Atomic Spectra/Flame Test Practical 2011
Introduction:
What do fireworks, lasers, and neon signs have in common? In each case, we see the brilliant
colours because the atoms and molecules are emitting energy in the form of visible light. The
chemistry of an element strongly depends on the arrangement of the electrons. Electrons in an
atom are normally found in the lowest energy level called the ground state. However, they can
be "excited" to a higher energy level if given the right amount of energy, usually in the form of
heat or electricity. Once the electron is excited to a higher energy level, it quickly loses the
energy and "relaxes" back to a more stable, lower energy level. If the energy released is the
same amount as the energy that makes up visible light, the element produces a colour.
Aim:




Observe/record the colours produced by metal salts in a flame
Use flame colours to identify unknown salts
Observe/record line spectra of hydrogen gas
Calculate frequency and energy values of specific wavelengths
Apparatus:
Bunsen burner, matches, cork mat, nichrome wire, watch glasses, 50ml beaker, Various salts,
spectrometer, apparatus for transferring power into hydrogen gas, fluorescent light and hydrochloric acid (4 molar).
Safety Considerations
For the flame test 4m HCl will be used to clean the nichrome wire after a test to clean off any
left-over salts. Thus participants require both safety goggles and lab coats. This being a practical
experiment involving the use of a Bunsen burner, participants must be mindful of correct
burner procedure and safety considerations including the use of the safety flame feature when
not actively conducting an experiment.
Some salts are hazardous to health and should be disposed of in waste bins provided and not
poured down the sink. Participants as per usual should take care to thoroughly wash hands
before and after the experiment.
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Atomic Spectra/Flame Test Practical 2011
Method 1#:
1. Clean a nichrome wire by dipping into HCl solution. Heat the wire in the heating flame
of a Bunsen until no more colour is produced.
2. Mix a little of the salt to be tested with the acid solution on a watch glass and dip the
end of the clean wire in the solution.
3. Hold this end of the wire in the outer blue part of the flame and note any flashes of
colour that may appear and the intensity. Note that these may fade quickly.
4. Test subsequent salts in the same manner, cleaning the wire between each test as
described in #1.
5. Tabulate your results making sure to include qualitative and quantitative data.
6. Use the information you have collected to identify the cations in each of the unknown
samples provided to you.
Results:
Table 1: Observation of Flame Tests
Salt
Strontium Chloride
Calcium Chloride
Lithium Chloride
Copper Chloride
Potassium Chloride
Baric Chloric
Magnesium Chloride
Iron Chloride
Sodium Chloride
Unknown 1
Unknown 2
Colour
Red
Orange
Red
Blue/White
Purple
Green
Blue
White
Yellow
Yellow
Yellow/Orange
Intensity
High
Very High
Medium
Very High
Medium
Low
Low
High
High
High
High
Duration (seconds)
2
1
2-3
4
<1
3-4
2
<1
2
2
<1
During the experiment I couldn’t discern any recognizable pattern or correlation between the
element’s colours and duration or intensity. The only noticeable pattern was that salts derived
from elements closer to Group 1 on the periodic table tended to combust faster and with
greater intensity. As far as the 2 unknown elements, Unknown 1 from my tests is Sodium
Chloride or plain salt. Unknown 2 matches the colour and intensity of Calcium Chloride.
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Atomic Spectra/Flame Test Practical 2011
Discuss:
The experiment only yielded one major pattern, that being that the elements that I knew to be
close to Group 1 of the periodic table of elements tended towards more rapid reactions. Both
unknown salts displayed clearly recognizable trends of other known salts tested beforehand.
With regards to why the salts composed of elements closer to Group 1 were shorter in reaction
rates, it is possibly because the elements have less energy levels for the electrons to pass
through once being “excited” by the heat or because they, being closer to group 1 thus being
naturally more reactive, reacting more violently than the other compounds to the energy
transfer.
The unknown salts displayed clear traits due to the fact that each element has its own unique
properties, though they may be grouped according to similarity each elements reacts to energy
transfer (heat) in a different way, ergo the salts can be identified fairly easily.
Evaluation:
Due to the fact that the flame test on each salt could only be done once it is possible that
several errors could alter the reactions observed. The wire may not have been properly cleaned
causing different colours of flame and duration of reaction. The wire could also have corrupted
the observation of the flame as it burns orange when red-hot thus it could make some salts
appear to be burning bright orange
A way to improve/prevent these issues could be to test each salt twice after burning off any
residue and soaking the wire in water before cleaning it in HCl. It might also be helpful to have
the accepted colours of such salts when burnt on hand to compare with the experiments
findings, although there cannot be much quantative error there could still be misleading results
due to impurities within the salts or just sloppy practice on the part of the experimenter.
