Tollens' reagent, used in what is known as the Tollens' test, or ‘silver mirror test’, is a special test chemical developed by Bernhard Tollens. It is used in organic chemistry as a distinguishing procedure
between organic ketones (RCOR’) and organic aldehydes (RCHO). With aldehydes
the reagent will react and deposit a silver coating on the glass surface of the
reaction apparatus, hence coining the ‘silver mirror’ name for this test.
However, with ketones, no reaction occurs and no silver is deposited.
The reaction behind the Tollens' test is generally an
alkali-catalysed redox reaction. To create Tollens' reagent, one first adds
concentrated ammonia to aqueous silver nitrate (AgNO3) in slight
excess to create the diamminesilver complex ion, [Ag(NH3)2]+
(aq).
During this procedure, insoluble silver oxide is first
precipitated, and with the addition of excess concentrated ammonia, the silver
oxide is 'ammoniated' to form the soluble silver complex ion, which is
indicated by the re-dissolving of the precipitate. After this, a little aqueous
sodium/potassium hydroxide is added to the reagent as a catalyst, and ammonia
added again to dissolve any precipitate.
Thus, in the reagent we now have an oxidizing agent, which
is the diamminesilver complex ion, and a catalyst, i.e. the hydroxide ion. However,
after a short period of time (roughly a day) the reagent will decompose to form
a very hazardous and extremely shock-sensitive explosive, silver nitride (Ag3N).
Thus, Tollens' reagent should be used immediately after preparation and
disposed of properly after use.
To run the test, one just has to add the organic compound that
is to be tested into the reagent and observe any changes.
For aldehydes, the –CHO group acts as a reducing agent,
reducing the diamminesilver ion to aqueous ammonia, and also elemental silver, which deposits on
the walls of the reaction vessel to form a silver mirror. In the process, the
aldehyde is oxidized to the corresponding carboxylic acid.
Half-reaction 1 (reduction)
Half-reaction 2 (oxidation)
Overall redox reaction
For ketones, however, the C=O group cannot reduce the
diamminesilver ion, and no reaction occurs.
Besides analytical purposes, Tollens' reagent was also used
before the advent of electroplating as a means of manufacturing mirrors.
However, due to its hazardous properties this method of production has fallen
out of use.
A distinguishing property of transition (d-block) metals is their ability to form
various coloured compounds with differing oxidation numbers, and in general a
colour corresponds to a specific oxidation number. Please refer to the
reduction reaction of manganese in the blog post: Chemical Chameleon for an
example.
Why transition metals exhibit so many colours in their
compounds is based on the quantum structure of the transition metal atom. All
transition metals have their valence electrons in the d-subshell, and in the isolated, atomic state the five d-orbitals have the same energy level,
i.e. they are degenerate. However, in a compound/complex ion the ligands
attached to the transition metal ion cause these d-orbitals to split in terms of energy, forming one group of d-orbitals with a lower energy and
another group of d-orbitals
with a higher energy. This happens due to the repulsion between the metal electrons and the electrons from the attached chemical.
Electrons in the lower energy d-orbitals can ‘jump’ to the higher energy d-orbitals via absorption of energy in the form of photons. To make
this jump the electron must generally absorb photons having energies of 2.8 to 5.0
x 10^-19 joules, which corresponds to a wavelength range of 400-700 nanometres,
the same wavelength range of the visible light spectrum. Thus, as visible light
passes through a transition metal compound, photons of a specific wavelength (and
thus colour) are absorbed by the compound, and the remaining photons pass
through the compound to give a coloured light.
For example, a solution
of copper (II) ions is blue as red light is absorbed by the d-electrons in the copper (II) ion (or
more accurately, the copper (II) hexahydrate ion). A solution of iron (II) ions
is green as in this case, the aqueous iron (II) ion absorbs reddish-blue light.
In this activity, CHS Science and Maths will experiment with
chromium compounds, creating as many types of coloured compounds as we can
through various physical and chemical reactions.
Follow the procedure here:
Starting with the commonly-used chromium (III) chloride (CrCl3) as a base chemical, the following reactions are conducted:
Dehydration - the dark green hydrated chromium (III) chloride hexahydrate is dehydrated via heat to become the violet anhydrous salt:
Amination - dark green aqueous chromium (III) chloride is aminated via concentrated ammonia solution to form the pink hexaammine chromium (III) complex ion, with chromium (III) hydroxide as an intermediate:
Oxidation (alkaline) - first the chromium (III) solution is made basic with the addition of sodium hydroxide solution. Excess sodium hydroxide solution is added to redissolve the insoluble chromium (III) hydroxide as shown below:
Following this hydrogen peroxide is added and the solution warmed. This oxidises the dark green chromate (III) ion to the yellow chromate (VI) ion:
Addition of acid - Adding acid to the alkaline chromate (VI) solution increases the concentration of H+ ions, turning the yellow chromate (VI) ion into the orange dichromate (VI) ion.
