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Modern Physics

Modern Physics. Quantum mechanics, Periodic Trends, atomic orbitals, electron configuration, By Zoe Poncher and Abby Zemach. Quantum Mechanical Model of the Atom.

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Modern Physics

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  1. Modern Physics Quantum mechanics, Periodic Trends, atomic orbitals, electron configuration, By Zoe Poncher and Abby Zemach

  2. Quantum Mechanical Model of the Atom • Summary: Theory that electrons have both wave like and particle properties, and reside in orbitals (each of which hold two electrons). • Names: • Heisenberg – Uncertainty Principle – both position and momentum cannot be known at the same time • Louis De Broglie – De Broglie Eq = λ = h/mv • Erwin Schrodinger—created equation that showed the quantum state, or wave function, of a particle as its position or time interval changed • Max Planck- E = hv=h(c/λ): Planck’s Constant = 6.63χ10-34

  3. Equations De Broglie’s Equation λ Plank’s Constant Avogadro's Number Speed of Light Mass of Electron

  4. History of Atom • Antoine Lavoisier – 1774 Law of Conservation of Matter • Joseph Proust – Law of Constant Composition • John Dalton - Atomic Theory & Law of Multiple Proportions • Michael Faraday-demonstrated electric nature of elements • Sir William Crookes – Cathode ray tube • JJ Thomson – Discovered Electrons- oil drop experiment • Ernest Rutherford – nuclear model of atom w/ Gold Foil exp. • Discovered Proton

  5. Quantum Model • Max Planck – German physicist (1858-1947) hypothesized that energy could be released or absorbed by atoms in discrete “chunks” of minimum size . ‘Quantum meant the smallest quantity of energy that could be emitted or absorbed as electromagnetic radiation. E = hv: Planks Constant:6.62606957 × 10-34J/ s

  6. Photoelectric Effect • Albert Einstein used the Quantum theory to explain why energy acts like a particle when it impacts metal, compared to wave property- otherwise known as the Photoelectric Effect! Ephoton –Ethreshold = KEelectron Ephoton= hf = hc/λ

  7. Principles You NEED to Know Principles: • Pauli Exclusion Principle –in a given atom no two electrons can have the same quantum #’s / can’t have same spin direction. • Aufbau Principle-electrons fill lowest energy orbitals first, • Hund’s Rule- the lowest energy state- which is the most stable –is the one with the greatest # of unpaired electrons,

  8. Definitions • Atomic size: Atomic radii between nuclei and outer electron shell. • Ionization Energy: Energy required to remove an electron from an atom in the gas phase • Electron Negativity: The ability of an atom to attract electrons, *** Trends develop from the distance which the outer shell is from the nuclei, and the ratio of protons to electrons. Ie Oxygen is smaller than Nitrogen because the effective nuclear charge is greater (more protons) increasing the pull on the outer orbitals.

  9. Definitions Valence Electrons – Outermost electrons, Isoelectric – Atoms having same electron configuration Paramagnetic- Atoms with UNPAIRED electrons which are attracted to magnetic fields, Diamagnetism – contains only paired electrons ,

  10. Periodic Trends Size dec. Ionization Energy inc. Electron Negativity inc. Smallest Element Largest Element Size inc. Ionization Energy dec. Electron Negativity dec. *Cations are usually SMALLER*

  11. Electron Negativity Mnemonic Skill Diagonal Trend: Estimate Electron Negativities by knowing EN increase as you go up and across the periodic table. Electron Negativity Differences and Polarities ΔEN <.4 = Covalent Bond – non polar .4 < ΔEN <1.7 = < Polar Covalent – Polar w/ Dipole pointing to higher negativity ΔEN > 1.7 = Ionic Bond –inherently polar

  12. Atomic Orbitals! Wherever the valance electrons lie from electron configuration shows what orbital the last electron Is in, • S orbital • P orbital • D orbital • F orbital Electrons may be anywhere in the orbital, However, never in the origin!

  13. Electron Configuration • Electron Configuration is the representation of the arrangement of electrons • Full extended version • All of the periods, orbitals, and electrons that are to the specific element, ie Se • Se: 1s22s22p63s23s64s23d104p4 • The principal energy levels • The orbitals used • Number of electrons filled in the orbital • Abbreviated version • Use previous Noble gas as place holder for all periods and orbitals in previous periods • Se: [Ar] 4s23d104p4

  14. Electron Configuration • Please Note Electrons which are Excited jump to a higher energy level, meaning electrons are promoted, • When you promote electrons you take from the inner orbitals first, • Before • After • When you remove electrons you take from the outermost shell, • Before • After

  15. Quantum Numbers • n = Principal Quantum # = Period / energy level • l = Angular Momentum (azimuthal) Quantum # = • s =0, p=1 , d=2 , f=3 • Ml = Magnetic Quantum # =- l - l value, • ie. P orbitals have -1,0,1 #’s • Ms = Spin Quantum # = -½ , ½ depends on electron configuration ** To find the Quantum numbers you use a combination of looking at the Periodic table and orbital diagram** No two different elements have the same set of Quantum Numbers

  16. Quantum Numbers • What are the Quantum Numbers for Carbon? Step One – notice Carbon is in the 2nd Period and P’ orbital n=2 & l=1 Step Two – sketch orbital diagram and add number of valence electrons -Be Carful to Follow Aufbau’s and Hund’s Principals Ml= 0 (lands in 0’s spot) & Ms = ½ (I chose to establish arrow’s as upwards first)

  17. Sources • Periodic Table Diagram: http://www.chemicalelements.com/ AP Chemistry Crash Course: by Michael D’Alessia, Barron’s AP Chemistry Exam Book Chemistry The Central Science 11e

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