Chapter 6

Hydrogen sulfide, \(\mathrm{H}_{2} \mathrm{~S}\), is a foul-smelling gas. It burns to form sulfur dioxide. $$ \begin{gathered} 2 \mathrm{H}_{2} \mathrm{~S}(g)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{SO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g) ; \\ \Delta H=-1036 \mathrm{~kJ} \end{gathered} $$ Calculate the enthalpy change to burn \(28.5 \mathrm{~g}\) of hydrogen sulfide.

Propane, \(\mathrm{C}_{3} \mathrm{H}_{8}\), is a common fuel gas. Use the following to calculate the grams of propane you would need to provide \(369 \mathrm{~kJ}\) of heat. $$ \begin{aligned} \mathrm{C}_{3} \mathrm{H}_{8}(g)+5 \mathrm{O}_{2}(g) & \longrightarrow 3 \mathrm{CO}_{2}(g)+4 \mathrm{H}_{2} \mathrm{O}(g) \\ \Delta H=-2043 \mathrm{~kJ} \end{aligned} $$

Ethanol, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\), is mixed with gasoline and sold as gasohol. Use the following to calculate the grams of ethanol needed to provide \(358 \mathrm{~kJ}\) of heat. $$ \begin{gathered} \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(l)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(g) ; \\ \Delta H=-1235 \mathrm{~kJ} \end{gathered} $$

You wish to heat water to make coffee. How much heat (in joules) must be used to raise the temperature of \(0.180 \mathrm{~kg}\) of tap water (enough for one cup of coffee) from \(19^{\circ} \mathrm{C}\) to \(96^{\circ} \mathrm{C}\) (near the ideal brewing temperature)? Assume the specific heat is that of pure water, \(4.18 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\).

An iron skillet weighing \(1.28 \mathrm{~kg}\) is heated on a stove to \(178^{\circ} \mathrm{C}\). Suppose the skillet is cooled to room temperature, \(21^{\circ} \mathrm{C}\). How much heat energy (in joules) must be removed to effect this cooling? The specific heat of iron is \(0.449 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\).

When steam condenses to liquid water, \(2.26 \mathrm{~kJ}\) of heat is released per gram. The heat from \(168 \mathrm{~g}\) of steam is used to heat a room containing \(6.44 \times 10^{4} \mathrm{~g}\) of air \((20 \mathrm{ft} \times 12 \mathrm{ft} \times 8 \mathrm{ft})\). The specific heat of air at normal pressure is \(1.015 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right) .\) What is the change in air temperature, assuming the heat from the steam is all absorbed by air?

When ice at \(0^{\circ} \mathrm{C}\) melts to liquid water at \(0^{\circ} \mathrm{C}\), it absorbs \(0.334 \mathrm{~kJ}\) of heat per gram. Suppose the heat needed to melt \(31.5 \mathrm{~g}\) of ice is absorbed from the water contained in a glass. If this water has a mass of \(0.210 \mathrm{~kg}\) and a temperature of \(21.0^{\circ} \mathrm{C}\), what is the final temperature of the water? (Note that you will also have \(31.5 \mathrm{~g}\) of water at \(0^{\circ} \mathrm{C}\) from the ice.)

When \(15.3 \mathrm{~g}\) of sodium nitrate, \(\mathrm{NaNO}_{3}\), was dissolved in water in a calorimeter, the temperature fell from \(25.00^{\circ} \mathrm{C}\) to \(21.56^{\circ} \mathrm{C}\). If the heat capacity of the solution and the calorimeter is \(1071 \mathrm{~J} /{ }^{\circ} \mathrm{C}\), what is the enthalpy change when \(1 \mathrm{~mol}\) of sodium nitrate dissolves in water? The solution process is $$ \mathrm{NaNO}_{3}(s) \longrightarrow \mathrm{Na}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) ; \Delta H=? $$

