Calculate ΔE and determine whether the process is endothermic or exothermic for the following cases: (a) q = 0.763 kJ and w = -840 J.

Verified step by step guidance

1. Convert the work (w) from joules to kilojoules to match the units of heat (q). Since 1 kJ = 1000 J, divide the work value by 1000: \( w = \frac{-840 \text{ J}}{1000} \text{ kJ} \).
2. Use the first law of thermodynamics, which states that the change in internal energy (\( \Delta E \)) is the sum of heat (q) and work (w): \( \Delta E = q + w \).
3. Substitute the given values into the equation: \( \Delta E = 0.763 \text{ kJ} + w \text{ (in kJ)} \).
4. Determine the sign of \( \Delta E \). If \( \Delta E \) is positive, the process is endothermic (absorbs energy). If \( \Delta E \) is negative, the process is exothermic (releases energy).
5. Interpret the result to conclude whether the process is endothermic or exothermic based on the sign of \( \Delta E \).

Key Concepts

First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed from one form to another. In a closed system, the change in internal energy (ΔE) is equal to the heat added to the system (q) minus the work done by the system (w). This principle is fundamental for understanding energy transfers in chemical processes.

Internal Energy Change (ΔE)

The change in internal energy (ΔE) of a system is calculated using the equation ΔE = q + w, where q represents heat transfer and w represents work done. A positive ΔE indicates that the system has gained energy, while a negative ΔE indicates a loss of energy. This concept is crucial for determining the nature of the process as either endothermic or exothermic.

Endothermic vs. Exothermic Processes

Endothermic processes absorb heat from the surroundings, resulting in a positive q value, while exothermic processes release heat, leading to a negative q value. The sign of ΔE helps classify the process: if ΔE is positive, the process is endothermic; if negative, it is exothermic. Understanding these distinctions is essential for predicting the thermal behavior of chemical reactions.