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Ch. 5 Atomic Models • Dalton – • Thomson – • Rutherford – Previous models could not explain the chemical properties of elements. • Why do different elements give off different colors when heated? • Why does the color change when more heat is added to an element? Bohr Model • Bohr proposed that an electron is found only in specific circular paths, or orbits around the nucleus. • Electrons are restricted to fixed (quantized) orbits • Electrons can jump orbits by absorbing or emitting a photon that has a specific wavelength of energy. 5 . 1 The Bohr Model Each possible electron orbit in Bohr’s model has a fixed energy. –The fixed energies an electron can have are called energy levels. –A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level. 5 The Bohr Model . 1Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced. The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level. The Quantum Mechanical Model –The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. The Quantum Mechanical Model • The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller. Modern Model of the Atom: Quantum Mechanical Model Erwin Schrodinger developed an equation to describe the energy of electrons “I don't like it, and I'm sorry I ever had anything to do with it.” (Erwin Schrodinger talking about Quantum Physics) 1887-1961 Atomic Orbital: a region of space in which there is a high probability of finding an electron. (S, P, D, F) S sublevel: sphere shaped Each S sublevel holds 1 orbitals P sublevel: dumbbell shaped Each P sublevel holds 3 orbitals D sublevel: clover leaf shape Each D sublevel holds 5 orbitals F sublevel: complex shape Each D sublevel holds 7 orbitals (Don’t worry, I’ll explain this later) 5 Atomic Orbitals . The numbers and kinds of atomic orbitals depend on 1 the energy sublevel. 5• The number of electrons allowed in each of . the first four energy levels are shown here. 1 Electron Arrangement in Atoms Chapter 5 section 2 Electron Configuration • The way electrons are arranged in various orbitals around the nuclei of an atom • Example: Arsenic 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3 Or (orbital diagram) 1s __ 2s __ 2p __ __ __3s___ 3p__ __ __ 4s__ 3d __ __ __ __ __ 4p__ __ __ 5 . 2 Electron Configurations • Orbital Filling Diagram 3 RULES FOR WRITING THE ELECTRON CONFIGURATION Aufbau Principle Electrons occupy the lowest energy level first. • The first energy level is the lowest energy level Pauli Exclusion Principle • An atomic orbital may describe and hold at most two electrons. • To occupy the same orbital electrons must have opposite spins; clockwise and counter clockwise • Example 1s___ Hund’s Rule • Electrons occupy orbitals of the same energy level so that one electron enters each orbital until all the orbitals contain one electron with the same spin. • Example 3p __ __ __ Exceptions to the Electron Configurations • Chromium • 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1 • Copper • 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1 Section 3 Physics and the Quantum Mechanical Model c=λf • The wavelength and frequency of light are inversely proportional to each other. • c = the speed of light 3.00 x 108 m/s • λ = wavelength in meters • f = frequency in Hertz (s-1 or cycles/second) All electromagnetic waves travel in a vacuum at 3.00 x 108 m/s What color of light has the highest frequency? • The amount of radiation emitted at each wavelength depends on the temperature of the object. Hot objects emit more of their light at short wavelengths, and cold objects emit more of their light at long wavelengths. 700 nm 400 nm • The amount of light produced at each wavelength depends on the temperature of the object producing the light.. Solid objects heated to 1,000 degrees C appear red but are putting out far more (invisible) infrared light. Max Planck analysis • The color of light from a hot glowing solid varies with temperature. • There is a relationship of energy in a solid and the wavelength of light emitted • Energy varies by small whole numbers • Any frequency of light should emit photons. (not so) Einstein • Light travels in packets of energy (photons) in waves • Light energy can be both particles and waves • Quantum – minimum amount of energy that can be lost or gained by an atom. Planck’s Equation (with the help of Einstein) The energy of a photon is described by the following equation. E=fh • E: joules • f : frequency • h : Planck’s constant 6.626x 10-34 j-s Each frequency of light has its own specific energy per photon and no other. Atomic spectra • Neils Bohr (1913) explains… • When atoms absorb energy, electrons move into higher energy levels. • These electrons then lose energy by emitting light when they return to lower energy levels. Bright line spectrum Consists of several distinct lines of color, each with its own frequency. Every element has a different bright line spectrum. Modern Model of H Atom Improving Bohr’s model Louis deBroglie (1923) Experimentally confirmed Einstein and Planck’s theory that particles (photons) could Have properties of waves. Won the Nobel Prize in 1929 The work done by de Broglie and Schrodinger became the branch of physics known as Quantum Mechanics It is often stated that of all the theories proposed in this century, the silliest is quantum theory. Some say the only thing that quantum theory has going for it, in fact, is that it is unquestionably correct. - R. Feynman Heisenberg Uncertainty Principle Werner Heisenberger It is impossible to know exactly both the velocity and position of a electron (particle) at the same time. Orbitals and Energy Levels Every energy level has a specific number of orbitals. Energy level # of orbitals 2 n n 1 1 2 4 3... 9... 4 quantum numbers used to describe an electron in an orbital 1. n = principle quantum number. It describes the size of the orbital. 2. l = designates the shape of the orbital. s = spherical p = double-lobed 4 quantum numbers used to describe an electron in an orbital 3. m = describes the orbitals orientation (around x-axis, y-axis, or z-axis). 4. Designates the spin of the electron.