As shown by the double-headed arrow, the chromate (VI) and dichromate (VI) ions exist in dynamic equilibrium, with their concentration determined by the acidity (pH) of the solution. By Le Chatelier's principle, a low pH (high H+ conc.) encourages the formation of the orange dichromate (VI) ion whereas a high pH (low H+ conc.) encourages the formation of the yellow chromate (VI) ion.
Addition of peroxide/Ligand exchange reaction
The addition of acidified hydrogen peroxide to a solution of either yellow chromate (VI) or orange dichromate (VI) ions yields the blue chromium (VI) oxide peroxide. It is unstable when dissolved in water, so to obtain a sample of the peroxide the reaction has to be carried out in an organic solvent. In this case, a chromate (VI) compound dissolved(solvated) in ethanol is used.
Addition of chlorine/Ligand exchange reaction
Dehydrated, strongly acidic orange dichromate (VI) ions react with concentrated hydrochloric acid to form the powerfully oxidising red compound of chromyl chloride. Due to the corrosive and toxic nature of the reactants and products involved this reaction is best carried out in a fume hood. Chromyl chloride reacts spontaneously with water, thereby requiring that it be stablised in a dry medium.
The Iodine Clock reaction, first discovered by Hans Heinrich Landolt in the 19th century, is what is known as a 'chemical clock reaction', and today is one of only few known to science. It is useful as a tool in determining the rate of a chemical reaction due to a marked change in the colour of the reacting solution when a particular reaction has been completed.
In general, two solutions containing iodine, starch, a reducing agent and an oxidising agent are prepared separately. When ready one solution is added to another and after a period of time the reacting solution always changes colour from colourless to a dark blue-black. The time taken for the solution to change colour can be recorded and the rate of reaction thus determined. To get an even more accurate reading colourimetry can be used as the reacting solution is adept at absorbing certain wavelengths of light when it changes colour.
Two reactions occur simultaneously in the iodine clock; in one iodide ions, I- are oxidised to form triiodide ions, (I3)-, and in the other the reverse occurs. This particular version of the iodine clock that will be demonstrated uses sodium thiosulphate (sodium metabisulfite), Na2S2O3 as a reducing agent for iodine whereas hydrogen peroxide, H2O2 is the oxidising agent. Sulphuric acid, H2SO4 is also used to acidify the solution.
In the first reaction iodide ions are oxidised to triiodide ions via the action of hydrogen peroxide:
In the second, triiodide ions are reduced back to iodide ions via the action of the thiosulphate anion:
Once the thiosulphate anions are depleted, an excess of triiodide ions is produced, which binds to the starch in the solution to form the blue-black triiodide-starch complex. At this point, the solution will hence turn blue-black.
Varying the concentrations of the various reactants involved will change the time taken for the colour change of the solution, t. Increasing the concentration of acid, hydrogen peroxide and iodide ions decreases the time taken for the colour change. Increasing the concentration of sodium thiosulphate increases the time taken for the colour change.
Taking rate of reaction to be directly proportional to 1/t, we can thus say that increasing the concentration of acid, hydrogen peroxide and iodide ions increases the rate of reaction, and increasing the concentration of sodium thiosulphate decreases the rate of reaction.
Feel free to watch the video below for the procedure.
Redox (reduction-oxidation) reactions play a major role in chemistry. In essence, redox reactions involve the transfer of one or more electrons from one reactant to another; the reactant that loses electrons is oxidized, and the reactant that gains electrons is reduced.
Students generally use mnemonics to remember the definitions of oxidation and reduction, and below are two more popular ones:
L - Loss of E - Electrons is O - Oxidation
G - Gain of E - Electrons is R - Reduction
(= Leo (lion) goes: 'GER' (grr))
the other being:
O - Oxidation I - Is L - Loss (of electrons)
R - Reduction I - Is G - Gain (of electrons)
(= oil rig).
How much an element/chemical can be reduced or oxidized is shown by its oxidation number, which generally ranges from -3 to +7. Some elements can only have two oxidation numbers (i.e. one oxidation state while bonded in a compound, and one neutral (charge 0) state), while others can have multiple oxidation numbers. These elements are generally the heavier elements starting from Period 3 of the Periodic Table as they have more electrons to 'play' with as they form compounds.
In general, the heavier Group 17 halogens (Cl, Br, I) and a class of metals known as the transition or d-block metals have multiple oxidation states. Many common metals such as iron and nickel are considered as transition metals, and as metals are all electron donors in chemical reactions, they can only be oxidized from their neutral state to form chemicals with a positive oxidation number. For ease of differentiating compounds with the same metal but with different oxidation states, Roman numerals (I to VII) are used to denote the oxidation state of transition metals in their compounds. For example, iron forms compounds with a +2 and +3 oxidation state; so iron chloride can either be iron (II) chloride (iron is +2) or iron (III) chloride (iron is +3).
One interesting aspect of transition metals is that their compounds are all colored, and the color of the compound depends on the oxidation state of the metal. In the illustration above, for example, iron (II) chloride and all other iron (II) compounds have a greenish color; iron (III) chloride and all other iron (III) compounds have a reddish-brown color similar to that of rust (which, by the way, is just iron (III) oxide).
