Chapter Questions
Explain the difference between kinetic and potential energy.
Why isn’t all energy available for work?
Explain what is meant by the term state function. Provide examples of state functions.
Which ball has the higher potential energy? Explain.
The law of conservation of energy means that energy is a state function. Explain why.
Explain the differences among heat, temperature, and thermal energy.
Do each of the following depend on the amount of substance you have? Explain.a. temperatureb. thermal energy
Provide a molecular-level explanation of why the temperatures of a cold soft drink and hot coffee in the same room will eventually be the same.
In which case is more heat involved: mixing 100.0-g samples of 90 °C water and 80 °C water or mixing 100.0-g samples of 60 °C water and 10 °C water? Assume no heat is lost to the environment.
If 100.0 g of water at 90 °C is added to 50.0 g of water at 10 °C, estimate the final temperature of the water. Explain your reasoning.
What is meant by potential energy in a chemical reaction? Where is it located?
What does it mean for a chemical to have a low potential energy?
In an endothermic reaction, do the reactants or the products have the lower potential energy?
Are the following processes exothermic or endothermic?a. When solid KBr is dissolved in water, the solution gets colder.b. Natural gas $\left(\mathrm{CH}_{4}\right)$ is burned in a furnace.c. When concentrated sulfuric acid is added to water, the solution gets very hot.d. Water is boiled in a teakettle.
What does it mean when energy is reported with a negative sign? A positive sign?
What does it mean when work is reported with a negative sign? A positive sign?
In thermodynamics, the chemist takes the system’s point of view. Explain.
How do engineers define energy and work differently from chemists? Why do engineers take this approach?
Calculate $\Delta E$ for each of the following cases.a. $q=+51 \mathrm{kJ}, w=-15 \mathrm{kJ}$b. $q=+100 . \mathrm{kJ}, w=-65 \mathrm{kJ}$c. $q=-65 \mathrm{kJ}, w=-20 . \mathrm{kJ}$
A gas absorbs 45 $\mathrm{kJ}$ of heat and does 29 $\mathrm{kJ}$ of work. Calculate $\Delta E .$
A system releases 125 $\mathrm{kJ}$ of heat, and 104 $\mathrm{kJ}$ of work is done on it. Calculate $\Delta E .$
Describe what happens to the molecules in a sample of ice as the sample is slowly heated until it liquefies and then vaporizes.
Metallic substances tend to have (lower/higher) specific heat capacities than nonmetallic substances.
If it takes 526 J of energy to warm 7.40 g of water by 17 °C, how much energy would be needed to warm 7.40 g of water by 55 °C?
Convert the following numbers of calories or kilocalories into joules or kilojoules.a. 7845 calb. $4.55 \times 10^{4}$ calc. 62.142 kcald. $43,024$ cal
If 7.24 kJ of heat is applied to a 952-g block of metal, the temperature increases by 10.7 °C. Calculate the specific heat capacity of the metal in J/g °C.
Three 75.0-g samples of copper, silver, and gold are available. Each of these samples is initially at 24.0°C, and then 2.00 kJ of heat is applied to each sample. Which sample will end up at the highest temperature?
A 35.2-g sample of metal X requires 1251 J of energy to heat the sample by 25.0 °C. Calculate the specific heat capacity of this metal.
The equation for the fermentation of glucose to alcohol and carbon dioxide is$$\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(a q) \rightarrow 2 \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(a q)+2 \mathrm{CO}_{2}(g)$$The enthalpy change for the reaction is $-67 \mathrm{~kJ} .$ Is the reaction endothermic or exothermic? Is energy, in the form of heat, absorbed or released as the reaction occurs?
For the reaction$$S(s)+O_{2}(g) \rightarrow \mathrm{SO}_{2}(g) \Delta H=-296 \mathrm{kJ} / \mathrm{mol}$$a. How much heat is released when 275 g of sulfur is burned in excess oxygen?b. How much heat is released when 25 mol of sulfur is burned in excess oxygen?c. How much heat is released when 150. g of sulfur dioxide is produced?
