Convert between UV-Vis wavelength, photon energy, frequency, and wavenumber. Enter a value in any unit to see all equivalent values and identify the spectral region (UV-C, UV-B, UV-A, visible color, or near IR).
Enter a wavelength, photon energy, frequency, or wavenumber to convert between common UV-Vis units and see which spectral region it falls into.
E = hc/λ (photon energy)
ν = c/λ (frequency)
ṽ = 1/λ (wavenumber)
This is an educational UV-Vis helper based on simple formulas. It is not a safety calculator or a substitute for instrument calibration.
Last Updated: November 17, 2025. This content is regularly reviewed to ensure accuracy and alignment with current spectroscopy principles and NIST constants.
Ultraviolet and visible (UV-Vis) light spans wavelengths from approximately 100 nm (ultraviolet) to 700 nm (visible red), covering a crucial region of the electromagnetic spectrum where electronic transitions in molecules occur. This region is fundamental in chemistry because molecules absorb UV-Vis light when electrons transition between energy levels, making it essential for spectroscopy, photochemistry, and understanding light-matter interactions. Understanding UV-Vis light is crucial for students studying general chemistry, physical chemistry, analytical chemistry, biochemistry, and materials science, as it explains how molecules interact with light, how spectroscopy works, and how photon energy relates to chemical processes. UV-Vis concepts appear on virtually every chemistry exam and are foundational to understanding electronic transitions, absorption spectroscopy, and photochemical reactions.
Photon energy is inversely proportional to wavelength: E = hc/λ, where E is energy, h is Planck's constant (6.626 × 10⁻³⁴ J·s), c is the speed of light (2.998 × 10⁸ m/s), and λ is wavelength. This means shorter wavelengths have higher energy—violet light (~400 nm) has nearly twice the photon energy of red light (~700 nm). The relationship E = hc/λ connects wavelength to energy, showing why UV light (shorter wavelength) has more energy than visible light, and why high-energy UV can break chemical bonds while visible light typically cannot. Understanding this relationship helps you see why different wavelengths have different effects on matter, why UV-Vis spectroscopy works, and how photon energy determines what electronic transitions are possible.
Frequency and wavenumber are alternative ways to describe light. Frequency (ν) is inversely proportional to wavelength: ν = c/λ, measured in Hz (cycles per second). Higher frequency means more wave cycles per second, corresponding to higher energy. Wavenumber (ṽ) is the reciprocal of wavelength: ṽ = 1/λ, typically expressed in cm⁻¹. Wavenumber is proportional to energy and frequency, making it convenient for spectroscopy—IR spectroscopists often prefer wavenumber because it scales linearly with energy, unlike wavelength. Understanding these relationships helps you convert between different ways of describing light and see why different units are used in different contexts (wavelength in UV-Vis, wavenumber in IR).
Energy units vary depending on context. Electron-volts (eV) describe the energy of a single photon and are convenient for atomic/molecular scale discussions (visible light is roughly 1.7–3.1 eV). Kilojoules per mole (kJ/mol) describe the energy of Avogadro's number of photons (one mole) and are useful when comparing to bond energies or reaction enthalpies, which are typically given in kJ/mol (visible light is roughly 170–300 kJ/mol). The conversion is: E(kJ/mol) = E(eV) × 96.485 kJ·mol⁻¹·eV⁻¹. Understanding these units helps you see why eV is used for single-photon interactions while kJ/mol is used for bulk processes, and how to convert between them.
Spectral regions are conventionally divided based on wavelength and biological/chemical effects. UV-C (~100–280 nm) is highest energy, germicidal, and absorbed by atmospheric ozone. UV-B (~280–315 nm) causes sunburn and vitamin D synthesis. UV-A (~315–400 nm) is lowest energy UV, used in "black lights" and tanning. Visible light (~400–700 nm) ranges from violet to red and is human-visible. Near IR (~700–2500 nm) is below visible and used in thermal imaging. Understanding these regions helps you identify where different wavelengths fall, understand their biological and chemical effects, and see why different regions are used for different applications.
