Chapter 5

Energy: Some Basic Principles Define the terms system and surroundings. What does it mean to say that a system and its surroundings are in thermal equilibrium?

Identify whether the following processes are exothermic or endothermic. Is the sign on \(q_{\text {sys }}\) positive or negative? (a) combustion of methane (b) melting of ice (c) raising the temperature of water from \(25^{\circ} \mathrm{C}\) to \(100^{\circ} \mathrm{C}\) (d) heating \(\mathrm{CaCO}_{3}(\mathrm{s})\) to form \(\mathrm{CaO}(\mathrm{s})\) and \(\mathrm{CO}_{2}(\mathrm{g})\)

Identify whether the following processes are exothermic or endothermic. Is the sign on \(q_{\mathrm{sys}}\) positive or negative? (a) the reaction of \(\mathrm{Na}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{g})\) (b) cooling and condensing gaseous \(\mathrm{N}_{2}\) to form liquid \(\mathrm{N}_{2}\) (c) cooling a soft drink from \(25^{\circ} \mathrm{C}\) to \(0^{\circ} \mathrm{C}\) (d) heating \(\mathrm{HgO}\) (s) to form \(\mathrm{Hg}(\ell)\) and \(\mathrm{O}_{2}(\mathrm{g})\)

The specific heat capacity of benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) is \(1.74 \mathrm{J} / \mathrm{g} \cdot \mathrm{K} .\) What is its molar heat capacity (in \(\mathrm{J} / \mathrm{mol} \cdot \mathrm{K}) ?\)

The specific heat capacity of copper metal is \(0.385 \mathrm{J} / \mathrm{g} \cdot \mathrm{K} .\) How much energy is required to heat 168 g of copper from \(-12.2^{\circ} \mathrm{C}\) to \(+25.6^{\circ} \mathrm{C} ?\)

The initial temperature of a 344 -g sample of iron is \(18.2^{\circ} \mathrm{C}\). If the sample absorbs \(2.25 \mathrm{kJ}\) of energy as heat, what is its final temperature?

After absorbing \(1.850 \mathrm{kJ}\) of energy as heat, the temperature of a 0.500 -kg block of copper is \(37^{\circ} \mathrm{C} .\) What was its initial temperature?

One beaker contains 156 g of water at \(22^{\circ} \mathrm{C},\) and a second beaker contains \(85.2 \mathrm{g}\) of water at \(95^{\circ} \mathrm{C}\) The water in the two beakers is mixed. What is the final water temperature?

A \(182-\mathrm{g}\) sample of gold at some temperature was added to \(22.1 \mathrm{g}\) of water. The initial water temperature was \(25.0^{\circ} \mathrm{C},\) and the final temperature was \(27.5^{\circ} \mathrm{C} .\) If the specific heat capacity of gold is \(0.128 \mathrm{J} / \mathrm{g} \cdot \mathrm{K},\) what was the initial temperature of the gold sample?

When 108 g of water at a temperature of \(22.5^{\circ} \mathrm{C}\) is mixed with \(65.1 \mathrm{g}\) of water at an unknown temperature, the final temperature of the resulting mixture is \(47.9^{\circ} \mathrm{C}\). What was the initial temperature of the second sample of water?

A 237 -g piece of molybdenum, initially at \(100.0^{\circ} \mathrm{C},\) is dropped into \(244 \mathrm{g}\) of water at \(10.0^{\circ} \mathrm{C} .\) When the system comes to thermal equilibrium, the temperature is \(15.3^{\circ} \mathrm{C}\). What is the specific heat capacity of molybdenum?

