It takes 348. kJ/mol to break a carbon-carbon single bond. Calculate the maximum wavelength of light for which a carbon-carbon single bond could be broken by absorbing a single photon. Round your answer to 3 significant digits. nm X
It takes 348. kJ/mol to break a carbon-carbon single bond. Calculate the maximum wavelength of light for which a carbon-carbon single bond could be broken by absorbing a single photon. Round your answer to 3 significant digits. nm X
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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![**Problem Statement:**
It takes 348 kJ/mol to break a carbon-carbon single bond. Calculate the maximum wavelength of light for which a carbon-carbon single bond could be broken by absorbing a single photon.
Round your answer to 3 significant digits.
**Input:**
- A text box is provided for the answer with units in nanometers (nm).
- There are buttons beneath the input box, including a reset button.
**Explanation:**
To solve this problem, you need to use the relationship between energy and wavelength of a photon, given by the equation:
\[ E = \frac{hc}{\lambda} \]
Where:
- \( E \) is the energy required to break the bond (in joules per photon).
- \( h \) is Planck's constant (\(6.626 \times 10^{-34} \, \text{Js}\)).
- \( c \) is the speed of light (\(3.00 \times 10^8 \, \text{ms}^{-1}\)).
- \( \lambda \) is the wavelength in meters.
Convert energy from kJ/mol to J/photon, and then solve for \( \lambda \):
1. Convert energy:
\[
\text{Energy (J/photon)} = \frac{348 \, \text{kJ/mol}}{6.022 \times 10^{23} \, \text{mol}^{-1}} \times 1000 \, \text{J/kJ}
\]
2. Rearrange the equation to find wavelength:
\[
\lambda = \frac{hc}{E}
\]
3. Convert wavelength from meters to nanometers by multiplying by \(10^9\).
Finally, the answer should be rounded to three significant digits and entered into the provided text box.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F9add9dee-1f9f-4ea0-8cd8-9e476bfd6ff4%2F38258323-1532-4160-bded-0a3a3a850509%2Ft4mdrzh_processed.png&w=3840&q=75)
Transcribed Image Text:**Problem Statement:**
It takes 348 kJ/mol to break a carbon-carbon single bond. Calculate the maximum wavelength of light for which a carbon-carbon single bond could be broken by absorbing a single photon.
Round your answer to 3 significant digits.
**Input:**
- A text box is provided for the answer with units in nanometers (nm).
- There are buttons beneath the input box, including a reset button.
**Explanation:**
To solve this problem, you need to use the relationship between energy and wavelength of a photon, given by the equation:
\[ E = \frac{hc}{\lambda} \]
Where:
- \( E \) is the energy required to break the bond (in joules per photon).
- \( h \) is Planck's constant (\(6.626 \times 10^{-34} \, \text{Js}\)).
- \( c \) is the speed of light (\(3.00 \times 10^8 \, \text{ms}^{-1}\)).
- \( \lambda \) is the wavelength in meters.
Convert energy from kJ/mol to J/photon, and then solve for \( \lambda \):
1. Convert energy:
\[
\text{Energy (J/photon)} = \frac{348 \, \text{kJ/mol}}{6.022 \times 10^{23} \, \text{mol}^{-1}} \times 1000 \, \text{J/kJ}
\]
2. Rearrange the equation to find wavelength:
\[
\lambda = \frac{hc}{E}
\]
3. Convert wavelength from meters to nanometers by multiplying by \(10^9\).
Finally, the answer should be rounded to three significant digits and entered into the provided text box.
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Transcribed Image Text:It takes 348. kJ/mol to break a carbon-carbon single bond. Calculate the maximum wavelength of light for which a carbon-carbon single bond could be broken
by absorbing a single photon.
Round your answer to 3 significant digits.
nm
x10
X
5
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