30 minutes reading time (5957 words)

    NSW Branch Nyholm Youth Lecture Oxygen - Friend or Foe?

    Peter Southwell-Keely
    University of New South Wales


    In this lecture I shall briefly summarise the history of the world from an oxygen point of view, say something of the discovery of oxygen, why we need it, its use as a chemotherapeutic agent and its toxicity. I shall then give some examples of reactions which demonstrate its reactivity and a number of examples of oxidation and reduction.

    It is often said that we live at the bottom of an ocean, an ocean of gas 20% of which is oxygen. However oxygen is not only present in the atmosphere but also makes up 89% of the hydrosphere (oceans, lakes, rivers etc. which cover the planet) and 50% of the lithosphere (the solid part of the earth). In short it is the most abundant element of the earth's crust and, under appropriate conditions, will combine with all the other elements except the inert gases.

    However it was not always so and for approximately the first 800 million years of its total 4.6 billion years existence the earth's atmosphere was anoxic and unfit for any living organism(1). About 2.8 billion years ago the simplest unicellular organisms had evolved to the stage of the blue-green algae which are photosynthetic organisms producing oxygen but the level of oxygen in the atmosphere would still have been less than one percent of its present value. It should be pointed out that oxygen would not have begun accumulating in the atmosphere until the various oxygen sinks such as sulfide gases, sulfide minerals, ferrous iron and reduced gases had become saturated.

    About 2 billion years ago the level of oxygen in the atmosphere had reached about 1% of its present value and a new form of metabolism had arisen to take advantage of it. Prior to this time cells had derived their energy from an anaerobic form of metabolism called "fermentation". However, by making use of oxygen, cells could switch over to an aerobic form of metabolism termed "respiration" and produce 16 times the energy for the same input. An important by-product of the increasing oxygen level was the formation of the ozone layer which absorbs short wavelength radiation from the sun and so protects life on this planet. By making use of the much more efficient "respiration', mode of energy production cells were able to become increasingly more complex with the first nucleated cells (eukaryotes) appearing 1.4 billion years ago followed by the multicellular organisms (metazoans) 670 million years ago. The appearance of the large fishes and land plants 400 million years ago marks the arrival of the fully oxic conditions that we experience today.

    Discovery of oxygen and the phlogiston theory

    Oxygen was actually discovered by the Swedish chemist Scheele in 1772 but the publication of his work did not appear until 1777. Meanwhile in England Joseph Priestley had independently discovered the gas in 1774 and because his work came to the attention of and therefore influenced other scientists before that of Scheele, it is to him that the discovery is usually attributed(2). The discovery of oxygen led to the overthrow of a long held theory, the phlogiston theory. The theory maintained that substances which burned did so because they contained something called phlogiston and when they burned they gave up phlogiston to the atmosphere. This was clear enough when applied to things such as charcoal but could also be applied to metals. A metal was believed to be made up from the metal calx (now known to be the metal oxide) and phlogiston and when it burned it gave up the phlogiston leaving the calx. 

     Similarly, when a metal calx was heated with charcoal it was converted into the metal by absorbing the phlogiston from the burning charcoal.

    The fact that when metals burned, they increased in weight, rather than decreased as the theory would have predicted, was not held to be a major problem with the theory. Thus when Priestley obtained his new gas by heating mercury calx (which had supposedly already lost its phlogiston) he called it dephlogisticated air (all gases were called some kind of air at that time).

    Antoine Lavoisier recognised that the gain rather than loss of weight of metals on heating was a major flaw in the phlogiston theory and he repeated Priestley's experiment applying very careful quantitative techniques. On heating liquid mercury in an atmosphere of air, he found that the gain in weight of mercury in forming its calx (mercuric oxide) was exactly equal to the loss of weight of the air. He then separated the mercury calx and, heating it more strongly, found that the calx was reconverted to liquid mercury and that the loss of weight was exactly equal to the weight of gas produced. Thus he had found the secret of combustion, i.e. that when substances burned, they did not lose phlogiston (or anything else) but combined with a component of air which he called oxygen. The name oxygen is interesting in that it means 'acid producer' and was so named because the first compounds of it known dissolved in water eg. carbon dioxide, sulfur dioxide, nitrogen dioxide, phosphorus pentoxide. Of course, as is now known, this name was premature and misleading since metals which combine with oxygen give basic rather than acidic solutions. Lavoisier's theory of combustion was the complete opposite of the phlogiston theory and took some time to gain widespread acceptance. It is interesting to note that Priestley remained a committed phlogistonist until the end of his life and that he called the final of his many books "The Doctrine of Phlogiston Established".

