Understanding Energy Transitions in Electromagnetic Radiation: UV vs Infrared

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Electromagnetic radiation in the ultraviolet (UV) region has shorter wavelengths, typically ranging from 10 to 400 nanometers, compared to infrared (IR) radiation, which ranges from 700 nanometers to 1 millimeter. The energy of a photon is calculated using the formula E = hc/λ, where E is energy, h is Planck's constant, c is the speed of light, and λ is the wavelength. Since UV radiation has shorter wavelengths, it corresponds to higher energy transitions than IR radiation. Understanding these differences is crucial in fields like chemistry and physics. This knowledge aids in grasping the fundamental principles of energy transitions in electromagnetic radiation.
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This was a problem on my Thanksgiving break chemistry homework. Some assistance would be greatly appreciated.

Why does electromagnetic radiation in the ulraviolet region represent a larger energy transition than does radiation in the infrared region?
 
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Hint: 1) What is the range of wavelengths for ultraviolet radiation? What is the range of wavelengths for infrared radiation?

2) How do you calculate the energy of a photon?
 
Thank you. I will give that some thought.
 
Thread 'Confusion regarding a chemical kinetics problem'

Thread 'Confusion regarding a chemical kinetics problem'

TL;DR Summary: cannot find out error in solution proposed. [![question with rate laws][1]][1] Now the rate law for the reaction (i.e reaction rate) can be written as: $$ R= k[N_2O_5] $$ my main question is, WHAT is this reaction equal to? what I mean here is, whether $$k[N_2O_5]= -d[N_2O_5]/dt$$ or is it $$k[N_2O_5]= -1/2 \frac{d}{dt} [N_2O_5] $$ ? The latter seems to be more apt, as the reaction rate must be -1/2 (disappearance rate of N2O5), which adheres to the stoichiometry of the...
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