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Lesson 3 - Electronic configuration

Lesson 3 - Electronic configuration. Definition. Shorthand notation for the amount of electrons in an orbital. It all tells the arrangement of electrons in the orbitals . How is it distributed?. Distribution of all electrons in an atom consist of: 1. number denoting the energy level

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Lesson 3 - Electronic configuration

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  1. Lesson 3 - Electronic configuration

  2. Definition • Shorthand notation for the amount of electrons in an orbital. It all tells the arrangement of electrons in the orbitals

  3. How is it distributed? • Distribution of all electrons in an atom consist of: 1. number denoting the energy level 2. Letter denoting the type of orbital • Superscript denoting the number of electrons in those orbitals.

  4. Electron Configuration 1s1 group # # valence e- possibilities are: s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- should equal Atomic # row # Energy level # possibilities are 1-7 7 rows Subshell/ orbital possibilities are s, p, d, or f 4 subshells What element has an electron configuration of 1s1?

  5. Practice: atomic # = 3 1s2 2s1 • Lithium: what is the electron configuration? • Boron: what is the electron configuration? 1s2 2s2 2p1 atomic # = 5 • What is the electron configuration for the following? • Sodium Sodium(11) • Chlorine  Chlorine (17) • Argon  Argon (18) • Calcium  Calcium (20) 1s2 2s22p63s1 1s2 2s22p63s23p5 1s2 2s22p63s23p6 1s2 2s22p63s23p64s2

  6. Order of Electron Subshell Filling:It does not go “in order” 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6

  7. 1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p

  8. Rules that govern distribution. • Pauli Exclusion Principle • 2. Aufbau Principle • 3. Hund’s Rule

  9. Pauli Exclusion Principle • No two electrons in the same atom can have exactly the same energy. • For example, no two electrons in the same atom can have identical sets of quantum numbers. • No two electrons can have the same spin

  10. Aufbau Principle • Build the electronic structure starting at the lowest energy level • s < p < d < f

  11. Hund’s Rule Spread electrons as much as possible before pairing When doing orbital diagrams

  12. Orbital Diagrams • Each box/ circle represents one orbital. • • Half-arrows represent the electrons. • • The direction of the arrow represents the spin of the electron.

  13. Orbital Diagram for A Nitrogen Atom N 1s2 2s2 2p3        F 1s2 2s2 2p5         

  14. Learning check • Draw orbital diagram for magnesium and iron

  15. Energy Diagram of orbitals

  16. Special Electronic Configurations • A half filled subshell and a full filled subshell lower the energy, gaining some stability. special electronic configurations. • Cr [Ar]4s1 3d5  All s and d subshells are half full • Cu [Ar]4s1 3d10 Prefers a filled d subshell, leaving s with 1

  17. 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.

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