This calculator is designed for educational exploration and practice. It helps students master UV-Vis conversions by converting between wavelength, photon energy, frequency, and wavenumber, identifying spectral regions, and understanding the relationships between these quantities. The tool provides step-by-step calculations showing how to use E = hc/λ, ν = c/λ, and ṽ = 1/λ, and how to convert between different units. For students preparing for chemistry exams, spectroscopy courses, or photochemistry labs, mastering UV-Vis conversions is essential—these calculations appear on virtually every chemistry assessment and are fundamental to understanding spectroscopy and light-matter interactions. The calculator supports multiple units and spectral region identification, helping students understand all aspects of UV-Vis light.
Critical disclaimer: This calculator is for educational, homework, and conceptual learning purposes only. It helps you understand UV-Vis theory, practice wavelength-energy conversions, and explore spectroscopy concepts. It does NOT provide instructions for actual UV safety assessments, laser system design, or phototherapy planning, which require proper training, calibrated equipment, safety protocols, and adherence to validated procedures. Never use this tool to determine UV exposure safety, design laser systems, plan phototherapy, or make safety or regulatory determinations. Real-world UV safety involves considerations beyond this calculator's scope: source intensity, exposure duration, beam geometry, refractive index effects, and proper safety standards (ACGIH, OSHA, ICNIRP). Use this tool to learn the theory—consult trained professionals and proper safety standards for practical applications.
Photon energy is inversely proportional to wavelength: E = hc/λ, where E is energy, h is Planck's constant (6.626 × 10⁻³⁴ J·s), c is the speed of light (2.998 × 10⁸ m/s), and λ is wavelength. This means shorter wavelengths have higher energy. For example, violet light (~400 nm) has energy E ≈ 3.1 eV, while red light (~700 nm) has energy E ≈ 1.8 eV—violet has nearly twice the energy of red. Understanding this relationship helps you see why UV light (shorter wavelength) has more energy than visible light, why high-energy UV can break chemical bonds, and how wavelength determines what electronic transitions are possible in molecules.
Use the equation E = hc/λ. For wavelength in nanometers (nm), a convenient approximation is: E(eV) ≈ 1240 / λ(nm). For example, 500 nm light: E ≈ 1240 / 500 = 2.48 eV. For exact calculations: (1) Convert wavelength to meters (nm → m: multiply by 10⁻⁹), (2) Calculate E = hc/λ in joules, (3) Convert to eV if needed (1 eV = 1.602 × 10⁻¹⁹ J). Understanding this conversion helps you determine photon energy from wavelength, compare energies of different wavelengths, and understand why shorter wavelengths have higher energy.
Electron-volts (eV) describe the energy of a single photon and are convenient for atomic/molecular scale discussions (visible light is roughly 1.7–3.1 eV). Kilojoules per mole (kJ/mol) describe the energy of Avogadro's number of photons (one mole) and are useful when comparing to bond energies or reaction enthalpies, which are typically given in kJ/mol (visible light is roughly 170–300 kJ/mol). The conversion is: E(kJ/mol) = E(eV) × 96.485 kJ·mol⁻¹·eV⁻¹. Understanding this distinction helps you see why eV is used for single-photon interactions while kJ/mol is used for bulk processes, and how to convert between them.
Frequency (ν) is the number of wave cycles per second, measured in Hz: ν = c/λ, where c is the speed of light. Higher frequency means more cycles per second, corresponding to higher energy. Wavenumber (ṽ) is the reciprocal of wavelength: ṽ = 1/λ, typically expressed in cm⁻¹. For wavelength in nm: ṽ (cm⁻¹) ≈ 10⁷ / λ (nm). Wavenumber is proportional to energy and frequency, making it convenient for spectroscopy—IR spectroscopists often prefer wavenumber because it scales linearly with energy. Understanding these relationships helps you convert between different ways of describing light and see why different units are used in different contexts.
