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ChemWiki: The Dynamic Chemistry Hypertext > Under Construction > Demonstrations > Additional Demos > LeChatelier's Principle-- The NO2/N2O4 Equilibrium

LeChatelier's Principle-- The NO2/N2O4 Equilibrium


Chemical Concept Demonstrated

  • LeChatelier's Principle


Two glass tubes are filled with gas, one with NO2 and one with NO and NO2.

Both of these tubes are immersed in liquid nitrogen.

As an alternative, fill two clear sealed syringes with NOgas, place both on an overhead projector, and press one of the plungers in.



The glass tube with NO and NOin it produces a blue liquid.   The flask with NOin it produces a white solid.  If the syringes are used, the compressed gas first becomes darker, then becomes lighter than the control syringe as the plunger is depressed.


NO and NOin the same tube are in equilibrium with the compound N2O3, which is a blue liquid.  At room temperature, the equilibrium is shifted far to the end of the gases.  However, when the temperature is lowered, LeChatelier's Principle comes into play.  Because the formation of N2Ois exothermic, lowering the temperature of the system makes the reaction more favorable.

A similar application of the principle explains why pure NOforms N2O(a white powder) at low temperatures.

In the syringe, the gas becomes darker at first simply because the gas is compressed into a smaller space.  The gas lightens because the resulting system is not in equilibrium.  Because different gases take up the same space with the same number of particles, it is advantageous for two NOmolecules to form a single molecule under high-pressure conditions.  N2Ois produced, fewer particles are in the syringe, and, as a result, the intensity of the color caused by the NOgoes down.



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Last modified
10:29, 2 Oct 2013



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This material is based upon work supported by the National Science Foundation under Grant Numbers 1246120, 1525057, and 1413739.

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