    Reason for oxygen requirement by aerobic organisms
    As aerobic organisms we have evolved to the point where we cannot exist for more than a very short time in the absence of or in a deficiency of oxygen. An example of our vestigial capacity for anaerobic metabolism can be experienced during severe exercise. Most people will have experienced muscular pain when they have run too far for too long. This is due to the build-up of lactic acid in the muscles and represents the switch-over to anaerobic metabolism in response to a deficiency of oxygen. Why do we require oxygen? When we sit in front of a log fire, we become warm due to the release of heat energy which accompanies the combustion (oxidation) of the carbohydrate cellulose (a polymer of glucose) by oxygen. The process can be represented by the following equation.C6H12O6 + 6O2 → 6CO2 + 6H2O + energy

    Cereals, bread, potatoes, flour and rice all contain the carbohydrate starch which is also a polymer of glucose. Most fruits contain significant amounts of glucose while sweets contain substantial amounts of sucrose, a combination of glucose and fructose. Thus when we eat any of the above we are essentially carrying out the same process as the burning of the log i.e. the oxidation of glucose (fructose is exactly equivalent to glucose) by oxygen which produces energy. Both the burning of the log and the eating of the food can be expressed by equation given above. The difference between the two processes is that all the energy produced by the log fire is released as heat whereas only a small amount of the energy which we derive from burning carbohydrate is released as heat to maintain our body temperature at 37°C while the bulk is stored away in the form of the high energy compound adenosine triphosphate (ATP). The reason for this storage of energy is that many of the chemical reactions which occur in living cells, particularly those which make new compounds, require energy input and this comes from ATP.

    Prior to the advent of the aerobic form of metabolism, organisms still used glucose as their major energy source for the production of ATP. Glucose can be oxidised without oxygen but it is only a partial oxidation and produces one sixteenth the amount of ATP produced in the aerobic mode. An example of an organism which can grow aerobically and anaerobically is brewers yeast. When it grows anaerobically the equation for the oxidation is

    C6H12O6 → 2C2H5OH + 2CO2 + energy

    Only two of the atoms of glucose are oxidised to CO2 compared with all six in the aerobic mode while the other 4 atoms appear as two molecules of ethanol.

    Another major energy source for the aerobic organism is fat and the combustion of fat by oxygen produces more energy on a weight-for-weight basis than does that of carbohydrate. Therefore the major, though not sole, use to which oxygen is put by the aerobic organism is the production of energy.

    Oxygen as a chemotherapeutic agent
    Apart from its ability to sustain the life of aerobic organisms through efficient energy production, oxygen (hyperbaric or high pressure oxygen) has become a weapon in the chemotherapeutic arsenal for the treatment of a variety of pathological conditions.

    A disease associated with warfare from the earliest times and still sometimes observed in automobile accidents is gas gangrene which is caused by organisms of the clostridium species. Because these organisms are strictly anaerobic, oxygen is toxic to them and also to many of the toxins which they produce.

    Often, non-healing wounds and diseases such as chronic osteomyelitis are due to a low level of oxygen in the tissues. Major agents in a body's defence mechanism, as we shall see later, are the phagocytic white cells which engulf and destroy foreign compounds and organisms. This defence process is stimulated by oxygen in the tissues and, conversely, retarded by low levels of oxygen. Therefore an effective treatment is to increase the oxygen in the tissues by exposure to hyperbaric oxygen.

    Hyperbaric oxygen has also been shown to aid in the process of skin grafting and it is now common to expose the affected skin for a week prior to and after the operation.

    Probably the most common use of hyperbaric oxygen is in the treatment of carbon monoxide poisoning. Although carbon monoxide binds more strongly than oxygen to haemoglobin, it can easily be displaced from haemoglobin by the levels of oxygen attained in the hyperbaric chamber.

    In Japan hyperbaric oxygen is used to treat cerebral oedema where pressure builds up in the head due to excessive fluid. Oxygen treatment brings about a reduction in pressure by causing a vasoconstriction of the blood vessels and thereby reducing the pressure from that source.

    Although of no value by itself, oxygen has also been shown to be a useful adjunct in the treatment of cyanide poisoning.

