Endothermic and Exothermic Processes
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AP Chemistry › Endothermic and Exothermic Processes
When 1.0 mol of calcium chloride, $\text{CaCl}_2(s)$, is added to water and stirred, the beaker becomes noticeably warm to the touch. Assuming pressure is constant, which statement best describes the process and the sign of $\Delta H$ for dissolving $\text{CaCl}_2(s)$?
The process is endothermic; heat flows from the solution to the surroundings and $\Delta H>0$.
The process is exothermic; heat flows into the solution from the surroundings and $\Delta H<0$.
The process is nonspontaneous; therefore $\Delta G>0$ and $\Delta H<0$.
The process is endothermic; the beaker feels warm so $\Delta H>0$.
The process is exothermic; heat flows from the solution to the surroundings and $\Delta H<0$.
Explanation
This question tests the skill of identifying exothermic processes in dissolution by observing heat transfer to the surroundings and assigning the sign of ΔH. When calcium chloride dissolves, the beaker becomes warm, meaning heat is released from the solution to the surroundings. Exothermic processes release heat, leading to a negative ΔH for the system. This matches choice B, as the warming indicates energy is given off during ion hydration exceeding lattice energy costs. A tempting distractor is choice A, which misclassifies the process as endothermic despite heat flowing from the system, arising from the misconception that heat transfer to surroundings implies absorption by the system. A transferable strategy is to define the system clearly and track heat flow direction to determine if ΔH is positive or negative.
A student mixes equal volumes of 1.0 M HCl(aq) and 1.0 M NaOH(aq) in a coffee-cup calorimeter. The temperature of the mixture increases. Which statement correctly describes the neutralization process and the sign of $\Delta H$ for the reaction?
The process is spontaneous; therefore $\Delta G<0$ and $\Delta H>0$.
The process is exothermic; heat flows out of the reaction system and $\Delta H<0$.
The process is exothermic; temperature increases so $\Delta H>0$.
The process is endothermic; temperature increases so $\Delta H<0$.
The process is endothermic; heat flows into the solution and $\Delta H>0$.
Explanation
This question examines identifying neutralization reactions as exothermic through temperature increases and assigning the sign of ΔH. Mixing HCl and NaOH causes the mixture's temperature to rise, showing heat release from the reaction. Exothermic processes transfer heat out of the system, leading to a negative ΔH. Choice B correctly explains this, as the temperature increase in the calorimeter indicates energy liberation. A tempting distractor is choice C, which states it's exothermic but with ΔH > 0 due to temperature rise, arising from the misconception that system warming means positive ΔH instead of negative. When evaluating reactions, monitor temperature changes in the surroundings to infer heat flow and ΔH sign.
A hand warmer contains iron powder that reacts with oxygen in air over time. When exposed to air, the packet warms up. Which statement best describes the oxidation process in the hand warmer?
The process is endothermic; warming indicates $\Delta G<0$ so $\Delta H>0$.
The process is endothermic; heat flows from the packet to the surroundings and $\Delta H>0$.
The process is spontaneous; therefore $\Delta G<0$ and $\Delta H>0$.
The process is exothermic; heat flows from the packet to the surroundings and $\Delta H<0$.
The process is exothermic; warming indicates $\Delta H>0$.
Explanation
This question evaluates classifying oxidation reactions as exothermic based on warming and determining the sign of ΔH. The iron powder in the hand warmer reacts with oxygen, causing the packet to warm up, indicating heat release to the surroundings. Exothermic processes give off heat, resulting in a negative ΔH for the reaction. This matches choice B, with the warming confirming energy flow from the system. A tempting distractor is choice D, which says it's exothermic but with ΔH > 0 due to warming, from the misconception that temperature increase in surroundings means positive ΔH for the system. To analyze such reactions, track if the process heats the environment, signifying exothermic nature and negative ΔH.
Hydrogen gas burns in oxygen according to $2\text{H}_2(g)+\text{O}_2(g)\rightarrow 2\text{H}_2\text{O}(l)$. The reaction mixture gives off heat and bright light. At constant pressure, which statement best describes the reaction and the sign of $\Delta H$?
The process is endothermic; light emission means $\Delta G>0$ so $\Delta H>0$.
