AP Chemistry Quiz: Endothermic And Exothermic Processes

What this quiz covers

This quiz focuses on Endothermic And Exothermic Processes, giving you a quick way to practice the rules, question types, and explanations that matter most for AP Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

Question 1

1. The reaction is exothermic; heat is released so $\Delta H<0$. (correct answer)
2. The reaction is endothermic; heat is absorbed so $\Delta H>0$.
3. The reaction is exothermic because a catalyst is present, so $\Delta H<0$.
4. The reaction is endothermic because the container warms, so $\Delta H<0$.
5. The reaction is spontaneous, so $\Delta G<0$ implies $\Delta H>0$.

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.

Question 2

1. The reaction is endothermic; heat is absorbed from the surroundings so $\Delta H>0$. (correct answer)
2. The reaction is exothermic; heat is released to the surroundings so $\Delta H<0$.
3. The reaction is endothermic because the beaker cools, so $\Delta H<0$.
4. The reaction is exothermic because the beaker cools, so $\Delta H>0$.
5. 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).

Question 3

1. The process is endothermic; heat is absorbed so $\Delta H>0$.
2. The process is exothermic; heat is released so $\Delta H<0$. (correct answer)
3. The process is exothermic because it is spontaneous, so $\Delta G<0$ means $\Delta H>0$.
4. The process is endothermic because the packet temperature increases, so $\Delta H<0$.
5. The process is endothermic because oxidation always requires heat input, so $\Delta H>0$.

Explanation: This question tests classifying oxidation reactions in hand warmers as endothermic or exothermic based on temperature effects. The packet becoming warm during the iron oxidation indicates heat release to the surroundings, confirming an exothermic process with ΔH < 0. Energy is released as iron forms stronger bonds with oxygen in rust, transferring heat outward. At constant pressure, this aligns with a negative enthalpy change. A tempting distractor is choice A, which says endothermic with ΔH > 0, arising from the misconception that warming means heat input rather than output from the system. Observe if the system generates heat (exothermic) or requires it (endothermic) to classify reactions.

Question 4

1. The reaction is exothermic because a gas forms, so $\Delta H<0$.
2. The reaction is spontaneous, so $\Delta G<0$ means $\Delta H<0$.
3. The reaction is endothermic; heat is absorbed so $\Delta H>0$. (correct answer)
4. The reaction is endothermic because it is slow, so $\Delta H<0$.
5. The reaction is exothermic; heat is released so $\Delta H<0$.

Explanation: This question tests identifying decomposition reactions needing heat as endothermic or exothermic. The requirement for continuous heating with a heat lamp to decompose CaCO3 indicates an endothermic process with ΔH > 0. Energy is absorbed to break bonds and form CO2. At constant pressure, heat input confirms positive ΔH. A tempting distractor is choice A, labeling it exothermic with ΔH < 0, arising from the misconception that gas formation means release. Reactions needing sustained heat are endothermic.

Question 5

1. The process is exothermic; heat is released as bonds form and $\Delta H<0$.
2. The process is endothermic; heat is absorbed to overcome attractions and $\Delta H>0$. (correct answer)
3. The process is exothermic; constant temperature means no heat transfer so $\Delta H=0$.
4. The process is endothermic; constant temperature means $\Delta H<0$.
5. The process is endothermic; because the temperature is constant, $\Delta G>0$.

Explanation: This question tests understanding of phase transitions and heat flow during melting. When KCl melts at constant temperature despite continuous heating, the supplied energy is being absorbed to break the ionic bonds in the solid lattice structure. Melting is endothermic because energy must be absorbed to overcome the attractive forces holding the solid together, making ΔH > 0. The constant temperature during melting occurs because all the absorbed energy goes into increasing potential energy (breaking bonds) rather than kinetic energy (temperature). Choice A incorrectly identifies melting as exothermic, confusing it with freezing which releases heat as bonds form. During any melting process, heat is absorbed (endothermic) even when temperature remains constant.

Question 6

1. The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
2. The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$. (correct answer)
3. The process is exothermic; the temperature drop means $\Delta G<0$.
4. The process is endothermic; the temperature drop means $\Delta H<0$.
5. The process is endothermic; heat is released to the surroundings and $\Delta H>0$.

Explanation: This question tests the ability to identify endothermic and exothermic processes based on temperature changes. When NH₄NO₃ dissolves in water and the temperature decreases from 22.0°C to 17.5°C, the solution is absorbing heat energy from its surroundings (the water and calorimeter). In an endothermic process, the system absorbs heat from the surroundings, causing the temperature of the surroundings to decrease, and ΔH > 0 because energy is added to the system. The dissolution of NH₄NO₃ requires energy to break apart the ionic lattice, and this energy requirement exceeds the energy released when water molecules hydrate the ions. Choice C incorrectly states the process is exothermic despite the temperature drop, confusing the sign of ΔH with ΔG. When you observe a temperature decrease in a calorimeter, always identify the process as endothermic with ΔH > 0.