The energy required to melt \(1.00 \mathrm{g}\) of ice at \(0^{\circ} \mathrm{C}\) is 333 J. If one ice cube has a mass of \(62.0 \mathrm{g}\) and a tray contains 16 ice cubes, what quantity of energy is required to melt a tray of ice cubes to form liquid water at \(0^{\circ} \mathrm{C} ?\)

How much energy is required to vaporize \(125 \mathrm{g}\) of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6,}\) at its boiling point, \(80.1^{\circ} \mathrm{C}\) ? (The heat of vaporization of benzene is \(30.8 \mathrm{kJ} / \mathrm{mol} .)\)

Chloromethane, \(\mathrm{CH}_{3} \mathrm{Cl},\) arises from microbial fermentation and is found throughout the environment. It is also produced industrially, is used in the manufacture of various chemicals, and has been used as a topical anesthetic. How much energy is required to convert \(92.5 \mathrm{g}\) of liquid to a vapor at its boiling point, \(-24.09^{\circ} \mathrm{C} ?\) (The heat of vaporization of \(\mathrm{CH}_{3} \mathrm{Cl}\) is \(21.40 \mathrm{kJ} / \mathrm{mol} .\) )

The freezing point of mercury is \(-38.8^{\circ} \mathrm{C} .\) What quantity of energy, in joules, is released to the surroundings if \(1.00 \mathrm{mL}\) of mercury is cooled from \(23.0^{\circ} \mathrm{C}\) to \(-38.8^{\circ} \mathrm{C}\) and then frozen to a solid? (The density of liquid mercury is \(13.6 \mathrm{g} / \mathrm{cm}^{3}\). Its specific heat capacity is 0.140 J/g \cdot K and its heat of fusion is \(11.4 \mathrm{J} / \mathrm{g} .\) )

What quantity of energy, in joules, is required to raise the temperature of \(454 \mathrm{g}\) of tin from room temperature, \(25.0^{\circ} \mathrm{C},\) to its melting point, \(231.9^{\circ} \mathrm{C},\) and then melt the tin at that temperature? (The specific heat capacity of tin is \(0.227 \mathrm{J} / \mathrm{g} \cdot \mathrm{K},\) and the heat of fusion of this metal is \(59.2 \mathrm{J} / \mathrm{g} .\) )

Ethanol, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH},\) boils at \(78.29^{\circ} \mathrm{C} .\) How much energy, in joules, is required to raise the temperature of \(1.00 \mathrm{kg}\) of ethanol from \(20.0^{\circ} \mathrm{C}\) to the boiling point and then to change the liquid to vapor at that temperature? (The specific heat capacity of liquid ethanol is \(2.44 \mathrm{J} / \mathrm{g} \cdot \mathrm{K},\) and its enthalpy of vaporization is \(855 \mathrm{J} / \mathrm{g} .\) )

A 25.0 -mL. sample of benzene at \(19.9^{\circ} \mathrm{C}\) was cooled to its melting point, \(5.5^{\circ} \mathrm{C},\) and then frozen. How much energy was given off as heat in this process? (The density of benzene is \(0.80 \mathrm{g} / \mathrm{mL},\) its specific heat capacity is \(1.74 \mathrm{J} / \mathrm{g} \cdot \mathrm{K}\) and its heat of fusion is \(127 \mathrm{J} / \mathrm{g} .\) )

Heat, Work, and Internal Energy As a gas cools, it is compressed from 2.50 L to 1.25 L under a constant pressure of \(1.01 \times 10^{5}\) Pa. Calculate the work (in J) required to compress the gas.

A balloon expands from 0.75 L. to 1.20 L.as it is heated under a constant pressure of \(1.01 \times 10^{5} \mathrm{Pa}\). Calculate the work (in J) done by the balloon on the environment.

A balloon does 324 J of work on the surroundings as it expands under a constant pressure of \(7.33 \times 10^{4} \mathrm{Pa}\). What is the change in volume (in L) of the balloon?

As the gas trapped in a cylinder with a movable piston cools, \(1.34 \mathrm{kJ}\) of work is done on the gas by the surroundings. If the gas is at a constant pressure of \(1.33 \times 10^{5} \mathrm{Pa}\), what is the change of volume (in L) of the gas?

