Can burn oxygen with a match

The role of oxygen in oxidations
Demonstrations 3 and 4 are only suitable for teachers
 
  
 
Fabrics Iron wool, oxygen, sulfur, flint made of cerium iron (from old gas lighters, available from welders), iron wire, quartz sand, potassium nitrate, granulated activated carbon, blue ink
equipment Beam balance, tripod, 4.5 volt battery, candle, wood chips, scotch tape, stopper, round bowl, combustion spoon, crucible tongs, several standing cylinders with cover, yoghurt jar with lid, burner, knife, 2 tripods, round filter, precision balance, hard-to-melt test tube 18 × 180 mm, 2 test tubes 20 × 180 mm with pierced stoppers, 2 long glass tubes, 2 beakers 250 ml
security The instructions for handling laboratory gases and explosives must be read. The respective safety precautions will be discussed separately in the demonstrations.

   
GBU Germany Burn substances in pure oxygen docx    pdf
GBU Germany Reactions in a potassium nitrate melt docx    pdf
SB Switzerland Burn substances in pure oxygen docx    pdf
SB Switzerland Reactions in a potassium nitrate melt docx    pdf
 

 
The demonstrations on the law of mass conservation showed that mass is not lost. A candle on a scale becomes lighter, but part of the mass is found in the combustion gases after the candle has burned down. The introduction to the subject of oxidation follows on from this experiment. This time a tuft of iron wool is hanging on a scale and the students are asked to guess what happens when the iron wool burns. However, the attempt initially provokes a contradiction.
 
 
Demonstration 1 Iron wool burns on a scale

This attempt must be on a fireproof pad as burning iron parts can fall down. Iron wool must never be stored with batteries, as these can ignite the iron wool.

Two equal pieces of iron wool are attached to a beam balance. Now you can show how a small piece of iron wool (without scales) burns by igniting it with a 4.5 volt battery. Have the students guess what will happen if one of the large pieces of iron wool on the beam balance is ignited. Various arguments are certainly put forward:
  • The iron wool becomes lighter because it burns, consequently the scales on the side of the iron wool that burns become lighter. This argument approximates the outdated "phlogiston theory", according to which it was believed that the combustible principle, the "phlogiston", escapes in the event of burns. (For more details see> Oxygen)
  • The scales remain balanced because the masses are preserved.
  • The iron wool becomes heavier because oxygen is chemically bound and something is added.
At this point, the students will hardly bring up the last argument unless a student has already dealt with it in detail. However, the result of the demonstration must not be revealed under any circumstances. You can have a vote carried out for the three positions.
  
 
 
Movie
 
 
Observations: To the amazement of some students, iron wool becomes heavier when burned. And this happens even though many glowing pieces of iron wool fall off and are not weighed with it. How can you explain that? When examining the reaction product, changes in color can be seen. You can no longer light it (if it has burned out completely).
 
 
Scales after burning the iron wool
 
Theory: Even if you as a teacher already know why the iron wool is heavier (iron reacts with oxygen to form iron oxide), a possible explanation should be withheld from the students for the time being. Instead, another attempt to clarify the problem follows.
 
 
Demonstration 2 A candle burns in a locked room

A candle burned on a scale during the Mass Preservation demonstration. This was an open system, the combustion gases could escape. The aim now is to show what happens when a tea light burns in a closed system.
 
A candle is attached to a wooden chip with scotch tape (see picture), the chip is put into a plug and this is attached to the bottom of a round bowl with wax. Then you fill the bowl 1 cm high with water and light the candle. You show the audience an upside-down cylinder and ask what will happen if you put the cylinder on the candle. Again, assumptions are made:
  • The candle flame goes out immediately.
  • The candle flame burns for a while and then goes out.
  • In doing so, water is drawn into the standing cylinder.
  • The water then completely fills the cylinder because all the air is used up.
  • The water only rises a little because not all of the air is consumed.
  • The water rises so and so many percent ...
Observations: As some have surely correctly suspected, the candle flame goes out after a while, while the water rises less than a fifth of the volume into the standing cylinder.
  
