The Hydrogen Spectrum

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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.)

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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.

amplitude
Wavelength - ?
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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)

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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.

Nucleus
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• 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!!

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n
IR
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visible
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UV
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The Equation for Lightc ??

• 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!!!

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• 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.

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Calculating Energy

• 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

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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

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A Sample CalculationPart 2
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• 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

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Rydbergs Equation
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Today in Lab
Part 1- We will use a optical bench to
determine the wavelength
We will measure the distance between the source
and location of the light. The source is
directly behind 50 cm on the meter stick We
will ignore the left of the stick Our measurement
will be from the source to the location of the
light. Example light is visible at 75 cm so the
distance is 75- 50 or 25cm.
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60
70
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How to determine Wavelength
? wavelength X distance from light to
line ( in cm) D diffraction grating. (1667
nm) L length of bar distance from
grating to source (100 cm)
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Wavelength Example
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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