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Chapter 6: Problem 19

Standard enthalpies of formation are relative values. What are \(\Delta H_{\mathrm{f}}^{\circ}\) values relative to?

Short Answer

Expert verified
The standard enthalpies of formation (\(\Delta H_{\mathrm{f}}^{\circ}\)) values are relative to the enthalpy of the pure elements in their standard states at a specific temperature and pressure, usually 25掳C (298.15 K) and 1 atm pressure. By convention, the enthalpy of formation for the standard state of an element is assigned a value of zero.

Step by step solution

01

Definition of Standard Enthalpy of Formation

The standard enthalpy of formation (\(\Delta H_{\mathrm{f}}^{\circ}\)) is the change in enthalpy during the formation of one mole of a substance in its standard state from its constituent elements in their standard states. It is a thermodynamic quantity that helps in understanding the stability of a compound and predicting the heat changes during chemical reactions.
02

The Reference State

The \(\Delta H_{\mathrm{f}}^{\circ}\) values are relative to the enthalpy of the pure elements in their standard states at a specific temperature and pressure, which is usually 25掳C (298.15 K) and 1 atm pressure. The standard state of an element is the most stable form of the element at the specified temperature and pressure. For example, the standard state of oxygen is O2 gas, the standard state of carbon is graphite, and the standard state of hydrogen is H2 gas. By convention, the enthalpy of formation for the standard state of an element is assigned a value of zero.
03

Interpretation of \(\Delta H_{\mathrm{f}}^{\circ}\) Values

When comparing \(\Delta H_{\mathrm{f}}^{\circ}\) values for different compounds, a more negative value indicates a more stable compound (with respect to its constituent elements in their standard states), as more energy is released during its formation. Conversely, a more positive value indicates a less stable compound. In conclusion, the standard enthalpy of formation (\(\Delta H_{\mathrm{f}}^{\circ}\)) values are relative to the enthalpy of the pure elements in their standard states and are a useful tool for understanding compound stability and predicting heat changes in chemical reactions.

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Most popular questions from this chapter

A sample of nickel is heated to \(99.8^{\circ} \mathrm{C}\) and placed in a coffeecup calorimeter containing \(150.0 \mathrm{~g}\) water at \(23.5^{\circ} \mathrm{C}\). After the metal cools, the final temperature of metal and water mixture is \(25.0^{\circ} \mathrm{C}\). If the specific heat capacity of nickel is \(0.444 \mathrm{~J} /{ }^{\circ} \mathrm{C} \cdot \mathrm{g}\), what mass of nickel was originally heated? Assume no heat loss to the surroundings.

Consider \(5.5 \mathrm{~L}\) of a gas at a pressure of \(3.0 \mathrm{~atm}\) in a cylinder with a movable piston. The external pressure is changed so that the volume changes to \(10.5 \mathrm{~L}\). a. Calculate the work done, and indicate the correct sign. b. Use the preceding data but consider the process to occur in two steps. At the end of the first step, the volume is \(7.0 \mathrm{~L}\). The second step results in a final volume of \(10.5\) L. Calculate the work done, and indicate the correct sign. c. Calculate the work done if after the first step the volume is \(8.0 \mathrm{~L}\) and the second step leads to a volume of \(10.5 \mathrm{~L}\). Does the work differ from that in part b? Explain.

A cubic piece of uranium metal (specific heat capacity \(=0.117\) \(\mathrm{J} /{ }^{\circ} \mathrm{C} \cdot \mathrm{g}\) ) at \(200.0^{\circ} \mathrm{C}\) is dropped into \(1.00 \mathrm{~L}\) deuterium oxide ("heavy water," specific heat capacity \(=4.211 \mathrm{~J} /{ }^{\circ} \mathrm{C} \cdot \mathrm{g}\) ) at \(25.5^{\circ} \mathrm{C}\). The final temperature of the uranium and deuterium oxide mixture is \(28.5^{\circ} \mathrm{C}\). Given the densities of uranium \(\left(19.05 \mathrm{~g} / \mathrm{cm}^{3}\right)\) and deuterium oxide (1.11 \(\mathrm{g} / \mathrm{mL}\) ), what is the edge length of the cube of uranium?

Some automobiles and buses have been equipped to burn propane \(\left(\mathrm{C}_{3} \mathrm{H}_{8}\right) .\) Compare the amounts of energy that can be obtained per gram of \(\mathrm{C}_{3} \mathrm{H}_{8}(g)\) and per gram of gasoline, assuming that gasoline is pure octane, \(\mathrm{C}_{8} \mathrm{H}_{18}(l) .\) (See Example \(6.11 .\) ) Look up the boiling point of propane. What disadvantages are there to using propane instead of gasoline as a fuel?

Consider the dissolution of \(\mathrm{CaCl}_{2}\) : $$ \mathrm{CaCl}_{2}(s) \longrightarrow \mathrm{Ca}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q) \quad \Delta H=-81.5 \mathrm{~kJ} $$ An \(11.0-\mathrm{g}\) sample of \(\mathrm{CaCl}_{2}\) is dissolved in \(125 \mathrm{~g}\) water, with both substances at \(25.0^{\circ} \mathrm{C}\). Calculate the final temperature of the solution assuming no heat loss to the surroundings and assuming the solution has a specific heat capacity of \(4.18 \mathrm{~J} /{ }^{\circ} \mathrm{C} \cdot \mathrm{g}\).
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