Endothermic and Exothermic Processes
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AP Chemistry › Endothermic and Exothermic Processes
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 exothermic; heat is released so $\Delta H<0$.
The reaction is endothermic; heat is absorbed so $\Delta H>0$.
The reaction is spontaneous, so $\Delta G<0$ implies $\Delta H>0$.
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.
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 exothermic 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 endothermic because the beaker cools, so $\Delta H<0$.
The reaction is endothermic; heat is absorbed from 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).
A hand warmer contains iron powder that reacts with oxygen in air, and the packet becomes warm. Which statement best describes the oxidation process in terms of enthalpy?
The process is endothermic because the packet temperature increases, so $\Delta H<0$.
The process is exothermic; heat is released so $\Delta H<0$.
The process is exothermic because it is spontaneous, so $\Delta G<0$ means $\Delta H>0$.
The process is endothermic; heat is absorbed so $\Delta H>0$.
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.
A student uses a heat lamp to decompose calcium carbonate: $\text{CaCO}_3(s)\rightarrow \text{CaO}(s)+\text{CO}_2(g)$. The reaction only proceeds while heating continues. Which statement best describes the process at constant pressure?
The reaction is exothermic; heat is released so $\Delta H<0$.
The reaction is exothermic because a gas forms, so $\Delta H<0$.
The reaction is endothermic because it is slow, so $\Delta H<0$.
The reaction is endothermic; heat is absorbed so $\Delta H>0$.
The reaction is spontaneous, so $\Delta G<0$ means $\Delta H<0$.
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.
A student heats solid potassium chloride until it melts. During melting, the temperature remains constant even though the hot plate continues to supply energy. Which statement correctly classifies the melting process?
The process is endothermic; constant temperature means $\Delta H<0$.
The process is exothermic; heat is released as bonds form and $\Delta H<0$.
The process is endothermic; because the temperature is constant, $\Delta G>0$.
The process is endothermic; heat is absorbed to overcome attractions and $\Delta H>0$.
The process is exothermic; constant temperature means no heat transfer so $\Delta H=0$.
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.
A student dissolves $\text{NH}_4\text{NO}_3(s)$ in water in a coffee-cup calorimeter. The temperature of the solution decreases from $22.0^\circ\text{C}$ to $17.5^\circ\text{C}$ while the salt dissolves completely. Which statement correctly describes the process?
The process is exothermic; the temperature drop means $\Delta G<0$.
The process is endothermic; heat is released to the surroundings and $\Delta H>0$.
The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$.
The process is endothermic; the temperature drop means $\Delta H<0$.
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.
A student mixes barium hydroxide octahydrate, $\text{Ba(OH)}_2\cdot 8\text{H}_2\text{O}(s)$, with ammonium chloride, $\text{NH}_4\text{Cl}(s)$, in a beaker. The beaker becomes very cold and can even stick to a damp surface. Which statement best describes the overall process?
The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$.
The process is spontaneous only if $\Delta G<0$, so $\Delta H$ must be negative.
The process has $\Delta H=0$ because solids mixing cannot change enthalpy.
The process is exothermic because the temperature decreases, so $\Delta H<0$ for the system.
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.
When methane burns in oxygen in a well-insulated metal can, the can becomes hot. Which statement best describes the combustion process in terms of $\Delta H$ and heat flow?
The process is endothermic because the can warms, meaning heat flows into the system.
The process is exothermic only if $\Delta G<0$, so $\Delta H$ must be negative.
The process is exothermic; heat is released by the system and $\Delta H<0$.
The process is endothermic; heat is absorbed by the system and $\Delta H>0$.
The process has $\Delta H=0$ because the system is insulated, so no enthalpy change occurs.
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.
When solid ammonium nitrate, $\text{NH}_4\text{NO}_3(s)$, is stirred into water in a coffee-cup calorimeter, the solution temperature decreases noticeably. Assuming the process occurs at approximately constant pressure, which statement best describes the heat flow for the dissolution?
The process is exothermic; $\Delta G<0$ so heat must be released.
The process is endothermic; heat is absorbed from the surroundings and $\Delta H>0$.
The process is exothermic; heat is released to the surroundings and $\Delta H<0$.
The process is endothermic; the temperature decreases so $\Delta H<0$.
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).
A hand warmer packet contains iron powder that reacts with oxygen in air. After exposure to air, the packet warms up. Which statement best describes the process?
The process is endothermic; heat flows into the packet from the surroundings and $\Delta H>0$.
The process is exothermic; heat flows out of the packet to the surroundings and $\Delta H<0$.
The process is endothermic; the temperature increase means $\Delta H<0$.
The process is exothermic; heat flows into the packet from the surroundings and $\Delta H>0$.
The process is exothermic; the temperature increase means $\Delta G>0$ 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.