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eV
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J
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kJ/mol
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THz
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cm⁻¹
Enter values to see results
—
eV
—
J
—
kJ/mol
—
THz
—
cm⁻¹
The Wavelength to Energy Calculator determines the photon energy corresponding to a given electromagnetic wavelength using the Planck-Einstein relation. This is one of the most frequently used conversions in spectroscopy, photochemistry, quantum physics, and materials science. Whether you are analyzing an absorption spectrum, designing a laser system, evaluating semiconductor band gaps, or studying photochemical reactions, knowing the energy associated with a particular wavelength is essential. The calculator accepts wavelengths in nanometers, micrometers, millimeters, centimeters, meters, or angstroms and returns energy in electron volts (eV), joules (J), and kilojoules per mole (kJ/mol). Additional outputs include frequency and wavenumber for comprehensive spectroscopic reference across all measurement conventions.
The photon energy is calculated from the Planck-Einstein relation:
$$E = \frac{hc}{\lambda}$$
where h = 6.626 × 10⁻³⁴ J·s is Planck's constant, c = 2.998 × 10⁸ m/s is the speed of light, and λ is the wavelength in meters. To express energy in different units:
$$E (\text{eV}) = \frac{hc}{\lambda \cdot e} = \frac{1239.84}{\lambda (\text{nm})}$$
$$E (\text{kJ/mol}) = \frac{hcN_A}{\lambda} \times 10^{-3} = \frac{119627}{\lambda (\text{nm})}$$
where e = 1.602 × 10⁻¹⁹ C is the elementary charge and N_A = 6.022 × 10²³ mol⁻¹ is Avogadro's number. The energy per mole represents the total energy when one mole (6.022 × 10²³) of photons is absorbed, which is the relevant quantity for chemical reactions where amounts are measured in moles.
The energy output reveals the quantum of energy carried by each photon at the specified wavelength. In chemistry, energies in kJ/mol are directly comparable to bond dissociation energies and activation energies. For instance, if a photon's energy exceeds a bond's dissociation energy, it can break that bond upon absorption. In physics, eV is the standard unit for comparing to electronic energy levels, ionization potentials, and work functions. The inverse relationship means that short wavelengths (UV, X-ray) carry much more energy per photon than long wavelengths (IR, microwave), explaining why UV light causes chemical damage while IR primarily causes heating.
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The 254 nm UV-C wavelength used in germicidal lamps carries 4.88 eV per photon (471 kJ/mol), enough to break C-C bonds (~347 kJ/mol) and damage DNA, explaining its effectiveness in sterilization.
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The 1550 nm wavelength is the standard for fiber-optic telecommunications because silica fiber has minimum attenuation at this wavelength. The photon energy is 0.80 eV, well below silicon's band gap.
They are inversely proportional: E = hc/λ. Doubling the wavelength halves the energy, and vice versa. This means UV photons carry significantly more energy than visible light photons, and visible photons carry more energy than IR photons.
Different fields prefer different units. Physicists use eV because it relates directly to voltage and electronic energy levels. Chemists use kJ/mol because chemical reactions involve moles of molecules. Joules (J) are the SI base unit. The calculator provides all three for convenience.
Typical bond dissociation energies include: C-H ≈ 413 kJ/mol, C-C ≈ 347 kJ/mol, C=C ≈ 614 kJ/mol, O-H ≈ 463 kJ/mol, C=O ≈ 745 kJ/mol, and N≡N ≈ 945 kJ/mol. Comparing these with photon energies helps predict which wavelengths can cause photodissociation.
Not through a single-photon process under normal conditions. However, multiphoton absorption (using intense laser light) can deliver enough energy through multiple simultaneous absorptions. Also, photons can excite molecules to reactive excited states that then undergo bond breaking through different pathways.
The Sun's spectrum peaks at approximately 500 nm (green light), corresponding to about 2.48 eV or 239 kJ/mol per photon. The solar spectrum ranges from about 250 nm (UV) to 2500 nm (near-IR), covering energies from 5.0 eV down to 0.5 eV.
In the photoelectric effect, a photon must have energy equal to or greater than the work function (φ) of the metal surface to eject an electron: E_photon ≥ φ. For example, the work function of sodium is 2.28 eV, so only wavelengths shorter than 544 nm can cause photoemission from sodium.
Thermal infrared radiation (10 μm peak at room temperature, per Wien's law) carries about 0.124 eV or 12 kJ/mol per photon. This is much less than typical bond energies, which is why IR radiation causes heating (molecular vibration) rather than chemical bond breaking.
LED wavelength is determined by the semiconductor band gap: λ = hc/E_gap. By entering the band gap energy, you can predict the emission wavelength. For example, GaN-based LEDs with a 2.64 eV band gap emit at ~470 nm (blue), while AlGaInP with 1.96 eV emits at ~633 nm (red).
Radio photons have extremely low energies. An FM radio photon at 100 MHz (3 m wavelength) has an energy of only 4.14 × 10⁻⁷ eV, which is about 10 million times less than a visible photon. This is why radio waves are non-ionizing and have no chemical effects.
The Planck-Einstein relation (E = hc/λ) applies only to massless particles (photons). For massive particles, the de Broglie relation is λ = h/(mv), and energy includes kinetic energy: E = p²/(2m) = h²/(2mλ²). A different calculator would be needed for particle wavelengths.
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