Spectral regions are conventionally divided: (1) UV-C (~100–280 nm): Highest energy UV, germicidal, absorbed by atmospheric ozone. (2) UV-B (~280–315 nm): Causes sunburn, vitamin D synthesis. (3) UV-A (~315–400 nm): Lowest energy UV, "black light," tanning, least harmful UV. (4) Visible (~400–700 nm): Human-visible light, violet to red. (5) Near IR (~700–2500 nm): Below visible, thermal imaging. These boundaries are approximate and vary slightly between references. Understanding these regions helps you identify where different wavelengths fall, understand their biological and chemical effects, and see why different regions are used for different applications.
The visible spectrum (~400–700 nm) ranges from violet at short wavelengths to red at long wavelengths. Approximate color bands: violet 400–450 nm, blue 450–495 nm, green 495–570 nm, yellow 570–590 nm, orange 590–620 nm, red 620–700 nm. Perceived color depends on many factors including intensity, surrounding colors, and individual perception. Understanding this relationship helps you identify the approximate color of light at a given wavelength and see how wavelength determines the color we perceive.
From E = hc/λ, energy is inversely proportional to wavelength. Shorter wavelength means more wave cycles fit in a given distance, corresponding to higher frequency (ν = c/λ) and higher energy (E = hν). This is a fundamental property of electromagnetic radiation: shorter wavelengths pack more energy into each photon. For example, violet light (~400 nm) has energy E ≈ 3.1 eV, while red light (~700 nm) has energy E ≈ 1.8 eV—violet has nearly twice the energy. Understanding this helps you see why UV light (shorter wavelength) has more energy than visible light, why high-energy UV can break chemical bonds, and how wavelength determines photon energy.
This interactive tool helps you convert between UV-Vis wavelength, photon energy, frequency, and wavenumber. Here's a comprehensive guide to using each feature:
Choose your preferred units for display:
Default Wavelength Unit
Select nm (nanometers), Å (angstroms), or m (meters). Most UV-Vis work uses nm.
Default Energy Unit
Select eV (electron-volts) or kJ/mol (kilojoules per mole). eV is common for single photons, kJ/mol for bulk processes.
For each entry you want to convert:
Entry Label
Enter a descriptive name (e.g., "500 nm visible light" or "Problem 1"). This helps you organize multiple entries.
What You Know
Select what quantity you're entering: wavelength, photon energy, frequency, or wavenumber.
Enter Value and Unit
Enter the value and select the appropriate unit. The calculator will convert to all other quantities automatically.
Click "Calculate" to see all conversions:
View All Quantities
The calculator shows: (a) Wavelength in nm, (b) Photon energy in eV and kJ/mol, (c) Frequency in Hz, (d) Wavenumber in cm⁻¹, (e) Spectral region (UV-C, UV-B, UV-A, visible, near IR), (f) Approximate color (for visible wavelengths).
Visualization
The calculator provides visualizations showing where each wavelength falls on the electromagnetic spectrum, helping you understand spectral regions visually.
Example: Convert 500 nm to energy
Input: Wavelength = 500 nm
Output: Energy = 2.48 eV (or 239 kJ/mol)
Frequency = 5.996 × 10¹⁴ Hz
Wavenumber = 20,000 cm⁻¹
Region: Visible (green–yellow)
Understanding the mathematics empowers you to convert between wavelength, energy, frequency, and wavenumber on exams, verify calculator results, and build intuition about UV-Vis light.
E = hc/λ
Where:
E = photon energy (J)
h = Planck's constant = 6.626 × 10⁻³⁴ J·s
c = speed of light = 2.998 × 10⁸ m/s
λ = wavelength (m)
Key insight: Energy is inversely proportional to wavelength. Shorter wavelengths have higher energy. For wavelength in nm, a convenient approximation is: E(eV) ≈ 1240 / λ(nm). Understanding this relationship helps you see why UV light (shorter wavelength) has more energy than visible light, and why high-energy UV can break chemical bonds.
