PAUL ROBERTS: Hello, my name is Paul Roberts. And I'm a chemistry teacher here at Andover College.
Today, we're going to be looking at the core practical identifying ions
in unknown salts. This part of the practical will involve identifying the cations, the positive ions
in the salts. In this experiment, you'll be given a range of known and unknown solutions.
These solutions will contain either aluminium 3-plus ions, calcium 2-plus ions,
copper 2-plus ions, iron 2-plus ions, or iron 3-plus ions.
However, in this demonstration, I will show you how to identify the ions individually. And when
you are given a unknown solution, you will be able to deduce what the iron is. Along the way,
I'll explain the chemistry of what's happening and give you some tips to get some good results.
Regarding safety, all the metal ion solutions should be considered to be harmful. Although
they are dilute and in solution, they will still cause damage to eyes and skin. Therefore, it is
important to wear safety glasses and laboratory coats throughout the duration of this experiment.
Particular care should be made when handling the 1 mole per deciliter cubed sodium hydroxide
solution. At this concentration, it is corrosive. So avoid skin contact. Clean
up any spills immediately. And if it does come in contact with the skin, rinse with plenty of water.
Although this practical is straightforward,
serious errors in observation can occur if there is cross-contamination between the solutions.
Therefore, I recommend writing on each pipette which solution it will be used for.
When measuring out volumes for this practical, there is no need to be super accurate. You only
require about 2 centimeters depth of the metal ion solutions in each test tube.
We're only looking for color changes or precipitates. So I wouldn't bother
wasting any time trying to measure out the perfect volumes, as this won't change results.
All the tests employ a common chemical property of metal ions. This is that metal hydroxides are
mostly insoluble in water. In fact, only group 1 metal hydroxides have good water solubility.
Therefore, when a soluble metal compound, typically a nitrate or a sulfate, is reacted
with sodium hydroxide, one of the few soluble hydroxides, a precipitate will be observed.
The color of the precipitate will depend on the metal ion being reacted.
First, I'll perform a test for aluminium 3-plus ions, which are essentially aluminium atoms
which have lost three electrons. Aluminium ions are notoriously insoluble. However, aluminium
nitrate is soluble. I have a 2-centimeter depth solution of aluminium nitrate here. And I can test
for the presence of aluminium by adding a few drops of sodium hydroxide. The hydroxide ions
react with the aluminium ions to produce aluminium hydroxide, which is an insoluble white compound.
Testing for calcium ions-- I have a solution of calcium chloride here. And once again,
I'm going to react the soluble metal ion solution with sodium hydroxide. This time, the precipitate
of calcium hydroxide is produced. You'll notice that it is also white. So we have a problem. When
we have two white precipitates, how do we know if it's calcium or aluminium? During the practical,
you're going to investigate this further by adding extra sodium hydroxide. You should
notice that the aluminium hydroxide precipitate dissolves when extra sodium hydroxide is added.
I am going to test for copper 2-plus ions in exactly the same way, 2 centimeters depth of
copper sulfate solution followed by a few drops of sodium hydroxide solution.
And as you can see, a gelatinous blue precipitate is formed.
Conveniently, it's the same color as the original copper sulfate solution.
Until now, all the ions we have tested have been the common ions for each element.
For example aluminium commonly forms 3-plus ions. Calcium commonly forms 2-plus ions.
And copper commonly forms 2-plus ions. Other charges of these cations are rare.
Now we're going to look at iron, Fe. I have two solutions here-- one containing iron 2-plus ions
or iron II and the second containing iron 3-plus ions or iron III. When I
add the same amount of sodium hydroxide to each solution, you should notice the iron 2-plus forms
a green precipitate of iron II hydroxide. And iron 3 forms a brown precipitate of iron III hydroxide.
After a few minutes of exposure to the air, the surface of the iron II hydroxide precipitate
has started to darken and become more brown. What's happening here is the
iron II hydroxide is being oxidized by the air to form the brown iron III hydroxide.
The second part of is practical is going to involve observing flame colors of
lithium plus ions, sodium plus ions, potassium plus ions, calcium 2-plus ions,
and copper 2-plus ions. This will involve using a Bunsen burner where particular care must be taken
in its operation, making sure long hair is tied back and there are no flammable materials nearby.
Safety glasses and lab coats will be worn throughout this practical.
I suggest when performing these flame tests,
concentrations of at least one mole per deciliter cubed are used. This will enable a clear,
positive result. We are using solutions of the ions. Please note that you can
also use solid samples containing the ions which will obtain similar results.
First of all, I will light the Bunsen with the hole closed
and then adjust the flame so it is roaring. I'm going to take a piece of nichrome wire attached
to a glass rod and immerse it in a 2 mole per decimeter cubed hydrochloric acid solution.
This will remove any impurities on the nichrome wire.
Next, I'll immerse a nichrome wire in the solution to be tested, in this case, lithium chloride,
and then I will place a wire in the side of a blue Bunsen flame. You can see here the lithium cation
is producing a crimson red flame color. The chloride in solution does not produce a flame
in the visible light range. Therefore, no color is observed from the chloride anion.
I will now repeat the process with sodium chloride.
This time, we can see an orange flame caused by the sodium ions.
Once again, the same process is repeated with potassium chloride.
And this time, the potassium ions are responsible for a lilac flame.
When the flame test on calcium chloride is carried out, the calcium 2-plus ions cause
a dark red flame to be observed. Be careful not to confuse this dark red color of calcium
with the crimson red color of lithium. The final flame test is on copper II chloride.
And this time, blue-green flames are observed. These are caused by the copper 2-plus ions.
