Final answer:
The Heisenberg Uncertainty Principle is a true fundamental concept in quantum mechanics that dictates it is impossible to precisely determine both an electron's position and momentum at the same time.
Step-by-step explanation:
The Heisenberg Uncertainty Principle
The statement that 'Heisenberg's uncertainty principle says that you cannot know the momentum or the position of an electron exactly' is true. Werner Heisenberg, a German physicist, introduced a fundamental concept in quantum mechanics that limits our ability to simultaneously determine both the precise position and momentum of particles at the atomic scale, such as electrons. He suggested that the more precisely we try to measure the position of a particle, the less precisely we can know its momentum, and vice versa. This is not a limitation of technology or measurement techniques, but rather an inherent property of the quantum world.
The principle can be quantitatively expressed as: Δx Δp ≥ ħ/2, where Δx represents the uncertainty in position and Δp is the uncertainty in momentum, while ħ denotes the reduced Planck constant. This inequality demonstrates that the product of the uncertainties is always at least half of the reduced Planck constant, indicating a fundamental limitation of precision.
However, it's important to note that the Heisenberg Uncertainty Principle predominantly applies to subatomic particles and becomes far less significant when dealing with macroscopic objects, due to their much larger mass and scale.