    Another potentially valuable use of oxygen therapy may be in the treatment of Stroke (cerebral haemorrhage). A new approach relies on the fact that the brain is bathed by spinal fluid inside and out. This spinal fluid serves to bring nutrients to the brain and also as its sewage system to carry the waste away. In this approach, the brain is injected with a fluid containing a much higher than normal level of dissolved oxygen while spinal fluid is drained out through a needle in the back. With this technique animals which had been legally dead (absence of an encephalogram for several hours, the absence of a response to pain stimuli and the absence of spontaneous movement, breathing or reflex action) for several hours were revived with no apparent side effects. It now appears that, under conditions of Stroke, the brain shuts down its level of activity and goes into a sort of hibernation but that it does not necessarily die.

    Oxygen toxicity
    So far the effects of oxygen on aerobic organisms have been described in a positive or beneficial sense. The fact that oxygen is essential for our existence and that we take it for granted tends to obscure the fact that it is a very reactive and potentially dangerous element. The reason that we are able to regard it as being so benign is that we, and all aerobic organisms, have developed mechanisms to guard against its potentially deleterious effects. However it is distinctly possible that oxygen and more particularly its closely related metabolites, superoxide, hydrogen peroxide and hydroxyl radical are intimately involved in the so-called autoimmune diseases, diseases in which the body reacts against one or more of its own tissues or metabolites. One of these diseases, rheumatoid arthritis, is probably well known to many people(3).

    In rheumatoid arthritis an abnormal immunoglobulin is synthesised and the body reacts by forming an immunoglobulin (antibody) against it. These two proteins combine to form what is termed a soluble-immune-complex (SIC). In a normal individual this SIC complex is ingested and destroyed by phagocytic leukocytes and poses no problems. In rheumatoid individuals, (the majority of whom are women), the phagocytic leukocytes fail to destroy the SIC and are themselves destroyed in the attempt. They break open and liberate their contents into the surrounding medium. In the absence of cellular control many of the liberated materials are capable of producing an inflammatory response. In the case of the rheumatoid sufferers the major inflammation is observed in the joints. This is because the SIC tend to be retained at sites of filtration such as the joints and the leukocytes are attracted to these sites and inflammation develops there.

    Of the compounds used by leukocytes to destroy invading species, whether they be foreign compounds or microorganisms, some are derived from the reduction of oxygen as follows:

    O2 O2- H2O2 *OH H2O
    oxygen superoxide hydrogen peroxide hydroxyl water

    Both superoxide and hydroxyl are species known as free-radicals and are very reactive. The hydroxyl radical is regarded as particularly dangerous since it can react quite indiscriminately with almost any compound, often with deleterious results. Thus, as part of their armoury for destroying foreign compounds etc, the leukocytes produce some dangerous compounds. I should point out that these oxygen free radicals are normal metabolites of oxygen and the cells which produce them have other defensive compounds which prevent excessive production of the radicals. However, too great a production of these radicals would be dangerous both for the cells producing them and also for the surrounding cells.

    For example, all cells are surrounded by a cell membrane which is made of a fatty material called phospholipid. It has been clearly demonstrated that both superoxide and hydroxyl can stimulate the oxidation and breakdown of this membrane which leads to the death of the cell.

    In all of the joints of the body there is a special lubricating substance known as hyaluronic acid which is related to common starch. Both of these compounds form thick, viscous solutions because they are very large, slow-moving molecules which tend to attract other molecules of the same kind into forming larger, viscous aggregates. If these molecules are chopped up into smaller molecules, they become much faster-moving and, as a result, less viscous, and less good lubricants. Both superoxide and hydroxyl radicals have been shown to break up hyaluronic acid into smaller pieces, with consequent destruction of its lubricating properties. For rheumatoid sufferers this is of critical importance since destruction of the lubricating substance means that the bones on each side of the joint start rubbing against each other, creating further inflammation and leading to the painful distortions of the joints which are such a common feature of this disease. Thus once the condition has developed it tends to become self-promoting.

    It should be clear from the above examples that the oxygen free radicals which are so important to the defensive capability of the properly functioning leukocyte are dangerous if produced in excess. It is thought that the production of oxygen radicals may well play a role not only in rheumatoid arthritis but in other autoimmune diseases.