When \(23.6 \mathrm{~g}\) of calcium chloride, \(\mathrm{CaCl}_{2}\), was dissolved in water in a calorimeter, the temperature rose from \(25.0^{\circ} \mathrm{C}\) to \(38.7^{\circ} \mathrm{C}\). If the heat capacity of the solution and the calorimeter is \(1258 \mathrm{~J} /{ }^{\circ} \mathrm{C}\), what is the enthalpy change when \(1.20 \mathrm{~mol}\) of calcium chloride dissolves in water? The solution process is $$ \mathrm{CaCl}_{2}(s) \longrightarrow \mathrm{Ca}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q) $$

A sample of ethanol, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\), weighing \(2.84 \mathrm{~g}\) was burned in an excess of oxygen in a bomb calorimeter. The temperature of the calorimeter rose from \(25.00^{\circ} \mathrm{C}\) to \(33.73^{\circ} \mathrm{C}\). If the heat capacity of the calorimeter and contents was \(9.63 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\), what is the value of \(q\) for burning \(1 \mathrm{~mol}\) of ethanol at constant volume and \(25.00^{\circ} \mathrm{C}\) ? The reaction is $$ \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(l)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) $$

A sample of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}\), weighing \(3.51 \mathrm{~g}\) was burned in an excess of oxygen in a bomb calorimeter. The temperature of the calorimeter rose from \(25.00^{\circ} \mathrm{C}\) to \(37.18^{\circ} \mathrm{C}\). If the heat capacity of the calorimeter and contents was \(12.05 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\), what is the value of \(q\) for burning \(1.25 \mathrm{~mol}\) of benzene at constant volume and \(25.00^{\circ} \mathrm{C} ?\) The reaction is $$ \mathrm{C}_{6} \mathrm{H}_{6}(l)+\frac{15}{2} \mathrm{O}_{2}(g) \longrightarrow 6 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) $$

Hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), is a colorless liquid used as a rocket fuel. What is the enthalpy change for the process in which hydrazine is formed from its elements? $$ \mathrm{N}_{2}(g)+2 \mathrm{H}_{2}(g) \longrightarrow \mathrm{N}_{2} \mathrm{H}_{4}(l) $$ Use the following reactions and enthalpy changes: \(\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{N}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) ; \Delta H=-622.2 \mathrm{~kJ}\) \(\mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) ; \Delta H=-285.8 \mathrm{~kJ}\)

Hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\), is a colorless liquid whose solutions are used as a bleach and an antiseptic. \(\mathrm{H}_{2} \mathrm{O}_{2}\) can be prepared in a process whose overall change is $$ \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}_{2}(l) $$ Calculate the enthalpy change using the following data: $$ \begin{gathered} 2 \mathrm{H}_{2} \mathrm{O}_{2}(l) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g) ; \Delta H=-196.0 \mathrm{~kJ} \\ \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) ; \Delta H=-285.8 \mathrm{~kJ} \end{gathered} $$

Ammonia will burn in the presence of a platinum catalyst to produce nitric oxide, \(\mathrm{NO}\). $$ 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) $$ What is the heat of reaction at constant pressure? Use the following thermochemical equations: $$ \begin{gathered} \mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}(g) ; \Delta H=180.6 \mathrm{~kJ} \\ \mathrm{~N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g) ; \Delta H=-91.8 \mathrm{~kJ} \\ 2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(g) ; \Delta H=-483.7 \mathrm{~kJ} \end{gathered} $$

Hydrogen cyanide is a highly poisonous, volatile liquid. It can be prepared by the reaction $$ \mathrm{CH}_{4}(g)+\mathrm{NH}_{3}(g) \longrightarrow \mathrm{HCN}(g)+3 \mathrm{H}_{2}(g) $$ What is the heat of reaction at constant pressure? Use the following thermochemical equations: $$ \begin{gathered} \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g) ; \Delta H=-91.8 \mathrm{~kJ} \\ \mathrm{C}(\text { graphite })+2 \mathrm{H}_{2}(g) \longrightarrow \mathrm{CH}_{4}(g) ; \Delta H=-74.9 \mathrm{~kJ} \\ \mathrm{H}_{2}(g)+2 \mathrm{C}(\text { graphite })+\mathrm{N}_{2}(g) \longrightarrow 2 \mathrm{HCN}(g) \\ \Delta H=270.3 \mathrm{~kJ} \end{gathered} $$

