Question

Examine the chart:.
Chart:
Metal | Specific Heat (J/g°C)
Calcium (Ca) | 0.647
Iron (Fe) | 0.449
Silver (Ag) | 0.235
Gold (Au) | 0.129
If equal heat amounts are added to 25.0 g samples of each metal shown in the chart, which metal will undergo the most significant temperature change?
A. Au
B. Fe
C. Ag
D. Ca

Answer 1

To determine which metal will undergo the most significant temperature change when equal heat amounts are added to 25.0 g samples of each metal, we can use the formula:

\[
Q = mc\Delta T
\]

where:
- \( Q \) is the heat added,
- \( m \) is the mass (25.0 g for each sample),
- \( c \) is the specific heat capacity (in J/g°C), and
- \( \Delta T \) is the change in temperature.

Rearranging this formula gives:

\[
\Delta T = \frac{Q}{mc}
\]

Since \( Q \) and \( m \) are equal for all samples, the temperature change \( \Delta T \) will be inversely proportional to the specific heat capacity \( c \):

\[
\Delta T \propto \frac{1}{c}
\]

Thus, the metal with the lowest specific heat will experience the largest temperature change.

From the chart provided:
- Calcium (Ca) has a specific heat of 0.647 J/g°C
- Iron (Fe) has a specific heat of 0.449 J/g°C
- Silver (Ag) has a specific heat of 0.235 J/g°C
- Gold (Au) has a specific heat of 0.129 J/g°C

Comparing the specific heats:
- Au (0.129) < Ag (0.235) < Fe (0.449) < Ca (0.647)

Gold has the lowest specific heat capacity, meaning it will undergo the most significant temperature change when equal heat amounts are added.

Therefore, the answer is:

**A. Au**