A chemical system produces 155 kJ of heat and does 22 Kj

Problem 1DE In this chapter, we have learned about the photoelectric effect and its impact on the formulation of light as photons. We have also seen that some anomalous electron configurations of the elements are particularly favorable if each atom has one or more half-filled shell, such as the case for the Cr atom with its [Ar]4s13d5 electron configuration. Let’s suppose it is hypothesized that it requires more energy to remove an electron from a metal that has atoms with one or more half-filled shells than from those that do not. (a) Design a series of experiments involving the photoelectric effect that would test the hypothesis, (b) What experimental apparatus would be needed to test the hypothesis? It’s not necessary that you name actual equipment but rather that you imagine how the apparatus would work—think in terms of the types of measurements that would be needed, and what capability you would need in your apparatus, (c) Describe the type of data you would collect and how you would analyze the data to see whether the hypothesis were correct, (d) Could your experiments be extended to test the hypothesis for other parts of the periodic table, such as the lanthanide or actinide elements?

Problem 1E What is thermochemistry? Why is it important?

Problem 2E What is energy? What is work? List some examples of each.

Problem 4E What is the law of conservation of energy? How does it relate to energy exchanges between a thermodynamic system and its surroundings?

Problem 5E What is the SI unit of energy? List some other common units of energy.

Problem 8E What is a state function? List some examples of state functions.

Problem 6E What is the first law of thermodynamics? What are its implications?

Problem 9E What is internal energy? Is internal energy a state function?

Problem 10E If energy flows out of a chemical system and into the surroundings, what is the sign of ?Esystem?

Problem 11E If the internal energy of the products of a reaction is higher than the internal energy of the reactants, what is the sign of ?E for the reaction? In which direction does energy flow?

The propane fuel (C3H8) used in gas barbeques burns according to the thermochemical equation: If a pork roast must absorb 1.6 x 103 kJ to fully cook, and if only 10% of the heat produced by the barbeque is actually absorbed by the roast, what mass of CO2 is emitted into the atmosphere during the grilling of the pork roast?

Charcoal is primarily carbon. Determine the mass of CO2 produced by burning enough carbon (in the form of charcoal) to produce 5.00 x 102 kJ of heat.

Two substances, A and B, initially at different temperatures, come into contact and reach thermal equilibrium. The mass of substance A is 6.15 g and its initial temperature is 20.5 °C. The mass of substance B is 25.2 g and its initial temperature is 52.7 °C. The final temperature of both substances at thermal equilibrium is 46.7 °C. If the specific heat capacity of substance B is 1.17 J/g °C , what is the specific heat capacity of substance A?

When 0.514 g of biphenyl (C12H10) undergoes combustion in a bomb calorimeter, the temperature rises from 25.8 °C to 29.4 °C. Find rxn for the combustion of biphenyl in kJ/mol biphenyl. The heat capacity of the bomb calorimeter, determined in a separate experiment, is 5.86 kJ/°C.

Mothballs are composed primarily of the hydrocarbon naphthalene (C10H8). When 1.025 g of naphthalene burns in a bomb calorimeter, the temperature rises from 24.25 °C to 32.33 °C. Find rxn for the combustion of naphthalene. The heat capacity of the calorimeter, determined in a separate experiment, is 5.11 kJ/°C.

Zinc metal reacts with hydrochloric acid according to the balanced equation: When 0.103 g of Zn(s) is combined with enough HCl to make 50.0 mL of solution in a coffee-cup calorimeter, all of the zinc reacts, raising the temperature of the solution from 22.5 °C to 23.7 °C. Find rxn for this reaction as written. (Use 1.0 g/mL for the density of the solution and 4.18 J/g °C as the specific heat capacity.)

Problem 43E The gas in a piston (defined as the system) warms and absorbs 655 J of heat. The expansion performs 344 J of work on the surroundings. What is the change in internal energy for the system?

Problem 44E The air in an inflated balloon (defined as the system) warms over a toaster and absorbs 115 J of heat. As it expands, it does 77 kJ of work. What is the change in internal energy for the system?

Problem 45E We pack two identical coolers for a picnic, placing twenty-four 12-ounce soft drinks and 5 pounds of ice in each. However, the drinks that we put into cooler A were refrigerated for several hours before they were packed in the cooler, while the drinks that we put into cooler B were at room temperature. When we open the two coolers 3 hours later, most of the ice in cooler A is still present, while nearly all of the ice in cooler B has melted. Explain this difference.

A kilogram of aluminum metal and a kilogram of water are each warmed to 75 °C and placed in two identical insulated containers. One hour later, the two containers are opened and the temperature of each substance is measured. The aluminum has cooled to 35 °C while the water has cooled only to 66 °C. Explain this difference.

How much heat is required to warm 1.50 L of water from 25.0 °C to 100.0 °C? (Assume a density of 1.0 g/mL for the water.)

How much heat is required to warm 1.50 kg of sand from 25.0 °C to 100.0 °C?

Suppose that 25 g of each substance is initially at 27.0 °C. What is the final temperature of each substance upon absorbing 2.35 kJ of heat? a. gold b. silver c. aluminum d. water

An unknown mass of each substance, initially at 23.0 !C, absorbs 1.95 x 103 J of heat. The final temperature is recorded as indicated. Find the mass of each substance. a. Pyrex glass (Tf = 55.4 ?) b. sand (Tf = 62.1 ?) c. ethanol (Tf = 44.2 ?) d. water (Tf = 32.4 ?)

Problem 51E How much work (in J) is required to expand the volume of a pump from 0.0 L to 2.5 L against an external pressure of 1.1 atm?

Problem 52E The average human lung expands by about 0.50 L during each breath. If this expansion occurs against an external pressure of 1.0 atm, how much work (in J) is done during the expansion?

Problem 53E The air within a piston equipped with a cylinder absorbs 565 J of heat and expands from an initial volume of 0.10 L to a final volume of 0.85 L against an external pressure of 1.0 atm. What is the change in internal energy of the air within the piston?

Problem 54E A gas is compressed from an initial volume of 5.55 L to a final volume of 1.22 L by an external pressure of 1.00 atm. During the compression the gas releases 124 J of heat. What is the change in internal energy of the gas?

Problem 55E When 1 mol of a fuel burns at constant pressure, it produces 3452 kJ of heat and does 11 kJ of work. What are ?E and ?H for the combustion of the fuel?

Problem 56E The change in internal energy for the combustion of 1.0 mol of octane at a pressure of 1.0 atm is 5084.3 kJ. If the change in enthalpy is 5074.1 kJ, how much work is done during the combustion?

Problem 57E Determine whether each process is exothermic or endothermic and indicate the sign of ?H . a. natural gas burning on a stove b. isopropyl alcohol evaporating from skin c. water condensing from steam

Problem 58E Determine whether each process is exothermic or endothermic and indicate the sign of ?H . a. dry ice evaporating b. a sparkler burning c. the reaction that occurs in a chemical cold pack used to ice athletic injuries

Consider the thermochemical equation for the combustion of acetone (C3H6O), the main ingredient in nail polish remover. If a bottle of nail polish remover contains 177 mL of acetone, how much heat is released by its complete combustion? The density of acetone is 0.788 g/mL.

What mass of natural gas (CH4) must burn to emit 267 kJ of heat?

Nitromethane (CH3NO2) burns in air to produce significant amounts of heat. How much heat is produced by the complete reaction of 5.56 kg of nitromethane?

Titanium reacts with iodine to form titanium(III) iodide, emitting heat. Determine the masses of titanium and iodine that react if 1.55 x 103 kJ of heat is emitted by the reaction.

A silver block, initially at 58.5 °C, is submerged into 100.0 g of water at 24.8 °C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 26.2 °C. What is the mass of the silver block?