Compounds with carbon-carbon double bonds, such as ethylene, \(\mathrm{C}_{2} \mathrm{H}_{4}\), add hydrogen in a reaction called hydrogenation. $$ \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g) $$ Calculate the enthalpy change for this reaction, using the following combustion data: $$ \begin{gathered} \mathrm{C}_{2} \mathrm{H}_{4}(g)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \\ \Delta H=-1411 \mathrm{~kJ} \end{gathered} $$ $$ \begin{gathered} \mathrm{C}_{2} \mathrm{H}_{6}(g)+\frac{7}{2} \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) \\ \Delta H=-1560 \mathrm{~kJ} \\ \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) ; \Delta H=-286 \mathrm{~kJ} \end{gathered} $$

Acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}\), is contained in vinegar. Suppose acetic acid was formed from its elements, according to the following equation: $$ 2 \mathrm{C} \text { (graphite) }+2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CH}_{3} \mathrm{COOH}(l) $$ Find the enthalpy change, \(\Delta H\), for this reaction, using the following data: $$ \begin{gathered} \mathrm{CH}_{3} \mathrm{COOH}(l)+2 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \\ \Delta H=-874 \mathrm{~kJ} \\ \mathrm{C}(\text { graphite })+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g) ; \Delta H=-394 \mathrm{~kJ} \\ \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) ; \Delta H=-286 \mathrm{~kJ} \end{gathered} $$

The cooling effect of alcohol on the skin is due to its evaporation. Calculate the heat of vaporization of ethanol (ethyl alcohol), \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\). $$ \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(l) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(g) ; \Delta H^{\circ}=? $$ The standard enthalpy of formation of \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(l)\) is \(-277.7\) \(\mathrm{kJ} / \mathrm{mol}\) and that of \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(g)\) is \(-235.4 \mathrm{~kJ} / \mathrm{mol}\).

The energy, \(E\), needed to move an object a distance \(d\) by applying a force \(F\) is \(E=F \times d\). What must be the SI unit of force if this equation is to be consistent with the SI unit of energy for \(E ?\)

The potential energy of an object in the gravitational field of the earth is \(E_{p}=m g h .\) What must be the SI unit of \(g\) if this equation is to be consistent with the SI unit of energy for \(E_{p} ?\)

Hydrogen is an ideal fuel in many respects; for example, the product of its combustion, water, is nonpolluting. The heat given off in burning hydrogen to gaseous water is \(5.16 \times 10^{4}\) Btu per pound. What is this heat energy in joules per gram? ( 1 Btu = 252 cal; see also Table \(1.4 .\) )

Niagara Falls has a height of \(167 \mathrm{ft}\) (American Falls). What is the potential energy in joules of \(1.00 \mathrm{lb}\) of water at the top of the falls if we take water at the bottom to have a potential energy of zero? What would be the speed of this water at the bottom of the falls if we neglect friction during the descent of the water?

Any object, be it a space satellite or a molecule, must attain an initial upward velocity of at least \(11.2 \mathrm{~km} / \mathrm{s}\) in order to escape the gravitational attraction of the earth. What would be the kinetic energy in joules of a satellite weighing \(2354 \mathrm{lb}\) that has the speed equal to this escape velocity of \(11.2 \mathrm{~km} / \mathrm{s}\) ?

When calcium carbonate, \(\mathrm{CaCO}_{3}\) (the major constituent of limestone and seashells), is heated, it decomposes to calcium oxide (quicklime). $$ \mathrm{CaCO}_{3}(s) \longrightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g) ; \Delta H=177.9 \mathrm{~kJ} $$ How much heat is required to decompose \(21.3 \mathrm{~g}\) of calcium carbonate?

Calcium oxide (quicklime) reacts with water to produce calcium hydroxide (slaked lime). $$ \mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(s) ; \Delta H=-65.2 \mathrm{~kJ} $$ The heat released by this reaction is sufficient to ignite paper. How much heat is released when \(24.5 \mathrm{~g}\) of calcium oxide reacts?

