\( \qquad \) (1) \( Q_{1 C F}=c \mathrm{mDt}=2.03 \mathrm{~J} / \mathrm{g} \cdot \mathrm{C} \cdot 95 \mathrm{~g} \cdot 12^{\circ} \mathrm{C}=2,314.2 \mathrm{~J}=2.3142 \mathrm{~kJ} \) (2) \( \Delta H_{f}=\frac{95 \mathrm{~g} H_{1} \mathrm{O}\left|\operatorname{lmol}_{18 \mathrm{D}}\right| 6.02 \mathrm{~kJ}}{18.02 \mathrm{~g} H_{50} \mid \mathrm{lmolH}_{2} \mathrm{D}}=31.7 \mathrm{~kJ}=31.7 \mathrm{~kJ} \) (3) \( Q_{\text {wattr }}=c m \Delta t=4.18 \mathrm{~J} / \mathrm{g} \cdot \mathrm{C} \cdot 95 \mathrm{~g} \cdot 25^{\circ} \mathrm{C}=9,927.5 \mathrm{~J}=9.92 \mathrm{~J} \mathrm{~kJ} \) \[ +43.9 \mathrm{~kJ} \] Rewatch Released (D) Mull 2) Is this energy released or absorbed? Absorbed Released

Energy changes in a chemical reaction can be quite fascinating! When a substance undergoes a phase change or reaction, it either absorbs or releases energy, which is quantified by enthalpy changes. If the enthalpy change (\( \Delta H \)) is positive, it indicates that energy is absorbed from the surroundings (endothermic). Conversely, a negative \( \Delta H \) signifies that energy is released (exothermic). Understanding these concepts is essential in fields like thermodynamics and chemistry. In practical terms, knowing whether a reaction absorbs or releases energy is crucial for various applications. For instance, in designing chemical processes or even cooking, endothermic reactions might be deliberately used to achieve desired temperatures or effects. Likewise, in thermal management systems, understanding these energy exchanges helps engineers develop more efficient cooling or heating systems. So, the next time you see a reaction, think about whether it's a cozy energy hug or a spirited energy shove!