A 32.5 g iron rod, initially at 22.7 °C, is submerged into an unknown mass of water at 63.2 °C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 59.5 °C. What is the mass of the water?

A 31.1 g wafer of pure gold, initially at 69.3 °C, is submerged into 64.2 g of water at 27.8 °C in an insulated container. What is the final temperature of both substances at thermal equilibrium?

A 2.85 g lead weight, initially at 10.3 °C, is submerged in 7.55 g of water at 52.3 °C in an insulated container. What is the final temperature of both substances at thermal equilibrium?

A 2.74 g sample of a substance suspected of being pure gold is warmed to 72.1 °C and submerged into 15.2 g of water initially at 24.7 °C. The final temperature of the mixture is 26.3 °C. What is the heat capacity of the unknown substance? Could the substance be pure gold?

Problem 71E Exactly 1.5 g of a fuel burns under conditions of constant pressure and then again under conditions of constant volume. In measurement A the reaction produces 25.9 kJ of heat, and in measurement B the reaction produces 23.3 kJ of heat. Which measurement (A or B) corresponds to conditions of constant pressure? Which one corresponds to conditions of constant volume? Explain.

Problem 72E In order to obtain the largest possible amount of heat from a chemical reaction in which there is a large increase in the number of moles of gas, should you carry out the reaction under conditions of constant volume or constant pressure? Explain.

Problem 13E How is the change in internal energy of a system related to heat and work?

Problem 14E Explain how the sum of heat and work can be a state function, even though heat and work are themselves not state functions.

Problem 15E What is heat capacity? Explain the difference between heat capacity and specific heat capacity.

Problem 16E Explain how the high specific heat capacity of water can affect the weather in coastal regions.

Problem 17E If two objects, A and B, of different temperature come into direct contact, what is the relationship between the heat lost by one object and the heat gained by the other? What is the relationship between the temperature changes of the two objects? (Assume that the two objects do not lose any heat to anything else.)

Problem 18E What is pressure–volume work? How is it calculated?

Problem 19E What is calorimetry? Explain the difference between a coffeecup calorimeter and a bomb calorimeter. What is each designed to measure?

Problem 20E What is the change in enthalpy (?H) for a chemical reaction? How is ?H different from ?E?

Problem 21E Explain the difference between an exothermic and an endothermic reaction. Give the sign of ?H for each type of reaction.

Problem 22E From a molecular viewpoint, where does the energy emitted in an exothermic chemical reaction come from? Why does the reaction mixture undergo an increase in temperature even though energy is emitted?

Problem 23E From a molecular viewpoint, where does the energy absorbed in an endothermic chemical reaction go? Why does the reaction mixture undergo a decrease in temperature even though energy is absorbed?

Problem 24E Is the change in enthalpy for a reaction an extensive property? Explain the relationship between ?H for a reaction and the amounts of reactants and products that undergo reaction.

Problem 25E Explain how the value of ?H for a reaction changes upon each operation: a. multiplying the reaction by a factor. b. reversing the reaction. Why do these relationships hold?

Problem 26E What is Hess’s law? Why is it useful?

Problem 27E What is a standard state? What is the standard enthalpy change for a reaction?

Problem 28E What is the standard enthalpy of formation for a compound? For a pure element in its standard state?

How do you calculate from tabulated standard enthalpies of formation?

Problem 30E What are the main sources of the energy consumed in the United States?

Problem 32E Explain global climate change. What causes global warming? What is the evidence that global warming is occurring?

Problem 31E What are the main environmental problems associated with fossil fuel use?

Convert between energy units: a. 534 kWh to J b. 215 kJ to Cal c. 567 Cal to J d. 2.85 x 103 J to cal

Convert between energy units: a. 231 cal to kJ b. 132 x 104 kJ to kcal c. 4.99 x 103 kJ to kWh d. 2.88 x 104 J to Cal

Problem 35E Suppose that a person eats 2387 Calories per day. Convert this amount of energy into each unit: a. J b. kJ c. kWh

Problem 36E A particular frost-free refrigerator uses about 745 kWh of electrical energy per year. Express this amount of energy in each unit: a. J b. kJ c. Cal

Problem 37E Which statement is true of the internal energy of a system and its surroundings during an energy exchange with a negative ?Esys? a. The internal energy of the system increases and the internal energy of the surroundings decreases. b. The internal energy of both the system and the surroundings increases. c. The internal energy of both the system and the surroundings decreases. d. The internal energy of the system decreases and the internal energy of the surroundings increases.

Problem 38E During an energy exchange, a chemical system absorbs energy from its surroundings. What is the sign of ?Esys for this process? Explain.

Problem 40E Identify each energy exchange as primarily heat or work and determine whether the sign of ?Eis positive or negative for the system. a. A rolling billiard ball collides with another billiard ball. The first billiard ball (defined as the system) stops rolling after the collision. b. A book is dropped to the floor. (The book is the system). c. A father pushes his daughter on a swing. (The daughter and the swing are the system).

Problem 41E A system releases 622 kJ of heat and does 105 kJ of work on the surroundings. What is the change in internal energy of the system?

Problem 42E A system absorbs 196 kJ of heat and the surroundings do 117 kJ of work on the system. What is the change in internal energy of the system?

A gas expands in volume from 26.7 mL to 89.3 mL at constant temperature. Calculate the work done (in joules) if the gas expands (a) against a vacuum, (b) against a constant pressure of 1.5 atm, and (c) against a constant pressure of 2.8 atm.

A gas expands and does P-V work on the surroundings equal to 325 J. At the same time, it absorbs 127 J of heat from the surroundings. Calculate the change in energy of the gas

The work done to compress a gas is 74 J. As a result, 26 J of heat is given off to the surroundings. Calculate the change in energy of the gas

Calculate the work done when 50.0 g of tin dissolves in excess acid at 1.00 atm and 258C: Sn(s) 1 2H1(aq) Sn21(aq) 1 H2(g) Assume ideal gas behavior.

Calculate the work done in joules when 1.0 mole of water vaporizes at 1.0 atm and 1008C. Assume that the volume of liquid water is negligible compared with that of steam at 1008C, and ideal gas behavior

Define these terms: enthalpy, enthalpy of reaction. Under what condition is the heat of a reaction equal to the enthalpy change of the same reaction?

In writing thermochemical equations, why is it important to indicate the physical state (that is, gaseous, liquid, solid, or aqueous) of each substance?

Explain the meaning of this thermochemical equation: 4NH3(g) 1 5O2(g) 4NO(g) 1 6H2O(g) H 5 2904 kJ/mol

Consider this reaction: 2CH3OH(l) 1 3O2(g) 4H2O(l) 1 2CO2(g) H 5 21452.8 kJ/mol What is the value of DH if (a) the equation is multiplied throughout by 2, (b) the direction of the reaction is reversed so that the products become the reactants and vice versa, (c) water vapor instead of liquid water is formed as the product?

The first step in the industrial recovery of zinc from the zinc sulfide ore is roasting, that is, the conversion of ZnS to ZnO by heating: 2ZnS(s) 1 3O2(g) 2ZnO(s) 1 2SO2(g) H 5 2879 kJ/mol Calculate the heat evolved (in kJ) per gram of ZnS roasted.

Determine the amount of heat (in kJ) given off when 1.26 3 104 g of NO2 are produced according to the equation 2NO(g) 1 O2(g) 2NO2(g) H 5 2114.6 kJ/mol

Consider the reaction 2H2O(g) 2H2(g) 1 O2(g) H 5 483.6 kJ/mol If 2.0 moles of H2O(g) are converted to H2(g) and O2(g) against a pressure of 1.0 atm at 1258C, what is DU for this reaction?

