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SCH4U G Ra de 12 C He mistr Y http://www.youtube.com/ watch?v=-d23GS56HjQ Quantum Mechanics The application of quantum theory to explain the properties of matter, particularly electrons in atoms Schrodinger’s Standing Waves Louis De Broglie developed a theory that matter can have wave-like properties Schrodinger extended this theory to electrons bound to a nucleus Postulated that electrons resembled a standing wave Certain orbitals exist at whole wavelengths of electron vibrations Orbitals - Redefined Orbital: region around the nucleus where there is a high probability of finding an electron As per wave model of Schrodinger – because things are vibrating Heisenberg Uncertainty Principle Heisenberg Uncertainty Principle Heisenberg studied statistics and developed matrix algebra Developed a statistical approach to explaining how electrons works and realized… IT IS IMPOSSIBLE TO KNOW THE EXACT POSITION AND SPEED OF ELECTRON AT A GIVEN TIME At best, we can describe the probability of finding it at a specific place Wave functions: the mathematical probability of finding an electron in a certain region of space Wave functions give us: Electron probability densities: the probability of finding an electron at a given location, derived from wave equations Homework Quantum Numbers Quantum Numbers: numbers that describe the quantum mechanical properties (energies) of orbitals From the solutions to Schrodinger’s equation The most stable energy states is called the ground state Principal Quantum Number (n) Integer number (n) used to level the main shell or energy level of the electron Describes size and energy of the atomic orbital Increase number = increase energy, bigger Secondary Quantum Number, l Describes the shape of the orbital within each shell Each energy level contains several sublevels Relates to the shape of the orbital Can be any integer from 0 to (n-1) Values of l Value 0 1 2 3 4 Letter Used s p d f g Name sharp principal diffuse fundamental Each orbital is given a code: Example If n = 1, l = 0 then we call it a 1s orbital If n = 3, l = 2 then we call it a 3d orbital Magnetic Quantum Number, ml Describes the orientation of the orbital in 3- space Can be whole number integers from – l to + l Example: if l = 1, then ml can be -1, 0, +1 There are 3 possible p orbitals px, py, and pz What are possible values for ml if l is: 0 1 2 3 Spin Quantum Number Electrons are basically little magnetics spin around when placed in magnetic fields, they can have spin ‘up’ or spin ‘down’ ms can be either +1/2 or – 1/2 Homework Electron Configurations and Energy Level Diagrams The four quantum numbers tell us about the energies of electrons in each atom Unless otherwise stated were are talking about ground state energies Energy Diagrams Describe how electrons fill orbitals using quantum numbers Electrons fill the lowest energy level orbitals first Each shell is (for the most part) filled before moving to higher shells Rules Use circles (or boxes) to represent each orbital in any given energy level and arrows for electrons Unoccupied circles imply that there are no electrons in it A circle can have at most two electrons in it; only if the arrows are pointing in opposite directions Rules to Remember Pauli exclusion Principle: no two electrons can have the same 4 quantum numbers. Electrons in the same orbital can’t have the same spin Hund’s Rule: One electron occupies each of sub-orbitals in the same energy level before a second can occupy the same sub-orbital Aufbau Principle: each electron is added to the lowest energy orbital available Building Orbital Diagrams Practice H, B, C, Ne Mg, P, Ar Ca, Mn, Zn, Ge, Kr Electron Configurations Condensed versions of orbital diagrams Write the electron configuration for each of the atoms above Exceptions to the Rules Examine the allowed charges for Chromium and Copper Write the electron configuration for chromium and copper What actually happens? Why? Evidence suggests that half-filled and filled orbitals are more stable than other orbitals, so electrons rearrange to give these configurations Explaining Ion Charges In order to get particular charges, entire energy levels or sublevels get cleared first. Use electron configuration theory to explain why: Zn Zn2+ Pb Pb2+ or Pb4+ Explaining Trends in the Periodic Table Atomic Radius: size of the atom Ionization Energy: energy needed to remove an electron from the outermost energy level from an electron in the gaseous state Electron Affinity: change in energy that occurs when an electron is added to a gaseous atom