 
 
 
Theory: By following on from the discussion, arguments can now be found for the exact description of the burning process in the candle: Air is required for combustion. However, it is only part of the air (approx. 20%), we call this proportion oxygen, the remaining gas (approx. 80%) is called nitrogen. When all the oxygen in the cylinder is used up, the candle flame goes out. Then there are nitrogen and carbon dioxide in the cylinder. Both gases do not sustain the combustion. A graphic of the air composition is shown to provide more information.
   
 
 
 
In connection with the experiment with iron wool, an explanation can now be sought: When iron wool is burned, oxygen is also required. This reacts with the iron to form iron oxide. The scale shows a higher mass because oxygen atoms are chemically bound with the iron. We call reactions with oxygen oxidations. Oxides are formed as reaction products:
 
Reaction scheme for burning the iron wool:

Iron + oxygen Iron oxide
 
Reaction scheme for burning the soot in the candle flame:
 
Carbon + oxygen carbon dioxide
 
 
Demonstration 3 substances react with pure oxygen
 
The test series must be carried out on a fireproof pad be performed. A safety goggles is necessary when reacting with cerium too Protective gloves made of leather and a Protective screen. One shouldn't look directly into the flame. When the oxygen reacts with sulfur, toxic sulfur dioxide is produced, so this experiment should be carried out in Deduction The remaining sulfur, which may still be burning, is burned over the roaring burner flame in the fume cupboard.

Several standing cylinders are filled with pure oxygen and made available. A little sand is put into each cylinder to protect the soil. Although the oxygen should remain in the cylinders due to its higher density than air, these are covered with round panes of glass so that the oxygen does not escape through air circulation.
 
a) A burning candle is immersed in a cylinder filled with oxygen on a combustion spoon.
b) A tuft of iron wool is held with crucible tongs, made to glow at one point and immersed in the standing cylinder filled with oxygen.
c) In the fume cupboard, half fill a combustion spoon with sulfur and ignite it with a burner flame. Then you hold the burning sulfur in the standing cylinder filled with oxygen.
 
   
 
d) A yoghurt glass is filled with pure oxygen and screwed shut. Then you wrap a wire around a flint made of iron and heat the flint until it glows in the non-luminous burner flame. The wire with the glowing flint is dipped into the yogurt glass with pure oxygen.
 
 
 
Cerium iron reacts with pure oxygen. >Movie
 
 
Observations:
a) The candle flame gets brighter and hotter so that wax drips onto the floor. A smoldering chip also glows brightly and even catches fire. The students practice the chip test in the student exercise on oxygen production with potassium permanganate at another point in time.
b) The iron wool burns up with a lively spark reaction.
c) The sulfur burns with a bright blue flame.
d) The Cereisen burns explosively with a very bright flame, with the yoghurt glass shattering due to the heat shock.
 
 
 
It is sufficient for the students to discuss the reaction scheme in words at this point. When cerium reacts with oxygen, temperatures of up to 4000 ° C arise. This will destroy the yogurt glass. This reaction is used in flints of lighters or gas lighters to ignite the gases:
  

Flint of a gas lighter.



Disposal: The remaining sulfur is carefully burned in the fume cupboard. This creates highly toxic sulfur dioxide! The remnants of the iron wool are only disposed of in the waste when it is certain that there is no longer any embers. Finely divided cerium and the product from the experiment with cerium iron can be pyrophoric. Since the reaction is usually almost complete, it is sufficient to rinse off the remains, the base and also the remains of the glass with plenty of water and then dispose of the glass in the container for broken glass. It is generally recommended to collect solid waste and glass separately in a metal bin with a metal lid.

The subsequent experiments are less suitable for school, as toxic or problematic waste is produced. Therefore, the films are ideal. The highest chemically achievable temperature is obtained when burning zirconium powder in pure oxygen (4660 ° C):


Zr + O2  ZrO2     ΔHR. = −1101 kJ / mol
 
 
40Zr
Zirconium
Movie
14 sec
A little zirconium wool in a porcelain bowl is touched with a flame.
 
 
Variations: There are films in which further examples of oxidations are shown.
 