Consider the reaction H2(g) 1 Cl2(g) 2HCl(g) H 5 2184.6 kJ/mol If 3 moles of H2 react with 3 moles of Cl2 to form HCl, calculate the work done (in joules) against a pressure of 1.0 atm at 258C. What is DU for this reaction? Assume the reaction goes to completion.

What is the difference between specific heat and heat capacity? What are the units for these two quantities? Which is the intensive property and which is the extensive property?

Define calorimetry and describe two commonly used calorimeters. In a calorimetric measurement, why is it important that we know the heat capacity of the calorimeter? How is this value determined?

Consider the following data: When these two metals are placed in contact, which of the following will take place? (a) Heat will flow from Al to Cu because Al has a larger specific heat. (b) Heat will flow from Cu to Al because Cu has a larger mass. (c) Heat will flow from Cu to Al because Cu has a larger heat capacity. (d) Heat will flow from Cu to Al because Cu is at a higher temperature. (e) No heat will flow in either direction

A piece of silver of mass 362 g has a heat capacity of 85.7 J/8C. What is the specific heat of silver?

A 6.22-kg piece of copper metal is heated from 20.58C to 324.38C. Calculate the heat absorbed (in kJ) by the metal

Calculate the amount of heat liberated (in kJ) from 366 g of mercury when it cools from 77.08C to 12.08C.

A sheet of gold weighing 10.0 g and at a temperature of 18.08C is placed flat on a sheet of iron weighing 20.0 g and at a temperature of 55.68C. What is the final temperature of the combined metals? Assume that no heat is lost to the surroundings. (Hint: The heat gained by the gold must be equal to the heat lost by the iron. The specific heats of the metals are given in Table 6.2.)

To a sample of water at 23.48C in a constant- pressure calorimeter of negligible heat capacity is added a 12.1-g piece of aluminum whose temperature is 81.78C. If the final temperature of water is 24.98C, calculate the mass of the water in the calorimeter. (Hint: See Table 6.2.)

A 0.1375-g sample of solid magnesium is burned in a constant-volume bomb calorimeter that has a heat capacity of 3024 J/8C. The temperature increases by 1.1268C. Calculate the heat given off by the burning Mg, in kJ/g and in kJ/mol.

A quantity of 85.0 mL of 0.900 M HCl is mixed with 85.0 mL of 0.900 M KOH in a constant-pressure calorimeter that has a heat capacity of 325 J/8C. If the initial temperatures of both solutions are the same at 18.248C, what is the final temperature of the mixed solution? The heat of neutralization is 256.2 kJ/mol. Assume the density and specific heat of the solutions are the same as those for water.

What is meant by the standard-state condition?

How are the standard enthalpies of an element and of a compound determined?

What is meant by the standard enthalpy of a reaction?

Write the equation for calculating the enthalpy of a reaction. Define all the terms.

State Hesss law. Explain, with one example, the usefulness of Hesss law in thermochemistry.

Describe how chemists use Hesss law to determine the DH8f of a compound by measuring its heat (enthalpy) of combustion.

Which of the following standard enthalpy of formation values is not zero at 258C? Na(s), Ne(g), CH4(g), S8(s), Hg(l), H(g).

The DH8f values of the two allotropes of oxygen, O2 and O3, are 0 and 142.2 kJ/mol, respectively, at 258C. Which is the more stable form at this temperature?

Which is the more negative quantity at 258C: DH8f for H2O(l) or DH8f for H2O(g)?

Predict the value of DH8f (greater than, less than, or equal to zero) for these elements at 258C: (a) Br2(g); Br2(l). (b) I2(g); I2(s).

In general, compounds with negative DH8f values are more stable than those with positive DH8f values. H2O2(l) has a negative DH8f (see Table 6.4). Why, then, does H2O2(l) have a tendency to decompose to H2O(l) and O2(g)?

Suggest ways (with appropriate equations) that would enable you to measure the DH8f values of Ag2O(s) and CaCl2(s) from their elements. No calculations are necessary.

Calculate the heat of decomposition for this process at constant pressure and 258C: CaCO3(s) CaO(s) 1 CO2(g) (Look up the standard enthalpy of formation of the reactant and products in Table 6.4.)

The standard enthalpies of formation of ions in aqueous solutions are obtained by arbitrarily assigning a value of zero to H1 ions; that is, Hf[H1 (aq)] 5 0. (a) For the following reaction HCl(g) H2O H1(aq) 1 Cl2(aq) H 5 274.9 kJ/mol calculate DH8f for the Cl2 ions. (b) Given that DH8f for OH2 ions is 2229.6 kJ/mol, calculate the enthalpy of neutralization when 1 mole of a strong monoprotic acid (such as HCl) is titrated by 1 mole of a strong base (such as KOH) at 258C.

Calculate the heats of combustion for the following reactions from the standard enthalpies of formation listed in Appendix 3: (a) 2H2(g) 1 O2(g) 2H2O(l) (b) 2C2H2(g) 1 5O2(g) 4CO2(g) 1 2H2O(l)

Calculate the heats of combustion for the following reactions from the standard enthalpies of formation listed in Appendix 3: (a) C2H4(g) 1 3O2(g) 2CO2(g) 1 2H2O(l) (b) 2H2S(g) 1 3O2(g) 2H2O(l) 1 2SO2(g)

Methanol, ethanol, and n-propanol are three common alcohols. When 1.00 g of each of these alcohols is burned in air, heat is liberated as shown by the following data: (a) methanol (CH3OH), 222.6 kJ; (b) ethanol (C2H5OH), 229.7 kJ; (c) n-propanol (C3H7OH), 233.4 kJ. Calculate the heats of combustion of these alcohols in kJ/mol.

The standard enthalpy change for the following reaction is 436.4 kJ/mol: H2(g) H(g) 1 H(g) Calculate the standard enthalpy of formation of atomic hydrogen (H).

From the standard enthalpies of formation, calculate DH8rxn for the reaction C6H12(l) 1 9O2(g) 6CO2(g) 1 6H2O(l) For C6H12(l), Hf 5 2151.9 kJ/mol

Pentaborane-9, B5H9, is a colorless, highly reactive liquid that will burst into flame when exposed to oxygen. The reaction is 2B5H9(l) 1 12O2(g) 5B2O3(s) 1 9H2O(l) Calculate the kilojoules of heat released per gram of the compound reacted with oxygen. The standard enthalpy of formation of B5H9 is 73.2 kJ/mol.

Determine the amount of heat (in kJ) given off when 1.26 3 104 g of ammonia are produced according to the equation N2(g) 1 3H2(g) 2NH3(g) Hrxn 5 292.6 kJ/mol Assume that the reaction takes place under standardstate conditions at 258C.

At 8508C, CaCO3 undergoes substantial decomposition to yield CaO and CO2. Assuming that the DH8f values of the reactant and products are the same at 8508C as they are at 258C, calculate the enthalpy change (in kJ) if 66.8 g of CO2 are produced in one reaction.

From these data, S(rhombic) 1 O2(g) SO2(g) Hrxn 5 2296.06 kJ/mol S(monoclinic) 1 O2(g) SO2(g) Hrxn 5 2296.36 kJ/mol calculate the enthalpy change for the transformation S(rhombic) S(monoclinic) (Monoclinic and rhombic are different allotropic forms of elemental sulfur.)

