The Iodine Clock

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.

Oxidation states and the Chemical Chameleon

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

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.

Of Gunpowder, Treason and Plot.. well maybe not the last two...

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!

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.

To navigate just refer to the Labels box in the sidebar.

Enjoy!