Frequency is inversely proportional to wavelength:
ν = c/λ
Where:
ν = frequency (Hz)
c = speed of light = 2.998 × 10⁸ m/s
λ = wavelength (m)
Key insight: Higher frequency means more wave cycles per second, corresponding to higher energy. Frequency and energy are directly proportional: E = hν. Understanding this helps you see how frequency relates to wavelength and energy.
Wavenumber is the reciprocal of wavelength:
ṽ = 1/λ
Where:
ṽ = wavenumber (typically cm⁻¹)
λ = wavelength (same units as ṽ)
For λ in nm: ṽ (cm⁻¹) ≈ 10⁷ / λ (nm)
Key insight: Wavenumber is proportional to energy and frequency, making it convenient for spectroscopy. IR spectroscopists often prefer wavenumber because it scales linearly with energy, unlike wavelength. Understanding this helps you see why wavenumber is used in spectroscopy.
Energy can be expressed in different units:
Joules to eV:
E(eV) = E(J) / (1.602 × 10⁻¹⁹)
eV to kJ/mol:
E(kJ/mol) = E(eV) × 96.485
Joules to kJ/mol:
E(kJ/mol) = E(J) × (6.022 × 10²³) / 1000
Given: λ = 500 nm
Find: Photon energy in eV
Method 1: Using approximation
E(eV) ≈ 1240 / λ(nm)
E ≈ 1240 / 500
E ≈ 2.48 eV
Method 2: Exact calculation
λ = 500 nm = 500 × 10⁻⁹ m = 5.00 × 10⁻⁷ m
E = hc/λ = (6.626 × 10⁻³⁴ × 2.998 × 10⁸) / (5.00 × 10⁻⁷)
E = 3.97 × 10⁻¹⁹ J
E = 3.97 × 10⁻¹⁹ / (1.602 × 10⁻¹⁹) = 2.48 eV
Given: E = 3.0 eV
Find: Wavelength in nm
Step 1: Convert eV to J
E = 3.0 × 1.602 × 10⁻¹⁹ = 4.81 × 10⁻¹⁹ J
Step 2: Rearrange E = hc/λ
λ = hc/E
Step 3: Calculate
λ = (6.626 × 10⁻³⁴ × 2.998 × 10⁸) / (4.81 × 10⁻¹⁹)
λ = 4.13 × 10⁻⁷ m = 413 nm
Or using approximation:
λ(nm) ≈ 1240 / E(eV) = 1240 / 3.0 = 413 nm
Understanding UV-Vis wavelength and energy conversions is essential for students across chemistry coursework. Here are detailed student-focused scenarios (all conceptual, not actual procedures):
Scenario: Your general chemistry homework asks: "What is the photon energy of 500 nm light in eV?" Use the calculator: enter wavelength = 500 nm. The calculator shows: E = 2.48 eV (or 239 kJ/mol). You learn: how to use E = hc/λ and convert between units. The calculator helps you check your work and understand the relationship between wavelength and energy.
Scenario: An exam asks: "What spectral region does 300 nm light fall into?" Use the calculator: enter wavelength = 300 nm. The calculator shows: UV-B region (280–315 nm). You learn: how to identify spectral regions and understand their characteristics. The calculator makes this relationship concrete—you see exactly where different wavelengths fall.
Scenario: Your analytical chemistry lab report asks: "A compound absorbs at 260 nm. What is the photon energy?" Use the calculator: enter wavelength = 260 nm. The calculator shows: E = 4.77 eV (or 460 kJ/mol), UV-C region. Understanding this helps explain why this wavelength is used for nucleic acid quantification (DNA/RNA absorb strongly at ~260 nm). The calculator helps you verify your calculations and understand how wavelength relates to electronic transitions.