From the following data, C(graphite) 1 O2(g) CO2(g) Hrxn 5 2393.5 kJ/mol H2(g) 1 1 2O2(g) H2O(l) Hrxn 5 2285.8 kJ/mol 2C2H6(g) 1 7O2(g) 4CO2(g) 1 6H2O(l) Hrxn 5 23119.6 kJ/mol calculate the enthalpy change for the reaction 2C(graphite) 1 3H2(g) C2H6(g)

From the following heats of combustion, CH3OH(l) 1 3 2O2(g) CO2(g) 1 2H2O(l) Hrxn 5 2726.4 kJ/mol C(graphite) 1 O2(g) CO2(g) Hrxn 5 2393.5 kJ/mol H2(g) 1 1 2O2(g) H2O(l) Hrxn 5 2285.8 kJ/mol calculate the enthalpy of formation of methanol (CH3OH) from its elements: C(graphite) 1 2H2(g) 1 1 2O2(g) CH3OH(l)

Calculate the standard enthalpy change for the reaction 2Al(s) 1 Fe2O3(s) 2Fe(s) 1 Al2O3(s) given that 2Al(s) 1 3 2O2(g) Al2O3(s) Hrxn 5 21669.8 kJ/mol 2Fe(s) 1 3 2O2(g) Fe2O3(s) Hrxn 5 2822.2 kJ/mol

Define the following terms: enthalpy of solution, heat of hydration, lattice energy, heat of dilution.

Why is the lattice energy of a solid always a positive quantity? Why is the hydration of ions always a negative quantity?

Consider two ionic compounds A and B. A has a larger lattice energy than B. Which of the two compounds is more stable?

Mg21 is a smaller cation than Na1 and also carries more positive charge. Which of the two species has a larger hydration energy (in kJ/mol)? Explain.

Consider the dissolution of an ionic compound such as potassium fluoride in water. Break the process into the following steps: separation of the cations and anions in the vapor phase and the hydration of the ions in the aqueous medium. Discuss the energy changes associated with each step. How does the heat of solution of KF depend on the relative magnitudes of these two quantities? On what law is the relationship based?

Why is it dangerous to add water to a concentrated acid such as sulfuric acid in a dilution process?

Which of the following does not have DH8f 5 0 at 258C? He(g) Fe(s) Cl(g) S8(s) O2(g) Br2(l)

Calculate the expansion work done when 3.70 moles of ethanol are converted to vapor at its boiling point (78.38C) and 1.0 atm

The convention of arbitrarily assigning a zero enthalpy value for the most stable form of each element in the standard state at 258C is a convenient way of dealing with enthalpies of reactions. Explain why this convention cannot be applied to nuclear reactions.

Given the thermochemical equations: Br2(l) 1 F2(g) 2BrF(g) H 5 2188 kJ/mol Br2(l) 1 3F2(g) 2BrF3(g) H 5 2768 kJ/mol calculate the DH8rxn for the reaction BrF(g) 1 F2(g) BrF3(g)

The standard enthalpy change DH8 for the thermal decomposition of silver nitrate according to the following equation is 178.67 kJ: AgNO3(s) AgNO2(s) 1 1 2O2(g) The standard enthalpy of formation of AgNO3(s) is 2123.02 kJ/mol. Calculate the standard enthalpy of formation of AgNO2(s).

Hydrazine, N2H4, decomposes according to the following reaction: 3N2H4(l) 4NH3(g) 1 N2(g) (a) Given that the standard enthalpy of formation of hydrazine is 50.42 kJ/mol, calculate DH8 for its decomposition. (b) Both hydrazine and ammonia burn in oxygen to produce H2O(l) and N2(g). Write balanced equations for each of these processes and calculate DH8 for each of them. On a mass basis (per kg), would hydrazine or ammonia be the better fuel?

A quantity of 2.00 3 102 mL of 0.862 M HCl is mixed with an equal volume of 0.431 M Ba(OH)2 in a constant-pressure calorimeter of negligible heat capacity. The initial temperature of the HCl and Ba(OH)2 solutions is the same at 20.488C, For the process H1(aq) 1 OH2(aq) H2O(l) the heat of neutralization is 256.2 kJ/mol. What is the final temperature of the mixed solution?

A 3.53-g sample of ammonium nitrate (NH4NO3) was added to 80.0 mL of water in a constantpressure calorimeter of negligible heat capacity. As a result, the temperature of the water decreased from 21.68C to 18.18C. Calculate the heat of solution (DHsoln) of ammonium nitrate.

Consider the reaction N2(g) 1 3H2(g) 2NH3(g) Hrxn 5 292.6 kJ/mol If 2.0 moles of N2 react with 6.0 moles of H2 to form NH3, calculate the work done (in joules) against a pressure of 1.0 atm at 258C. What is DU for this reaction? Assume the reaction goes to completion

Calculate the heat released when 2.00 L of Cl2(g) with a density of 1.88 g/L react with an excess of sodium metal at 258C and 1 atm to form sodium chloride.

Venuss atmosphere is composed of 96.5 percent CO2, 3.5 percent N2, and 0.015 percent SO2 by volume. Its standard atmospheric pressure is 9.0 3 106 Pa. Calculate the partial pressures of the gases in pascals.

A 2.10-mole sample of crystalline acetic acid, initially at 17.08C, is allowed to melt at 17.08C and is then heated to 118.18C (its normal boiling point) at 1.00 atm. The sample is allowed to vaporize at 118.18C and is then rapidly quenched to 17.08C, so that it recrystallizes. Calculate DH8 for the total process as described.

Calculate the work done in joules by the reaction 2Na(s) 1 2H2O(l) 2NaOH(aq) 1 H2(g) when 0.34 g of Na reacts with water to form hydrogen gas at 08C and 1.0 atm.

You are given the following data: H2(g) 2H(g) H 5 436.4 kJ/mol Br2(g) 2Br(g) H 5 192.5 kJ/mol H2(g) 1 Br2(g) 2HBr(g) H 5 272.4 kJ/mol Calculate DH8 for the reaction H(g) 1 Br(g) HBr(g)

A gaseous mixture consists of 28.4 mole percent of hydrogen and 71.6 mole percent of methane. A 15.6-L gas sample, measured at 19.48C and 2.23 atm, is burned in air. Calculate the heat released.

When 2.740 g of Ba reacts with O2 at 298 K and 1 atm to form BaO, 11.14 kJ of heat are released. What is DH8f for BaO?

Methanol (CH3OH) is an organic solvent and is also used as a fuel in some automobile engines. From the following data, calculate the standard enthalpy of formation of methanol: 2CH3OH(l) 1 3O2(g) 2CO2(g) 1 4H2O(l) Hrxn 5 21452.8 kJ/mol

A 44.0-g sample of an unknown metal at 99.08C was placed in a constant-pressure calorimeter containing 80.0 g of water at 24.08C. The final temperature of the system was found to be 28.48C. Calculate the specific heat of the metal. (The heat capacity of the calorimeter is 12.4 J/8C.)

Using the data in Appendix 3, calculate the enthalpy change for the gaseous reaction shown here. (Hint: First determine the limiting reagent.) CO NO CO2 N2

Producer gas (carbon monoxide) is prepared by passing air over red-hot coke: C(s) 1 1 2O2(g) CO(g) Water gas (mixture of carbon monoxide and hydrogen) is prepared by passing steam over red- hot coke: C(s) 1 H2O(g) CO(g) 1 H2(g) For many years, both producer gas and water gas were used as fuels in industry and for domestic cooking. The large-scale preparation of these gases was carried out alternately, that is, first producer gas, then water gas, and so on. Using thermochemical reasoning, explain why this procedure was chosen.

Compare the heat produced by the complete combustion of 1 mole of methane (CH4) with a mole of water gas (0.50 mole H2 and 0.50 mole CO) under the same conditions. On the basis of your answer, would you prefer methane over water gas as a fuel? Can you suggest two other reasons why methane is preferable to water gas as a fuel?

