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Introduction

Introduction. Purpose: To observe the spectra of elements and relate the wavelengths to energy and energy levels of electrons.Spectroscope: Spectra

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Introduction

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    2. Introduction Purpose: To observe the spectra of elements and relate the wavelengths to energy and energy levels of electrons. Spectroscope: Spectra color, scope inspect thus inspect colors. Contains a prism, which separates emitted light into its constituent wavelengths. (red, greenetc.)

    3. Light Light: Electromagnetic energy (or combination of electric and magnetic fields) can be described by frequency and wavelength. Wavelength (?) distance between two peaks Frequency (?) Cycles (Wavelengths) per second.

    4. Electron Trends Electrons want to be as stable as possible, so initially they are in the lowest energy level possible. When electrons are heated they absorb energy. They travel to a higher energy level, and are now less stable. Electrons will release energy (light) to become more stable. An element when heated to its gaseous state, produces an emission line spectrum which we can observe by using a spectroscope. (Finger print)

    5. Bohrs Theory diagram of a Cl atom Electrons revolve around the nucleus in specific energy levels called orbits. Principle energy level (n): 1, 2, 3, n The greater the value of n the further away from the nucleus the electron is.

    6. Light emitted from hydrogen atom. We will observe energy being emitted as electrons drop from higher energy levels to lower ones. Electrons that fall to the 2nd energy level can been seen by us. !!Light is the disposal of energy!!

    7. The Equation for Light c= ?*? Speed of light (c) in a vacuum 3.0 x 108 m/s This in an inversely proportional relationship. If the wavelength increases, the frequency decreases. Note: 1 nm = 10-9 m!!!

    8. Since energy emitted depends on the size of the energy level drop, atoms may emit visible or non-visible light. Note: For hydrogen, each electron drop to n = 2 will result in the emission of visible light. nf will be 2 for our experiment.

    9. The energy evolved (absorbed or emitted) from an electrons transition is called a photon (discrete packet of energy). ?E = h? Where h = 6.63 x 10-34 Js (Plancks constant), and ? = frequency (sec-1or s-1) NOTE: ?E ? Negative value during emission ?E ? Positive value during absorption

    10. A Sample Calculation Part #1 A Hydrogen spectral line is observed at 486 nm. Find ?, and E, ? = c/? You must first convert nanometers to meters Where 1nm = 10-9 m 486nm = 4.86 x 107 m c = 3.0 x 108 m/s

    11. A Sample Calculation Part #2

    12. We can determine initial and final location of an electron, or a change in energy. Relates energy emitted to an electron shift. ?E ? Energy emitted, Joules Rh ? 2.18 x 1018 Joules Ni ? Initial energy level Nf ? Final energy level

    13. Rydbergs Equation

    14. Today in Lab

    15. How to determine Wavelength

    16. Wavelength Example

    18. Due Next Lab pg83 # 1 Flame test Demo, # 2 Visible Spectrum pgs 85-87 4A-4C Section C pg 87 identify unknown element Show all calculations for full credit

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