Scenario: Problem: "Compare the photon energies of 400 nm (violet) and 700 nm (red) light." Use the calculator: analyze both wavelengths. 400 nm: E = 3.10 eV. 700 nm: E = 1.77 eV. You learn: violet has nearly twice the energy of red, demonstrating the inverse relationship between wavelength and energy. The calculator helps you compare wavelengths and understand energy differences.
Scenario: Your physical chemistry homework asks: "Why do molecules absorb UV-Vis light?" Use the calculator: analyze typical UV-Vis wavelengths (200–700 nm). The calculator shows: energies range from ~1.8 eV (red) to ~6.2 eV (UV-C). Understanding this helps explain why these energies match electronic transition energies in molecules (typically 1–10 eV), and why UV-Vis spectroscopy works. The calculator makes this relationship concrete—you see exactly how photon energies relate to molecular energy levels.
Scenario: Problem: "What is the frequency of 600 nm light?" Use the calculator: enter wavelength = 600 nm. The calculator calculates: ν = c/λ = 4.997 × 10¹⁴ Hz. This demonstrates how to convert between wavelength and frequency, which is useful for understanding wave properties of light.
Scenario: Your instructor recommends practicing different types of UV-Vis conversion problems. Use the calculator to work through: (1) Wavelength to energy, (2) Energy to wavelength, (3) Wavelength to frequency, (4) Wavelength to wavenumber, (5) Identifying spectral regions. The calculator helps you practice all conversion types, identify common mistakes, and build confidence. Understanding how to convert between different quantities prepares you for exams where you might encounter various scenarios.
UV-Vis conversion problems involve wavelength, energy, frequency, wavenumber, and unit conversions that are error-prone. Here are the most frequent mistakes and how to avoid them:
Mistake: Using wavelength in nm directly in E = hc/λ without converting to meters, or mixing units (e.g., using nm with meters in the same calculation).
Why it's wrong: The equation E = hc/λ requires consistent units. If λ is in nm, you must convert to meters (multiply by 10⁻⁹) before using the equation, or use the approximation E(eV) ≈ 1240 / λ(nm). Using wrong units gives wrong energy values. For example, using 500 nm directly (without conversion) in E = hc/λ gives completely wrong results.
Solution: Always check units. For wavelength in nm, either: (1) Convert to meters (multiply by 10⁻⁹), or (2) Use the approximation E(eV) ≈ 1240 / λ(nm). The calculator handles unit conversions automatically—observe it to reinforce correct unit handling.
Mistake: Thinking that longer wavelengths have higher energy, or confusing the relationship between wavelength and energy.
Why it's wrong: From E = hc/λ, energy is inversely proportional to wavelength. Shorter wavelengths have higher energy, not longer wavelengths. For example, violet light (400 nm) has higher energy than red light (700 nm), not lower. Confusing this relationship leads to wrong energy comparisons and misunderstandings about UV vs visible light.
Solution: Always remember: shorter wavelength = higher energy. A helpful mnemonic: "Short and strong" (short wavelength, strong/high energy). The calculator shows this relationship—use it to reinforce the inverse relationship.
Mistake: Using eV when kJ/mol is appropriate (or vice versa), or not understanding when to use each unit.
Why it's wrong: eV describes single-photon energy, while kJ/mol describes energy per mole of photons. Using the wrong unit for the context can lead to wrong comparisons. For example, comparing a single-photon energy in eV to a bond energy in kJ/mol requires conversion. Using eV directly gives wrong comparisons.
Solution: Always identify the context: single-photon interactions use eV, bulk processes use kJ/mol. When comparing to bond energies or reaction enthalpies (typically in kJ/mol), convert photon energy to kJ/mol. The calculator shows both units—use them to reinforce when each is appropriate.
Mistake: Using incorrect values for Planck's constant (h), speed of light (c), or conversion factors (eV to J).