    Two other massive areas of research of great interest to us all, and which I shall mention only in passing, are those of ageing and of cancer(4). During the course of a lifetime unsaturated fatty substances in the body are subject to continuous attack by oxygen and give rise to fatty oxygen radicals which are also indiscriminantly reactive and dangerous. Compounds known as antioxidants (vitamins E and C, β-carotene, glutathione, glutathione peroxidase, superoxide dismutase etc) protect us from the harmful effects of these radicals. However over many years deleterious reactions gradually build up to the point that the body can no longer repair and this in turn is believed to promote ageing and cancer.

    To summarise so far, I would remind you that the title of this lecture is a question not a statement. While it is obvious that we cannot survive without oxygen and that we tend to regard it as innocuous, it is in reality a very reactive element and may in fact contain the seeds of our ultimate demise.

    Properties of oxygen
    The power of oxygen to support combustion is readily demonstrated by burning two wax tapers side by side, one in an inverted conical flask (500 mL) filled with air and the second in an identical flask filled with pure oxygen. If the room lights are extinguished it can be seen clearly that the taper burning in oxygen burns with a brilliant, white, crackling flame for about one minute while the taper burning in air has a yellowish flame which lasts only 10-20 sec. In addition the heat of the oxygen flame is much greater and results in significant wax melt and loss of length (3-4 cm) of the taper. The taper burning in air does not show much wax melt at all and loses hardly any of its length.

    A much more spectacular (and dangerous) demonstration of the ability of oxygen to support combustion is to pour some liquid oxygen onto a small piece of burning cotton wool (previously doused in a little methanol) in a 1L beaker contained in a metal tray for safety. What was previously a small yellow flame becomes a large, crackling, violent conflagration. Liquid oxygen is easily generated by leading gaseous oxygen into a gas trap which is cooled in liquid nitrogen. Since oxygen boils at a slightly higher temperature (-183°C) than nitrogen (-196°C), it is condensed in the trap. In 30 - 40 minutes several hundred mL of oxygen can be collected. When sufficient liquid oxygen has been collected, it is advisable to stopper (glass) the flask until use as an open flask will dilute itself by condensing water out of the atmosphere. I cannot emphasise strongly enough how dangerous liquid oxygen is and that it must be kept away from oils, greases and other contaminants.

    If the liquid oxygen is collected in a colourless glass container, the pale blue colour of the element is clearly visible. The blue colour of oxygen is due to the fact that it contains two unpaired electrons and is therefore a free radical and all long-lived radicals are coloured.

    The unpaired electrons of oxygen cause it to be paramagnetic (weakly magnetic). This may be demonstrated by putting liquid oxygen into a test tube which is suspended between the poles of a fixed magnet. If the magnet is mounted on a 'lazy susan', its poles may be turned towards the test tube. As the poles of the magnet are moved closer to the test tube, it swings towards them. Although oxygen is magnetic, it has no pole itself and will therefore move towards either the north or south pole of the magnet.

    Oxidation and reduction
    Oxygen gives its name to the process known as oxidation. I shall now define oxidation and give a number of examples which lend themselves to demonstration.

    When substances burn they are actually reacting with oxygen and being converted into different species. This type of reaction is called an oxidation. Many, possibly the majority, of the reactions which we encounter in chemistry, in biology and in our daily life are oxidation reactions. Oxidation means different things to different people. For an organic chemist oxidation occurs when a compound either combines with oxygen or, alternatively, loses hydrogen. Of course if one compound is combining with oxygen, another compound must be losing oxygen. Similarly, if one compound is losing hydrogen another compound must be gaining hydrogen. The process in which a compound either loses oxygen or, alternatively, gains hydrogen is called reduction and it is impossible to have an oxidation without an accompanying reduction.

    The simplest oxidation/reduction (redox) reaction, and also one of the most spectacular, is that between hydrogen and oxygen:

    2H2 + O2 → 2H2O + energy (286 kJ/mol)

    The two gases do not react at all unless a source of heat is supplied and then they react with a very loud bang and the evolution of a great deal of energy. The hydrogen combines with oxygen to form dihydrogen-oxide (water). Thus, according to the definition, hydrogen has been oxidised. By contrast, oxygen has combined with hydrogen to form oxygen dihydride (water) and therefore the oxygen has been reduced. As a demonstration it is useful to have 3 balloons, one filled with balloon gas (helium), one filled with hydrogen and one filled with a mixture of two parts hydrogen and one part oxygen. For best effect the balloons should be suspended by string several metres above the demonstrator and ignited successively in the above order with a lighted wax taper attached at the end of a long rod.

