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Chapter 4 ARRANGEMENT OF ELECTRONS IN ATOMS THE DEVELOPMENT OF A NEW ATOMIC MODEL Light has characteristics of both particles and waves Electromagnetic radiation – a form of energy that exhibits wave-like behavior as it moves through space Electromagnetic spectrum ELECTROMAGNETIC SPECTRUM = wavelength (units of meters) = frequency (units of Hertz, Hz, c = speed of light (3.00 x 108 m/s) h = Planck’s constant (6.626 x 10-34 Js) 1 seconds 1 c E h ) EMISSION AND ABSORPTION Ground state – the lowest energy state of an atom Excited state – when an atom has a higher potential energy than it does at its ground state Balmer series – represents visible light region Lyman series – represents the UV region Paschen series – represents the infrared series EMISSION LINE SPECTRA When a narrow beam of light is passes through a prism, it separates into specific colors in the visible spectrum. Each element has a signature spectrum. EMISSION SPECTRA In addition to indicators in the visible spectrum, energy can be detected in the UV region as well as the infrared region. This has been scientifically observed. What wavelengths represent indicators in the UV range? What about the IR range? EMISSION SPECTRA Scientists had predicted that emission spectra would be on a continuous spectrum. Is this what was scientifically observed? BOHR’S MODEL AND TRANSITION STATES Bohr’s Model helped to explain the quantum energy levels of the atom When a photon is absorbed, the electron gains enough energy to move to an outer energy level. When an electron loses energy (in the form of a photon), energy is released. ELECTRON ENERGY TRANSITIONS PHOTOELECTRIC EFFECT Recap – what is the relationship between energy and frequency of a wave? What is the visible region of light? Which colors in the visible spectrum carry more energy? http://www.youtube.com/watch?v=0qKrOFgJZ4 HOMEWORK P.97, #1-6 (2 days to complete) QUANTUM MODEL OF THE ATOM Electrons have wavelike properties Investigations from the photoelectric effect and hydrogen’s emission line spectra determined that light acts as both a wave and a particle. ELECTRONS Electrons have interference patterns Constructive interference Destructive interference DOUBLE SLIT EXPERIMENT http://www.youtube.com/watch?v=DfPeprQ7o Gc HEISENBERG Heisenberg’s Uncertainty Principle – you cannot know the instantaneous position and velocity of an electron (or any other particle) x – represents position p – represents momentum (velocity multiplied by mass) - represents a constant Δ – in this case, delta represents the uncertainty. When you have extremely precise information about x, your uncertainty for p goes through the roof. ORBITALS Orbitals indicate probable electron locations Schrodinger’s equation! This formed the foundation of modern quantum theory Quantum theory – Quantum numbers – PRINCIPLE QUANTUM NUMBER Symbolized by n Indicates the main energy level occupied by an electron. True or false – these can only be integers The total number of orbitals in a given shell is equal to n2! ANGULAR MOMENTUM QUANTUM NUMBER Different orbital shapes can exist for each principle quantum number. Represented by l What orbital shapes could you find at n=3? What about when n=3 and l=2? MAGNETIC QUANTUM NUMBER Represented by “m” Ranges from –l to l Indicates the orientation of an orbital around the nucleus SPIN QUANTUM NUMBER Represented by “s” Has only two possible states A single orbital can hold a maximum of two different electrons SMARTBOARD DEMONSTRATION LET’S PRACTICE! CLASS WORK How many electrons could be represented by n=2? How many electrons could be represented with a principle quantum number of 3 and an angular momentum quantum number of 0? How many electrons could be represented with a principle quantum number of 1 and a spin number of ½? How many electrons could be represented with a principle quantum number of 3 and a spin number of ½? How many electrons could be represented by l=3?