Heisenberg's Uncertainty Principle
Definition and meaning of Heisenberg's Uncertainty Principle in chemistry.
Heisenberg's uncertainty principle states that the exact position and exact momentum of a particle, such as an electron, cannot both be known simultaneously with arbitrary precision.
In more detail
The principle reflects a fundamental property of quantum systems arising from wave-particle duality, not a limitation of measuring instruments. The uncertainties in position (Δx) and momentum (Δp) are related by Δx·Δp ≥ ħ/2, where ħ is the reduced Planck constant: the more precisely one quantity is known, the less precisely the other can be. This principle justified replacing the Bohr model's fixed electron orbits with the quantum mechanical model, in which electrons occupy orbitals describing regions of probability rather than definite paths.
Key facts
| Formula | Δx·Δp ≥ ħ/2 |
|---|---|
| Proposed by | Werner Heisenberg (1927) |
| Reduced Planck constant (ħ) | 1.055 × 10⁻³⁴ J·s |
| Field | Physical Chemistry |
An electron in a hydrogen atom cannot be assigned both a precise location and a precise velocity at once; instead, its behavior is described by an atomic orbital, a probability distribution showing where the electron is likely to be found.
Frequently asked questions
Does the uncertainty principle apply to everyday objects?
In principle yes, but for macroscopic objects the effect is far too small to detect because ħ is extremely tiny; it becomes significant only at atomic and subatomic scales.
Why did the uncertainty principle end the Bohr model?
The Bohr model assumed electrons travel in fixed circular orbits with precisely known position and momentum, which contradicts the uncertainty principle; the quantum mechanical model of orbitals replaced it.