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The graph shows a calibration curve with the equation \( y = 2755x \), where \( y \) is the absorbance, and \( x \) is the concentration (in molarity, \( M \)).

You are asked to calculate the molar concentration for a solution with a measured absorbance of 0.212.

We can use the given linear equation to find the concentration:

\[
y = 2755x
\]

Substitute \( y = 0.212 \) for the absorbance:

\[
0.212 = 2755x
\]

Now solve for \( x \):

\[
x = \frac{0.212}{2755}
\]

Let me calculate that for you.The concentration \( x \) is approximately \( 7.70 \times 10^{-5} \, M \).

Thus, the molar concentration is:

\[
\text{Concentration} = 7.70 \times 10^{-5} \, M
\]

Would you like further details or clarifications on this?

Here are some related questions:
1. What is the significance of the calibration curve in experiments?
2. How can we improve the accuracy of absorbance measurements?
3. How does Beer's Law relate to the equation of the calibration curve?
4. What happens if the absorbance exceeds the linear range of the calibration curve?
5. How is molar absorptivity determined in spectrophotometry?

**Tip:** Always check that the absorbance values fall within the linear region of the calibration curve to ensure accurate concentration calculations.

Using the calibration curve, calculate the molar concentration for a solution with a measured absorbance of 0.212.

Learn how to calculate the molar concentration for a solution with a measured absorbance of 0.212 using a calibration curve. This example utilizes the linear equation y = 2755x and applies the Beer-Lambert Law for spectrophotometry experiments. Suitable for high school students in Grades 9-12.

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