What is the wavelength of light needed to break double bonded carbons if double bonded carbons have an enthalpy of 602 kJ/mol?
602 kJ/mol x (1 mol/6.02E23) = ?kJ/particle
Eparticle = hc/wavelength
Plug in and solve for wavelength. Post your work if you get stuck.
To determine the wavelength of light needed to break double bonded carbons, we need to use the relationship between the energy of light and its wavelength. This is described by the equation E = hc/λ, where E represents energy, h is Planck's constant (6.626 x 10^-34 J·s), c is the speed of light in a vacuum (3.00 x 10^8 m/s), and λ is the wavelength of light.
First, we need to convert the given enthalpy value of the double bonded carbons to energy per mole. We can use the equation E = ΔH * N, where E represents energy, ΔH is the enthalpy change (602 kJ/mol in this case), and N is Avogadro's number (6.022 x 10^23 mol^-1).
E = (602 kJ/mol) * (6.022 x 10^23 mol^-1)
E = 3.627224 x 10^26 J
Next, we can rearrange the equation E = hc/λ to solve for λ:
λ = hc/E
Substituting the values, we have:
λ = (6.626 x 10^-34 J·s * 3.00 x 10^8 m/s) / (3.627224 x 10^26 J)
λ ≈ 5.46 x 10^-7 m or 546 nm
Therefore, the wavelength of light needed to break double bonded carbons with an enthalpy of 602 kJ/mol is approximately 546 nanometers.