One major web source describes rubidium superoxide as being dark brown on one page and orange on another! The rubidium doesn't show a clear flame colour in this video, although the caesium does show traces of blue-violet.īoth superoxides are described in most sources as being either orange or yellow. There is a bit of video from the Royal Society of Chemistry showing the two metals burning on exposure to air. The equations are the same as the equivalent potassium one. The formula for a superoxide always looks wrong! There is more about these oxides later on.īoth metals catch fire in air and produce superoxides, RbO 2 and CsO 2. The formula for a peroxide doesn't look too stange, because most people are familiar with the similar formula for hydrogen peroxide. I don't recall ever seeing it yellow or orange! I have a bit of a problem with this, because over my teaching career I have heated potassium in air many times and, if memory serves correctly, it always leaves a greyish white film on the bit of porcelain you are heating it on. Note: Potassium peroxide and superoxide are described as being somewhere between yellow and orange depending on what source you look at. The equation for the formation of the peroxide is just like the sodium one above: Larger pieces of potassium burn with a lilac flame. Small pieces of potassium heated in air tend to just melt and turn instantly into a mixture of potassium peroxide and potassium superoxide without any flame being seen. The equation for the formation of the simple oxide is just like the lithium one. You get a white solid mixture of sodium oxide and sodium peroxide. Using larger amounts of sodium or burning it in oxygen gives a strong orange flame. Small pieces of sodium burn in air with often little more than an orange glow. Use the BACK button on your browser to return to this page from either of these links. You will find this discussed on the page about electronegativity. There is a diagonal relationship between lithium and magnesium. Lithium's reactions are often rather like those of the Group 2 metals. You will find what you want about 3/4 of the way down that page. Note: You will find the reason why lithium forms a nitride on the page about reactions of Group 2 elements with air or oxygen. Lithium is the only element in this Group to form a nitride in this way. With pure oxygen, the flame would simply be more intense.įor the record, it also reacts with the nitrogen in the air to give lithium nitride. It reacts with oxygen in the air to give white lithium oxide. Lithium burns with a strongly red-tinged flame if heated in air. Lithium is unique in the Group because it also reacts with the nitrogen in the air to form lithium nitride (again, see below). Reaction with oxygen is just a more dramatic version of the reaction with air. The tubes are broken open when the metal is used.ĭepending on how far down the Group you are, different kinds of oxide are formed when the metals burn (details below). They are stored either in a vacuum or in an inert atmosphere of, say, argon. Rubidium and caesium are normally stored in sealed glass tubes to prevent air getting at them. It is, anyway, less reactive than the rest of the Group.) (Lithium in fact floats on the oil, but there will be enough oil coating it to give it some protection. Lithium, sodium and potassium are stored in oil. Reactivity increases as you go down the Group. These are all very reactive metals and have to be stored out of contact with air to prevent their oxidation. It also deals very briefly with the reactions of the elements with chlorine. This page mainly looks at the reactions of the Group 1 elements (lithium, sodium, potassium, rubidium and caesium) with oxygen - including the simple reactions of the various kinds of oxides formed. REACTIONS OF THE GROUP 1 ELEMENTS WITH OXYGEN AND CHLORINE Reactions of the Group 1 elements with oxygen and chlorine
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