To determine the concentration of an unknown solution with an absorbance of 0.60 using extrapolation, you typically need a standard curve based on Beer-Lambert Law, which states:

$A = \varepsilon \cdot c \cdot l$

Where:

• $A$ is the absorbance,
• $\varepsilon$ is the molar absorptivity (L/mol·cm),
• $c$ is the concentration (mol/L),
• $l$ is the path length of the cell (usually 1 cm).

Steps:

1. Create a Standard Curve: Measure the absorbance for several known concentrations of the solution. The plot should be absorbance (A) vs. concentration (M).

2. Linear Relationship: If Beer-Lambert Law holds, the standard curve will be a straight line, typically described by $y = mx + b$, where:

   • $y$ is the absorbance,
   • $m$ is the slope (related to $\varepsilon$),
   • $x$ is the concentration,
   • $b$ is the intercept (which should ideally be close to 0).

3. Extrapolation: With an absorbance of 0.60, extrapolate on the standard curve to find the corresponding concentration.

Example (Hypothetical Data):

• Assume a linear equation for the standard curve is: $A = 2.0 \cdot c$ Where $A$ is the absorbance and $c$ is the concentration in M.

4. Solve for the Concentration: Given $A = 0.60$, solve for $c$: $0.60 = 2.0 \cdot c$ $c = \frac{0.60}{2.0} = 0.30 \, \text{M}$

Thus, the concentration of the unknown solution is 0.30 M.

Would you like more details or have any questions?

---

Related Questions:

1. How does temperature affect the absorbance reading?
2. What is the role of molar absorptivity in Beer-Lambert Law?
3. How can I create a standard curve for absorbance measurements?
4. What causes deviations from Beer-Lambert Law in real-world experiments?
5. How can you determine molar absorptivity ($\varepsilon$) for a substance?

Tip: Always zero the spectrophotometer with a blank solution before taking absorbance readings to ensure accuracy.