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Chapter 6: Problem 44

Are the following processes exothermic or endothermic? a. the combustion of gasoline in a car engine b. water condensing on a cold pipe c. \(\mathrm{CO}_{2}(s) \longrightarrow \mathrm{CO}_{2}(g)\) d. \(\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{~F}(g)\)

Short Answer

Expert verified
a. Exothermic b. Exothermic c. Endothermic d. Endothermic

Step by step solution

01

Analyze combustion of gasoline in a car engine

The combustion of gasoline in a car engine is a chemical reaction that occurs when gasoline is burned. This reaction produces heat and releases energy, which is used to power the car. As a result, this process is exothermic since energy is released as heat to the surroundings.
02

Analyze water condensing on a cold pipe

In this process, water vapor molecules lose energy when they come into contact with a cold pipe and change from the gas phase to the liquid phase. This energy loss is manifested as a release of heat, as the molecules come together and result in a lower energy state. Since energy is released to the surroundings during this process, it is an exothermic process.
03

Analyze \(\mathrm{CO}_{2}(s) \longrightarrow \mathrm{CO}_{2}(g)\)

The given process describes the sublimation of solid carbon dioxide (CO2) into gaseous CO2. In this process, CO2 molecules in the solid phase gain energy to break the bonds holding them together, allowing them to enter the gas phase. Because energy is absorbed from the surroundings to facilitate this phase change, the process is endothermic.
04

Analyze \(\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{~F}(g)\)

This process represents the dissociation of fluorine gas (F2) into individual fluorine atoms (F). To break the bond between the two fluorine atoms in the F2 molecule, energy must be absorbed from the surroundings. As a result, this process is endothermic. To summarize: a. Exothermic b. Exothermic c. Endothermic d. Endothermic

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Most popular questions from this chapter

For the process \(\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g)\) at \(298 \mathrm{~K}\) and \(1.0 \mathrm{~atm}\), \(\Delta H\) is more positive than \(\Delta E\) by \(2.5 \mathrm{~kJ} / \mathrm{mol}\). What does the \(2.5 \mathrm{~kJ} / \mathrm{mol}\) quantity represent?

A system undergoes a process consisting of the following two steps: Step \(1:\) The system absorbs \(72 \mathrm{~J}\) of heat while \(35 \mathrm{~J}\) of work is done on it. Step 2: The system absorbs \(35 \mathrm{~J}\) of heat while performing \(72 \mathrm{~J}\) of work. Calculate \(\Delta E\) for the overall process.

One way to lose weight is to exercise! Walking briskly at \(4.0\) miles per hour for an hour consumes about 400 kcal of energy. How many hours would you have to walk at \(4.0\) miles per hour to lose one pound of body fat? One gram of body fat is equivalent to \(7.7\) kcal of energy. There are \(454 \mathrm{~g}\) in \(1 \mathrm{lb}\).

Given the following data $$ \begin{aligned} \mathrm{Ca}(s)+2 \mathrm{C}(\text { graphite }) & \longrightarrow \mathrm{CaC}_{2}(s) & \Delta H &=-62.8 \mathrm{~kJ} \\ \mathrm{Ca}(s)+\frac{1}{2} \mathrm{O}_{2}(g) & \longrightarrow \mathrm{CaO}(s) & \Delta H &=-635.5 \mathrm{~kJ} \\ \mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) & \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q) & \Delta H &=-653.1 \mathrm{~kJ} \\ \mathrm{C}_{2} \mathrm{H}_{2}(g)+\frac{5}{2} \mathrm{O}_{2}(g) & \longrightarrow 2 \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) & \Delta H &=-1300 . \mathrm{kJ} \\ \mathrm{C}(\text { graphite })+\mathrm{O}_{2}(g) & \mathrm{CO}_{2}(\mathrm{~g}) & \Delta H &=-393.5 \mathrm{~kJ} \end{aligned} $$ calculate \(\Delta H\) for the reaction $$ \mathrm{CaC}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)+\mathrm{C}_{2} \mathrm{H}_{2}(g) $$

The specific heat capacity of silver is \(0.24 \mathrm{~J} /{ }^{\circ} \mathrm{C} \cdot \mathrm{g}\). a. Calculate the energy required to raise the temperature of \(150.0 \mathrm{~g}\) Ag from \(273 \mathrm{~K}\) to \(298 \mathrm{~K}\). b. Calculate the energy required to raise the temperature of \(1.0\) mole of Ag by \(1.0^{\circ} \mathrm{C}\) (called the molar heat capacity of silver). c. It takes \(1.25 \mathrm{~kJ}\) of energy to heat a sample of pure silver from \(12.0^{\circ} \mathrm{C}\) to \(15.2^{\circ} \mathrm{C}\). Calculate the mass of the sample of silver.
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