The so-called hydrogen economy is based on hydrogen produced from water using solar energy. The gas may be burned as a fuel: 2H2(g) 1 O2(g) 2H2O(l) A primary advantage of hydrogen as a fuel is that it is nonpolluting. A major disadvantage is that it is a gas and therefore is harder to store than liquids or solids. Calculate the volume of hydrogen gas at 258C and 1.00 atm required to produce an amount of energy equivalent to that produced by the combustion of a gallon of octane (C8H18). The density of octane is 2.66 kg/gal, and its standard enthalpy of formation is 2249.9 kJ/mol.

Ethanol (C2H5OH) and gasoline (assumed to be all octane, C8H18) are both used as automobile fuel. If gasoline is selling for $4.50/gal, what would the price of ethanol have to be in order to provide the same amount of heat per dollar? The density and DH8f of octane are 0.7025 g/mL and 2249.9 kJ/mol and of ethanol are 0.7894 g/mL and 2277.0 kJ/mol, respectively. 1 gal 5 3.785 L.

The combustion of what volume of ethane (C2H6), measured at 23.08C and 752 mmHg, would be required to heat 855 g of water from 25.08C to 98.08C?

If energy is conserved, how can there be an energy crisis?

The heat of vaporization of a liquid (DHvap) is the energy required to vaporize 1.00 g of the liquid at its boiling point. In one experiment, 60.0 g of liquid nitrogen (boiling point 21968C) are poured into a Styrofoam cup containing 2.00 3 102 g of water at 55.38C. Calculate the molar heat of vaporization of liquid nitrogen if the final temperature of the water is 41.08C.

Explain the cooling effect experienced when ethanol is rubbed on your skin, given that C2H5OH(l) C2H5OH(g) H 5 42.2 kJ/mol

For which of the following reactions does DH8rxn 5 DH8f ? (a) H2(g) 1 S(rhombic) H2S(g) (b) C(diamond) 1 O2(g) CO2(g) (c) H2(g) 1 CuO(s) H2O(l) 1 Cu(s) (d) O(g) 1 O2(g) O3(g)

Calculate the work done (in joules) when 1.0 mole of water is frozen at 08C and 1.0 atm. The volumes of one mole of water and ice at 08C are 0.0180 L and 0.0196 L, respectively.

A quantity of 0.020 mole of a gas initially at 0.050 L and 208C undergoes a constant-temperature expansion until its volume is 0.50 L. Calculate the work done (in joules) by the gas if it expands (a) against a vacuum and (b) against a constant pressure of 0.20 atm. (c) If the gas in (b) is allowed to expand unchecked until its pressure is equal to the external pressure, what would its final volume be before it stopped expanding, and what would be the work done?

Calculate the standard enthalpy of formation for diamond, given that C(graphite) 1 O2(g) CO2(g) H 5 2393.5 kJ/mol C(diamond) 1 O2(g) CO2(g) H 5 2395.4 kJ/mol

(a) For most efficient use, refrigerator freezer compartments should be fully packed with food. What is the thermochemical basis for this recommendation? (b) Starting at the same temperature, tea and coffee remain hot longer in a thermal flask than chicken noodle soup. Explain.

Calculate the standard enthalpy change for the fermentation process. (See Problem 3.72.)

Portable hot packs are available for skiers and people engaged in other outdoor activities in a cold climate. The air-permeable paper packet contains a mixture of powdered iron, sodium chloride, and other components, all moistened by a little water. The exothermic reaction that produces the heat is a very common onethe rusting of iron: 4Fe(s) 1 3O2(g) 2Fe2O3(s) When the outside plastic envelope is removed, O2 molecules penetrate the paper, causing the reaction to begin. A typical packet contains 250 g of iron to warm your hands or feet for up to 4 hours. How much heat (in kJ) is produced by this reaction? (Hint: See Appendix 3 for DH8f values.)

A person ate 0.50 pound of cheese (an energy intake of 4000 kJ). Suppose that none of the energy was stored in his body. What mass (in grams) of water would he need to perspire in order to maintain his original temperature? (It takes 44.0 kJ to vaporize 1 mole of water.)

The total volume of the Pacific Ocean is estimated to be 7.2 3 108 km3 . A medium-sized atomic bomb produces 1.0 3 1015 J of energy upon explosion. Calculate the number of atomic bombs needed to release enough energy to raise the temperature of the water in the Pacific Ocean by 18C.

A 19.2-g quantity of dry ice (solid carbon dioxide) is allowed to sublime (evaporate) in an apparatus like the one shown in Figure 6.5. Calculate the expansion work done against a constant external pressure of 0.995 atm and at a constant temperature of 228C. Assume that the initial volume of dry ice is negligible and that CO2 behaves like an ideal gas.

The enthalpy of combustion of benzoic acid (C6H5COOH) is commonly used as the standard for calibrating constant-volume bomb calorimeters; its value has been accurately determined to be 23226.7 kJ/mol. When 1.9862 g of benzoic acid are burned in a calorimeter, the temperature rises from 21.848C to 25.678C. What is the heat capacity of the bomb? (Assume that the quantity of water surrounding the bomb is exactly 2000 g.)

The combustion of a 25.0-g gaseous mixture of H2 and CH4 releases 2354 kJ of heat. Calculate the amounts of the gases in grams.

Calcium oxide (CaO) is used to remove sulfur dioxide generated by coal-burning power stations: 2CaO(s) 1 2SO2(g) 1 O2(g) 2CaSO4(s) Calculate the enthalpy change for this process if 6.6 3 105 g of SO2 are removed by this process every day.

Glaubers salt, sodium sulfate decahydrate (Na2SO4 ? 10H2O), undergoes a phase transition (that is, melting or freezing) at a convenient temperature of about 328C: Na2SO4 ? 10H2O(s) Na2SO4 ? 10H2O(l) H 5 74.4 kJ/mol As a result, this compound is used to regulate the temperature in homes. It is placed in plastic bags in the ceiling of a room. During the day, the endothermic melting process absorbs heat from the surroundings, cooling the room. At night, it gives off heat as it freezes. Calculate the mass of Glaubers salt in kilograms needed to lower the temperature of air in a room by 8.28C at 1.0 atm. The dimensions of the room are 2.80 m 3 10.6 m 3 17.2 m, the specific heat of air is 1.2 J/g ? 8C, and the molar mass of air may be taken as 29.0 g/mol.

A balloon 16 m in diameter is inflated with helium at 188C. (a) Calculate the mass of He in the balloon, assuming ideal behavior. (b) Calculate the work done (in joules) during the inflation process if the atmospheric pressure is 98.7 kPa.

Acetylene (C2H2) can be hydrogenated (reacting with hydrogen) first to ethylene (C2H4) and then to ethane (C2H6). Starting with one mole of C2H2, label the diagram shown here analogous to Figure 6.10. Use the data in Appendix 3. Enthalpy

Calculate the DH8 for the reaction Fe31(aq) 1 3OH2(aq) Fe(OH)3(s)

An excess of zinc metal is added to 50.0 mL of a 0.100 M AgNO3 solution in a constant-pressure calorimeter like the one pictured in Figure 6.9. As a result of the reaction Zn(s) 1 2Ag1(aq) Zn21(aq) 1 2Ag(s) the temperature rises from 19.258C to 22.178C. If the heat capacity of the calorimeter is 98.6 J/8C, calculate the enthalpy change for the above reaction on a molar basis. Assume that the density and specific heat of the solution are the same as those for water, and ignore the specific heats of the metals

(a) A person drinks four glasses of cold water (3.08C) every day. The volume of each glass is 2.5 3 102 mL. How much heat (in kJ) does the body have to supply to raise the temperature of the water to 378C, the body temperature? (b) How much heat would your body lose if you were to ingest 8.0 3 102 g of snow at 08C to quench thirst? (The amount of heat necessary to melt snow is 6.01 kJ/mol.)