   
 
Demonstration of 4 reactions in a potassium nitrate melt
 
Burns are well known to students from various fields. Probably the most fascinating are the burns that occur with explosives or rocket propellants. The reaction between hydrogen and oxygen found a technical application in the American space shuttle. The space shuttle, which was retired in 2011, carried a hydrogen and oxygen tank with it. The two gases were ignited in the engine and generated the energy required for locomotion.
  
 
 
 
When the space shuttle started, the solid rocket rockets on the side ignited. Solid rocket propellants carry substances with them that contain chemically bound oxygen atoms and act as oxidizing agents. This experience is known from the student exercise on the production of oxygen from potassium permanganate. You also need a fuel, such as coal or sulfur, which provides the necessary energy when burned with the oxygen. First a (possibly burned out) fireworks rocket is shown and the structure is demonstrated. Attempts to cut open fireworks are not recommended.
 
Burning black powder produces large amounts of gas that blow up a container. This creates the bang at the firecrackers. If there is a small hole in the container of a rocket, the gas flows out of this hole and a forward thrust is generated so that the rocket flies. The chemical reaction equation for burning black powder is so complicated that it should not be stated. Combustion gases such as carbon monoxide, carbon dioxide and nitrogen are produced. Burning 10 g of black powder produces around 2.8 liters of these gases (see also information on the demonstrations on explosives).
 
Attempt: Are there safety goggles and Protective gloves to wear, the goggles requirement also applies to spectators, unless there is one Protective screen used. The aim is to show that the components of black powder can react violently with one another. To do this, put 2cm high potassium nitrate into a difficult-to-melt test tube (18x180mm) and heat it with the rustling burner flame until a clear melt is formed. Then you take the burner away. In a series of experiments, a small paper ball, a small piece of wood (half a match without a fuse head), 3 grains of activated charcoal and a spatula tip of sulfur are thrown into the melt one after the other. You wait after each reaction and heat again until the melt is clear again. In all cases the substances burn vigorously with phenomena of fire. The demonstration shows the oxidative effect of potassium nitrate.
 
 
 
The grain glows and bounces around violently.
 
The sulfur burns with bright phenomena of fire.
 
Movie
 
 
Disposal: Small test tube residues of the potassium nitrate melt can be diluted with plenty of water and thrown down the drain after they have cooled down completely. If nitrate solutions are collected in a container for disposal, it must be ensured that the pH value is alkaline (pH = 8), as hydrogen cyanide can form in acidic solutions.

Note: Due to current law, black powder cannot be made in class. Instead, you can show films about it.
 
 
 
Movie
 
 
Demonstration 5 The rusting process as an example of slow oxidation
 
Slow oxidations are also common in nature. During the respiratory and digestive process, glucose is transported to the muscles and the brain via the bloodstream. There, the glucose oxidizes to carbon dioxide with the help of the blood oxygen that comes from breathing. This releases energy. Dextrose contains chemically bonded carbon atoms that oxidize to carbon dioxide, which is released into the bloodstream and the lungs.
 
Execution: Another form of slow oxidation is the rusting process. Two test tubes are filled with iron wool. The iron wool is additionally moistened with water in a test tube. The two test tubes are closed with a pierced stopper and long tubes are provided. These are immersed in water colored with blue ink. The entire apparatus is left to stand for several weeks and observed.
 
 
 
 
Observation: After two to three weeks, the iron wool moistened with water shows clear signs of rust. The colored water in this tube has risen, while there are no changes in the other test tube.
  
 

Rust efflorescence on an iron plate.
 
 
Theory: During the rusting process, oxygen is consumed in a complex chemical reaction. Rusting is also an oxidation, which releases heat. The reaction scheme shows the complicated rusting process in a very simplified way:
  
Iron + water + oxygen Iron hydroxide (+ energy)
 
The iron hydroxide then reacts with the oxygen in the air to form iron (III) oxide (Fe2O3) further. Other iron oxides are also formed, for example iron (II, III) oxide (Fe3O4) or iron (II) oxide (FeO).
  
Outlook: Demonstrations and fire-fighting experiments are suitable for deepening the topic. The topic of reductions also follows. Variations: If you put an iron nail in an open vessel with cold, oxygenated water, the nail rusts faster than if it is placed in a closed vessel with boiled water.