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Thread: caesium in water

  1. #1 caesium in water 
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    Caesium in water by popular demand for The Periodic Table of Videos

    http://www.youtube.com/watch?v=zCARhVfeX5U


    Check out scientists doing their stuff
    http://www.test-tube.org.uk/
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  3. #2  
    Forum Bachelors Degree Demen Tolden's Avatar
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    From what I've heard of the reactions alkali metals have with water, I was expecting a much more dangerous reaction. Maybe my teacher was being overly cautious.


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  4. #3  
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    Not all alkali metals react dangerously with water. The reaction only becomes more vigorous as you go down the alkali-metal group of the periodic table. Lithium, at the top of the group, merely fizzes on the surface of the water. The next element, sodium, reacts with a small, delayed explosion, while potassium reacts with bigger and more immediate explosions. Itís when you get to rubidium that the experiment becomes dangeous.

    http://youtube.com/watch?v=GVoJZkmAAfA

    The theory is that the further down the group you go, the further the outermost electron is from the nucleus of the atom, hence the more easily it can escape the electrostatic attraction of the nucleus to participate in reactions. Hence the reactivity of the elements increases as you go down the group.

    Someone here recently suggested dropping francium (below caesium in the group) in water and seeing what happens. Iím not sure anyone has done the experiment yet. When they do, Iíll be interested to see what happens.
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  5. #4  
    Forum Sophomore oceanwave's Avatar
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    just a question I had after watching the very entertaining vid (i wish MY chem teachers were that interested in proving a theory) is:
    Yes, we all know that the reactions get more explosive down the group due to the ease of the electron parting from the atom of the metal. Thus a law of conservation of energy comes into mind: Why doesnt the caesium (or any other reactive metal) react just nicely AND save the energy seeing as the electron "flies" away so easily? Why the need for the explosion? The prof in the vid said that less energy required to remove the atom translates into more energy for an explosive reaction. But based on the law of conservation of energy, why doesnt the reaction just happen without the additional heat / flame being liberated? (or as I was typing this: why does any reaction at all need to liberate that much energy at all?)

    My current answer (yes, to my own question) is that when ionic bonds are formed (in the current case, caesium hydroxide) a certain amount of energy is released. Before that, energy is required to extract the electron from the atom and that energy is "supplied" by the willingness to form a bond (ionic) which liberates energy and thus stabilises the new compound as a whole. And since the caesium atom is pretty big (helps to allow lose of electron) and the hydroxide ion is pretty small (helps electronegativity) that the strength of the new ionic bond is very strong and the resultant energy seen is the energy released as a result of the bond minus the energy required to take the electron from the caesium atom in the first place.

    And erm, sounds weird but is the above answer to my own question correct? Anything more to add or correct? (My school - in general my WHOLE country - doesnt really place emphasis on the thought process, just forcing everyone to swallow what's in the notes and telling those who bother to ask that its "out-of-the-syllabus".
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