When \(745 \mathrm{J}\) of energy in the form of heat is transferred from the environment to a gas, the expansion of the gas does 312 J of work on the environment. What is the change in internal energy of the gas?

The internal energy of a gas decreases by \(1.65 \mathrm{kJ}\) when it transfers \(1.87 \mathrm{kJ}\) of energy in the form of heat to the surroundings. (a) Calculate the work done by the gas on the surroundings. (b) Does the volume of gas increase or decrease?

A volume of 1.50 L of argon gas is confined in a cylinder with a movable piston under a constant pressure of \(1.22 \times 10^{5}\) Pa. When \(1.25 \mathrm{kJ}\) of energy in the form of heat is transferred from the surroundings to the gas, the internal energy of the gas increases by 1.11 kJ. What is the final volume of argon gas in the cylinder?

Nitrogen gas is confined in a cylinder with a movable piston under a constant pressure of \(9.95 \times 10^{4}\) Pa. When \(695 \mathrm{J}\) of energy in the form of heat is transferred from the gas to the surroundings, its volume decreases by 1.88 L. What is the change in internal energy of the gas?

Enthalpy Changes Nitrogen monoxide, a gas recently found to be involved in a wide range of biological processes, reacts with oxygen to give brown \(\mathrm{NO}_{2}\) gas. $$ \begin{aligned} 2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow & 2 \mathrm{NO}_{2}(\mathrm{g}) \\ \Delta_{r} H^{\circ} &=-114.1 \mathrm{kJ} / \mathrm{mol}-\mathrm{pxn} \end{aligned} $$ Is this reaction endothermic or exothermic? What is the enthalpy change if \(1.25 \mathrm{g}\) of \(\mathrm{NO}\) is converted completely to \(\mathrm{NO}_{2} ?\)

Calcium carbide, \(\mathrm{CaC}_{2}\), is manufactured by the reaction of CaO with carbon at a high temperature. (Calcium carbide is then used to make acetylene. \(\mathrm{CaO}(\mathrm{s})+3 \mathrm{C}(\mathrm{s}) \rightarrow \mathrm{CaC}_{2}(\mathrm{s})+\mathrm{CO}(\mathrm{g})\) $$ \Delta_{1} H^{\circ}=+464.8 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ Is this reaction endothermic or exothermic? What is the enthalpy change if \(10.0 \mathrm{g}\) of \(\mathrm{CaO}\) is allowed to react with an excess of carbon?

Isooctane \((2,2,4-\) trimethylpentane), one of the many hydrocarbons that make up gasoline, burns in air to give water and carbon dioxide. $$ 2 \mathrm{C}_{8} \mathrm{H}_{18}(\ell)+25 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \begin{array}{l} 16 \mathrm{CO}_{2}(\mathrm{g})+18 \mathrm{H}_{2} \mathrm{O}(\ell) \\ \Delta_{i} H^{\circ}=-10,922 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} \end{array} $$ What is the enthalpy change if you burn 1.00 L of isooctane \((d=0.69 \mathrm{g} / \mathrm{mL}) ?\)

Calorimetry Assume you mix 100.0 \(\mathrm{mL}\) of \(0.200 \mathrm{M}\) CsOH with \(50.0 \mathrm{mL}\) of \(0.400 \mathrm{M} \mathrm{HCl}\) in a coffee-cup calorimeter. The following reaction occurs: $$ \mathrm{CsOH}(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{CsCl}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell) $$ The temperature of both solutions before mixing was \(22.50^{\circ} \mathrm{C},\) and it rises to \(24.28^{\circ} \mathrm{C}\) after the acid-base reaction. What is the enthalpy change for the reaction per mole of CsOH? Assume the densities of the solutions are all \(1.00 \mathrm{g} / \mathrm{mL}\) and the specific heat capacities of the solutions are \(4.2 \mathrm{J} / \mathrm{g} \cdot \mathrm{K}.\)