Question 7

1. The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
2. The process has $\Delta H=0$ because solids mixing cannot change enthalpy.
3. The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$. (correct answer)
4. The process is exothermic because the temperature decreases, so $\Delta H<0$ for the system.
5. The process is spontaneous only if $\Delta G<0$, so $\Delta H$ must be negative.

Explanation: This question tests understanding of endothermic reactions between solids. When Ba(OH)₂·8H₂O and NH₄Cl react and the beaker becomes very cold, this indicates that the reaction absorbs heat from the surroundings, making it endothermic with ΔH > 0. The reaction produces NH₃ gas, water, and BaCl₂, but the energy required to break the reactant bonds and vaporize ammonia exceeds the energy released from forming new bonds, resulting in net heat absorption. The dramatic cooling can even freeze water between the beaker and a surface, demonstrating significant heat absorption. Choice D incorrectly concludes that temperature decrease means ΔH < 0 for the system, confusing the direction of heat flow - when the beaker cools, it's because heat flows INTO the reacting system. To identify strongly endothermic reactions, look for dramatic cooling effects that can freeze water or make containers painfully cold to touch.

Question 8

1. The process is endothermic; heat is absorbed by the system and $\Delta H>0$.
2. The process has $\Delta H=0$ because the system is insulated, so no enthalpy change occurs.
3. The process is exothermic; heat is released by the system and $\Delta H<0$. (correct answer)
4. The process is endothermic because the can warms, meaning heat flows into the system.
5. The process is exothermic only if $\Delta G<0$, so $\Delta H$ must be negative.

Explanation: This question tests understanding of combustion reactions as exothermic processes. When methane burns in oxygen and the can becomes hot, this indicates that the combustion reaction releases heat, making it exothermic with ΔH < 0. The reaction CH₄ + 2O₂ → CO₂ + 2H₂O releases substantial energy because the bonds formed in the products (C=O and O-H bonds) are much stronger than the bonds broken in the reactants (C-H and O=O bonds). This excess energy is released as heat, warming the can even though it's insulated. Choice D incorrectly interprets the can warming as meaning heat flows INTO the system, when actually the can warms because heat flows OUT of the reacting system. To identify combustion reactions thermodynamically, remember that all combustion reactions are exothermic because they release energy stored in chemical bonds as heat.

Question 9

1. The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
2. The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$. (correct answer)
3. The process is exothermic; $\Delta G<0$ so heat must be released.
4. The process is endothermic; the temperature decreases so $\Delta H<0$.
5. The process is neither; a temperature change does not indicate heat flow.

Explanation: This question tests the ability to identify endothermic and exothermic processes based on temperature changes. When ammonium nitrate dissolves in water and the temperature decreases, heat is being absorbed from the surroundings (the water) to break apart the ionic lattice and hydrate the ions. Since heat flows from the surroundings into the system, the process is endothermic and ΔH > 0. Choice D incorrectly states that a temperature decrease means ΔH < 0, confusing the sign convention—when the surroundings cool down, it's because heat has been absorbed by the system, making ΔH positive. A key strategy is to remember that if the surroundings get colder, heat has been absorbed by the system (endothermic, ΔH > 0), while if the surroundings get warmer, heat has been released by the system (exothermic, ΔH < 0).

Question 10

1. The process is endothermic; heat flows into the packet from the surroundings and $\Delta H>0$.
2. The process is exothermic; heat flows out of the packet to the surroundings and $\Delta H<0$. (correct answer)
3. The process is endothermic; the temperature increase means $\Delta H<0$.
4. The process is exothermic; the temperature increase means $\Delta G>0$ and $\Delta H<0$.
5. The process is exothermic; heat flows into the packet from the surroundings and $\Delta H>0$.

Explanation: This question tests the ability to classify oxidation reactions based on temperature changes. When iron powder reacts with oxygen and the packet warms up, the oxidation reaction is releasing heat to the surroundings, making it exothermic with ΔH < 0. The formation of iron oxide releases energy as new Fe-O bonds form, and this energy release exceeds any energy required to break bonds in the reactants. The heat flows from the reacting system (iron and oxygen) outward to warm your hands. Choice A incorrectly identifies the process as endothermic despite the temperature increase, reversing the relationship between heat flow and exothermic reactions. When a chemical reaction causes warming, it is exothermic with heat flowing out and ΔH < 0.

Question 11

1. The process is endothermic; heat flows from the solution to the surroundings and $\Delta H>0$.
2. The process is exothermic; heat flows from the solution to the surroundings and $\Delta H<0$. (correct answer)
3. The process is endothermic; the beaker feels warm so $\Delta H>0$.
4. The process is exothermic; heat flows into the solution from the surroundings and $\Delta H<0$.
5. The process is nonspontaneous; therefore $\Delta G>0$ 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.