A drivers manual states that the stopping distance quadruples as the speed doubles; that is, if it takes 30 ft to stop a car moving at 25 mph then it would take 120 ft to stop a car moving at 50 mph. Justify this statement by using mechanics and the first law of thermodynamics. [Assume that when a car is stopped, its kinetic energy (1 2mu2 ) is totally converted to heat.]

At 258C, the standard enthalpy of formation of HF(aq) is given by 2320.1 kJ/mol; of OH2(aq), it is 2229.6 kJ/mol; of F2(aq), it is 2329.1 kJ/mol; and of H2O(l), it is 2285.8 kJ/mol. (a) Calculate the standard enthalpy of neutralization of HF(aq): HF(aq) 1 OH2(aq) F2(aq) 1 H2O(l) (b) Using the value of 256.2 kJ as the standard enthalpy change for the reaction H1 (aq) 1 OH2(aq) H2O(l) calculate the standard enthalpy change for the reaction HF(aq) H1 (aq) 1 F2(aq)

Why are cold, damp air and hot, humid air more uncomfortable than dry air at the same temperatures? (The specific heats of water vapor and air are approximately 1.9 J/g ? 8C and 1.0 J/g ? 8C, respectively.)

From the enthalpy of formation for CO2 and the following information, calculate the standard enthalpy of formation for carbon monoxide (CO). CO(g) 1 1 2O2(g) CO2(g) H 5 2283.0 kJ/mol Why cant we obtain it directly by measuring the enthalpy of the following reaction? C(graphite) 1 1 2O2(g) CO(g)

A 46-kg person drinks 500 g of milk, which has a caloric value of approximately 3.0 kJ/g. If only 17 percent of the energy in milk is converted to mechanical work, how high (in meters) can the person climb based on this energy intake? [Hint: The work done in ascending is given by mgh, where m is the mass (in kilograms), g the gravitational acceleration (9.8 m/s2 ), and h the height (in meters).]

The height of Niagara Falls on the American side is 51 m. (a) Calculate the potential energy of 1.0 g of water at the top of the falls relative to the ground level. (b) What is the speed of the falling water if all of the potential energy is converted to kinetic energy? (c) What would be the increase in temperature of the water if all the kinetic energy were converted to heat? (See Problem 6.121 for suggestions.)

In the nineteenth century two scientists named Dulong and Petit noticed that for a solid element, the product of its molar mass and its specific heat is approximately 25 J/8C. This observation, now called Dulong and Petits law, was used to estimate the specific heat of metals. Verify the law for the metals listed in Table 6.2. The law does not apply to one of the metals. Which one is it? Why?

Determine the standard enthalpy of formation of ethanol (C2H5OH) from its standard enthalpy of combustion (21367.4 kJ/mol).

Acetylene (C2H2) and benzene (C6H6) have the same empirical formula. In fact, benzene can be made from acetylene as follows: 3C2H2(g) C6H6(l) The enthalpies of combustion for C2H2 and C6H6 are 21299.4 kJ/mol and 23267.4 kJ/mol, respectively. Calculate the standard enthalpies of formation of C2H2 and C6H6 and hence the enthalpy change for the formation of C6H6 from C2H2

Ice at 08C is placed in a Styrofoam cup containing 361 g of a soft drink at 238C. The specific heat of the drink is about the same as that of water. Some ice remains after the ice and soft drink reach an equilibrium temperature of 08C. Determine the mass of ice that has melted. Ignore the heat capacity of the cup. (Hint: It takes 334 J to melt 1 g of ice at 08C.)

After a dinner party, the host performed the following trick. First, he blew out one of the burning candles. He then quickly brought a lighted match to about 1 in above the wick. To everyones surprise, the candle was relighted. Explain how the host was able to accomplish the task without touching the wick.

How much heat is required to decompose 89.7 g of NH4Cl? (Hint: You may use the enthalpy of formation values at 258C for the calculation.)

A gas company in Massachusetts charges $1.30 for 15 ft3 of natural gas (CH4) measured at 208C and 1.0 atm. Calculate the cost of heating 200 mL of water (enough to make a cup of coffee or tea) from 208C to 1008C. Assume that only 50 percent of the heat generated by the combustion is used to heat the water; the rest of the heat is lost to the surroundings.

Calculate the internal energy of a Goodyear blimp filled with helium gas at 1.2 3 105 Pa. The volume of the blimp is 5.5 3 103 m3 . If all the energy were used to heat 10.0 tons of copper at 218C, calculate the final temperature of the metal. (Hint: See Section 5.7 for help in calculating the internal energy of a gas. 1 ton 5 9.072 3 105 g.)

Decomposition reactions are usually endothermic, whereas combination reactions are usually exothermic. Give a qualitative explanation for these trends

Acetylene (C2H2) can be made by reacting calcium carbide (CaC2) with water. (a) Write an equation for the reaction. (b) What is the maximum amount of heat (in joules) that can be obtained from the combustion of acetylene, starting with 74.6 g of CaC2?

The average temperature in deserts is high during the day but quite cool at night, whereas that in regions along the coastline is more moderate. Explain

When 1.034 g of naphthalene (C10H8) are burned in a constant-volume bomb calorimeter at 298 K, 41.56 kJ of heat are evolved. Calculate DU and DH for the reaction on a molar basis

From a thermochemical point of view, explain why a carbon dioxide fire extinguisher or water should not be used on a magnesium fire.

Calculate the DU for the following reaction at 298 K: 2H2(g) 1 O2(g) 2H2O(l)

Lime is a term that includes calcium oxide (CaO, also called quicklime) and calcium hydroxide [Ca(OH)2, also called slaked lime]. It is used in the steel industry to remove acidic impurities, in airpollution control to remove acidic oxides such as SO2, and in water treatment. Quicklime is made industrially by heating limestone (CaCO3) above 20008C: CaCO3(s) CaO(s) 1 CO2(g) H 5 177.8 kJ/mol Slaked lime is produced by treating quicklime with water: CaO(s) 1 H2O(l) Ca(OH)2(s) H 5 265.2 kJ/mol The exothermic reaction of quicklime with water and the rather small specific heats of both quicklime (0.946 J/g ? 8C) and slaked lime (1.20 J/g ? 8C) make it hazardous to store and transport lime in vessels made of wood. Wooden sailing ships carrying lime would occasionally catch fire when water leaked into the hold. (a) If a 500-g sample of water reacts with an equimolar amount of CaO (both at an initial temperature of 258C), what is the final temperature of the product, Ca(OH)2? Assume that the product absorbs all of the heat released in the reaction. (b) Given that the standard enthalpies of formation of CaO and H2O are 2635.6 kJ/mol and 2285.8 kJ/mol, respectively, calculate the standard enthalpy of formation of Ca(OH)2.

A 4.117-g impure sample of glucose (C6H12O6) was burned in a constant-volume calorimeter having a heat capacity of 19.65 kJ/8C. If the rise in temperature is 3.1348C, calculate the percent by mass of the glucose in the sample. Assume that the impurities are unaffected by the combustion process. See Appendix 3 for thermodynamic data

Construct a table with the headings q, w, DU, and DH. For each of the following processes, deduce whether each of the quantities listed is positive (1), negative (2), or zero (0). (a) Freezing of benzene. (b) Compression of an ideal gas at constant temperature. (c) Reaction of sodium with water. (d) Boiling liquid ammonia. (e) Heating a gas at constant volume. (f) Melting of ice.