Why it's wrong: Physical constants must be accurate for correct calculations. Using wrong values gives wrong results. For example, using h = 6.6 × 10⁻³⁴ instead of 6.626 × 10⁻³⁴ gives slightly wrong energies. Using very wrong values (e.g., c = 3 × 10¹⁰) gives completely wrong results.
Solution: Always use correct constants: h = 6.626 × 10⁻³⁴ J·s, c = 2.998 × 10⁸ m/s, 1 eV = 1.602 × 10⁻¹⁹ J. The calculator uses standard CODATA values—use them to verify your constants.
Mistake: Not verifying that calculated values are reasonable (e.g., visible light energy 1.7–3.1 eV, UV-C energy > 4 eV).
Why it's wrong: If you calculate visible light energy as 10 eV or UV-C energy as 0.5 eV, something is wrong. Visible light is roughly 1.7–3.1 eV, UV-C is roughly 4.4–12.4 eV. Not checking reasonableness means you might accept wrong answers. For example, if you get E = 0.1 eV for 400 nm light, you made an error (should be ~3.1 eV).
Solution: Always check: visible light ≈ 1.7–3.1 eV, UV-A ≈ 3.1–3.9 eV, UV-B ≈ 3.9–4.4 eV, UV-C ≈ 4.4–12.4 eV. The calculator shows spectral regions—use them to verify your answers make sense.
Mistake: Using wrong units for wavenumber (e.g., m⁻¹ instead of cm⁻¹) or calculating wavenumber incorrectly.
Why it's wrong: Wavenumber is typically expressed in cm⁻¹. For wavelength in nm: ṽ (cm⁻¹) ≈ 10⁷ / λ (nm). Using wrong units or wrong calculation gives wrong wavenumber values. For example, if you calculate ṽ = 1 / 500 nm = 0.002 nm⁻¹, you're using wrong units—should be ṽ ≈ 10⁷ / 500 = 20,000 cm⁻¹.
Solution: Always use cm⁻¹ for wavenumber. For wavelength in nm: ṽ (cm⁻¹) ≈ 10⁷ / λ (nm). The calculator shows wavenumber in cm⁻¹—use it to reinforce correct units and calculations.
Mistake: Using calculator results for light in a medium (glass, water) without accounting for refractive index.
Why it's wrong: This calculator assumes light travels in vacuum (n = 1). In a medium with refractive index n, wavelength shortens to λ/n while frequency stays constant. For example, in glass (n ≈ 1.5), 500 nm light has wavelength ≈ 333 nm in the medium, but frequency (and energy per photon) stays the same. Using vacuum calculations for medium calculations gives wrong wavelengths.
Solution: Always remember: this calculator assumes vacuum. For light in a medium, wavelength changes but frequency and photon energy stay constant. The calculator notes this limitation—use it to reinforce when vacuum approximation is valid.
Once you've mastered basics, these advanced strategies deepen understanding and prepare you for complex spectroscopy and photochemistry problems:
Conceptual insight: From E = hc/λ, energy is inversely proportional to wavelength because shorter wavelengths pack more wave cycles into a given distance, corresponding to higher frequency (ν = c/λ) and higher energy (E = hν). This is a fundamental property of electromagnetic radiation: shorter wavelengths carry more energy per photon. Understanding this helps you see why UV light (shorter wavelength) has more energy than visible light, why high-energy UV can break chemical bonds, and how wavelength determines what electronic transitions are possible. This provides deep insight beyond memorization: wavelength and energy are fundamentally linked through the wave nature of light.
Quantitative insight: UV-Vis photon energies (roughly 1.7–12.4 eV) match typical electronic transition energies in molecules (typically 1–10 eV). This is why molecules absorb UV-Vis light—the photon energy matches the energy difference between electronic states. For example, π→π* transitions in conjugated systems typically require 3–6 eV, corresponding to UV-Vis wavelengths. Understanding this helps you see why UV-Vis spectroscopy works and how photon energy determines what transitions are possible.
