Development of the Periodic Table- Dmitri Mendeleev. Elements on the Periodic Table are arranged by: (2) increasing atomic #, or protons
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1 Chapter 8: Electron Configuration and Chemical Periodicity Development of the Periodic Table- Dmitri Mendeleev Elements on the Periodic Table are arranged by: (1) similar chemical properties (2) increasing atomic #, or protons
2 Nucleus (center of the atom) contains protons and neutrons. Electrons are located in orbitals outside the nucleus within energy levels, n =1,2,3..
3 Energy levels contain sub-energy levels: s orbitals (one per energy level) p orbitals (three per energy level for > n=2) d orbitals (five per energy level for > n=3) f orbitals (seven per energy level for > n=4) Each orbital can hold a maximum of 2 electrons
4 Review: - Quantum number n defines the principal energy level - Quantum number l defines the sublevel type - Quantum number m l defines the orientation of the orbital - Quantum number m s defines the spin
5 Let s look at another feature of the periodic table The blocks/groupings of the elements.
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8 Consider the number of electrons in an oxygen atom. There are 8 electrons. Where is each one located in terms of energy levels and orbitals? Consider Silver, it has 47 electrons, where are they located?
9 Electron Configurations allow us to predict where the electrons of a ground state atom are located in terms of energy levels and orbital types. Let s begin with the following: H, He, Li, Be, O, Ne, Mg, S Your turn, on the board: Ca, Mn, Ge, Ag
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11 A few other pieces of information you ll need: (1) Valence Electrons: Electrons in the outermost energy level; the electrons involved in compound formation. (2) Inner or Core Electrons: Electrons that fill all of the energy levels of an atom present in the previous Noble Gas. (3) Condensed Electron Configuration: Use the previous Noble Gas to represent the core electrons and finish with the rest of the electrons. Examples....
12 F 1s 2 2s 2 2p 5 [He] 2s 2 2p 5 Condensed or Abbreviated Electron Configurations Ga 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 [Ar] 4s 2 3d 10 4p 1 Te 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 4 [Kr] 5s 2 4d 10 5p 4
13 Some specific rules....
14 The Pauli Exclusion Principle: No two electrons in the same atom have the same set of four quantum numbers (n, l, m l, m s ). Aufbau Principle: Electrons fill energy levels and orbitals (sublevels) with the lowest orbitals filling first and proceeding to the next lowest orbital (sublevel). Hund s Rule: Electrons fill equal energy orbitals one at a time before pairing.
15 Energy level n=1 1H 1s 1 1s 2He 1s 2 1s
16 3Li 4Be 1s 2s 8O 10Ne 1s 2s 1s 2s 2p 1s 2s 2p
17 12Mg 3s 3p 16S 3s 3p
18 20Ca 25Mn 32Ge 47Ag 4s 3d 4p 4s 3d 4p 4s 3d 4p 5s 4d 5p
19 Let s go back through the previous orbital diagrams and assign quantum numbers.
20 What is the set of quantum numbers of the last entering electron of Molybdenum? What is the set of quantum numbers of the first entering electron in the p orbitals of Arsenic? What is the set of quantum numbers of the fifth entering electron of the highest energy d orbitals of Arsenic?
21 Atoms with completely filled energy levels are very stable (Noble gases) Atoms with completely filled sub-levels have greater stability than those with incomplete sub-levels Atoms with half-filled sub-levels have relatively greater stability than incomplete sub-levels, but not as stable as completely filled sub-levels.
22 Increasing Stability Completely filled energy level Completely filled energy sub-level Half filled energy sub-level Incomplete energy sub-level
23 Cr Predicted: [Ar] 4s 2 3d 4 Actual: [Ar] 4s 1 3d 5 Mo Predicted: [Kr] 5s 2 4d 4 Actual: [Kr] 5s 1 4d 5 Cu Predicted: [Ar] 4s 2 3d 9 Actual: [Ar] 4s 1 3d 10 Ag Predicted: [Kr] 5s 2 4d 9 Actual: [Kr] 5s 1 4d 10 Note: These deviations only occur in the d & f orbitals and not in the s or p orbitals.
24 Paramagnetic: The tendency for a particle with at least one unpaired electron to be attracted to a magnetic field. Diamagnetic: The tendency for a particle not to be attracted toward a magnetic field. This is caused by an atom having only paired electrons.
25 Draw an orbital diagram of the outermost electrons of the following elements: H He S Ni Se Ag Determine if the atoms above are paramagnetic or diamagnetic.
26 Write the (a) full electron configuration, (b) the orbital diagram, (c) magnetic property designation for each of the following: S 2- Ca 2+ N 3- Cu + Cu 2+ Hint: Electrons leave from the highest energy sublevels first.
27 Similar electron configurations Group I --- ns 1 Group II --- ns 2 Group VII --- ns 2, np 5 Group VIII --- ns 2, np 6
28 GENERAL TRENDS in the Periodic Table Atomic size decreases left to right in a period and increases going down in a group.
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31 Ionization Energy (IE) increases left to right in a period and decreases going down in a group. A (g) + energy A + (g) + e - IE 1 < IE 2 < IE 3, etc.
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33 Compare IE 1 (MJ/mol) for second period atoms Li (0.52) Be (0.90) B (0.80) C (1.09) N (1.40) O (1.31) F (1.68) Ne (2.08)
34 Electron Affinity (EA) is irregular in groups and periods. A (g) + e - A - (g) + energy Metallic character of elements increases from right to left and down the periodic table.
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