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Unit 3-3 Electron configuration

Unit 3-3 Electron configuration. Chap 8. Electron Configurations. Electron configurations tells us in which orbitals the electrons for an element are located. Three rules: electrons fill orbitals starting with lowest energy level (n=1) and moving upwards (aufbau principle);

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Unit 3-3 Electron configuration

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  1. Unit 3-3 Electron configuration Chap 8

  2. Electron Configurations • Electron configurations tells us in which orbitals the electrons for an element are located. • Three rules: • electrons fill orbitals starting with lowest energy level (n=1) and moving upwards (aufbau principle); • no two electrons can fill one orbital with the same spin (Pauli’s exclusion); • for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule). • The orbitals in the same subshell have equal energy . Thus they are called degenerate orbitals.

  3. Pauli Exclusion Principle • no two electrons in an atom may have the same set of 4 quantum numbers • therefore no orbital may have more than 2 electrons, and they must have with opposite spins • knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel • s sublevel has 1 orbital, therefore it can hold 2 electrons • p sublevel has 3 orbitals, therefore it can hold 6 electrons • d sublevel has 5 orbitals, therefore it can hold 10 electrons • f sublevel has 7 orbitals, therefore it can hold 14 electrons Tro, Chemistry: A Molecular Approach

  4. Spin Quantum Number, ms • spin quantum number describes how the electron spins on its axis • clockwise or counterclockwise • spin up or spin down • spins must cancel in an orbital • paired • mscan have values of ±½

  5. Aufbau Principle an electron occupies the lowest-energy orbital that can receive it. Finding the order of Subshell Fillingin Ground State Electron Configurations

  6. 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Energy • Notice the following: • The orbitals within a sublevel have the same energy (degenerate). But sublevels within an energy level are NOT degenerate. • From the 4th and higher energy levels, their s sublevel is lower in energy than the d sublevel of the previous energy level • the energy difference between levels becomes smaller for higher energy levels

  7. Finding the order of Subshell Fillingin Ground State Electron Configurations Start by drawing a diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) next, draw arrows through the diagonals, looping back to the next diagonal each time You must know how to draw the filling order!

  8. Hund’s Rule • “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.” • Within a subshell, electrons fill each orbital singly before any orbital gets a second electron

  9. Electron Configurations notation • A shorthand method of writing the location of electrons by sublevel. • The sublevel is written followed by a superscript with the number of electrons in the sublevel. • If the 2p sublevel contains 2 electrons, it is written 2p2

  10. Writing Electron Configurations practice Write e- configurations of Na and Mg: • First, determine how many electrons are in the atom. • Arrange the energy sublevels according to increasing energy: • 1s 2s 2p… • Fill each sublevel with electrons until you have used all the electrons in the atom: • Na: 1s2 2s2 2p6 3s1 • The sum of the superscripts equals the atomic number of the element

  11. Noble Gas Notation • short hand for larger atoms • configuration for the last noble gas is abbreviated by the noble gas’s symbol in brackets

  12. Orbital diagram Write e- configuration diagram of Si: • number of electrons: 14 • last electron is in sublevel: 3p 1s 2s 2p 3s 3p • Valence Electrons- the electrons in the outermost energy level

  13. Electron Configuration of Atoms in their Ground State Tro, Chemistry: A Molecular Approach

  14. Electron configuration of the elements of the first three series

  15. Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table, then correspond to different types of orbitals that are not fulfilled by e-.

  16. Valence Electrons • the electrons in the highest principal energy shell and are furthest away from the nucleus are called the valence electrons • electrons in lower energy shells are called core electrons • When an atom undergoes a chemical reaction, only the outermost electrons are involved. • For the representative elements (s and p blocks), the valence electrons are the s and p electrons beyond the noble gas core.

  17. Predicting Valence Electrons • The Roman numeral in the American convention indicates the number of valence electrons. • Group IA elements have 1 valence electron • Group VA elements have 5 valence electrons • When using the IUPAC designations for group numbers, the last digit indicates the number of valence electrons. • Group 14 elements have 4 valence electrons • Group 2 elements have 2 valence electrons

  18. Lewis Electron Dot Formulas • An electron dot formula of an elements shows the symbol of the element surrounded by its valence electrons. • We use one dot for each valence electron. • Consider phosphorous, P, which has 5 valence electrons. Here is the method for writing the electron dot formula.

  19. Ion Electron Configurations in their ground states • When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na+ + e- 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 • When we write the electron configuration of a negative ion, we add one electron for each negative charge: O + e-→ O2- 1s2 2s2 2p4→ 1s2 2s2 2p6 • Write Lewis dot structures of Na+ and O2-

  20. Ionic Charge • Recall, that atoms lose or gain electrons to form ions. • Atom lose e-  positive charge  cation • Atom gain e-  negative charge  anion • The charge of an ion is related to the number of valence electrons on the atom. • Group IA/1 metals lose their one valence electron to form 1+ ions. • Na → Na+ + e- • Metals lose their valence electrons to form cation. Metals can NOT gain e- , therefore cannot form anion! • Nonmetals gain e- and become monatomic anion when forming binary ionic compound (XY) with metal

  21. Predicting Ionic Charge • Group IA/1 metals form 1+ ions, group IIA/2 metals form 2+ ions, group IIIA/13 metals form 3+ ions, and group IVA/14 metals from 4+ ions. • By losing their valence electrons, they achieve a noble gas configuration. • Similarly, nonmetals can gain electrons to achieve a noble gas configuration. • Group VA/15 elements form -3 ions, group VIA/16 elements form -2 ions, and group VIIA/17 elements form -1 ions.

  22. Homework • Write both the electron configurations and orbital diagrams of all the atoms from period 1 to 3. • Page 359 46 a, c and d, 48 a, 50 b, 52, 53, 54

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