    The burning of the wax taper is typical of an organic oxidation which we can illustrate with the equation for the burning of methane (the simplest hydrocarbon) gas as follows:

    CH4 + 2O2 → CO2 + 2H2O + energy (891 kJ/mol)

    In this reaction, which also requires heat, the carbon atom of methane (CH4) has combined with oxygen from the air to form carbon dioxide and therefore has become oxidised. The hydrogen of the methane has combined with oxygen to form dihydrogen oxide (water) and has also been oxidised. However it is impossible to have oxidation without reduction and oxygen has combined with hydrogen to form oxygen dihydride (water) and has been reduced.

    Inorganic chemists accept the above definition of oxidation/reduction but require a more general one to cover similar types of reactions which take place without any involvement of oxygen or hydrogen. Thus oxidation may also be defined as the loss of electrons and reduction as the gain of erections. The reaction between iodine (I2) and sodium thiosulfate (Na2S2O3) is a very important analytical oxidation/reduction reaction which does not involve the gain or loss of either oxygen or hydrogen:

    I2 + 2Na2S2O3 → 2NaI + Na2S4O6

    To see what has been oxidised and what reduced, we can separate this equation into two half equations:

    I2 + 2e → 2I-
    2S2O32- -2e → S4O62-

    Iodine has gained electrons and has been reduced to iodide (I-) while thiosulfate has lost electrons and has been oxidised to tetrathionate (S4O62-).

    Oxygenation of haemoglobin - a Clayton's oxidation?
    The major function of blood is to transport oxygen to the tissues of the body, all of which need oxygen for their metabolism. Blood is mostly water and water is capable of dissolving very little oxygen. However the body can increase the amount of oxygen in water more than fifty times by binding it to a red, iron containing protein called haemoglobin. Thus blood is red because of the haemoglobin in it. Blood which issues from a wound near the surface of the body or is donated at the Blood Bank is venous (i.e. from veins), is devoid of oxygen and is a deep red, almost purple colour. When venous blood is pumped through the lungs it picks up oxygen and changes to a bright, mail box red as the oxygen binds with the haemoglobin. The reaction may be expressed as follows (where Hb = haemoglobin)

    or Hb(FeII) + O2 Hb(FeII).O2
    (dark red) (bright orange red)

    Since oxygen binds to the iron of the haemoglobin, it can be said to be oxidised according to the organic chemical view of things and indeed the colour change of the solution tells us that a reaction has occurred. However the iron does not change its oxidation state, which remains at +II. In the inorganic chemical view, oxidation has not occurred and in this sense it is a Clayton's oxidation - the oxidation you have when you're not having an oxidation. In fact it is most important that iron is not oxidised in this reaction for, if it were, a new compound, methaemoglobin, would be formed which is unable to bind and transport oxygen. The ability of the FeII in haemoglobin to react with oxygen without becoming oxidised is a particularly clever trick on the part of the haemoglobin when it is realised that FeII salts are very readily oxidised by oxygen in neutral and alkaline solutions. It is very easy to demonstrate the oxygenation of haemoglobin by putting venous blood into two equivalent, large test tubes and swirling one of them by hand for several minutes. This is all the time needed for the bright red oxyhaemoglobin to form.

    Browning of fruits
    Most people are aware that, when the surface of a fruit is exposed either by peeling or chewing, it will, if left uneaten, turn brown some little time later. The speed of browning is determined by the fruit concerned. Apples that have been kept in cold storage seem to brown rather quickly (within 5-15 minutes). This reaction (actually a series of reactions) is an enzyme-catalysed oxidation which leads to the formation of brown pigments called melanins. This sequence of reactions is the same as that which causes us to become brown and may be abbreviated as follows

