S1.
INTRODUCTION TO SPECTROSCOPY
PURPOSE
1.
2.
3.
4.
To use hand held spectroscopes to identify continuous and line spectra.
To observe the spectra emitted by elements in a flame and by various light sources
To use emission spectra for qualitative and quantitative purposes
To record simple absorption spectra.
PREWORK
1. List the chemicals you will be working with in this practical.
2. Note any safety requirements as indicated by the MSDS
3. Give reasons for not looking directly at the sun.
INTRODUCTION
A spectroscope is an instrument that splits up light into a spectrum, just as a prism does.
Using a spectroscope it is possible to determine whether the light contains a “continuous
spectrum” where the light appears to contain all wavelengths more or less equally, or whether
the light consists of just a few “bands” of colour, or perhaps just a few really narrow bands
known as “lines”.
Light is emitted when an outer electron takes in heat energy, jumps to a higher energy
position, then gives out, emits, the energy in the form of light as it falls back to its original
position.
PROCEDURE
I. Use of hand-held Spectroscope
• Take the spectroscope out of its box and practise by looking at the lights in the
laboratory. Find out how to increase and decrease the amount of light coming into the
instrument by narrowing and widening the slit.
II. Emission Spectra of Elements in a Flame
1. Light the flame as directed by your teacher, and aspirate the series of metal ion solutions
provided into the flame. Observe the colour by eye, and using the spectroscope, determine
the spectrum of lines emitted by each ion. Record your observations in the results section.
Between each solution, rinse the system by aspirating de-ionised water into the flame until
any unusual colour disappears. Compare your results to the spectrum shown on the wall
chart.
2. Obtain the set of unknown solutions, which are mixtures of two of the metals ions
observed in step 1. Noting the colour of the flame by eye and by using the spectroscope
identify the elements in the unknown solutions.
3. Pour about 50 mL of de-ionised water into a small beaker, and stir with two fingers for
about 1 minute.
4. Aspirate the three Na standard solutions (5, 25 and 50 mg/L) and a sample of tap water.
Estimate the Na content of the two water samples, recording your results.
III. Emission Spectra of Light Sources
1. Using the spectroscope , record the spectrum obtained from:
a. sunlight (by projecting the light onto a piece of white paper…do not look directly at
the sun)
b. an incandescent light
c. a fluorescent light
d. light shining through a piece of coloured glass
From the spectrum chart in the laboratory, identify the element responsible for the light from
the fluorescent tube.
IV. Basic Absorption Spectroscopy
1. Using a scanning spectrophotometer, record the absorbance of a 50 mg/L permanganate
solution every 20 nm between 400 and 700 nm using a clean plastic cell. It will be
necessary to run a blank of distilled water across the same wavelength range. Note the
wavelength of maximum absorbance of the permanganate and draw a sketch of the scan.
2. Using the scanning spectrophotometer, record the spectrum of a didymium filter between
400 and 700 nm. Note the wavelengths of the peaks. Note the appearance of the spectrum
on the wallchart for didymium.
3. Using the maximum absorption wavelength for permanganate, measure the absorbance of
the 50 mg/L solution in TEN different plastic cells.
CALCULATIONS
1. Plot a graph of absorbance (y-axis) against wavelength (x-axis) for the permanganate
readings.
2. Calculate the average and relative precision for the absorbances from the different cells.
DISCUSSION
• explain how you identified the element emitting light in the fluorescent tube, and the
identity of the two elements in the unknown mixtures.
• comment on the difference between your manual spectrum and the scanned spectrum
of permanganate
• explain how you could make your permanganate spectrum more similar to the scanned
one
• explain how the two spectra of didymium (scanned and wallchart) are providing the
same information
• comment on the variation in absorbance values for the different cells
QUESTIONS
1. How does the measurement of absorption and emission spectra differ?
2. What is the difference between line and band spectra?
3. Explain how spectroscopy can be used for quantitative and/or qualitative analysis
4. What causes the series of dark lines in the sunlight spectrum?
5. Copper (II) ions give a blue-green colour in solution. This means they are absorbing most
of the radiation from about 600-650 nm. If you looked at copper (II) solution through a
hand held spectroscope, what would you see?
6. Compare the emission spectrum of copper obtained in this experiment with the absorption
spectrum described in Question 5. Why are they different?
S1 RESULTS SHEET
Date of analysis
Instrument identity
II. Emission spectra of Elements in a Flame
Element
Copper
Lithium
Potassium
Calcium
Strontium
Sodium
Barium
Unknown 1
Identity
of
elements
Unknown 2
Identity
of
elements
Sample
Tap water
Skin
Flame Colour
2
2
Approx. Conc
III. Emission of Spectra of Light Sources
Appearance of sunlight spectrum
Appearance of incandescent light spectrum
Appearance of fluorescent light spectrum
Appearance of coloured light spectrum
Colour of Spectral Lines
IV. Basic Absorption Spectra.
Manual spectrum of permanganate
Wavelength (nm)
400
420
440
460
480
500
520
540
Absorbance
Wavelength (nm)
560
580
600
620
640
660
680
700
Absorption
Peaks for didymium (nm)
Spectrophotometer
Wallchart
Absorbance of permanganate solution in different cells
Have you?
Completed the instrument log
Completed the sample register
Completed the standard register
Teachers signature
Date
Signature