Question

Can someone explain the Heisenberg uncertainty principle in laymans terms

Answer 1

Really really simple (too simple if you ask me) In quantum physics it says that 2 physical properties of an object cannot be measured precisely (for example actual position and future position because of the momentum).... It also claims that the more precisely you measure one, the less precise the other shall be

Answer 2

So this is how I learned it, because I want to know the WHY one measurement gets less certain as you reach certainty on the other property. To measure the position of an electron, you need to make it visible; this requires bouncing a photon of light off the electron, and viewing the scattered light through a microscope. But photons carry momentum, so when it impacts the electron, it changes the electrons momentum. So while you learn the position, you no longer know its momentum. If you increase the wavelength of light, thereby decreasing the momentum that the photon will transfer to the photon, you decrease the resolution of your microscope and thereby get a poor measurement of its position. So whichever measurement you're going for, you lose information about the other property by your disturbance.

Answer 3

you increase the speed of an electron, its mass increases. crazy, right?!

Answer 4

rbolio is close, but not precisely spot on. The equation is as follows: \[\Delta x \Delta p \ge h \div2\] Where, roughly, dx is the range of possible positions of the, say electron, in question, and dp is the range of possible momenta (which, in a more basic sense, is a way of saying the trajectory or the velocity) of the electron in question. By more precisely measuring the exact position (by making dx smaller) you necessarily alter your ability to determine the future trajectory of the electron. To maintain the truth of the inequality, by reducing the value of dx, the value of dp must increase. Correspondingly, if you measure the exact momentum, dp, of the electron in question, you necessarily compromise your ability to determine the present or future position of the electron, and must rely on something called a wave function to find the probabilities that the electron will arrive at a certain place.

The reasons that account for the uncertainty principle's accuracy are subtle. The most frequently cited "reason" I believe is that dealing with atomic and subatomic subjects, measuring a specific value necessarily compromises your ability to measure another. It's like this: we "see" a basketball's position because photons bounce off of them and towards our eyes. But the thing is, when dealing with very, very, very small entities, it isn't like bouncing photons off of the objects, it would be like bouncing baseballs off of the basketball to determine the location of the basketball. The thing is, by bouncing a baseball off of the basketball, you would alter the trajectory. The only way to "see" where the basketball was after that would be to bounce more baseballs off of it, but this would alter the basketball's trajectory some more.

Forgive my sordid analogy. Hope this helped clear stuff up.

Answer 5

I completely forgot to mention h is a constant (specifically, it's Planck's constant with value \[6.63\times10^{-34} J \times s\].