Download Spectrophotometry Chapter 18

Spectrophotometry Chapter 18
Review of Electron Energies
• The electrons of an atom have different energies.
• Not all energies exist, only certain allowed energy levels.
• Electrons with more energy are able to get farther away from the nucleus and its + charges.
• Therefore, electrons in higher energy levels spend more time farther away from the nucleus.
• The higher energy levels are larger so they can hold more electrons.
• Electrons are not orbiting the nucleus like planets!
• The energy levels of the electrons are divided into sublevels and orbitals.
Light as a Wave
• Light is also known as electromagnetic radiation.
– energy radiated in the form of a WAVE caused by an electric field interacting with a magnetic
field
– result of the acceleration of a charged particle
– does not require a material medium and can travel through a vacuum
Properties of Waves
All waves have:
• wavelength (, “lambda”).
– The distance between 2 crests of a wave.
– In units of length (m, cm, nm)
• frequency (, “nu”).
– The # of waves that oscillate per second.
– In units of 1/sec or Hertz, Hz
• speed (for light waves speed = c)
– In units of m/s
• Amplitude (A)
– The “height” of the wave
The Wave Nature of Light (electromagnetic radiation or EMR)
• The speed of a wave, v, is given by its frequency multiplied by its wavelength:
•
v =   (units are m/s)
• For light in a vacuum, speed v = c =2.9979 x 108 m/s and this is the fastest anything could
travel.
• c =   (units are m/s)
– Note that the longer the wavelength the smaller frequency.
– Or the shorter the wavelength the higher (larger) the frequency.
Visible Light:
The Light our Eyes Can Detect
• Wavelengths: ~ 400 to 750 nanometers.
– Each wavelength corresponds to a different color of light.
– Red light: ~ 620 - 750 nm
– Violet light: ~ 390 – 450 nm
• White light is a mixture of all the colors of the rainbow.
– A prism can split white light into its colors.
Electromagnetic Radiation (EMR) Spectrum
• Visible light is only one type of EMR.
• Other types include:
– Gamma rays
– X-rays
– Ultraviolet
– Infrared
– Microwaves
– Radio and TV waves
• All of these have different wavelengths.
What is the frequency of blue light with  = 440 nm?
• c= 
•  = c/ 
•  = (3.00 x 108 m/s)/(440 x 10-9 m)
•  = 6.8 x 1014 1/s
•  = 6.8 x 1014 Hz
Wave Number
• The relationship between energy and frequency of light is: E = h
where h is Planck’s constant (6.626  10-34 Js).
• Using the above equation and the relationship between wavelength and frequency we get:
E = h c/ 
• Since h and c are constants, it follows that 1/  is proportional to E.
• The quantity 1/  is given the symbol  and is called the wave number.
Line Spectra
• Each element emits (gives off) its own, individual set of colors when energized.
• This is called its atomic emission spectrum. (or line spectrum)
– The colors show up as a series of lines when viewed through a prism.
– Called spectral lines.
• Like a fingerprint for that element.
Ground and Excited States
• Ground State: The state of least possible energy in a physical system, as of elementary particles.
– Also called ground level.
• Excited State: Being at an energy level higher than the ground state
Absorption
• A photon of light hits a molecule or atom.
• If the energy of the photon is “right”, it is absorbed.
•
•
An electron excited to a higher energy level.
• The energy of the photon must be exactly equal to the difference of the electron energy levels.
Photons that aren’t absorbed are scattered (or transmitted).
Types of Transitions
• Light in the visible and UV may cause electronic transitions.
• IR radiation effects vibrational states of molecules (more later).
• Microwave radiation effects the rotational states of molecules (more later).
Emission
• An electron is excited to a higher energy level by an outside energy.
– Heat, electricity, etc.
• The electron returns to a lower energy level.
• A photon (tiny packet) of light is emitted.
• The energy of the photon is equal to the difference between the energy levels of the electron.
Spectrophotometry
• In a spectrophotometry experiment, we measure the light absorbed or emitted.
• This week in lab: absorbance in the visible region.
Absorption Methods: Transmittance
T = P/Po
where
T => transmittance
P => power of transmitted radiation
Po => power of incident radiation
%T = (P/Po)*100
where
%T => percent transmittance
Absorption Methods, Absorbance
A = - log10T = - log10 (P/Po)
where
A => absorbance
Relation Between Transmittance and Absorbance
P/Po
%T
A
1
100
0
0.1
10
1
0.01
1
2
Absorption Methods, Beer’s Law
A = abc = bc
where
a => absorptivity
b => path length
c => concentration
 => molar absorptivity
Homework Chp 18
Exercises: 18A, B, C
Problems: 1 – 4, 6 – 12, 16, 18