Question 12

1. The process is endothermic; heat flows into the solution and $\Delta H>0$.
2. The process is exothermic; heat flows out of the reaction system and $\Delta H<0$. (correct answer)
3. The process is exothermic; temperature increases so $\Delta H>0$.
4. The process is endothermic; temperature increases so $\Delta H<0$.
5. The process is spontaneous; therefore $\Delta G<0$ 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.

Question 13

1. The process is endothermic; heat flows from the packet to the surroundings and $\Delta H>0$.
2. The process is spontaneous; therefore $\Delta G<0$ and $\Delta H>0$.
3. The process is exothermic; warming indicates $\Delta H>0$.
4. The process is endothermic; warming indicates $\Delta G<0$ so $\Delta H>0$.
5. The process is exothermic; heat flows from the packet to the surroundings and $\Delta H<0$. (correct answer)

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.

Question 14

1. The process is endothermic; heat flows into the system and $\Delta H>0$.
2. The process is exothermic; products have higher enthalpy so $\Delta H>0$.
3. The process is exothermic; heat flows out of the system and $\Delta H<0$. (correct answer)
4. The process is spontaneous; therefore $\Delta G<0$ and $\Delta H>0$.
5. The process is endothermic; light emission means $\Delta G>0$ 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.

Question 15

1. The process is exothermic; heat is released to the solution and $\Delta H<0$.
2. The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$. (correct answer)
3. The process is exothermic; the temperature decreases so $\Delta H<0$.
4. The process is endothermic; the temperature decreases so $\Delta H<0$.
5. The process is spontaneous; therefore $\Delta G<0$ and $\Delta H<0$.

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.

Question 16

1. The process is exothermic; heat flows from the surroundings into the pack and $\Delta H<0$.
2. The process is endothermic; heat flows from the surroundings into the pack and $\Delta H>0$. (correct answer)
3. The process is exothermic; a lower temperature means $\Delta H<0$.
4. The process is endothermic; the pack is cold so $\Delta H<0$.
5. The process is nonspontaneous; therefore $\Delta G>0$ 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.

Question 17

1. The process is spontaneous, so $\Delta G<0$ implies $\Delta H<0$.
2. The process is exothermic; heat flows out so $\Delta H<0$.
3. The process is exothermic because a gas forms, so $\Delta H>0$.
4. The process is endothermic because it increases disorder, so $\Delta H<0$.
5. The process is endothermic; heat flows in so $\Delta H>0$. (correct answer)

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.

Question 18

1. The process is exothermic; heat is released to the surroundings so $\Delta H<0$.
2. The process is endothermic; heat is absorbed from the surroundings so $\Delta H>0$. (correct answer)
3. The process is endothermic because the temperature decreases, so $\Delta H<0$.
4. The process is exothermic because the temperature decreases, so $\Delta H>0$.
5. The process is spontaneous, so $\Delta H<0$ even if the cup cools.

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.

Question 19

1. The reaction is endothermic; heat is absorbed so $\Delta H>0$. (correct answer)
2. The reaction is endothermic because activation energy is required, so $\Delta H<0$.
3. The reaction is spontaneous, so $\Delta G<0$ means $\Delta H>0$.
4. The reaction is exothermic because products are higher energy, so $\Delta H<0$.
5. The reaction is exothermic; heat is released so $\Delta H<0$.

Explanation: This question tests interpreting energy diagrams and heat requirements for endothermic or exothermic classification. Products at higher potential energy than reactants require net energy input, making the reaction endothermic with ΔH > 0. Continuous heat supply confirms absorption to reach higher energy state. The diagram shows positive ΔH. A tempting distractor is choice A, labeling it exothermic with ΔH < 0, stemming from the misconception that heat supply implies release rather than absorption. If products have higher energy, the reaction is endothermic.

Question 20

1. The process is exothermic; heat is released so $\Delta H<0$. (correct answer)
2. The process is endothermic; heat is absorbed so $\Delta H>0$.
3. The process is exothermic because temperature is constant, so $\Delta H=0$.
4. The process is endothermic because a solid forms, so $\Delta H<0$.
5. The process is spontaneous, so $\Delta G<0$ means $\Delta H>0$.

Explanation: This question tests identifying freezing as an endothermic or exothermic phase change. Cooling liquid ethanol to freeze it at constant temperature releases heat as molecules form a ordered solid, making it exothermic with ΔH < 0. The constant temperature means heat is expelled during lattice formation. At constant pressure, this is the negative enthalpy of fusion. A tempting distractor is choice C, labeling it endothermic with ΔH > 0, arising from the misconception that constant temperature implies absorption instead of release. Remember, transitions to more ordered states (liquid to solid) are exothermic.