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Electron Configuration Ch 11

Electron Configuration Ch 11. review. What do we know about electrons? How many electrons does a neutral atom contain? To learn: How are these electrons arranged around the nucleus?. Development of Atomic Theory. Elements (Boyle – 1600’s)

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Electron Configuration Ch 11

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  1. Electron ConfigurationCh 11

  2. review • What do we know about electrons? • How many electrons does a neutral atom contain? • To learn: • How are these electrons arranged around the nucleus?

  3. Development of Atomic Theory • Elements (Boyle – 1600’s) • Elements are made up of atoms, which cannot be further broken down. All atoms of elements are exactly alike. (Dalton early 1800’s) • Atoms with positive & negative charges mixed together (JJ Thompson – mid 1800’s) • Atoms with solid nucleus & electrons on outside (Rutherford – later 1800’s) • Electrons in concentric orbits, like planets, around nucleus (Bohr – early 1900’s)

  4. Electron locations – Bohr model • First electrons closest to nucleus • Only room for two – why? • Draw the Bohr model for hydrogen, helium • Next “orbit” holds 8 electrons • Draw the Bohr model for nitrogen, neon • 3rd “orbit” holds 8 electrons • Draw Bohr model for magnesium, sulfur • What is last atom with three levels?

  5. Development of Atomic Theory – cont’d • Wave-mechanical model (deBroglie – 1920’s) • Firefly analogy – orbitals are not orbits • Orbitals = 3-dimensional region in which there is a high probability of finding an electron in an atom

  6. Electron Configuration – rulesSee Ch 11, section 4 • Rules of filling: • a maximum of 2 electrons per orbital • Electrons in any one orbital spin in opposite directions. • Electrons “fill in” from the lowest energy level to the highest. • Periods correspond to energy levels for s & p blocks. The d block energy level = period -1. Format: 3d4, where the 3 = energy level, • d = orbital shape, • 4 = number of electrons in the orbital

  7. Orbital configuration • Much like electron configuration, but also: • Shows spin • Shows individual orbitals within blocks • Shows how electrons fill in • 3 p orbitals in each energy level: x, y & z • First: px1 py1 pz1 • Then: px2 py2 pz2

  8. Valence electrons • Valence electrons = electrons at the highest energy level • Why aren’t d orbitals ever valence electrons? • b/c they always fill in at one lower energy level than the valence s orbital for that element (period – 1) • What is the maximum number of valence electrons an element can ever have? • 8 (two from s + 6 from p)

  9. Lewis Dot Structures • Named after G.N Lewis (1902) • Shorthand way of showing valence electrons • Dots around element symbol • s orbital first (two dots on top) • p orbitals next, one at a time going around symbol. HW: draw the lewis dot structures for all the elements 11-18 & 19 - 36

  10. Electrons and Energy • Energy as light – waves or particles? • Both or either • Called wave-particle duality • Photon = “particle” of light (electromagnetic radiation) • When electrons absorb energy, they get “excited” and jump to a higher energy state • When the electrons go to lower energy state: • Emit photons of light

  11. Electromagnetic Spectrum • Different elements emit different visible colors • Why? • Because the different energy levels correspond to different wavelengths

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