You mix \(125 \mathrm{mL}\) of \(0.250 \mathrm{M}\) CsOH with \(50.0 \mathrm{mL}\) of 0.625 M HF in a coffee-cup calorimeter, and the temperature of both solutions rises from \(21.50^{\circ} \mathrm{C}\) before mixing to \(24.40^{\circ} \mathrm{C}\) after the reaction. $$ \mathrm{CsOH}(\mathrm{aq})+\mathrm{HF}(\mathrm{aq}) \rightarrow \mathrm{CsF}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell) $$ What is the enthalpy of reaction per mole of CsOH? Assume the densities of the solutions are all \(1.00 \mathrm{g} / \mathrm{mL},\) and the specific heat capacities of the solutions are \(4.2 \mathrm{J} / \mathrm{g} \cdot \mathrm{K}\)

A piece of titanium metal with a mass of \(20.8 \mathrm{g}\) is heated in boiling water to \(99.5^{\circ} \mathrm{C}\) and then dropped into a coffee-cup calorimeter containing \(75.0 \mathrm{g}\) of water at \(21.7^{\circ} \mathrm{C}\). When thermal equilibrium is reached, the final temperature is \(24.3^{\circ} \mathrm{C}\) Calculate the specific heat capacity of titanium.

A piece of chromium metal with a mass of \(24.26 \mathrm{g}\) is heated in boiling water to \(98.3^{\circ} \mathrm{C}\) and then dropped into a coffee-cup calorimeter containing 82.3 g of water at \(23.3^{\circ}\) C. When thermal equilibrium is reached, the final temperature is \(25.6^{\circ} \mathrm{C} .\) Calculate the specific heat capacity of chromium.

You should use care when dissolving \(\mathrm{H}_{2} \mathrm{SO}_{4}\) in water because the process is highly exothermic. To measure the enthalpy change, \(5.2 \mathrm{g}\) of concentrated \(\mathrm{H}_{2} \mathrm{SO}_{4}(\ell)\) was added (with stirring) to 135 g of water in a coffee-cup calorimeter. This resulted in an increase in temperature from \(20.2^{\circ} \mathrm{C}\) to \(28.8^{\circ} \mathrm{C} .\) Calculate the enthalpy change for the process \(\mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}),\) in kJ/mol.

Suppose you burned 0.300 g of \(\mathrm{C}(\mathrm{s})\) in an excess of \(\mathbf{O}_{2}(g)\) in a constant-volume calorimeter to give \(\mathrm{CO}_{2}(\mathrm{g})\) $$ \mathrm{C}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{CO}_{2}(\mathrm{g}) $$ The temperature of the calorimeter, which contained 775 g of water, increased from \(25.00^{\circ} \mathrm{C}\) to \(27.38^{\circ} \mathrm{C} .\) The heat capacity of the bomb is \(893 \mathrm{J} / \mathrm{K} .\) Calculate \(\Delta U\) per mole of carbon.

A \(0.692-\mathrm{g}\) sample of glucose, \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6},\) was burned in a constant-volume calorimeter. The temperature rose from \(21.70^{\circ} \mathrm{C}\) to \(25.22^{\circ} \mathrm{C}\). The calorimeter contained \(575 \mathrm{g}\) of water, and the bomb had a heat capacity of \(650 \mathrm{J} / \mathrm{K}\). What is \(\Delta U\) per mole of glucose?

An "ice calorimeter" can be used to determine the specific heat capacity of a metal. A piece of hot metal is dropped onto a weighed quantity of ice. The energy transferred from the metal to the ice can be determined from the amount of ice melted. Suppose you heated a 50.0 -g piece of silver to \(99.8^{\circ} \mathrm{C}\) and then dropped it onto ice. When the metal's temperature had dropped to \(0.0^{\circ} \mathrm{C},\) it is found that \(3.54 \mathrm{g}\) of ice had melted. What is the specific heat capacity of silver?