The process is exothermic; heat flows out of the system and $\Delta H<0$.
The process is spontaneous; therefore $\Delta G<0$ and $\Delta H>0$.
The process is endothermic; heat flows into the system and $\Delta H>0$.
The process is exothermic; products have higher enthalpy so $\Delta H>0$.
Explanation
This question tests classifying combustion reactions as exothermic based on heat and light emission and determining ΔH's sign. The reaction of hydrogen with oxygen produces water, releasing heat and bright light, indicating energy is given off. Exothermic reactions release heat to the surroundings, resulting in a negative ΔH. This is described in choice B, with the heat flow out of the system confirming the classification. A tempting distractor is choice A, which mislabels it endothermic with heat flowing into the system, from the misconception that light emission requires heat absorption rather than release. A useful strategy is to observe if the surroundings gain energy, signaling an exothermic process with negative ΔH.
A student dissolves 10.0 g of ammonium nitrate, $\text{NH}_4\text{NO}_3(s)$, in 100 mL of water in a foam cup. The temperature of the solution drops from $22.0^\circ\text{C}$ to $16.5^\circ\text{C}$ while the salt dissolves. Which statement correctly classifies the dissolving process and the sign of $\Delta H$ for the process?
The process is endothermic; the temperature decreases so $\Delta H<0$.
The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$.
The process is spontaneous; therefore $\Delta G<0$ and $\Delta H<0$.
The process is exothermic; heat is released to the solution and $\Delta H<0$.
The process is exothermic; the temperature decreases so $\Delta H<0$.
Explanation
This question tests the skill of classifying dissolution processes as endothermic or exothermic based on observed temperature changes and determining the sign of ΔH. The dissolution of ammonium nitrate causes the solution temperature to drop from 22.0°C to 16.5°C, indicating that heat is absorbed from the solution by the dissolving process. In endothermic processes, the system absorbs heat from the surroundings, resulting in a positive ΔH value. This aligns with choice B, as the temperature decrease shows energy is taken in to break solute-solute and solvent-solvent interactions. A tempting distractor is choice C, which incorrectly states the process is exothermic with ΔH < 0 due to the temperature decrease, stemming from the misconception that cooling always indicates heat release rather than absorption by the system. To classify such processes reliably, always consider whether the system is gaining or losing heat based on the surroundings' temperature change.
A chemical cold pack is activated by breaking an inner pouch, allowing a solid to dissolve in water. The pack quickly becomes cold. Which statement best describes the dissolving process occurring in the cold pack?
The process is endothermic; the pack is cold so $\Delta H<0$.
The process is nonspontaneous; therefore $\Delta G>0$ and $\Delta H>0$.
The process is exothermic; a lower temperature means $\Delta H<0$.
The process is exothermic; heat flows from the surroundings into the pack and $\Delta H<0$.
The process is endothermic; heat flows from the surroundings into the pack and $\Delta H>0$.
Explanation
This question tests classifying dissolution in cold packs as endothermic based on cooling effects and determining ΔH's sign. Activating the cold pack by dissolving the solid makes it cold, meaning heat is absorbed from the surroundings into the pack. Endothermic processes absorb heat, resulting in a positive ΔH for the system. This is captured in choice B, with the cooling confirming energy intake during dissolution. A tempting distractor is choice C, which calls it exothermic due to lower temperature implying ΔH < 0, based on the misconception that system cooling means heat release rather than absorption. A transferable approach is to identify if the process cools the surroundings, indicating endothermic absorption with positive ΔH.
Hydrogen peroxide decomposes in the presence of a catalyst: $2\text{H}_2\text{O}_2(aq)\rightarrow 2\text{H}_2\text{O}(l)+\text{O}_2(g)$. The container warms during the reaction. Which statement is correct?
The reaction is endothermic because the container warms, so $\Delta H<0$.
The reaction is spontaneous, so $\Delta G<0$ implies $\Delta H>0$.
The reaction is exothermic; heat is released so $\Delta H<0$.
The reaction is endothermic; heat is absorbed so $\Delta H>0$.
The reaction is exothermic because a catalyst is present, so $\Delta H<0$.