To demonstrate this phenomenon an experiment called the 'Chemical Chameleon' is conducted. The transition metal used is manganese, Mn; it has the distinction of being one of a few metals with oxidation states up to +7. The manganese compound used is potassium permanganate, aka potassium manganate (VII), chemical formula KMnO4. In solution form it is a purple colour.
In the experiment we add KMnO4 solution to a reducing agent, to reduce it until Mn has a +2 oxidation state. In between are a few transition stages, each with its own colors. Watch the video below to see what they are.
To summarize, the glucose reduces the manganate (VII) ion (MnO4-) with the aid of sodium hydroxide as a catalyst. Manganese undergoes the following reduction stages:
MnO4- + e- ---> MnO4 2- (+6)
MnO4 2- + 2 e- ---> MnO2 (+4) + O2
MnO2 + 2e- ---> Mn 2+ (+2) + O2
And that is the mechanism of the chemical chameleon, a simple experiment explaining a major concept in chemistry.
Flame tests are conducted to determine the identity of the metal present in a salt sample. Different metal cations produce different coloured flames, and their identity can be determined by looking up a reference table.
Please refer to the video for the procedure.
The theory behind the flame test is the same one that explains emission spectra and is an extension of the quantum theory which explains the quantum structure of atoms and ions.
When the salt is heated, the electrons in the metal ions in the salt sample are 'excited', i.e. they gain energy and are promoted to a higher energy level in the ion's electronic structure. These electrons however, are not in a stable state due to their high energy and they consequently release that extra energy after a brief moment, and return back to a lower energy level. The energy released leaves the ions in the form of photons, i.e. light, which causes the salt sample to emit a coloured flame.
The higher the energy released, the higher the frequency of the light emitted, with the lowest frequency being red and the highest violet (or even ultraviolet), following the sequence of the colours of the rainbow.
Different amounts of energy are released depending on the electronic structure of the ion, which is unique to the element of that metal ion. Hence, each metal ion produces its own characteristic colour, which allows us to use the flame test to identify them.
Some common metal ions and their flame colours are listed below:
Lithium - red
Sodium - orange
Potassium - lilac
Calcium - brick-red
Barium - pale green/apple green
Copper (II) - bluish-green with white flame centre if heated at high temperature
Lead - greyish-white
Magnesium - white
Zinc - blue
Some metal ions do not produce flame colours; rather, they produce 'sparks'.
One example is iron (III), which produces gold sparks.
Gunpowder is an ancient Chinese invention, reported to have been first created during the Song dynasty while alchemists (there were no chemists then) were trying to find an elixir of immortality. We highly doubt that gunpowder was able to serve this purpose.
Gunpowder is flammable and explosive, due to the ingredients present in its 'recipe'. Original gunpowder was made from charcoal bits, sulphur and saltpeter. This combination of ingredients creates what is known as a 'self-oxidising fuel', due to the presence of potassium nitrate (KNO3) in the saltpeter. Once ignited the carbon in the charcoal combusts very rapidly, as the nitrate in the saltpeter decomposes to give nitrogen, its oxides, and oxygen, which increase the rate of combustion. At the same time sulphur reduces the temperature needed for ignition, which helps initiate the reaction and raises the combustion rate as well. The end result is a potent mix of solid powders that burn rapidly, creating large amounts of heat and smoke.
As such, gunpowder is used as a propellant, making rockets fly and bullets zip.
In this activity we will explore the science behind the synthesis of gunpowder.
Original-grade gunpowder is first made using the following laboratory chemicals:
Carbon powder 0.15 g
Potassium nitrate 0.75 g
Sulphur 0.10 g
for a mixture totaling 1.0 grams. It is not advised to synthesise more than this amount as a large amount of gunpowder will produce too much thermal energy for lab apparatus to cope with safely. At the same time, larger amounts of gunpowder increase the risk of violent explosion if ignited all at once.
The above chemicals are weighed out and mixed together in a crucible. A glass rod is used to break up any large particles in the mixture to ensure an even distribution of all ingredients. Once mixed properly, the crucible containing the gunpowder mixture is placed in a glass basin, or on an asbestos tile, in a well-ventilated area clear of any flammable materials and ignited by means of a match, and the resulting combustion is observed.
After this, a stronger oxidising agent can be substituted in place of potassium nitrate to alter the rate and intensity of combustion. In this experiment potassium chlorate (V), KClO3, is used. As it is a very strong oxidising agent, ignition is now carried out using a long burning splint to avoid burns.
Welcome to the official weblog of the Catholic High School Science and Maths Society!
Every week we'll be posting about the activities that have been conducted by the society for our visitors' perusal, or (hopefully), viewing pleasure.
Various demonstration videos of experiments conducted by our committee members as well as some in-depth explanations regarding the science behind the activities will be included in the posts, to further make an experience on this blog an interesting one. As such, we do hope that our followers will gain much from reading our posts, and have the interest in Science and Maths sparked in them.
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