Hess's Law The enthalpy changes for the following reactions can be measured: \(\mathrm{CH}_{4}(\mathrm{g})+2 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) $$ \Delta_{r} H^{\circ}=-802.4 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ \(\mathrm{CH}_{3} \mathrm{OH}(\mathrm{g})+3 / 2 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) \(\Delta_{1} H^{\circ}=-676 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn}\) (a) Use these values and Hess's law to determine the enthalpy change for the reaction \(\mathrm{CH}_{4}(\mathrm{g})+1 / 2 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{CH}_{3} \mathrm{OH}(\mathrm{g})\) (b) Draw an energy level diagram that shows the relationship between the energy quantities involved in this problem.

The enthalpy changes of the following reactions can be measured: \(\mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) $$ \Delta_{i} H^{\circ}=-1411.1 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(\ell)+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CO}_{2}(\mathrm{g})+3 \mathrm{H}_{2} \mathrm{O}(\ell)\) \(\Delta, H^{\circ}=-1367.5 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn}\) (a) Use these values and Hess's law to determine the enthalpy change for the reaction \(\mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(\ell)\) (b) Draw an energy level diagram that shows the relationship between the energy quantities involved in this problem.

Enthalpy changes for the following reactions can be determined experimentally: \(\mathrm{N}_{2}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{NH}_{3}(\mathrm{g})\) $$ \Delta_{r} H^{\circ}=-91.8 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ \(4 \mathrm{NH}_{3}(\mathrm{g})+5 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{NO}(\mathrm{g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) $$ \Delta_{\mathrm{r}} H^{\circ}=-906.2 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ \(\mathrm{H}_{2}(\mathrm{g})+1 / 2 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) $$ \Delta_{i} H^{\circ}=-241.8 \mathrm{kJ} / \mathrm{mol}-\mathrm{v} \mathrm{xn} $$ Use these values to determine the enthalpy change for the formation of \(\mathrm{NO}(\mathrm{g})\) from the elements (an enthalpy change that cannot be measured directly because the reaction is reactant-favored). $$ ^{1 / 2} \mathrm{N}_{2}(\mathrm{g})+1 / 2 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{NO}(\mathrm{g}) \quad \Delta_{r} H^{\circ}=? $$

You wish to know the enthalpy change for the formation of liquid \(\mathrm{PCl}_{3}\) from the elements. $$ \mathrm{P}_{4}(\mathrm{s})+6 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{PCl}_{3}(\ell) \quad \Delta_{r} H^{\circ}=? $$ The enthalpy change for the formation of \(\mathrm{PCl}_{5}\) from the elements can be determined experimentally, as can the enthalpy change for the reaction of \(\mathrm{PCl}_{3}(\ell)\) with more chlorine to give \(\mathrm{PCl}_{5}(\mathrm{s}):\) \(\mathrm{P}_{4}(\mathrm{s})+10 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{PCl}_{5}(\mathrm{s})\) \(\Delta_{i} H^{\circ}=-1774.0 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn}\) \(\mathrm{PCl}_{3}(\ell)+\mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{PCl}_{5}(\mathrm{s})\) $$ \Delta_{r} H^{\circ}=-123.8 \mathrm{kJ} / \mathrm{mol}-\mathrm{Dxn} $$ Use these data to calculate the enthalpy change for the formation of 1.00 mol of \(\mathrm{PCl}_{3}(\ell)\) from phosphorus and chlorine.

Write a balanced chemical equation for the formation of \(\mathrm{CaCO}_{3}(\mathrm{s})\) from the elements in their standard states. Find the value for \(\Delta_{f} H^{\circ}\) for \(\mathrm{CaCO}_{3}(\mathrm{s})\) in Appendix L.