Calculate the enthalpy change when 1.00 g of methane is burned in excess oxygen according to the reaction$$\begin{array}{c}{\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)} \\ {\Delta H=-891 \mathrm{kJ} / \mathrm{mol}}\end{array}$$
Given the following data:$$S(s)+\frac{3}{2} O_{2}(g) \rightarrow S O_{3}(g) \quad \Delta H=-395.2 \mathrm{kJ}$$$$2 \mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{SO}_{3}(g) \quad \Delta H=-198.2 \mathrm{kJ}$$calculate $\Delta H$ for the reaction$$\mathrm{S}(s)+\mathrm{O}_{2}(g) \rightarrow \mathrm{SO}_{2}(g)$$
Given the following data:$$\mathrm{C}_{2} \mathrm{H}_{2}(g)+\frac{5}{2} \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)\Delta H=-1300 . \mathrm{kJ}$$$$\mathrm{C}(s)+\mathrm{O}_{2}(g) \quad \rightarrow \mathrm{CO}_{2}(g) \Delta H=-394 \mathrm{kJ}$$$$\mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \quad \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \quad \Delta H=-286 \mathrm{kJ}$$calculate $\Delta H$ for the reaction$$2 \mathrm{C}(s)+\mathrm{H}_{2}(g) \rightarrow \mathrm{C}_{2} \mathrm{H}_{2}(g)$$
Given the following data:$$\begin{aligned} 2 \mathrm{O}_{3}(g) \rightarrow 3 \mathrm{O}_{2}(g) & \Delta H=-427 \mathrm{kJ} \\ \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{O}(g) & \Delta H=+495 \mathrm{kJ} \end{aligned}$$$$\mathrm{NO}(g)+\mathrm{O}_{3}(g) \rightarrow \mathrm{NO}_{2}(g)+\mathrm{O}_{2}(g)\Delta H=-199 \mathrm{kJ}$$calculate $\Delta H$ for the reaction$$\mathrm{NO}(g)+\mathrm{O}(g) \rightarrow \mathrm{NO}_{2}(g)$$
Given the following data:$$\begin{array}{c}{\mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{CO}(g) \rightarrow 2 \mathrm{Fe}(s)+3 \mathrm{CO}_{2}(g)} \\ {\Delta H^{\circ}=-23 \mathrm{kJ}}\end{array}$$$$\begin{array}{c}{3 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+\mathrm{CO}(g) \rightarrow 2 \mathrm{Fe}_{3} \mathrm{O}_{4}(s)+\mathrm{CO}_{2}(g)} \\ {\Delta H^{2}=-39 \mathrm{kJ}}\end{array}$$$$\begin{array}{c}{\mathrm{Fe}_{3} \mathrm{O}_{4}(s)+\mathrm{CO}(g) \rightarrow 3 \mathrm{FeO}(s)+\mathrm{CO}_{2}(g)} \\ {\Delta H^{\circ}=-18 \mathrm{kJ}}\end{array}$$calculate $\Delta H^{\circ}$ for the reaction$$\mathrm{FeO}(\mathrm{s})+\mathrm{CO}(g) \rightarrow \mathrm{Fe}(s)+\mathrm{CO}_{2}(g)$$
What is the difference between the quality of energy and the quantity of energy? Which is decreasing?
Why can no work be done when everything in the universe is at the same temperature?
What was the advantage of using tetraethyl lead in gasoline? What were two disadvantages?
Why do we need some greenhouse gases? What is the problem with having too much of the greenhouse gases?
Which energy sources used in the United States have declined the most in the last 150 years? Which have increased the most?
Why can’t the first law of thermodynamics explain why a ball doesn’t spontaneously roll up a hill?
What is meant by the term driving force?
What is the driving force of an exothermic reaction—matter spread or energy spread? Explain.
Exothermic reactions have a driving force. Nevertheless, water melting into a liquid is endothermic and this process occurs at room conditions. Explain why.
Consider a sample of steam (water in the gaseous state) at 150 °C. Describe what happens to the molecules in the sample as the sample is slowly cooled until it liquefies and then solidifies.
Convert the following numbers of kilojoules into kilocalories. (Remember: kilo means 1000.)a. 462.4 kJb. 18.28 kJc. 1.014 kJd. 190.5 kJ
Perform the indicated conversions.a. 45.62 kcal into kilojoulesb. 72.94 kJ into kilocaloriesc. 2.751 kJ into caloriesd. 5.721 kcal into joules
Calculate the amount of energy required (in calories) to heat 145 g of water from 22.3 °C to 75.0 °C.
It takes 1.25 kJ of energy to heat a certain sample of pure silver from 12.0 °C to 15.2 °C. Calculate the mass of the sample of silver.
If 50. J of heat is applied to 10. g of iron, by how much will the temperature of the iron increase? (See Table 10.1.)
The specific heat capacity of gold is 0.13 J/g °C. Calculate the specific heat capacity of gold in cal/g °C.
Calculate the amount of energy required (in joules) to heat 2.5 kg of water from 18.5 °C to 55.0 °C.