Explanation
This question tests classifying catalyzed decompositions based on temperature changes. The container warming during H2O2 decomposition indicates heat release, making it exothermic with ΔH < 0. Breaking O-O bonds and forming stronger ones releases energy. At constant pressure, this is negative ΔH. A tempting distractor is choice B, saying endothermic with ΔH > 0, based on the misconception that catalysts imply absorption. Warming indicates exothermic regardless of catalyst.
Solid iodine is heated gently and purple iodine vapor forms: $\text{I}_2(s)\rightarrow \text{I}_2(g)$. Which statement best describes the sublimation at constant pressure?
The process is endothermic; heat flows in so $\Delta H>0$.
The process is exothermic; heat flows out so $\Delta H<0$.
The process is spontaneous, so $\Delta G<0$ implies $\Delta H<0$.
The process is endothermic because it increases disorder, so $\Delta H<0$.
The process is exothermic because a gas forms, so $\Delta H>0$.
Explanation
This question tests the identification of sublimation as an endothermic or exothermic phase change. Heating solid iodine to form purple vapor requires energy input to overcome intermolecular forces, making sublimation endothermic with ΔH > 0. The process absorbs heat to transition from solid to gas without a liquid phase, increasing potential energy. At constant pressure, this heat absorption corresponds to the positive enthalpy of sublimation. A tempting distractor is choice A, labeling it exothermic with ΔH < 0, due to the misconception that gas formation implies energy release instead of absorption. For phase changes, remember that increasing molecular freedom (solid to gas) is endothermic, while decreasing it is exothermic.
A student dissolves solid $\text{NH}_4\text{NO}_3$ in water in an open cup and observes that the solution temperature decreases noticeably. Assuming pressure is approximately constant, which statement best describes the process?
The process is endothermic because the temperature decreases, so $\Delta H<0$.
The process is exothermic because the temperature decreases, so $\Delta H>0$.
The process is spontaneous, so $\Delta H<0$ even if the cup cools.
The process is exothermic; heat is released to the surroundings so $\Delta H<0$.
The process is endothermic; heat is absorbed from the surroundings so $\Delta H>0$.
Explanation
This question tests the skill of identifying endothermic and exothermic processes based on observed temperature changes in chemical systems. When solid NH4NO3 dissolves in water and the solution temperature decreases, it indicates that the system is absorbing heat from the surroundings to facilitate the dissolution process. This heat absorption means the process is endothermic, as energy is required to break the ionic bonds in the solid and hydrate the ions, resulting in a positive enthalpy change (ΔH > 0). The open cup at constant pressure allows us to equate the heat flow with ΔH, confirming that the surroundings cool because heat is transferred into the system. A tempting distractor is choice A, which incorrectly labels the process as exothermic with ΔH < 0, stemming from the misconception that a decrease in temperature always means heat is released by the system rather than absorbed. To distinguish between endothermic and exothermic processes, always consider the direction of heat flow: if the system absorbs heat (surroundings cool), it is endothermic; if the system releases heat (surroundings warm), it is exothermic.
In a lab demonstration, barium hydroxide octahydrate reacts with ammonium chloride and the beaker becomes very cold, sometimes freezing water beneath it. Which statement best describes the reaction's enthalpy change?
The reaction is endothermic; heat is absorbed from the surroundings so $\Delta H>0$.
The reaction is exothermic because the beaker cools, so $\Delta H>0$.
The reaction is endothermic because the beaker cools, so $\Delta H<0$.
The reaction is exothermic; heat is released to the surroundings so $\Delta H<0$.
The reaction is spontaneous, so $\Delta G<0$ ensures $\Delta H<0$.
Explanation
This question tests identifying reactions that cool surroundings as endothermic or exothermic. The beaker becoming very cold during the reaction of barium hydroxide octahydrate and ammonium chloride indicates heat absorption from the surroundings, making it endothermic with ΔH > 0. The process requires energy to break bonds and form new species, cooling the system significantly. At constant pressure, this corresponds to a positive enthalpy change. A tempting distractor is choice B, labeling it exothermic with ΔH < 0, stemming from the misconception that extreme cooling means heat release rather than intense absorption. To evaluate, check if surroundings lose heat (endothermic) or gain heat (exothermic).