Formic acid, \(\mathrm{HCHO}_{2}\), was first discovered in ants ( formica is Latin for "ant"). In an experiment, \(5.48 \mathrm{~g}\) of formic acid was burned at constant pressure. $$ 2 \mathrm{HCHO}_{2}(l)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) $$ If \(30.3 \mathrm{~kJ}\) of heat evolved, what is \(\Delta H\) per mole of formic acid?

Acetic acid, \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\), is the sour constituent of vinegar (acetum is Latin for "vinegar"). In an experiment, \(3.58 \mathrm{~g}\) of acetic acid was burned. $$ \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}(l)+2 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) $$ If \(52.0 \mathrm{~kJ}\) of heat evolved, what is \(\Delta H\) per mole of acetic acid?

Suppose you mix \(21.0 \mathrm{~g}\) of water at \(52.7^{\circ} \mathrm{C}\) with \(54.9 \mathrm{~g}\) of water at \(31.5^{\circ} \mathrm{C}\) in an insulated cup. What is the maximum temperature of the solution after mixing?

Suppose you mix \(20.5 \mathrm{~g}\) of water at \(66.2^{\circ} \mathrm{C}\) with \(45.4 \mathrm{~g}\) of water at \(35.7^{\circ} \mathrm{C}\) in an insulated cup. What is the maximum temperature of the solution after mixing?

A piece of lead of mass \(121.6 \mathrm{~g}\) was heated by an electrical coil. From the resistance of the coil, the current, and the time the current flowed, it was calculated that \(235 \mathrm{~J}\) of heat was added to the lead. The temperature of the lead rose from \(20.4^{\circ} \mathrm{C}\) to \(35.5^{\circ} \mathrm{C}\). What is the specific heat of the lead?

The specific heat of copper metal was determined by putting a piece of the metal weighing \(35.4 \mathrm{~g}\) in hot water. The quantity of heat absorbed by the metal was calculated to be \(47.0 \mathrm{~J}\) from the temperature drop of the water. What was the specific heat of the metal if the temperature of the metal rose \(3.45^{\circ} \mathrm{C}\) ?

A \(50.0\) -g sample of water at \(100.00^{\circ} \mathrm{C}\) was placed in an insulated cup. Then \(25.3 \mathrm{~g}\) of zinc metal at \(25.00^{\circ} \mathrm{C}\) was added to the water. The temperature of the water dropped to \(96.68^{\circ} \mathrm{C}\). What is the specific heat of zinc?

A 19.6-g sample of a metal was heated to \(61.67^{\circ} \mathrm{C}\). When the metal was placed into \(26.7 \mathrm{~g}\) of water in a calorimeter, the temperature of the water increased from \(25.00^{\circ} \mathrm{C}\) to \(30.00^{\circ} \mathrm{C}\). What is the specific heat of the metal?

A 14.1-mL sample of \(0.996 \mathrm{M} \mathrm{NaOH}\) is mixed with \(32.3 \mathrm{~mL}\) of \(0.905 M \mathrm{HCl}\) in a coffee-cup calorimeter (see Section \(6.6\) of your text for a description of a coffee-cup calorimeter). The enthalpy of the reaction, written with the lowest wholenumber coefficients, is \(-55.8 \mathrm{~kJ} .\) Both solutions are at \(21.6^{\circ} \mathrm{C}\) prior to mixing and reacting. What is the final temperature of the reaction mixture? When solving this problem, assume that no heat is lost from the calorimeter to the surroundings, the density of all solutions is \(1.00 \mathrm{~g} / \mathrm{mL}\), the specific heat of all solutions is the same as that of water, and volumes are additive.

A 29.1-mL sample of \(1.05 \mathrm{M}\) KOH is mixed with \(20.9 \mathrm{~mL}\) of \(1.07 M \mathrm{HBr}\) in a coffee-cup calorimeter (see Section \(6.6\) of your text for a description of a coffee-cup calorimeter). The enthalpy of the reaction, written with the lowest wholenumber coefficients, is \(-55.8 \mathrm{~kJ} .\) Both solutions are at \(21.8^{\circ} \mathrm{C}\) prior to mixing and reacting. What is the final temperature of the reaction mixture? When solving this problem, assume that no heat is lost from the calorimeter to the surroundings, the density of all solutions is \(1.00 \mathrm{~g} / \mathrm{mL}\), and volumes are additive.

