The reactions of the Group 2 metals with air rather than oxygen is complicated by the fact that they all react with nitrogen to produce nitrides. In each case, you will get a mixture of the metal oxide and the metal nitride. The general equation for the Group is:. For example, the familiar white ash you get when you burn magnesium ribbon in air is a mixture of magnesium oxide and magnesium nitride.
There are no simple patterns in the way the metals burn. While it would be tempting to say that the reactions get more vigorous as you go down the Group, but it is not true.
The overall amount of heat evolved when one mole of oxide is produced from the metal and oxygen also shows no simple pattern:. If anything, there is a slight tendency for the amount of heat evolved to decrease as you go down the Group. But how reactive a metal seems to be depends on how fast the reaction happens i.
The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. The activation energy will fall because the ionization energies of the metals fall.
In this case, though, the effect of the fall in the activation energy is masked by other factors - for example, the presence of existing oxide layers on the metals, and the impossibility of controlling precisely how much heat you are supplying to the metal in order to get it to start burning.
Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do. There is an increase in the tendency to form the peroxide as you go down the Group. The peroxide ion, O 2 2- looks like this:. The covalent bond between the two oxygen atoms is relatively weak.
Electrons in the peroxide ion will be strongly attracted towards the positive ion. This is then well on the way to forming a simple oxide ion if the right-hand oxygen atom as drawn below breaks off. We say that the positive ion polarizes the negative ion.
This works best if the positive ion is small and highly charged - if it has a high charge density. Ions of the metals at the top of the Group have such a high charge density because they are so small that any peroxide ion near them falls to pieces to give an oxide and oxygen.
See Section 13 for disposal information. Information about protection against explosions and fires: No special measures required.
Conditions for safe storage, including any incompatibilities: Requirements to be met by storerooms and receptacles: No special requirements. Information about storage in one common storage facility: Not required. Further information about storage conditions: Keep receptacle tightly sealed. Store in cool, dry conditions in well sealed receptacles. Specific end use s No data available. Additional information about design of technical systems: No further data; see item 7.
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Instead, YouTube currently has three short video clips showing barium metal burning. All of these show intense white flames with no convincing trace of green.
You might possibly be able to imagine a trace of very pale greenish colour surrounding the white flame in the third video, but to my eye, they all count as a white flame. Anything else that I could find in a short clip from YouTube involved a flame test for a barium compound, irrespective of how it was described in the video.
Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating in oxygen. Mixtures of barium oxide and barium peroxide will be produced. The reactions of the Group 2 metals with air rather than oxygen is complicated by the fact that they all react with nitrogen to produce nitrides.
In each case, you will get a mixture of the metal oxide and the metal nitride. The familiar white ash you get when you burn magnesium ribbon in air is a mixture of magnesium oxide and magnesium nitride despite what you might have been told when you were first learning Chemistry! There are no simple patterns. It would be tempting to say that the reactions get more vigorous as you go down the Group, but it isn't true. The overall amount of heat evolved when one mole of oxide is produced from the metal and oxygen shows no simple pattern:.
If anything, there is a slight tendency for the amount of heat evolved to get less as you go down the Group. But how reactive a metal seems to be depends on how fast the reaction happens - not the overall amount of heat evolved.
The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster.
The activation energy will fall because the ionisation energies of the metals fall. Note: This has been argued through in detail on the page about the reactions of these metals with water or steam.
If you need to know about the reactions with oxygen, you will almost certainly need to know about the reactions with water as well. In this case, though, the effect of the fall in the activation energy is masked by other factors - for example, the presence of existing oxide layers on the metals, and the impossibility of controlling precisely how much heat you are supplying to the metal in order to get it to start burning.
Note: It is interesting to look at what happens if you heat a very reactive metal like potassium in air. The potassium melts at a low temperature and almost instantly turns into a pool of molten potassium oxide. The activation energy is so low that the reaction happens very quickly at quite a low temperature. There is often no trace of flame. It can be fairly boring! Magnesium, on the other hand, has to be heated to quite a high temperature before it will start to react.
The activation energy is much higher. There are also problems with surface coatings. It is then so hot that it produces the typical intense white flame. It would obviously be totally misleading to say that magnesium is more reactive than potassium on the evidence of the bright flame. You haven't had to heat them by the same amount to get the reactions happening. Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do.
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