1. Lab 12 Atomic spectra and atomic structure

2. amplitude
The Hydrogen Spectrum
Wavelength - 
Experiment 6

3. Introduction
Purpose: To observe the spectra of elements and relate the wavelengths to energy and energy levels of electrons.
Spectroscope: Contains a prism, which separates emitted light into its constituent wavelengths. (red, green…etc.)

4. Anatomy of a Light Wave
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 - 

5. 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!!!

6. Bohr’s Theory
An element, when heated to its gaseous state, produces an emission line spectrum which we can observe by using a spectroscope. (Finger print)
What is atomic structure and how is light emitted? Electrons in an atom exist in specific regions at various distances from the nucleus. (eg. Planets circling the sun.)

7. Bohr’s Theory (2) diagram of a Cl atom
Nucleus x
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. Cl

8. Bohr’s Theory (3) diagram of Cl atom with energy shells
Usually electrons exist at their lowest energy level or “ground state” where they are most stable, but if the atom is heated, exposed to light or electricity, the electrons can absorb energy and move to a higher energy level (Excited State).
Electrons can only absorb energy in specific amounts. They absorb the entire amount first then instantaneously jump to a higher level(s).
Eventually electrons emit all or some of the absorbed energy and drop from the higher energy level to a lower energy level and emit a photon for each drop in energy level.
Nucleus

9. The Experiment
Light emitted from hydrogen atom. We will observe energy being emitted as electrons drop from higher energy levels to lower ones. 6 5 4 n= 3 IR 2
visible 1 UV

10. The Experiment (2)
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.

11. The Experiment (3)
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 J•s (Planck’s constant), and  = frequency (sec-1or s-1)
NOTE:
E  Negative value during emission
E  Positive value during absorption

12. The Experiment (4) a sample calculation
A Hydrogen spectral line is observed at 486 nm. Find , E, and Ni:
 = c/
You must first convert nanometers to meters  486nm = 4.86 x 10–7 m
c = 3.0 x 108 m/s

13. Calibration of Mercury
Look through spectroscope to determine emission lines of mercury
Compare the observed  with the theoretical  of these colored lines
Prepare a calibration curve by plotting your experimental wavelengths to the theoretical ones given on page 126
Graph needs to be computer generated with a trend line and a equation of the line

14. Part A. Emission spectrum of Hydrogen
Look at hydrogen through a spectroscope and determine the colors that the emission spectrum produces
Use Rydberg’s Equation found on page 5 and 6 on pg 127 to calculate wavelength in nanometers for the electron transitions listed on pg 129.
Look at the chart in the classroom, and determine which colors correlate to which assignment

15. Rydberg’s Equation
Relates energy emitted to an electron shift in a hydrogen atom.
E  Energy emitted, Joules
Rh  2.18 x 10–18 Joules
Ni  Initial energy level
Nf  Final energy level
nf will be 2 for our experiment!

16. Rydberg’s Equation (2)
Example: (green) nf = 2; ni = 4

17. Part B Emission Spectra of Group 1 A and 2A elements
When solutions of metals are heated in a Bunsen burner flame, they give off characteristic colours. For example, sodium makes the flame turn bright orange –

18. Part B
Using a nichrome wire place a small sample of the known metal into the flame and record the gross color.
Use HCL to clean the wire and repeat with the next metal
Determine the unknowns based on your standards

19. Due next week
Pgs 25-27
Questions 1-10 pg 27
Study for next weeks quiz