A 2.20-g sample of phenol (C₆H₅OH) was burned in a bomb calorimeter whose total heat capacity is 11.90 kJ/°C. The temperature of the calorimeter plus contents increased from 21.50 to 27.50 °C. (b) What is the heat of combustion per mole of phenol?

Under constant-volume conditions, the heat of combustion of benzoic acid (C₆H₅O₆) is 15.57 kJ/g. A 3.500-g sample of sucrose is burned in a bomb calorimeter. The temperature of the calorimeter increases from 20.94 to 24.72 °C. (a) What is the total heat capacity of the calorimeter?

Under constant-volume conditions, the heat of combustion of benzoic acid (C₆H₅O₆) is 15.57 kJ/g. A 3.500-g sample of sucrose is burned in a bomb calorimeter. The temperature of the calorimeter increases from 20.94 to 24.72 °C. (b) If the size of the sucrose sample had been exactly twice as large, what would the temperature change of the calorimeter have been?

Under constant-volume conditions, the heat of combustion of naphthalene (C₁₀H₈) is 40.18 kJ/g. A 2.50-g sample of naphthalene is burned in a bomb calorimeter. The temperature of the calorimeter increases from 21.50 to 28.83 °C. (c) Suppose that in changing samples, a portion of the water in the calorimeter were lost. In what way, if any, would this change the heat capacity of the calorimeter?

Consider the following hypothetical reactions: A → B ΔH_I = +60 kJ B → C ΔH_II = -90 kJ (a) Use Hess’s law to calculate the enthalpy change for the reaction A → C.

Consider the following hypothetical reactions: A → B ΔH_I = +60 kJ B → C ΔH_II = -90 kJ (b) Construct an enthalpy diagram for substances A, B, and C, and show how Hess's law applies.

Calculate the enthalpy change for the reaction P₄O₆(s) + 2 O₂(g) → P₄O₁₀(s) given the following enthalpies of reaction: P₄(s) + 3 O₂(g) → P₄O₆(s) ΔH = -1640.1 kJ P₄(s) + 5 O₂(g) → P₄O₁₀(s) ΔH = -2940.1 kJ

From the enthalpies of reaction 2 C(s) + O₂(g) → 2 CO(g) ΔH = -221.0 kJ 2 C(s) + O₂(g) + 4 H₂(g) → 2 CH₃OH(g) ΔH = -402.4 kJ Calculate ΔH for the reaction CO(g) + 2 H₂(g) → CH₃OH(g)

From the enthalpies of reaction H₂(g) + F₂(g) → 2 HF(g) ΔH = -537 kJ C(s) + 2 F₂(g) → CF₄(g) ΔH = -680 kJ 2 C(s) + 2 H₂(g) → C₂H₄(g) ΔH = +52.3 kJ Calculate H for the reaction of ethylene with F₂: C₂H₄(g) + 6 F₂(g) → 2 CF₄(g) + 4 HF(g)

Given the data N₂(g) + O₂(g) → 2 NO(g) ΔH = +180.7 kJ 2 NO(g) + O₂(g) → 2 NO₂(g) ΔH = -113.1 kJ 2 N₂O(g) → 2 N₂(g) + O₂(g) ΔH = -163.2 kJ use Hess's law to calculate ΔH for the reaction N₂O(g) + NO₂(g) → 3 NO(g)

(c) What is meant by the term standard enthalpy of formation?

(a) Why does the standard enthalpy of formation of both the very reactive fluorine (F₂) and the almost inert gas nitrogen (N₂) both read zero?

For each of the following compounds, write a balanced thermochemical equation depicting the formation of one mole of the compound from its elements in their standard states and then look up ΔH°_f for each substance in Appendix C. (a) NO₂(g) (b) SO₃(g) (c) NaBr(s) (d) Pb(NO₃)₂(s).

Write balanced equations that describe the formation of the following compounds from elements in their standard states, and then look up the standard enthalpy of formation for each substance in Appendix C: (a) CH₃OH(l)

Acetylene (C₂H₂(g)) is used for welding because oxyacetylene is the hottest burning common fuel gas. Using standard enthalpies of formation, calculate the quantity of heat produced when 10 g of acetylene is completely combusted in air under standard conditions.

Using values from Appendix C, calculate the value of H for each of the following reactions: (a) CaO(s) + 2 HF(g) → CaF₂(s) + H₂O(g) (b) Fe₂O₃(s) + 3 C(s) → 2 Fe(s) + 3CO(g) (c) 2 CO(g) + 2 NO(g) → N₂(s) + 2 CO₂(g) (d) 4 NH₃(g) + 5 O₂(g) → 4 NO(g) + 6 H₂Og)

Complete combustion of 1 mol of acetone (C₃H₆O) liberates 1790 kJ: C₃H₆O(l) + 4 O₂(g) → 3 CO₂(g) + 3 H₂O(l) ΔH° = -1790 kJ Using this information together with the standard enthalpies of formation of O₂(g), CO₂(g), and H₂O(l) from Appendix C, calculate the standard enthalpy of formation of acetone.

Calcium carbide (CaC₂) reacts with water to form acetylene (C₂H₂) and Ca(OH)₂. From the following enthalpy of reaction data and data in Appendix C, calculate H°_f for CaC₂(s): CaC₂(s) + 2 H₂O(l) → Ca(OH₂)(s) + C₂H₂(g) ΔH° = -127.2 kJ

Gasoline is composed primarily of hydrocarbons, including many with eight carbon atoms, called octanes. One of the cleanest–burning octanes is a compound called 2,3,4- trimethylpentane, which has the following structural formula: The complete combustion of one mole of this compound to CO₂(g) and H₂O(g) leads to ΔH° = -5064.9 kJ. (b) By using the information in this problem and data in Table 5.3, calculate H°f for 2,3,4-trimethylpentane.

Diethyl ether, C₄H₁₀O(l), a flammable compound that was once used as a surgical anesthetic, has the structure The complete combustion of 1 mol of C₄H₁₀O(l) to CO₂(g) and H₂O(l) yields ΔH° = -2723.7 kJ. (a) Write a balanced equation for the combustion of 1 mol of C₄H₁₀O(l).

Ethanol (C₂H₅OH) is blended with gasoline as an automobile fuel. (b) Calculate the standard enthalpy change for the reaction, assuming H₂O(g) as a product.

Ethanol (C₂H₅OH) is blended with gasoline as an automobile fuel. (c) Calculate the heat produced per liter of ethanol by combustion of ethanol under constant pressure. Ethanol has a density of 0.789 g/mL.

Ethanol (C₂H₅OH) is blended with gasoline as an automobile fuel. (d) Calculate the mass of CO₂ produced per kJ of heat emitted.

Methanol (CH₃OH) is used as a fuel in race cars. (b) Calculate the standard enthalpy change for the reaction, assuming H₂O(g) as a product.

Methanol (CH₃OH) is used as a fuel in race cars. (c) Calculate the heat produced by combustion per liter of methanol. Methanol has a density of 0.791 g/mL.

Methanol (CH₃OH) is used as a fuel in race cars. (d) Calculate the mass of CO₂ produced per kJ of heat emitted.