Questions: 1. Why was it not possible to determine the charge of the electron by this method at the time when Thomson was doing his experiment? 2. A constant current of 0.800 A was used to deposit copper at the cathode of a cell. Calculate the grams of Cu deposited in minutes. 3. If the time and current are each uncertain by 1.0% and the mass of copper is uncertain by 0.10%, what is the uncertainty?

\(\boxed{\text{Thomson could not determine the charge of the electron directly due to the lack of technology and methods at the time.}}\)

To calculate the mass of copper deposited, we use Faraday's laws of electrolysis. The mass \( m \) of a substance deposited is given by: \[ m = \frac{Q \cdot M}{n \cdot F} \] where:

• \( Q \) is the total charge passed through the electrolyte,
• \( M \) is the molar mass of the substance (Cu = 63.546 g/mol),
• \( n \) is the number of electrons involved in the reaction (for Cu, \( n = 2 \)),
• \( F \) is Faraday's constant (\( F = 96485 \, \text{C/mol} \)).

First, calculate the total charge \( Q \): \[ Q = I \cdot t \] where \( I = 0.800 \, \text{A} \) and \( t \) is the time in seconds. Convert minutes to seconds: \[ t = 60 \, \text{s/min} \]

Thus, \[ Q = 0.800 \, \text{A} \cdot 60 \, \text{s} = 48.00 \, \text{C} \]

Now, calculate the mass \( m \): \[ m = \frac{48.00 \, \text{C} \cdot 63.546 \, \text{g/mol}}{2 \cdot 96485 \, \text{C/mol}} \] \[ m = \frac{3049.008 \, \text{g} \cdot \text{C/mol}}{192970 \, \text{C/mol}} \] \[ m \approx 0.0158 \, \text{g} \]

\(\boxed{0.0158 \, \text{g}}\)

Given:

• Uncertainty in time (\( t \)) = \( 1.0 \% \)
• Uncertainty in current (\( I \)) = \( 1.0 \% \)
• Uncertainty in mass of copper (\( m \)) = \( 0.10 \% \)

The total uncertainty in the mass of copper deposited can be calculated using the formula for combined uncertainties: \[ \text{Total Uncertainty} = \sqrt{(\text{Uncertainty in } I)^2 + (\text{Uncertainty in } t)^2 + (\text{Uncertainty in } m)^2} \]

Substitute the given values: \[ \text{Total Uncertainty} = \sqrt{(1.0\%)^2 + (1.0\%)^2 + (0.10\%)^2} \] \[ \text{Total Uncertainty} = \sqrt{1.00\% + 1.00\% + 0.01\%} \] \[ \text{Total Uncertainty} = \sqrt{2.01\%} \] \[ \text{Total Uncertainty} \approx 1.4177\% \]

\(\boxed{1.4177\%}\)