# Electron Configurations Notes

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```					Electron Configurations Notes
Homework (page 166) – 35, 37, 38, 46, 47, 48, 53, 60, 66, 67, 70, 72, 77, 78, 79, 81, 82, 83, 85, 86, 87, 88, 90, 91, 92,

1. See section 5.1 in book to know the following terms and conditions. a. Wavelength (400nm-700nm is the visible spectrum) b. Speed of Light is 3.00 x 108 m/s c. Frequency is waves/second and measured in hertz d. Relationship between wavelength and frequency (c = wavelength / frequency) e. Frequency of waves and their ability to travel through materials f. ROYGBIV and the electromagnetic spectrum

2. Quantum Theory a. Planck’s Theory or Quantum Theory i. There is a fundamental restriction on the amount of energy that an object can emit or absorb ii. Energy is found in tiny packets called quanta iii. Quantum of energy is different for different frequencies iv. Relationship between frequency and energy 1. E = hν, where h = Planck’s constant (6.6262 x 10-34J.s) 2. As frequency goes up, energy goes up 3. Photoelectric Effect a. When light of certain wavelength hits a metal, the metal will emit electrons i. All metals have a threshold of the amount of energy they need ii. Light is made of particles called photons, which each have a certain amount of energy iii. “All or none” response iv. Dark rooms, X rays, old TV rays, cell phones 4. Dual Nature of Light a. Photons exhibit a wavelike pattern. Therefore, it acts as a wave. b. If photon of light hits and electron, it bounces off. Therefore it is a particle 5. Spectra – Atomic emission spectra a. Continuous spectrum – All wavelengths of light are present. True white light. b. Line spectra – Only certain wavelengths of light are present.

6. Three theories dealing with electrons a. Bohr model of the atom (quantizing of atom) i. Electrons are found in shells ii. Lowest state an electron is normally found is called the ground state iii. When electron gains energy, it is in an excited state iv. When electron goes from excited state to ground state, it releases E=hv b. DeBroglie’s Matter Waves (matter waves) i. All matter moves in a wavelike pattern, but only small matter can be measured c. Heisenberg’s Uncertainty Principle (momentum and position of an electron) i. Momentum and position of an electron cannot BOTH be determined d. Schrodinger Wave Equation i. Took Bohr’s model and applied the other two theories to come up with the quantum mechanical model. 7. Quantum Mechanical Model of the Atom a. Electrons are found in shells (principle energy levels) and are called quantum numbers (n) b. Shells are broken down into sublevels (l) – s=0, p=1, d=2, f=3 i. Number of sublevels = quantum number ii. s, p, d, and f are the levels 1. n = 3 can have s, p, and d c. Sublevels are broken down into orbitals (ml) +/- l i. The sublevel determines the number of orbitals (d = +2, +1, 0, -1, -2 or 5 orbitals) ii. s = 1, p = 3, d = 5, f = 7 d. Three theories that govern electron orbitals i. Aufbau Principle – electrons start filling an atom at its lowest level of energy and build up to the highest. ii. Pauli Exclusion Principle (ms)– 2 electrons per orbital and they spin in opposite directions to counteract magnetic spin. iii. Hund’s Rule – Electrons must fill up equal energy orbitals before pairing up in the same sublevel. e. Example – Theoretical amount of electrons that can fit into the 3rd shell i. N = 3  s, p, and d  1, 3, and 5 orbitals  2, 6, and 10 electrons = 18 electrons

8. Electron configurations - A way to show where every electron is found in an atom a. 1s2 – quantum number, sublevel, number of electrons in that sublevel b. Diagonal Rule – Shows the order of sublevels following the Aufbau Principle c. Example i. Mg = 12 protons and 12 electrons 1. 1s22s22p63s2 ii. O = 6 protons and 6 electrons 1. 1s22s22p2 d. Exceptions i. It’s more stable for atoms to have the last sublevel full or half full than to be almost full. ii. Chromium – expect 4s23d4, but it is actually 4s13d5 9. Orbital Diagrams a. Only show the s and p sublevels of the last shell…show them both though b. If electron configuration ends in a d or f sublevel, show only the d or f sublevel c. Example i. O = 1s22s22p2 ii. 2s _____ 2p _____ 10. Noble Gas Notation a. Replace all electron configurations that lead up to a noble gas with [Noble Gas] b. Labeling the following: i. [Ne]3s2  [Ne] = noble gas inner core, 3 = energy level, s = sublevel, 2 = electrons 11. Electron Dot Structures a. Place dots around chemical symbol to show number of valence electrons _____ _____

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