The combustion of 0.4196 g of a hydrocarbon releases 17.55 kJ of heat. The masses of the products are CO2 5 1.419 g and H2O 5 0.290 g. (a) What is the empirical formula of the compound? (b) If the approximate molar mass of the compound is 76 g, calculate its standard enthalpy of formation

Metabolic activity in the human body releases approximately 1.0 3 104 kJ of heat per day. Assuming the body is 50 kg of water, how much would the body temperature rise if it were an isolated system? How much water must the body eliminate as perspiration to maintain the normal body temperature (98.68F)? Comment on your results. The heat of vaporization of water may be taken as 2.41 kJ/g

Give an example for each of the following situations: (a) Adding heat to a system raises its temperature, (b) adding heat to a system does not change (raise) its temperature, and (c) a systems temperature is changed even though no heat is added or removed from it.

From the following data, calculate the heat of solution for KI: NaCl NaI KCl KI Lattice energy 788 686 699 632 (kJ/mol) Heat of solution 4.0 25.1 17.2 ?

Starting at A, an ideal gas undergoes a cyclic process involving expansion and compression, as shown here. Calculate the total work done. Does your result support the notion that work is not a state function? P (atm) 1 2 V (L) 2 1 B C A D

For reactions in condensed phases (liquids and solids), the difference between DH and DU is usually quite small. This statement holds for reactions carried out under atmospheric conditions. For certain geochemical processes, however, the external pressure may be so great that DH and DU can differ by a significant amount. A well-known example is the slow conversion of graphite to diamond under Earths surface. Calculate (DH 2 DU) for the conversion of 1 mole of graphite to 1 mole of diamond at a pressure of 50,000 atm. The densities of graphite and diamond are 2.25 g/cm3 and 3.52 g/cm3 , respectively.

The diagrams shown here represent various physical and chemical processes. (a) 2A(g) A2(g). (b) MX(s) M1(aq) 1 X2(aq). (c) AB(g) 1 C(g) AC(g) 1 B(g). (d) B(l) B(g). Predict whether the situations shown are endothermic or exothermic. Explain why in some cases no clear conclusions can be made. (a) (b) (c) (d)

A 20.3-g sample of an unknown metal and a 28.5-g sample of copper, both at 80.68C, are added to 100 g of water at 11.28C in a constant-pressure calorimeter of negligible heat capacity. If the final temperature of the metals and water is 13.78C, determine the specific heat of the unknown metal.

For most biological processes, DH < DU. Explain

Estimate the potential energy expended by an average adult male in going from the ground to the top floor of the Empire State Building using the staircase.

The fastest serve in tennis is about 150 mph. Can the kinetic energy of a tennis ball traveling at this speed be sufficient to heat 1 mL of water by 308C?

Can the total energy output of the sun in one second be sufficient to heat all of the ocean water on Earth to its boiling point?

It has been estimated that 3 trillion standard cubic feet of methane is released into the atmosphere every year. Capturing that methane would provide a source of energy, and it would also remove a potent greenhouse gas from the atmosphere (methane is 25 times more effective at trapping heat than an equal number of molecules of carbon dioxide). Standard cubic feet is measured at 608F and 1 atm. Determine the amount of energy that could be obtained by combustion of the methane that escapes each year

Biomass plants generate electricity from waste material such as wood chips. Some of these plants convert the feedstock to ethanol (C2H5OH) for later use as a fuel. (a) How many grams of ethanol can be produced from 1.0 ton of wood chips, if 85 percent of the carbon is converted to C2H5OH? (b) How much energy would be released by burning the ethanol obtained from 1.0 ton of wood chips? (Hint: Treat the wood chips as cellulose.)

Suppose an automobile carried hydrogen gas in its fuel tank instead of gasoline. At what pressure would the hydrogen gas need to be kept for the tank to contain an equivalent amount of chemical energy as a tank of gasoline?

A press release announcing a new fuel-cell car to the public stated that hydrogen is relatively cheap and some stations in California sell hydrogen for $5 a kilogram. A kg has the same energy as a gallon of gasoline, so its like paying $5 a gallon. But you go two to three times as far on the hydrogen. Analyze this claim.

We hear a lot about how the burning of hydrocarbons produces the greenhouse gas CO2, but what about the effect of increasing energy consumption on the amount of oxygen in the atmosphere required to sustain life. The figure shows past and projected energy world consumption. (a) How many moles of oxygen would be required to generate the additional energy expenditure for the next decade? (b) What would be the resulting decrease in atmospheric oxygen? 400 2005 World energy consumption (1015 kJ) Year 500 600 700 800 2010 2015 2020 2025 2030

Problem 1SAQ A chemical system produces 155 kJ of heat and does 22 Kj of work. What is ?E for the surroundings ? a) 177 kJ b) -177 kJ c) 133 kJ d) -133 kJ

Problem 2SAQ Which sample is most likely to undergo the smallest change in temperature upon the absorption of 100 kJ of heat? a) 15 g water b) 15 g lead c) 50 g water d) 50 g lead

How much heat must be absorbed by a 15.0 g sample of water to raise its temperature from 25.0 °C to 55.0 °C? (For water, Cs = 4.18 J/g °C.) a) 1.57 kJ b) 1.88 kJ c) 3.45 kJ d) 107 J

Problem 3E What is kinetic energy? What is potential energy? List some examples of each.

A 12.5 g sample of granite initially at 82.0 °C is immersed into 25.0 g of water initially at 22.0 °C. What is the final temperature of both substances when they reach thermal equilibrium? (For water, Cs = 4.18 J/g °C and for granite, Cs = 0.790 J/g °C.) a) 52.0 °C b) 1.55 x 103 °C c) 15.7 °C d) 27.2 °C

Problem 5SAQ A cylinder with a moving piston expands from an initial volume of 0.250 L against an external pressure of 2.00 atm. The expansion does 288 J of work on the surroundings. What is the fi nal volume of the cylinder? a) 1.42 L b) 1.17 L c) 144 L d) 1.67 L

When a 3.80 g sample of liquid octane (C8H18) is burned in a bomb calorimeter, the temperature of the calorimeter rises by 27.3 °C. The heat capacity of the calorimeter, measured in a separate experiment, is 6.18 kJ/ °C. Determine the enthalpy of combustion for octane in units of kJ/mol octane. a) -5.07 x 103 kJ/mol b) 5.07 x 103 kJ/mol c) -44.4 x 103 kJ/mol d) -16.7 x 103 kJ/mol

A friend claims to have constructed a machine that creates electricity but requires no energy input. Explain why you should be suspicious of your friend’s claim.

Hydrogen gas reacts with oxygen to form water. Determine the minimum mass of hydrogen gas required to produce 226 kJ of heat. a) 8.63 g b) 1.88 g c) 0.942 g d) 0.935 g

Manganese reacts with hydrochloric acid to produce manganese(II) chloride and hydrogen gas. \(\mathrm{Mn}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{MnCl}_2(a q)+\mathrm{H}_2(g)\) When \(0.625 \mathrm{~g} \mathrm{Mn}\) is combined with enough hydrochloric acid to make \(100.0 \mathrm{~mL}\) of solution in a coffee-cup calorimeter, all of the Mn reacts, raising the temperature of the solution from \(23.5^{\circ} \mathrm{C}\) to \(28.8^{\circ} \mathrm{C}\). Find \(\Delta H_{\mathrm{rxn}}\) for the reaction as written. (Assume that the specific heat capacity of the solution is \(4.18 \mathrm{~J} / \mathrm{g}{ }^{\circ} \mathrm{C}\) and the density is \(1.00 \mathrm{~g} / \mathrm{mL}\).) (a) \(-195 \mathrm{~kJ}\) (b) \(-3.54 \mathrm{~kJ}\) (c) \(-1.22 \mathrm{~kJ}\) (d) \(-2.21 \mathrm{~kJ}\)

Consider the reactions: What is for the reaction 2 B ? 3 C ? a) ?H1 + ?H2 b) ?H1 - ?H2 c) ?H2 - ?H1 d) 2 x (?H1 + ?H2)

Use standard enthalpies of formation to determine for the reaction: a) -541.2 kJ b) -2336 kJ c) -541.2 kJ d) -24.8 kJ

Problem 11SAQ Two substances, A and B, of equal mass but at different temperatures come into thermal contact. The specific heat capacity of substance A is twice the specific heat capacity of substance B. Which statement is true of the temperature of the two substances when they reach thermal equilibrium? (Assume no heat loss other than the thermal transfer between the substances.) a) The final temperature of both substances will be closer to the initial temperature of substance A than the initial temperature of substance B. b) The final temperature of both substances will be closer to the initial temperature of substance B than the initial temperature of substance A. c) The final temperature of both substances will be exactly midway between the initial temperatures of substance A and substance B. d) The final temperature of substance B will be greater than the final temperature of substance A.

