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Bohr Model Nucleus: Neutrons and Protons Orbits: Electrons We know both specific energy and location of each electron Nucleus Energy Levels Electrons orbit the nucleus in certain fixed energy levels (or shells) Bohr Model Bohr’s Atomic Model of Hydrogen Bohr - electrons exist in energy levels AND defined orbits around the nucleus. Each orbit corresponds to a different energy level. The further out the orbit, the higher the energy level Bohr’s Model The Photoelectric Effect Light releases electrons Not all colors work Atomic Emission Spectra Hydrogen gas emitted specific bands of light Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 65 4 3 2 1 Electromagnetic Radiation Photoelectric Effect – There is a minimum frequency to eject the electron Electromagnetic Radiation Photoelectric Effect Only explained by “energy packets” of light called a quantum Quantum - minimum amount of energy that can be gained or lost by an atom Photons are massless particles of light of a certain quantum of energy Based on the frequency and wavelength of the photon Bohr’s Model Excited electrons Energy added to atom – electrons “jump” up energy levels When the atom relaxes electron “falls” to lower energy levels and emits photon Bohr Model of hydrogen Reference Sheets!!!!! Electromagnetic Radiation Atomic Line Spectra Electrons in an atom add energy to go to an “excited state”. When they relax back to the ground state, they emit energy in specific energy quanta Electromagnetic Radiation These observations suggested that electrons must exist in defined energy levels First, the electron absorbs energy and jumps from the ground state to an excited state 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ hv Next, the excited electron relaxes to a lower excited state or ground state 5 ______ 5 ______ 4 ______ 4 ______ 3 ______ 3 ______ 2 ______ 2 ______ 1 ______ 1 ______ hv Electromagnetic Radiation Wave Nature of Light Property of Waves Frequency No. of waves per second Wavelength ▪ Distance between corresponding points in a wave Amplitude ▪ Size of the wave peak ▪ Electromagnetic Radiation Mathematical Relations C=λf • C = speed of light = 3.0 x 108 m/s • λ (lamda) = wavelength (m) • f = frequency (Hz or s-1) This is how we know what color light is emitted! Frequency is inversely proportional to wavelength If λ increases, f decreases. If f increases, λ decreases. Speed of the wave is always constant at 3.0 x 108 m/s. Electromagnetic Radiation Particle Nature of Light Wave nature could not explain all observations (Planck & Einstein) Photoelectric Effect E = hf When light strikes a metal electrons are ejected Atomic Line Spectra ▪ When elements are heated, they emit a unique set of frequencies of visible and non-visible light. Other Scientists Contributions de Broglie Heisenberg Modeled electrons as waves Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron Electrons exist in orbital’s of probability Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron Other Scientists Contributions Schrödinger Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom Quantum Mechanical Model of the atom – current model of the atom treating electrons as waves. Quantum Mechanical Model Nucleus: Neutrons and protons Orbitals: region in space surrounding the nucleus where there is a 95% probability of finding an electron. We know either energy or location of each electron. Reflection Who contributed to the modern model of the atom? How is it different from Bohr’s? Why do atoms give unique atomic line spectra? What are ground and excited states?