    Tyrosine is a colourless amino acid which is widely distributed in the body, mainly in proteins. It is oxidised (adds an atom of oxygen) by oxygen in the presence of the enzyme tyrosinase to dihydroxyphenylalanine (DOPA), also a colourless compound. DOPA is then oxidised (in this case it loses two hydrogen atoms) by tyrosinase to DOPAquinone, a red compound. A number of transformations then occur, which are beyond this discussion, resulting in the brown melanin pigments. Many people will be aware that this browning reaction in fruit and vegetables can be prevented by squeezing a little citrus (i.e. orange or lemon) juice over the item in question. The citrus juice probably functions in two ways, one by being acidic and two by being a reducing agent. All enzymes work best at a certain acidity (pH) termed the pH optimum which is a different value for each enzyme. At pH values either less than or greater than the pH optimum the efficiency of the enzyme drops dramatically. Since citrus juices are quite acidic, they would certainly lower the pH below the pH optimum for tyrosinase, thereby inhibiting the oxidation. However the citrus juices also contain the strong reducing agent, ascorbic acid (vitamin C). Ascorbic acid can reduce (add two hydrogen atoms to) DOPAquinone back to DOPA thereby inhibiting the reaction. It is important to note that the process cannot be reversed once the brown pigments have been formed.

    Red wine into white wine
    This actually happened to me once when I was making red wine. The colour of red wine is due to the anthocyanin pigments in the skins of the red grapes. Below is the structure of a typical anthocyanin, pelargonin. The casks or containers in which the fermentation takes place must be sterilised of all stray microorganisms or the wine may develop an unpleasant taste or odour due to by-products of these stray organisms. The sterilising agent most often used is potassium metabisulfite (K2S2O5) which turns into two molecules of potassium bisulfite (KHSO3) when dissolved in water and which, amongst other things, is a very good reducing agent. My problem was that I used too much metabisulfite and, as a result, converted the anthocyanin pigments to their reduced form which is colourless. Consequently the wine turned out white rather than red. This reaction can be demonstrated very simply using inexpensive red wine from a cask or, for a clearer result, with Ribena (5.1%) blackcurrant juice drink. The pigments in Ribena and red wine are the same. For demonstration purposes, sodium dithionite (Na2S2O4) gives a faster and clearer reaction than metabisulfite. A few crystals of dithionite will decolourise (reduce) Ribena within seconds.

     Anthocyanin pigments occur combined with sugars. In the structures above the sugar is glucose (Gl). The reduction of an anthocyanin pigment by bisulfite can be represented by the above equation with pelargonin as the representative anthocyanin. The structure on the left is a red colour while that on the right is colourless.The reaction can be best understood as the reduction of the positively charged ring oxygen and the oxidation of the negatively charged sulfur in the bisulfite.

    Bleaching of blue cotton
    Another true life adventure. Recently my daughter bought a white cotton jacket and attempted to dye it a blue colour. The dyeing process worked but my daughter was disappointed in the shade of blue. My wife had recently bought a new bleach and offered to remove the blue colour from the jacket. She soaked the jacket in a solution of the bleach for an hour or two. The colour came out very well and we thought we had achieved a brilliant success. However, after the excess bleach had been washed out of the jacket and it had been drying for a relatively short time, the blue colour gradually began to return and was soon as blue as it had originally been. The reason for this is that there are two types of bleach, an oxidising type (such as sodium hypochlorite, (NaOCl) and a reducing type, normally sodium dithionite (mentioned above). We had used a reducing bleach which had added two hydrogen atoms to the blue pigment molecule and turned it, not exactly colourless, but a pale yellow. However this particular pigment can be easily reoxidised to the blue form by oxygen simply by allowing it to stand in air. Had we used an oxidising bleach the colour would have been removed and would not have returned because the blue pigment molecule would have been broken into smaller, colourless fragments which could not recombine. I am not absolutely certain that the pigment in this reaction was indigo but the following equation using indigo will demonstrate the point.

     Indigo is very easily reduced (adds two hydrogen atoms) by sodium dithionite and just as easily reoxidised by the oxygen in air.

