is quantized and can only exist at certain wavelengths. 7 -

2. 2 Bohr Model Niels Bohr hypothesized that electrons orbit the nucleus just as the planets orbit the sun (planetary model). He labeled the electron orbits with a number, starting with 1 closest to the nucleus and increasing as the orbits get further away from the nucleus. –The number is known as the Principal Quantum Number (n). 7 -

3. 3 Bohr Model Orbits have a fixed radius. The orbit with the lowest energy is closest to the nucleus. The energy of each orbit increases as the orbits get further away from the nucleus. When an electron jumps from one orbit to another, it absorbs or emits energy according to the equation: ∆E = E f – E i 7 -

4. 4 - If you excite an atom with heat, then electrons absorb energy & jump to higher energy levels (shells further away from the nucleus. When the electrons return to ground state, they release the excess energy in the form of light. - Atoms have characteristic, quantized energy levels which results in specific wavelengths of light being released.

5. 5 Modern Model of the Atom The modern model of the atom is based on Schrodinger’s mathematical model of waves This model describes electrons as occupying orbitals, not orbits. –Orbitals Three dimensional regions in space where electrons are likely to be found, not a circular pathway –Principal energy level Orbitals of similar size 7 - Figure 7.11

6. 6 Modern Model of the Atom 7 - Figure 7.12

7. 7 Orbitals Come in different shapes and sizes. –Lower energy orbitals are smaller. –Higher energy orbitals are larger and extend further away from the nucleus. Four most common types are s, p, d, and f. 7 - Figure 7.13

8. 8 s Orbitals 7 - Figure 7.14 Maximum number of electrons allowed = ______

9. 9 p Orbitals 7 - Maximum number of electrons allowed = ______

10. 10 d Orbitals 7 - Maximum number of electrons allowed = ______

11. 11 Orbital Diagram Rules Two principles and 1 rule determine how the electrons are filled in the principal energy levels and sublevels. –Aufbau principle Electrons fill orbitals starting with the lowest-energy orbitals. –Pauli exclusion principle A maximum of two electrons can occupy each orbital, and they must have opposite spins. –Hund’s rule Electrons are distributed into orbitals of identical energy (same sublevel) in such a way as to give the maximum number of unpaired electrons. 7 -

12. 12 Did you really understand that? Electrons are always filled in their ground state, or ____________ energy state 1.Aufbau (building up) principle. Electrons occupy the _______________ subshell available. 2. Hund’s rule. In orbitals of equal energy, each orbital gets _______ electron before any orbital gets a second electron. 3. Pauli exclusion principle. An orbital can only hold _____electrons which must have _______________ spins.

13. 13 Did you really understand that? Electrons are always filled in their ground state, or LOWEST energy state 1.Aufbau (building up) principle. Electrons occupy the LOWEST ENERGY subshell available. 2. Hund’s rule. In orbitals of equal energy, each orbital gets ONE electron before any orbital gets a second electron. 3. Pauli exclusion principle. An orbital can only hold __2___electrons which must have OPPOSITE spins.

14. 14 ELECTRON SPIN PAULI EXCLUSION PRINCIPLE -- In a given atom, no two electrons can have the same set of quantum numbers (address). -- An orbital can hold only two electrons which must have opposite spins, either +1/2 or –1/2.

15. 15 Electronic Energy Levels - The electronic energy levels are classified as: 1) Main levels: n = 1, 2, 3, 4, 5, 6, 7 2) Sublevels: s p d f 3) Orbitals: final resting place for up to two electrons 4) Spin States: 2 electrons of opposite spin per orbital - We will concentrate on the Main & the Sublevels. - The electrons fill the atom starting with the lowest main energy level & lowest sublevel.

16. 16 III. Electronic Energy Levels - Main Energy Levels & Accompanying Sublevels: n = 1 has an s sublevel (1 sublevel) n = 2 has s & p sublevels(2 sublevels) n = 3 has s, p & d sublevels(3 sublevels) n = 4 has s, p, d & f sublevels(4 sublevels) - Maximum electrons in each sublevel: s holds up to 2 electrons1 orbital p holds up to 6 electrons3 orbitals d holds up to 10 electrons5 orbitals f holds up to 14 electrons7 orbitals

17. 17 Summary: Principle Levels & Sublevels for a Neutral Zn atom (30 electrons). s p n = 2 s d p n = 3 f s d p n = 4 s n = 1 Energy 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 When writing out electron configurations - you must write them in the order of lowest energy to highest energy orbitals.

18. 18 Electron Configurations Shorthand notation which shows the distribution of electrons among sublevels When we write electron configurations, we write the number of the principal quantum number followed by a symbol (s,p,d or f) for the sublevel, and then add a superscript to each sublevel symbol to designate the number of electrons in that sublevel. Carbon has 6 electrons. Therefore, using the orbital diagram we obtain: 1s 2 2s 2 2p 2 7 -

19. 19 Practice – Electron Configurations Write electron configurations for the following: –Al –Sc –K –Br –Hg

20. 20 Practice – Electron Configurations Write electron configurations for the following: –Al 1s 2 2s 2 2p 6 3s 2 3p 1 –Sc1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 –K1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 –Br1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 –Hg 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10

21. 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d6f 7s7p7d7f 1s 2 1s 2 2s 2 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 etc

22. 22

23. 23 Periodicity of Electron Configurations Can you identify the following elements? –Alkali Metals ____1s 2 2s 1 ____ 1s 2 2s 2 2p 6 3s 1 –Alkaline Earth Metals ____ 1s 2 2s 2 2p 6 3s 2 ____ 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 –Halogens ____ 1s 2 2s 2 2p 6 3s 2 3p 5 ____ 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 –Noble Gases ____ 1s 2 2s 2 2p 6 ____ 1s 2 2s 2 2p 6 3s 2 3p 6 7 -

24. 24 Periodicity of Electron Configurations Can you tell the patterns among the following groups of elements? –Alkali Metals (Group IA (1)) Li1s 2 2s 1 Na1s 2 2s 2 2p 6 3s 1 –Alkali Earth Metals (Group IIA (2)) Mg1s 2 2s 2 2p 6 3s 2 Ca1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 –Halogens (Group VIIA (17)) Cl1s 2 2s 2 2p 6 3s 2 3p 5 Br1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 –Noble Gases (Group VIIIA (18)) Ne1s 2 2s 2 2p 6 Ar1s 2 2s 2 2p 6 3s 2 3p 6 7 -

25. 25 Using the Periodic Table for Electron Configurations The periodic table can be used to fill orbital diagrams or to find electron configurations. –First, we need to separate the periodic table into blocks. Blocks contain elements with the same highest-energy sublevel. 7 -

26. 26 The Principal Quantum Number and Sublevel on the Periodic Table 7 - Figure 7.21