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Chapter 9: Problem 53

What is the sign of \(\Delta H\) for an exothermic reaction? For an endothermic reaction?

Short Answer

Expert verified
For exothermic reactions, 螖H is negative; for endothermic reactions, 螖H is positive.

Step by step solution

01

Understanding Exothermic Reactions

An exothermic reaction releases energy to the surroundings, usually in the form of heat. This means that the energy of the products is lower than the energy of the reactants.
02

Identifying Sign of 螖H for Exothermic

Since energy is released, the change in enthalpy (螖H) for an exothermic reaction is negative, indicating a loss of energy from the system to the surroundings.
03

Understanding Endothermic Reactions

An endothermic reaction absorbs energy from the surroundings, usually as heat. As a result, the energy of the products is higher than the energy of the reactants.
04

Identifying Sign of 螖H for Endothermic

Since energy is absorbed, the change in enthalpy (螖H) for an endothermic reaction is positive, indicating a gain of energy by the system from the surroundings.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Exothermic Reaction
In chemical terms, an exothermic reaction is one that releases energy into its surroundings. This energy is often released in the form of heat but can also include light or sound. During these reactions, the energy of the products is less than the energy of the reactants, because energy is given off during the reaction process.
One characteristic feature of exothermic reactions is that they often result in an increase in the temperature of the surroundings.
  • Example: Combustion reactions, such as burning wood or gasoline, are typical exothermic processes.
  • In these processes, the chemical bonds in the reactants break and new bonds form in the products, releasing heat.
Thus, for an exothermic reaction, the enthalpy change (螖H) is negative, as energy is lost from the system to the surroundings.
Endothermic Reaction
Conversely, an endothermic reaction is a process in which energy is absorbed from the surroundings into the system. This usually occurs in the form of heat uptake. Consequently, the energy of the products in an endothermic reaction is greater than the energy of the reactants. This intake of energy is essential to drive the reaction forward.
During endothermic reactions, you might notice a drop in temperature in the surroundings because the reaction absorbs heat.
  • Example: Photosynthesis in plants is a classic endothermic reaction, where energy from sunlight is absorbed to convert carbon dioxide and water into glucose and oxygen.
  • Another common example is the reaction of barium hydroxide with ammonium thiocyanate, where the system absorbs energy, causing the container to feel cold.
For endothermic reactions, the enthalpy change (螖H) is positive, indicating an energy gain by the system from its surroundings.
Energy Absorption and Release
Understanding energy absorption and release is crucial in thermodynamics and helps us predict the behavior of chemical reactions.
  • For exothermic reactions, energy is released, often increasing the surrounding temperature. This is due to the fact that the energy required to break the bonds in the reactants is less than the energy released when new bonds form in the products.
  • In endothermic reactions, energy is absorbed, often resulting in a cooling effect on the surroundings as the system takes in energy to facilitate the reaction.
The sign of the enthalpy change (\(\Delta H\)) reflects these energy changes:
  • Negative \(\Delta H\) indicates energy release (exothermic).
  • Positive \(\Delta H\) indicates energy absorption (endothermic).
Understanding these concepts helps in determining the energy benefits or needs of a reaction, which is especially important in fields ranging from industrial chemistry to environmental science.

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Most popular questions from this chapter

When a solution containing \(8.00 \mathrm{~g}\) of \(\mathrm{NaOH}\) in \(50.0 \mathrm{~g}\) of water at \(25.0^{\circ} \mathrm{C}\) is added to a solution of \(8.00 \mathrm{~g}\) of \(\mathrm{HCl}\) in \(250.0 \mathrm{~g}\) of water at \(25.0^{\circ} \mathrm{C}\) in a calorimeter, the temperature of the solution increases to \(33.5^{\circ} \mathrm{C}\). Assuming that the specific heat of the solution is \(4.18 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\) and that the calorimeter itself absorbs a negligible amount of heat, calculate \(\Delta H\) in kilojoules \(/ \mathrm{mol}\) for the reaction $$\mathrm{NaOH}(a q)+\mathrm{HCl}(a q) \longrightarrow \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$$ When the experiment is repeated using a solution of \(10.00 \mathrm{~g}\) of HCl in \(248.0 \mathrm{~g}\) of water, the same temperature increase is observed. Explain.

Citric acid has three dissociable hydrogens. When \(5.00 \mathrm{~mL}\) of 0.64 M citric acid and \(45.00 \mathrm{~mL}\) of \(0.77 \mathrm{M} \mathrm{NaOH}\) are mixed at an initial temperature of \(26.0^{\circ} \mathrm{C}\), the temperature rises to \(27.9^{\circ} \mathrm{C}\) as the citric acid is neutralized. The combined mixture has a mass of \(51.6 \mathrm{~g}\) and a specific heat of \(4.0 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\). Assuming that no heat is transferred to the surroundings, calculate the enthalpy change for the reaction of \(1.00 \mathrm{~mol}\) of citric acid in \(\mathrm{kJ}\). Is the reaction exothermic or endothermic?

What is the difference between the internal energy change \(\Delta E\) and the enthalpy change \(\Delta H ?\) Which of the two is measured at constant pressure, and which at constant volume?

When \(1.045 \mathrm{~g}\) of \(\mathrm{CaO}\) is added to \(50.0 \mathrm{~mL}\) of water at \(25.0^{\circ} \mathrm{C}\) in a calorimeter, the temperature of the water increases to \(32.3^{\circ} \mathrm{C}\). Assuming that the specific heat of the solution is \(4.18 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\) and that the calorimeter itself absorbs a negligible amount of heat, calculate \(\Delta H\) in kilojoules \(/ \mathrm{mol} \mathrm{Ca}(\mathrm{OH})_{2}\) for the reaction $$\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)$$

Calculate the amount of heat required to raise the temperature of \(250.0 \mathrm{~g}\) (approximately 1 cup) of hot chocolate from \(25.0^{\circ} \mathrm{C}\) to \(80.0^{\circ} \mathrm{C}\). Assume hot chocolate has the same specific heat as water \(\left[4.18 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\right]\)
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