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Chapter 4 Notes Sec 4-1 Sec. 4-1 The Rutherford model (or the planetary model) of the atom was an improvement over the previous models, but it was still incomplete. It didn’t explain where the e- were located in the space surrounding the nucleus. Before we can understand what came next, we must first learn about light. Electromagnetic Spectrum Electromagnetic radiation is a form of energy that travels through space. All electromagnetic radiation (or waves) travel at 3.0 x 108 m/s. This is called the speed of light. In the early 1900’s, scientists conducted 2 experiments involving light and matter. The first involved an phenomenon known as the photoelectric effect, where metals will release e- when a light shines on them. Ex. Solar cells, photovoltaic cells, etc. In 1900, a German physicist, Max Plank, suggested that the object emits energy in small, specific amounts of quanta. A quantum is the minimum quantity of energy that can be lost or gained by an atom. (Quantum Mechanics is now a branch of physics) In 1905, Albert Einstein expanded on Plank’s theory and said that any electromagnetic radiation exhibited properties of waves and particles. This was called the Wave-Particle Duality. Light now behaved as a wave and as a particle. Each particle of light carries a quantum of energy called a photon. Einstein explained the photoelectric effect by proposing the electromagnetic radiation is absorbed in whole #’s of photons. In 1913, Niels Bohr came to work with Rutherford. Rutherford knew that there were errors in his model. Bohr then proposed the atom’s e- was linked to the photon emissions. He included that the ecan circle the nucleus in allowed paths or orbits. E1 is called ground state E3 E1 E2 nucleus When 1 photon hits an atom, it becomes excited and moves to a higher energy state. We call this absorption. E3 E1 E2 When an electron gets so excited that it “pops” off of the atom it becomes free and creates electricity. Emission is when an excited electron falls back down. E3 E1 E2 Excited neon atoms emit light when falling back to ground state or to a lower energy level. Each element gives off a distinct band of light when it is excited. Sec 4.2 Because the wave – particle theory was still being debated, Bohr’s model of the atom wasn’t readily accepted. Investigations into Einstein’s idea of the Wave-Particle Duality led to a revolution of our basic idea of matter. In 1924, Louis de Broglie showed that the e- were waves by showing that they exhibited diffraction and interference. In 1927, Werner Heisenberg came up with his uncertainty principle. He said it is impossible to determine simultaneously both the velocity and position of an e- . This was very hard to swallow In 1926, Erwin Schrodinger developed an equation that treated e- as waves. Together, the Heisenberg uncertainty principle and the Schrodinger equation laid the foundation for the Quantum theory. Heisenberg uncertainty principle gives us a probability of where to find the e-. The e- does not travel around the nucleus in neat orbital’s (like Bohr thought). They exist in certain regions called orbital's. We still use the planetary model to show two dimensions, but the orbital’s are actually 3 dimensional. Quantum #’s – Specify the properties of the atomic orbital's and the properties of the e- in the orbital's. 4 quantum #’s – n, l, ml, ms “n” - Principal Quantum #– indicated the main energy level occupied by the e-. “l” – Angular Momentum # - indicated the shape of the orbit. There are 4 shapes – s, p, d, and f. “ml” – magnetic quantum # -orientation of the orbital. Is it located in the x, y, z position, etc. “ms”– Spin - shows the spin of the electron Can be +1/2 or -1/2 for clockwise and counterclockwise (overheads) Sec 4.3 Pauli’s Exclusion Principle – e- carry a spin, 1 cw and 1 ccw, which create a magnetic field. Aufbau’s Principle – e- must occupy the lowest energy level possible. Hund’s Rule – e- fill up each orientation (x,y,z, etc) before they pair up. Electron Configuration -