The Heisenberg Uncertainty Principle, stated in a way that's useful for chemists, is that it's impossible to determine accurately both the momentum and position of an electron simultaneously.
More broadly, the Heisenberg Uncertainty Principle applies to pairs of complementary variables. Position and momentum form one such pair, while energy and time form another. The Heisenberg Uncertainty principle is correct for objects of any mass but is significant only for extremely small objects such as the particles which form atoms.
Consider the question: What is the momentum, p, and position, x, of a particle at time t?
To know the position, the particle must be localized at x, with uncertainty of Δx. Therefore the de Broglie wave associated with it must also be localized. A localized wave is formed from a superposition of many waves of different wavelength, but this implies many different values of momentum, because p=h/λ. (h is Planck's Constant and λ is wavelength.)
Therefore we cannot associate a definite momentum with such a localized particle and hence Δp is large.
If, on the other hand, momentum is definite, the de Broglie wave would have a definite wavelength and be spread out through space - i.e. non-localized and large Δx.
It can be shown that:
δpδx ≥ h/4π
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