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 Bohr’s single quantum number (n) was expanded to a total of four quantum numbers n, l, m l, and m s  These four quantized values describe an electron.

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Presentation on theme: " Bohr’s single quantum number (n) was expanded to a total of four quantum numbers n, l, m l, and m s  These four quantized values describe an electron."— Presentation transcript:

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2  Bohr’s single quantum number (n) was expanded to a total of four quantum numbers n, l, m l, and m s  These four quantized values describe an electron in an atom (quantized values are restricted to certain discrete values)  These values add order to our description of the electron in the atom

3 Table 1: Summary of Quantum Numbers

4 Table 2: Values and Letters for the Secondary Quantum Numbers

5 OrbitsOrbitals 2-D path3-D region in space Fixed distance from nucleus Variable distance from nucleus Circular or elliptical path No path; varied shape or region 2n 2 electrons per orbit2 electrons per orbital

6 Table 4: Energy Levels, Orbitals, and Shells  The first two quantum numbers (n and l) describe electrons with different energies under NORMAL circumstances  The last two quantum numbers (m l and m s ) describe electrons with different energies under SPECIAL conditions (e.g. magnetic field)

7  Moving forward, we will be focusing on the electrons position in space (not energy), the language will change ◦ Main (principal) energy level = shell ◦ Energy sublevel = subshell ◦ WHY? ◦ Its easier! ◦ 1s orbital can be communicated as n=1, l=0 ◦ 2p orbital can be communicated as n=2, l=1

8 Table 5: Classification of Energy Sublevels (subshells)

9  Figure 1: Energy- level diagrams show the relative energies of electrons in various orbitals under normal conditions (each orbital can hold a maximum of 2 e - )

10  The energy of an electron increases with an increasing value of principal quantum number, n  For a given number of n, the subshells increase in energy, in order, s

11  Pauli Exclusion Principle – no two electrons in an atom can have the same four quantum numbers; no two electrons in the same atomic orbital can have the same spin, only two electrons with opposite spins can occupy any one orbital  What order do we fill the orbitals?  Aufbau Principle – each electron is added to the lowest energy orbital available in an atom or ion  An energy sublevel must be filled before moving onto the next higher sublevel

12 Figure 3: In this aufbau diagram, start at the bottom (1s) and add electrons in the order shown by the diagonal arrows. You work your way from the bottom left corner to the top right corner.

13 Figure 4: Classification of elements by the sublevels that are being filled

14  Hund’s Rule – one electron occupies each of the several orbitals at the same energy before a second electron can occupy the same orbital SEATWORK  Read pp. 189 – Drawing energy-level diagrams for atoms, anions, cations  Practice p. 191 UC # 3, 4

15  If you’re thinking this is too easy to be true, you’re right!  There are a few complications as the atoms get larger  As the energy level gets farther from the nucleus, the distance between energy levels decreases  As a matter of fact, it is believed that the energy levels actually overlap

16  Therefore, some energy levels start filling orbitals before the previous energy level is finished filling its subshell  The first time this is encountered is with potassium, in which the 4s starts to fill before the 3d

17  The second complication has to do with a variation of Hund’s Rule that takes into account the minimizing of the electron- electron repulsion  It states, the most stable arrangement of electrons is the arrangement with the maximum number of unpaired electrons.  So, when the transition metals’ orbitals are filling with electrons, at d 4 and d 9, an electron from the s JUMPS up into the d 5 and d 10

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19  Overall energy state of the atom is lower after the promotion of the electrons  Half-filled and filled subshells are more stable (lower energy) than unfilled subshells

20  A method of communicating the location and number of electrons in electron energy levels (presents same information as energy-level diagrams BUT much more concise) Figure 5: Example of electron configuration

21 Writing Electron Configurations The electron configuration below represents a boron atom in its ground state. The superscripts indicate the number of electrons occupying each sublevel.

22  Writing out electron configurations can become awkward as the atoms increase in the number of electrons  The shorthand involves using the abbreviation of the last noble gas (placed in square brackets) to indicate that all the orbitals to that point are full. Then the configuration is continued as usual.

23 Nitrogen 1s 2 2s 2 2p 3 [He] 2s 2 2p 3 Chromium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 [Ar] 4s 2 3d 4 Strontium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 [Kr] 5s 2

24  Learning Checkpoint  Read pp. 192 – 193 ◦ Understand FULL electron configuration and Shorthand (NOBLE GAS CORE) electron configurations ◦ Add the summary for “Procedure for Writing an Electron Configuration” on p. 193  Complete “Electron Configuration” worksheet  Practice Questions p. 194 UC # 6, 8, 9, 10  Section 3.6 Questions p. 197 UC # 2, 3, 4, 5, 6, 7, 8, 9 10, 11, 12, 13, 14 24


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