    Making a silver mirror
    Excellent silver mirrors can be made using the Tollens reaction in which aldehydes reduce silver ions to silver metal. The secret of a good mirror is clean glassware. Wash the vessel in concentrated nitric acid, rinse thoroughly with water and distilled water and allow to drain and air dry. In a 15 x 2.5 cm test tube mix 1 mL of Tollens reagent A (10% silver nitrate) and 1 mL of Tollens reagent B (10% sodium hydroxide) to form insoluble, brown silver oxide. Add 2M aqueous ammonia sufficient to not quite dissolve the silver oxide. Then add to the tube sufficient glucose to cover three thumb nails to a depth of about 3 mm and swirl at room temperature with a circular motion. Almost immediately black, finely divided silver metal begins to precipitate and after several minutes a beautiful silver mirror begins to appear on the surface of the glass. The mirror should have covered the sides of the tube to a height of about 3-5 cm within 10 minutes. Glucose is a particularly good aldehyde to use in this test because it deposits the silver slowly and therefore allows a fine uniform mirror to develop. Other aldehydes tend to precipitate the silver very quickly and often give a blotchy mirror. The reason that glucose reacts more slowly than other aldehydes is that glucose is really five compounds in equilibrium with each other. The three major components of the equilibrium are the α-D-glucopyranose (36%), β-D-glucopyranose (64%) and straight chain glucose (0.02%) and, of these, only straight chain glucose (an aldehyde) has reducing properties. As the small amount of straight chain glucose reacts and is, thereby, removed from the equilibrium, more of the α- and β-D- glucopyranoses undergo ring opening to form straight chain glucose and reestablish the equilibrium allowing more straight chain glucose to react and so on. Eventually all of the glucose reacts thus behaving as if it had been 100% straight chain glucose in the first place. This is an excellent example of Le Chatelier's principle. It should be noted that Tollens reagent (ammoniacal silver nitrate) becomes explosive on standing and should be disposed of as soon as it has been used and washed down the sink with plenty of water.

    Copper into silver into goldThis was the dream of the Alchemists and is a popular demonstration particularly with younger people(5). This demonstration works best with old pennies (pre 1966) which, for optimum effect, should be polished before use to remove surface grime and copper oxide which would inhibit the reaction. Place 25g granulated zinc in a 150 mL beaker and add 50 mL of 3M NaOH. Heat (Bunsen burner) until the solution begins to boil and place two clean copper coins on the zinc. Leave in the hot solution for 3-4 minutes with occasional turning. The coins should be silver by this time and can be removed from the solution, washed with water and blotted dry. Place one of the coins in a Bunsen flame and heat carefully until it turns gold. It is very important not to heat too long or the coin becomes a dirty yellow colour. Then cool under a tap or in a beaker of water. Display an original copper coin together with the newly formed 'silver'and 'gold' coins.

    This is a very interesting reaction and not completely understood. The silver colour which is observed first is due to the deposition of zinc metal on the surface of the copper coin. When the zinc is heated in the presence of copper it forms the alloy brass which is a yellow colour.

    Zinc will dissolve in alkaline solution forming sodium zincate (Na2ZnO2) and liberating hydrogen gas

    Zn(s) + 2OH- → ZnO22-(aq) + H2(g)

    Thus zinc is a reducing agent and reduces protons in the water to hydrogen while being itself oxidised to Zn(II). It seems that, in addition to donating electrons to protons, the zinc passes electrons to the copper and then on to the zincate which is thereby reduced back to zinc metal.


    1. Cloud, P., "The biosphere". Sci. Am. 1983, 249, 132-144.
    2. Scott, A.F., "The invention of the balloon and the birth of modern chemistry". Sci. Am., 1984, 250, 102-111.
    3. Halliwell, B. and Gutteridge, J.M.C., "Free radicals in biology and medicine". 2nd edit 1989, Clarendon Press, Oxford.
    4. Harman, D., in "Free radicals, aging and degenerative diseases" (Johnson, J., Harman, D., Walford, R. and Miguel, J. eds) 1986, pp 3-49.
    5. Summerlin, L.R., and Ealy, J.L., "Chemical demonstrations. A sourcebook for teachers". vol.1, pp 137-138.

    About the author

    Peter Southwell-Keely obtained his B.Sc. from the University of Sydney in 1964 and his Ph.D. from the University of N.S.W. in 1968. He then did post-doctoral work at the Veterans Administration Hospital, Long Beach, California with Prof. Klaus Schwarz on the nutritional biochemistry of selenium and at the Max Planck Institute fur Zellchemie in Munich with Prof. Feodor Lynen on fatty acid biosynthesis. He returned to the University of N.S.W in 1972 as a Lecturer and is currently Senior Lecturer in Organic Chemistry. His research interests include free radical reactions of organosulfur compounds, the synthesis of organoselenium compounds for use as antifungal reagents, the biochemistry of rheumatoid inflammation and the mode of action of vitamin E. He enjoys doing chemical demonstrations on Open Days and for high school students and had great fun doing the Nyholm youth lectures in 1993 ably supported by Chris Owens (UWS Nepean). He is now retired. 

    NSW Branch Nyholm Youth Lecture and SA Branch D.R....


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