If 10. J of heat is applied to 5.0-g samples of each of the substances listed in Table 10.1, which substance’s temperature will increase the most? Which substance’s temperature will increase the least?
A 50.0-g sample of water at 100. °C is poured into a 50.0-g sample of water at 25 °C. What will be the final temperature of the water?
A 25.0-g sample of pure iron at 85 °C is dropped into 75 g of water at 20. °C. What is the final temperature of the water–iron mixture?
If it takes 4.5 J of energy to warm 5.0 g of aluminum from 25 °C to a certain higher temperature, then it will take _________ J to warm 10. g of aluminum over the same temperature interval.
For each of the substances listed in Table 10.1, calculate the quantity of heat required to heat 150. g of the substance by 11.2 °C.
Suppose you have 10.0-g samples of each of the substances listed in Table 10.1 and that 1.00 kJ of heat is applied to each of these samples. By what amount would the temperature of each sample be raised?
Calculate $\Delta E$ for each of the following.a. $q=-47 \mathrm{kJ}, w=+88 \mathrm{kJ}$b. $q=+82 \mathrm{kJ}, w=+47 \mathrm{kJ}$c. $q=+47 \mathrm{kJ}, w=0$d. In which of these cases do the surroundings do work on the system?
Are the following processes exothermic or endothermic?a. the combustion of gasoline in a car engineb. water condensing on a cold pipec. $\mathrm{CO}_{2}(s) \rightarrow \mathrm{CO}_{2}(g)$d. $\mathrm{F}_{2}(g) \rightarrow 2 \mathrm{F}(g)$
The overall reaction in commercial heat packs can be represented as$$4 \mathrm{Fe}(s)+3 \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s) \Delta H=-1652 \mathrm{kJ}$$a. How much heat is released when 4.00 mol iron is reacted with excess $\mathrm{O}_{2} ?$b. How much heat is released when 1.00 mol $\mathrm{Fe}_{2} \mathrm{O}_{3}$ is produced?c. How much heat is released when 1.00 $\mathrm{g}$ iron is reacted with excess $\mathrm{O}_{2} ?$d. How much heat is released when 10.0 $\mathrm{g}$ Fe and 2.00 $\mathrm{g} \mathrm{O}_{2}$ are reacted?
Consider the following equations:$$\begin{array}{ll}{3 \mathrm{A}+6 \mathrm{B} \rightarrow+3 \mathrm{D}} & {\Delta H=-403 \mathrm{kJ} / \mathrm{mol}} \\ {\mathrm{E}+2 \mathrm{F} \rightarrow \mathrm{A}} & {\Delta H=-105.2 \mathrm{kJ} / \mathrm{mol}} \\ {\mathrm{C} \rightarrow \mathrm{E}+3 \mathrm{D}} & {\Delta H=+64.8 \mathrm{kJ} / \mathrm{mol}}\end{array}$$Suppose the first equation is reversed and multiplied by $\frac{1}{6},$ the second and third equations are divided by 2, and the three adjusted equations are added. What is the net reaction and what is the overall heat of this reaction?
It has been determined that the body can generate 5500 kJ of energy during one hour of strenuous exercise. Perspiration is the body’s mechanism for eliminating this heat. How many grams and how many liters of water would have to be evaporated through perspiration to rid the body of the heat generated during two hours of exercise? (The heat of vaporization of water is 40.6 kJ/mol.)
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 miles 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 g in 1 lb.
You are given a metal and asked to determine its identity. You are to do this by determining the specific heat of the metal. You place the metal in a boiling water bath for a few minutes and then transfer the metal to a 100.0 g sample of water at a measured temperature. You then record the highest temperature of the water. The data are given in the table below.$$\begin{array}{|l|l|l|}\hline & {\text { Data }} \\ \hline \text { Mass water } & {100.0 g} \\ \hline \text { Initial temperature of the water } & {21.31^{\circ} \mathrm{C}} \\ \hline \text { Final temperature of the water } & {24.80^{\circ} \mathrm{C}} \\ \hline \text { Mass metal } & {50.0 {g}} \\ \hline\end{array}$$a. Given that the specific heat capacity of water is 4.184 J/g °C, calculate the amount of heat that is absorbed by the water.b. How does the heat absorbed by the water compare to the heat lost by the metal?c. What is the initial temperature of the “hot” metal? What is the final temperature of the metal?d. Calculate the specific heat capacity of the metal.e. Given the following specific heat capacities, determine the identity of your metal.Tin: 0.227 J/g °CZinc: 0.388 J/g °CAluminum: 0.891 J/g °C