In a calorimetric experiment, \(6.48 \mathrm{~g}\) of lithium hydroxide, LiOH, was dissolved in water. The temperature of the calorimeter rose from \(25.00^{\circ} \mathrm{C}\) to \(36.66^{\circ} \mathrm{C}\). What is \(\Delta H\) for the solution process? $$ \mathrm{LiOH}(s) \longrightarrow \mathrm{Li}^{+}(a q)+\mathrm{OH}^{-}(a q) $$ The heat capacity of the calorimeter and its contents is \(547 \mathrm{~J} /{ }^{\circ} \mathrm{C}\).

When \(21.45 \mathrm{~g}\) of potassium nitrate, \(\mathrm{KNO}_{3}\), was dissolved in water in a calorimeter, the temperature fell from \(25.00^{\circ} \mathrm{C}\) to \(14.14^{\circ} \mathrm{C}\). What is the \(\Delta H\) for the solution process? $$ \mathrm{KNO}_{3}(s) \longrightarrow \mathrm{K}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) $$ The heat capacity of the calorimeter and its contents is \(682 \mathrm{~J} /{ }^{\circ} \mathrm{C} .\)

A \(10.00-\mathrm{g}\) sample of acetic acid, \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\), was burned in a bomb calorimeter in an excess of oxygen. $$ \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}(l)+2 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) $$ The temperature of the calorimeter rose from \(25.00^{\circ} \mathrm{C}\) to \(35.84^{\circ} \mathrm{C}\). If the heat capacity of the calorimeter and its contents is \(13.43 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\), what is the enthalpy change for the reaction?

The sugar arabinose, \(\mathrm{C}_{5} \mathrm{H}_{10} \mathrm{O}_{5}\), is burned completely in oxygen in a calorimeter. $$ \mathrm{C}_{5} \mathrm{H}_{10} \mathrm{O}_{5}(s)+5 \mathrm{O}_{2}(g) \longrightarrow 5 \mathrm{CO}_{2}(g)+5 \mathrm{H}_{2} \mathrm{O}(l) $$ Burning a \(0.548-\mathrm{g}\) sample caused the temperature to rise from \(20.00^{\circ} \mathrm{C}\) to \(20.54^{\circ} \mathrm{C}\). The heat capacity of the calorimeter and its contents is \(15.8 \mathrm{~kJ} /{ }^{\circ} \mathrm{C} .\) Calculate \(\Delta H\) for the combustion reaction per mole of arabinose.

Hydrogen sulfide, \(\mathrm{H}_{2} \mathrm{~S}\), is a poisonous gas with the odor of rotten eggs. The reaction for the formation of \(\mathrm{H}_{2} \mathrm{~S}\) from the elements is $$ \mathrm{H}_{2}(g)+\frac{1}{8} \mathrm{~S}_{8}(\text { rhombic }) \longrightarrow \mathrm{H}_{2} \mathrm{~S}(g) $$ Use Hess's law to obtain the enthalpy change for this reaction from the following enthalpy changes: $$ \begin{gathered} \mathrm{H}_{2} \mathrm{~S}(g)+\frac{3}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{SO}_{2}(g) ; \Delta H=-518 \mathrm{~kJ} \\\ \mathrm{H}_{2}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g) ; \Delta H=-242 \mathrm{~kJ} \\ \frac{1}{8} \mathrm{~S}_{8}(\text { rhombic })+\mathrm{O}_{2}(g) \longrightarrow \mathrm{SO}_{2}(g) ; \Delta H=-297 \mathrm{~kJ} \end{gathered} $$

Ethylene glycol, \(\mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{OH}\), is used as antifreeze. It is produced from ethylene oxide, \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}\), by the reaction $$ \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{OH}(l) $$ Use Hess's law to obtain the enthalpy change for this reaction from the following enthalpy changes: $$ \begin{gathered} 2 \mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{CO}_{2}(g)+4 \mathrm{H}_{2} \mathrm{O}(l) \\ \Delta H=-2612.2 \mathrm{~kJ} \\ \mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{OH}(l)+\frac{5}{2} \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) ; \\ \Delta H=-1189.8 \mathrm{~kJ} \end{gathered} $$

Describe the physical characteristics of white phosphorus. Is it found in any modern matches? Why or why not?