The standard enthalpy of formation of solid barium oxide, \(\mathrm{BaO},\) is \(-553.5 \mathrm{kJ} / \mathrm{mol},\) and the standard enthalpy of formation of barium peroxide, \(\mathrm{BaO}_{2},\) is \(-634.3 \mathrm{kJ} / \mathrm{mol}\) (a) Calculate the standard enthalpy change for the following reaction. Is the reaction exothermic or endothermic? \(2 \mathrm{BaO}_{2}(\mathrm{s}) \rightarrow 2 \mathrm{BaO}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g})\) (b) Draw an energy level diagram that shows the relationship between the enthalpy change of the decomposition of \(\mathrm{BaO}_{2}\) to \(\mathrm{BaO}\) and \(\mathrm{O}_{2}\) and the enthalpies of formation of \(\mathrm{BaO}(\mathrm{s})\) and \(\mathrm{BaO}_{2}(\mathrm{s})\)

The enthalpy change for the oxidation of naphthalene, \(\mathrm{C}_{10} \mathrm{H}_{8},\) is measured by calorimetry. $$ \begin{aligned} \mathrm{C}_{10} \mathrm{H}_{g}(\mathrm{s})+12 \mathrm{O}_{2}(\mathrm{g}) \rightarrow & 10 \mathrm{CO}_{2}(\mathrm{g})+4 \mathrm{H}_{2} \mathrm{O}(\ell) \\\ \Delta_{i} H^{\circ} &=-5156.1 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} \end{aligned} $$ Use this value, along with the standard enthalpies of formation of \(\mathrm{CO}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2} \mathrm{O}(\ell),\) to calculate the enthalpy of formation of naphthalene, in kJ/mol.

The enthalpy change for the oxidation of styrene, \(\mathrm{C}_{8} \mathrm{H}_{8},\) is measured by calorimetry. \(\mathrm{C}_{\mathrm{s}} \mathrm{H}_{\mathrm{s}}(\ell)+10 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 8 \mathrm{CO}_{2}(\mathrm{g})+4 \mathrm{H}_{2} \mathrm{O}(\ell)\) $$ \Delta_{r} H^{\circ}=-4395.0 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} $$ Use this value, along with the standard enthalpies of formation of \(\mathrm{CO}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2} \mathrm{O}(\ell),\) to calculate the enthalpy of formation of styrene, in \(\mathrm{kJ} / \mathrm{mol}\).

These questions are not designated as to type or location in the chapter. They may combine several concepts. The following terms are used extensively in thermodynamics. Define each and give an example. (a) exothermic and endothermic (b) system and surroundings (c) specific heat capacity (d) state function (e) standard state (f) enthalpy change, \(\Delta H\) (g) standard enthalpy of formation

For each of the following tell whether the process is exothermic or endothermic. (No calculations are required.) (a) \(\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{s})\) (b) \(2 \mathrm{H}_{2}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) (c) \(\mathrm{H}_{2} \mathrm{O}\left(\ell, 25^{\circ} \mathrm{C}\right) \rightarrow \mathrm{H}_{2} \mathrm{O}\left(\ell, 15^{\circ} \mathrm{C}\right)\) (d) \(\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\)

For each of the following, define a system and its surroundings, and give the direction of energy transfer between system and surroundings. (a) Methane burns in a gas furnace in your home. (b) Water drops, sitting on your skin after a swim, evaporate. (c) Water, at \(25^{\circ} \mathrm{C},\) is placed in the freezing compartment of a refrigerator, where it cools and eventually solidifies. (d) Aluminum and \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})\) are mixed in a flask sitting on a laboratory bench. A reaction occurs, and a large quantity of energy is evolved as heat.

What does the term standard state mean? What are the standard states of the following substances at \(298 \mathrm{K}: \mathrm{H}_{2} \mathrm{O}, \mathrm{NaCl}, \mathrm{Hg}, \mathrm{CH}_{4} ?\)

You have a large balloon containing 1.0 mol of gaseous water vapor at \(80^{\circ} \mathrm{C}\). How will each step affect the internal energy of the system? (a) The temperature of the system is raised to \(90^{\circ} \mathrm{C}\) (b) The vapor is condensed to a liquid, at \(40^{\circ} \mathrm{C}.\)