Practical framework: For wavelength in nm, the approximation E(eV) ≈ 1240 / λ(nm) is very convenient and accurate enough for most purposes. This comes from: E = hc/λ with hc ≈ 1240 eV·nm. For quick estimates: 400 nm ≈ 3.1 eV, 500 nm ≈ 2.5 eV, 600 nm ≈ 2.1 eV, 700 nm ≈ 1.8 eV. These mental shortcuts help you quickly estimate energies on multiple-choice exams and check calculator results. Understanding approximate relationships builds intuition about UV-Vis energies.
Unifying concept: UV-Vis photon energies (roughly 170–1200 kJ/mol) can be compared to chemical bond energies (typically 200–500 kJ/mol for single bonds). This helps explain why high-energy UV can break chemical bonds (UV-C at ~4.4–12.4 eV ≈ 420–1200 kJ/mol can break many bonds), while visible light typically cannot (visible at ~1.7–3.1 eV ≈ 170–300 kJ/mol is usually below bond energies). Understanding this connection helps you see why UV light is more chemically active than visible light and how photon energy relates to chemical reactivity.
Exam technique: For quick identification: < 280 nm = UV-C, 280–315 nm = UV-B, 315–400 nm = UV-A, 400–700 nm = visible, > 700 nm = near IR. Approximate energies: visible ≈ 1.7–3.1 eV, UV-A ≈ 3.1–3.9 eV, UV-B ≈ 3.9–4.4 eV, UV-C ≈ 4.4–12.4 eV. These mental shortcuts help you quickly identify spectral regions on multiple-choice exams and check calculator results. Understanding approximate relationships builds intuition about UV-Vis regions.
Advanced consideration: This calculator assumes vacuum (n = 1) and single wavelengths. Real systems show: (a) Refractive index effects (wavelength changes in media, frequency and photon energy stay constant), (b) Spectral width (real light sources have finite bandwidth, not single wavelengths), (c) Instrumental response (detectors and optics have wavelength-dependent efficiency), (d) Absorption coefficients (real absorption depends on concentration and path length, not just wavelength). Understanding these limitations shows why empirical measurements may differ from calculated values, and why advanced spectroscopic techniques are needed for accurate work in research and industry, especially for precise measurements or non-vacuum conditions.
Advanced consideration: Light exhibits both wave and particle properties. The wave nature is described by wavelength, frequency, and wavenumber. The particle nature is described by photon energy (E = hν). The relationships E = hc/λ and ν = c/λ connect these descriptions, showing how wave properties (wavelength, frequency) relate to particle properties (photon energy). Understanding this duality helps you see why different quantities (wavelength, energy, frequency, wavenumber) are all valid ways to describe the same light, and why conversions between them are possible.
• Vacuum Wavelength: Conversions use vacuum speed of light (c = 2.998 × 10⁸ m/s). In media with refractive index n > 1, wavelength changes (λ_medium = λ_vacuum/n) while frequency and photon energy remain constant.
• Monochromatic Light: Equations apply to single wavelengths. Real light sources have finite bandwidth—spectral width affects apparent wavelength and energy distribution. Broad-band sources require integration over the spectrum.
• No Absorption Information: Wavelength-energy conversion says nothing about absorption probability. Beer-Lambert law (A = εcl) requires molar absorptivity (ε) from experimental data to relate wavelength to measured absorbance.
• Energy per Photon vs. Total Energy: E = hν gives energy of one photon. Total radiant energy depends on photon flux (number of photons). High-energy UV photons at low intensity may deliver less total energy than lower-energy visible light at high intensity.
Important Note: This calculator is strictly for educational and informational purposes only. It demonstrates wavelength-energy relationships for learning. For spectroscopic analysis, photochemistry research, or laser applications, use calibrated instrumentation and appropriate spectral databases.
The UV-Vis spectroscopy principles and wavelength-energy relationships referenced in this content are based on authoritative chemistry sources:
Calculations assume vacuum conditions (n = 1). Wavelength values in media require refractive index corrections.
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