Question
It takes 163kJ/mol to break an nitrogen-nitrogen single bond calculate the maximum wavelength of light for which nitrogen-nitrogen single bond could be broken by absorbing a single photon
Be sure your answer has correct number of significant digits in nm
Be sure your answer has correct number of significant digits in nm
Answers
GPT 3.5
To calculate the maximum wavelength of light required to break a nitrogen-nitrogen single bond, we will use the equation:
E = hc/λ
Where:
E = energy required to break the bond = 163 kJ/mol = 163 x 10^3 J/mol
h = Planck's constant = 6.626 x 10^-34 J·s
c = speed of light = 2.998 x 10^8 m/s
λ = wavelength of light required
Rearranging the equation, we get:
λ = hc/E
Substituting the given values:
λ = (6.626 x 10^-34 J·s * 2.998 x 10^8 m/s) / (163 x 10^3 J/mol)
λ = (1.983 x 10^-25 J·m) / (163 x 10^3 J/mol)
λ ≈ 1.22 x 10^-28 m
To express the answer in nm, we convert meters to nanometers by multiplying by 10^9.
λ ≈ 1.22 x 10^-19 nm
The maximum wavelength of light required to break a nitrogen-nitrogen single bond is approximately 1.22 x 10^-19 nm.
E = hc/λ
Where:
E = energy required to break the bond = 163 kJ/mol = 163 x 10^3 J/mol
h = Planck's constant = 6.626 x 10^-34 J·s
c = speed of light = 2.998 x 10^8 m/s
λ = wavelength of light required
Rearranging the equation, we get:
λ = hc/E
Substituting the given values:
λ = (6.626 x 10^-34 J·s * 2.998 x 10^8 m/s) / (163 x 10^3 J/mol)
λ = (1.983 x 10^-25 J·m) / (163 x 10^3 J/mol)
λ ≈ 1.22 x 10^-28 m
To express the answer in nm, we convert meters to nanometers by multiplying by 10^9.
λ ≈ 1.22 x 10^-19 nm
The maximum wavelength of light required to break a nitrogen-nitrogen single bond is approximately 1.22 x 10^-19 nm.