What is the phosphorus compound used in "strike anywhere" matches. What is the chemical equation for the burning of this compound in air?

How fast (in meters per second) must an iron ball with a mass of \(56.6 \mathrm{~g}\) be traveling in order to have a kinetic energy of \(15.75 \mathrm{~J} ?\) The density of iron is \(7.87 \mathrm{~g} / \mathrm{cm}^{3} .\)

Sulfur dioxide gas reacts with oxygen, \(\mathrm{O}_{2}(\mathrm{~g})\), to produce \(\mathrm{SO}_{3}(g)\). This reaction releases \(99.0 \mathrm{~kJ}\) of heat (at constant pressure) for each mole of sulfur dioxide that reacts. Write the thermochemical equation for the reaction of 2 mol of sulfur dioxide, and then also for the decomposition of \(3 \mathrm{~mol}\) of sulfur trioxide gas into oxygen gas and sulfur dioxide gas. Do you need any other information to answer either question?

How many grams of oxygen gas are required to produce \(7.60 \mathrm{~kJ}\) of heat when hydrogen gas burns at constant pressure to produce gaseous water? $$ 2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(g) ; \Delta H=-484 \mathrm{~kJ} $$ Liquid water has a heat of vaporization of \(44.0 \mathrm{~kJ}\) per mole at \(25^{\circ} \mathrm{C}\).

A piece of iron was heated to \(95.4^{\circ} \mathrm{C}\) and dropped into a constant-pressure calorimeter containing \(284 \mathrm{~g}\) of water at \(32.2^{\circ} \mathrm{C}\). The final temperature of the water and iron was \(51.9^{\circ} \mathrm{C}\). Assuming that the calorimeter itself absorbs a negligible amount of heat, what was the mass (in grams) of the piece of iron? The specific heat of iron is \(0.449 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\), and the specific heat of water is \(4.18 \mathrm{~J} /\left(\mathrm{g} \cdot{ }^{\circ} \mathrm{C}\right)\).

The enthalpy of combustion, \(\Delta H\), for benzoic acid, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\), is \(-3226 \mathrm{~kJ} / \mathrm{mol}\). When a sample of benzoic acid was burned in a calorimeter (at constant pressure), the temperature of the calorimeter and contents rose from \(23.44^{\circ} \mathrm{C}\) to \(27.65^{\circ} \mathrm{C}\). The heat capacity of the calorimeter and contents was \(12.41 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\). What mass of benzoic acid was burned?

Given the following (hypothetical) thermochemical equations: $$ \begin{aligned} &\mathrm{A}+\mathrm{B} \longrightarrow 2 \mathrm{C} ; \Delta H=-447 \mathrm{~kJ} \\ &\mathrm{~A}+3 \mathrm{D} \longrightarrow 2 \mathrm{E} ; \Delta H=-484 \mathrm{~kJ} \\ &2 \mathrm{D}+\mathrm{B} \longrightarrow 2 \mathrm{~F} ; \Delta H=-429 \mathrm{~kJ} \end{aligned} $$ Calculate \(\Delta H\), in \(\mathrm{kJ}\), for the equation $$ 4 \mathrm{E}+5 \mathrm{~B} \longrightarrow 4 \mathrm{C}+6 \mathrm{~F} $$

What will be the final temperature of a mixture made from \(25.0 \mathrm{~g}\) of water at \(15.0^{\circ} \mathrm{C}\), from \(45.0 \mathrm{~g}\) of water at \(50.0^{\circ} \mathrm{C}\), and from \(15.0 \mathrm{~g}\) of water at \(37.0^{\circ} \mathrm{C}\) ?