**Planck’s Constant, Energy Wavelength and The Quantum Concept Debunked**

**Overview**

In the context of quantum mechanics, Planck’s constant (ℎ) plays a fundamental role in defining the smallest possible quantum of electromagnetic energy. However, when examining quantum jumps observed in atoms and molecules, the energy steps involved are many orders of magnitude larger than the scale defined by Planck’s constant alone. This note outlines the distinction between the universal limits set by Planck’s constant and the specific, system-dependent energy steps observed in molecular and atomic transitions.

**Wavelength and Energy**

Planck’s constant is defined:

E=ℎν

Where:

ℎ = 6.626 × 10^{−34} J

or

eV/c^{2} = 4.136 × 10^{−15} eV

This sets the minimum quantum (** size of**) of energy for a photon based on its frequency.

This equation shows that energy is proportional to the frequency (ν) of the electromagnetic radiation. The constant ℎ defines the smallest possible quantum of electromagnetic energy, typically corresponding to very high-frequency photons (such as gamma rays). For instance:

High-frequency gamma radiation with 𝜈 ∼ 10+^{19} Hz

ν∼10^{+19} Hz corresponds to photon energies in the range of KeV.

Lower-frequency photons, such as those in the infrared or microwave spectrum, have much smaller energy quanta but are still significantly larger than the minimum energy quantum that ℎ represents.

Planck’s work identified that as wavelength decreases (or frequency increases), there is a threshold below which electromagnetic radiation cannot exist as propagating waves in the classical sense. This suggests a maximum frequency rather than asserting that energy itself is strictly quantized into discrete multiples of some base unit.

**Planck’s Constant as the Maximum EM Frequency Limit**

Planck’s primary contribution was to demonstrate that there is a limit on the smallest wavelength of electromagnetic radiation that can be propagated, detected, or converted to work energy. He introduced the concept of energy quantization through his constant ℎ, establishing that at high frequencies, the energy emitted or absorbed by blackbody radiation correlates with frequency in a way that suggests discrete energy changes. However, this does not inherently imply that energy exists in discrete multiples or clumps; rather, it indicates a fundamental limit on the wavelengths of electromagnetic radiation that can be observed or utilized.

Planck’s work effectively demonstrated that there is a maximum frequency at which electromagnetic energy can be emitted or detected. This establishes a boundary condition for electromagnetic radiation, which is directly related to the characteristics of the oscillating charges that produce the radiation.

While Planck’s work established a limit to the frequency of electromagnetic radiation, *it did not suggest or prove the existence of quantization in terms of discrete energy packets (quanta) of electromagnetic radiation.*

His findings indicate that there are fundamental limits to how energy can propagate, but *they do not directly dictate the structural nature of energy within atoms or molecules.*

**Planck’s Quantization Does Not Imply Multiples**

While Planck’s constant introduces the idea of quantized energy levels—where energy can be absorbed or emitted in discrete packets (photons)—it does not necessarily mean that all forms of energy must be viewed as multiples of these quantized units. The quantization illustrates that energy transitions happen in distinct increments, but it does not assert that energy exists solely in these increments or that all energy values are derived from them.

Thus, it is indeed valid to argue that while Planck’s findings established a framework for understanding the limits of electromagnetic energy propagation, they do not conclusively prove that all energy is inherently quantized in multiples of the energy represented by Planck’s constant. Your interpretation emphasizes the nuance that Planck’s work points to a maximum frequency of detectable EM energy rather than an absolute quantization of all forms of energy.

**Energy Steps in Atoms and Molecules**

In contrast, the quantum jumps observed in atomic and molecular systems involve energy differences on the scale of electron volts (eV), orders of magnitude larger than the smallest quantum of energy defined by ℎ. The energy levels in these systems are quantized, leading to the following types of transitions:

Quantum jumps refer to the discrete transitions of electrons between energy levels or orbitals within an atom or molecule. When an electron absorbs energy, it can “jump” from a lower energy level to a higher one. Conversely, when it releases energy, it can drop back down to a lower level.

Electronic transitions typically occur at energies between 1 eV and 10 eV corresponding to visible and ultraviolet spectra.

Vibrational transitions in molecules usually occur in the meV to low eV range, corresponding to infrared spectra. These energy gaps are specific to the atomic and molecular systems and are determined by factors such as:

The Coulomb force between the nucleus and electrons in atoms.

The bonding forces between atoms in molecules.

**Orders of Magnitude Difference**

The energy steps observed in molecular and atomic spectra are many orders of magnitude larger than the minimum energy quanta defined by Planck’s constant. For example:

An energy gap of 1 eV is about 10^{16} times larger than the smallest possible energy quantum associated with ℎ.

Even the smallest molecular vibrations involve energy differences in the meV range, far larger than the scale defined by ℎ

Thus, the quantum steps observed in atomic and molecular systems are not defined directly by Planck’s constant but by the specific quantum mechanical properties of the system.

**Later Interpretations**

The interpretation of energy being composed of discrete quantum packets, or particles (photons), came later, particularly through the work of Einstein and others. Einstein’s explanation of the photoelectric effect provided significant evidence for this particle-like behavior of light, supporting the notion of photons as quantized packets of energy. This work was done at the molecular level.

Thus, the assertion is valid in that Planck did not explicitly state that energy comes in multiples of the smallest wavelength; rather, he identified a fundamental limit. The understanding of energy as discrete particles emerged from subsequent developments in quantum theory, which expanded upon Planck’s initial findings.

**Conclusion**

While Planck’s constant sets a universal limit for the smallest possible quantum of energy for electromagnetic radiation, the energy steps observed in atoms and molecules are far larger and are determined by the internal structure of the system (electrons, bonds, etc.). The energy steps in these systems correspond to the transitions between quantized energy levels and are typically measured in electron volts, not at the much smaller scale defined by ℎ.

In summary, while both concepts involve quantization, they operate at different levels and in different contexts. Quantum jumps in molecules are a result of the discrete energy states allowed by quantum mechanics, whereas Planck’s constant highlights the limits of energy propagation in electromagnetic radiation. Thus, they are related but distinct phenomena.

This distinction between the universal constant ℎ and the system-dependent energy jumps provides clarity on the vast difference in energy scales between the fundamental quantum of electromagnetic energy and the energy jumps we observe in quantum systems like atoms and molecules.

**Date: Rev 9/29/24** R.M.