Just to clarify, the chloride ions dissolved in all these solutions have not been responsible for
any of the flame colors. The different flame colors have been caused by the different metal
ions dissolved in solution. After the flame tests, it is important the Bunsen is turned off promptly.
The third part of this practical involves testing for ammonia, a neutrally-charged molecule,
and ammonium, a positively-charged ion. During your test for ammonium,
you will be heating a solution in a test tube over a gentle Bunsen flame.
Now, during this process, it is important not to point the mouth of the test tube
towards another person. Ammonia and H3 is a gas at room temperature and pressure. However,
it is highly soluble in water and is often used in its aqueous form. Ammonia has a characteristic
pungent, fishy smell and is often produced as a byproduct of biological matter decaying.
Ammonia is a basic compound. So if a dampened piece of red litmus paper is held over the
mouth of the test tube containing ammonia, the red litmus should turn from red to blue.
This is particularly easy to identify if the ammonia solution is of a high concentration.
However, if the ammonia solution is of a low concentration,
it can be gently warmed in a test tube to encourage the evolution of the gas.
Ammonium NH4+ is the neutralized cation of ammonia. Ammonium does not have the characteristic
smell of ammonia. And all ammonium compounds are solid at room temperature and pressure.
In order to test for ammonium, we need to change it to ammonia. And we do this
by removing the extra H+ from the cation. And this is done by using sodium hydroxide.
I'll start off by adding 2 centimeters depth of ammonium chloride solution to a test tube.
To this test tube, I will add 1 centimeter depth of one mole per decimeter cubed sodium hydroxide.
I will now apply gentle heating with a Bunsen burner once again being careful not to point
the mouth of the test tube at another person. The sodium hydroxide will react with the ammonium ion
to produce ammonia. And this gas can be once again detected using dampened red litmus paper.
So in summary, we have now gone through the procedures for identifying common metal ions,
also the ammonium ion and the ammonia molecule. The procedures we've used have been
metal hydroxide precipitate tests, flame test to identify metal ions,
and the red litmus test to identify the ammonium ion and the ammonia molecule.
I'm going to be demonstrating the core practical testing for anions. The anions
we're going to be looking at in particular are chloride, bromide, iodide, sulfate,
and carbonate. In this demonstration, I will show you how to identify these anions individually,
and then you'll be able to use these results to identify anions in an unknown solution.
Regarding safety in this experiment, assume all unknown solutions to be harmful. When trying to
identify sulfate anions, you will be using barium chloride. Any barium compounds either reacted
or unreacted should be disposed of separately in the inorganic waste and not down the sink. Barium
compounds are particularly harmful and will damage aquatic life. You will also be handling dilute
hydrochloric acid and dilute nitric acid. Both will cause skin irritation and damage to the eyes.
Therefore, it is important to wear safety glasses and lab coats throughout this experiment.
Although this practical is fairly straightforward,
serious errors in observation can occur if there is cross-contamination between the solutions.
Therefore, I recommend writing on each pipette which solution it will be used for. When measuring
out volumes for this practical, there is no need to be super accurate. You only require about
2 centimeters depth of each of the solutions to be tested in each test tube.
The area students often make mistakes is with observations. When you are adding chemicals to
your unknown solutions, carefully observe what is happening. Don't add the chemical
and then decide to look 30 seconds later, as the reaction may have already happened.
Resist the temptation to carry out batches of reactions, as by the time you've finished
adding the last chemical to the last test tube, the first test tube may have finished reacting.
First of all, we're going to test for the presence of halides. Halides are the group 7 anions.
And the specific halides we're going to test for are chloride, bromide, and iodide. Initially,
I'm going to add 2 centimeters depth of potassium chloride solution to a test tube.
To this, I will add a few drops of dilute nitric acid
and then add a few drops of silver nitrate solution. You can see the chloride anions
have reacted with the silver cations and have produced a white precipitate of silver chloride.
I will repeat the same procedure with the other halide solutions sodium bromide and
sodium iodide. In comparison to the chloride reaction, the bromide anions have reacted
with the silver cations and produced a cream precipitate of silver bromide. And the iodide
ions have reacted with the silver cations and produced a yellow precipitate of silver iodide.
Next, I'm going to test for the presence of sulfate anions, assay 4 2-. Here I have a
sample of sodium sulfate. And once again, I'm adding it to a test tube to a depth
of 2 centimeters. I will add a few drops of hydrochloric acid,
briefly shake, and then add a few drops of barium chloride. You'll notice almost immediately a
heavy white precipitate which will settle fairly quickly. This has been produced by the reaction
between the barium cations and the sulfate anions hence barium sulfate.
The final anion I'm going to identify is carbonate. A sodium carbonate solution is
added once again to a 2 centimeter depth in this test tube. To the carbonate solution,
I'm going to add a few drops of either nitric acid or hydrochloric acid.
With careful inspection, you should notice small bubbles forming in the test tube.
These are bubbles of carbon dioxide, which are being produced as a result
of an acid carbonate neutralization. To confirm the bubbles of carbon dioxide,
you could apply a line water test. Ask your teacher for further guidance.
Regarding the disposal of chemicals, particularly care should be made not to dispose of any solution
containing barium ions down the sink, as barium ions pose a threat to aquatic life.
They should be disposed of in a separate inorganic waste container.
The contents of the other test tubes can be disposed of down the sink with plenty of water.
In this practical, we have shown you how to identify the main anions found in GCSE chemistry,
namely chloride, bromide, iodide, sulfate, and carbonate ions. Your teacher
may now give you some unknown solutions for you to test to see which anions they contain.