Now having completed the experiment I would like the opportunity to test several salts from
different groups in order to see if the flame test will yield more identifiable patterns according
to Groups in the periodic table. Perhaps an element maybe identified in terms of which group it
belongs to just by its colour or the duration of its combustion?
Method 2#: Atomic Spectra
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Atomic Spectra/Flame Test Practical 2011
1. Go to the station where a hydrogen gas discharge tube is set up. Here, an induction coil
is being used to provide energy to “excite” the gas atoms.
2. Using the spectrometer, look at one of the fluorescent light in the room. You will see
coloured lines corresponding to specific wavelengths emitted by the white light.
3. Record the values for the violet and green lines. These should be 4360 Å and 5460 Å
respectively. Note any error in these readings so that adjustments can be made in
subsequent readings.
4. Turn on the electricity to the hydrogen gas tube (for no more than 30 sec at one time to
prevent burn out).
5. Look through the spectrometer and observe/record the coloured lines that are
produced.
6. Tabulate your results, including qualitative and quantitative data.
7. Process your data thoroughly, providing as much information as possible about each line
found including error and uncertainty.
Results 2#: Fluorescent vs Hydrogen
1: Fluorescent Bulb
Red
Yellow
Green
Blue
Purple
6200
5850
5460
4200
4360
Red
Yellow
Green
Blue
Purple
6700
5800
5400
4400
4410
2: Hydrogen
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Atomic Spectra/Flame Test Practical 2011
8000
7000
6000
5000
Hydrogen
4000
Fluorescent Light
3000
2000
1000
0
0
1
2
3
4
5
6
Figure 1, 1 = Red, 2= Yellow, 3 = Green, 4=Blue, 5= Violet
From the use of spectrometers we can see from observation that the hydrogen when excited by
electricity deviates from the fluorescent bulb’s spectrum by no more than 500 in each of the
separate colours. In fact both gases exhibit similar light both with and without observation
through the spectrometer.
Discussion
From the spectra test I found that hydrogen’s spectra erred towards the extreme ends of the
spectra, having greater values than the fluorescent light in all but two of the spectrum colours.
This may be related to the fact that hydrogen gas is a much simpler element than the
compound of gases found in fluorescent bulbs, (Krypton and Fluorine to name a couple).
Why the fluorescent bulb produced lesser values of each spectrum colour is unclear to me, not
having an extensive knowledge of both the compound present in the bulbs or even a basic
understanding of the characteristics of hydrogen as opposed to such a compound. However I
would say that a more complex gas such as the fluorescent bulb contains would be less reactive
to electricity as all the respective elements within the compound would experience changes in
different ways. All that I could suggest would be that hydrogen as a gas is more reactive and
sensitive to energy than the fluorescent lighting.
Evaluation:
Any errors or inconsistencies with the data would be largely down to my own inexperience with
spectrometers as they are particularly difficult to use effectively without having to estimate the
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Atomic Spectra/Flame Test Practical 2011
actual value being displayed on the spectrum. The data collected is “approximate” at best as
attempting to remember a value then record it after putting down the spectrometer is
unwieldy at best. However the data, though not necessarily exact still displays certain trends
which makes reaching a conclusion no more difficult than if the readings had been taken down
reliably.
The method could be improved if it was done by two people as one could take a reading and
report it to another who could then record it, thus allowing the participant viewing the
hydrogen to remain focused on the behaviour of the gas and the spectrum colours.
Having observed hydrogen while being stimulated by electricity I’m forced to wonder what
properties other gases have when stimulated by electricity and what readings they produce on
the emissions spectrum. The only way to fully understand how elements behave is to observe
their behaviour in action.
Conclusion:
From the flame test I learned that each element produces differing results when subjected to
heat as the electrons jumping from energy level to energy level traveled differently with vary
degrees of light intensity, colour and duration of reaction.
While conducting the atomic spectra observations I found that each element has its own
distinct emission spectra when subjected to electric stimulation, having compared only two
gases, both of which were vastly different from each other (one a element, the other a
compound) I found that the base elemental gas (H) had mostly greater values on the emission
spectrum of colours when compared to the emission readings of the common fluorescent light
bulb.
These experiments have stimulated a greater curiousity in regards to individual characteristics
of elements as their electrons react to energy and how chemists can record and draw
conclusions from such complex processes occurring on such a small scale.
Marks
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