Problem 13SAQ Which fuel is not a fossil fuel? a) coal b) hydrogen c) natural gas d) petroleum

Problem 12SAQ Which process is endothermic? a) The evaporation of water from the skin. b) The burning of candle wax. c) The oxidation of iron in a chemical hand warmer. d) The combustion of natural gas in a stove.

The standard enthalpy of formation for glucose [C6H12O6( s )] is -1273.3 kJ/mol. What is the correct formation equation corresponding to this ?

Natural gas burns in air to form carbon dioxide and water, releasing heat. What minimum mass of CH4 is required to heat 55 g of water by 25 °C? (Assume 100% heating efficiency.) a) 0.115 g b) 2.25 x 103 g c) 115 g d) 8.70 g

Identify each energy exchange as primarily heat or work and determine whether the sign of \(\Delta E\) is positive or negative for the system. a. Sweat evaporates from skin, cooling the skin. (The evaporating sweat is the system.) b. A balloon expands against an external pressure. (The contents of the balloon is the system.) c. An aqueous chemical reaction mixture is warmed with an external flame. (The reaction mixture is the system.)

Instant cold packs used to ice athletic injuries on the field contain ammonium nitrate and water separated by a thin plastic divider. When the divider is broken, the ammonium nitrate dissolves according to the endothermic reaction: In order to measure the enthalpy change for this reaction, 1.25 g of NH4NO3 is dissolved in enough water to make 25.0 mL of solution. The initial temperature is 25.8 °C and the final temperature (after the solid dissolves) is 21.9 °C. Calculate the change in enthalpy for the reaction in kJ. (Use 1.0 g/mL as the density of the solution and 4.18 J/g °C as the specific heat capacity.)

For each generic reaction, determine the value of 2 in terms of 1.

Consider the generic reaction: \(\mathrm{A}+2 \mathrm{~B} \longrightarrow \mathrm{C}+3 \mathrm{D} \quad \Delta H=155 \mathrm{~kJ}\) Determine the value of \(\Delta H\) for each related reaction: a. \(3 \mathrm{~A}+6 \mathrm{~B} \longrightarrow 3 \mathrm{C}+9 \mathrm{D}\) b. \(\mathrm{C}+3 \mathrm{D} \longrightarrow \mathrm{A}+2 \mathrm{~B}\) c. \(1 / 2 \mathrm{C}+3 / 2 \mathrm{D} \longrightarrow 1 / 2 \mathrm{~A}+\mathrm{B}\)

Calculate rxn for the reaction: Use the following reactions and given ’s.

Calculate \(\Delta H_{\mathrm{rxn}}\) for the reaction: \(\mathrm{CaO}(s)+\mathrm{CO}_2(g) \longrightarrow \mathrm{CaCO}_3(s)\) Use the following reactions and given \(\Delta H\) 's. \(\begin{array}{r} \mathrm{Ca}(s)+\mathrm{CO}_2(g)+1 / 2 \mathrm{O}_2(g) \longrightarrow \mathrm{CaCO}_3(s) \\ \Delta H=-812.8 \mathrm{~kJ} \end{array}\) \(2 \mathrm{Ca}(s)+\mathrm{O}_2(g) \longrightarrow 2 \mathrm{CaO}(s) \quad \Delta H=-1269.8 \mathrm{~kJ}\)

Calculate rxn for the reaction: Use the following reactions and given ’s.

Calculate rxn for the reaction: Use the following reactions and given ’s.

Write an equation for the formation of each compound from its elements in their standard states , and find for each from Appendix IIB . a. NH3(g) b. CO2(g) c. Fe2O3(s) d. CH4(g)

Hydrazine \(\left(\mathrm{N}_2 \mathrm{H}_4\right)\) is a fuel used by some spacecraft. It is normally oxidized by \(\mathrm{N}_2 \mathrm{O}_4\) according to the equation: \(\mathrm{N}_2 \mathrm{H}_4(l)+\mathrm{N}_2 \mathrm{O}_4(g) \longrightarrow 2 \mathrm{~N}_2 \mathrm{O}(g)+2 \mathrm{H}_2 \mathrm{O}(g)\) Calculate \(\Delta H_{\mathrm{rxn}}^{\circ}\) for this reaction using standard enthalpies of formation.

Write an equation for the formation of each compound from its elements in their standard states, and find \(\Delta H_{\mathrm{rxn}}^{\circ}\) for each from Appendix IIB. a. \(\mathrm{NO}_{2}(g)\) b. \(\mathrm{MgCO}_{3}(s)\) c. \(\mathrm{C}_2\mathrm{H}_4(\mathrm{g})\) d. \(\mathrm{CH}_{3} \mathrm{OH}(l)\)

Pentane (C5H12) is a component of gasoline that burns according to the following balanced equation: Calculate for this reaction using standard enthalpies of formation. (The standard enthalpy of formation of liquid pentane is -146.8 kJ/mol.)

Use standard enthalpies of formation to calculate for each reaction:

Use standard enthalpies of formation to calculate for each reaction:

During photosynthesis, plants use energy from sunlight to form glucose (C6H12O6) and oxygen from carbon dioxide and water. Write a balanced equation for photosynthesis and calculate .

Ethanol (C2H5OH) can be made from the fermentation of crops and has been used as a fuel additive to gasoline. Write a balanced equation for the combustion of ethanol and calculate .

Top fuel dragsters and funny cars burn nitromethane as fuel according to the balanced combustion equation: \(2\mathrm{\ CH}_3\mathrm{NO}_2(l)+3/2\mathrm{O}_2(g)\ \longrightarrow\ 2\mathrm{\ CO}_2(g)+3\mathrm{\ H}_2\mathrm{O}(l)+\mathrm{N}_2(g)\ \Delta H_{\mathrm{rxn}}^{\circ}=-1418\mathrm{\ kJ}\) The enthalpy of combustion for nitromethane is -709.2 kJ/mol. Calculate the standard enthalpy of formation \(\left(\Delta H_{\mathrm{f}}^{\circ}\right)\) for nitromethane.

The explosive nitroglycerin \(\left(\mathrm{C}_3\mathrm{H}_5\mathrm{N}_3\mathrm{O}_9\right)\) decomposes rapidly upon ignition or sudden impact according to the balanced equation: \(4\mathrm{\ C}_3\mathrm{H}_5\mathrm{N}_3\mathrm{O}_9(l)\longrightarrow12\ \mathrm{CO}_2(\mathrm{g})+10\mathrm{\ H}_2\mathrm{O}(\mathrm{g})+6\mathrm{\ N}_2(g)+\mathrm{O}_2(g)\) \(\Delta H_{\mathrm{rxn}}^{\circ}=-5678 \mathrm{\ kJ}\) Calculate the standard enthalpy of formation \(\left(\Delta H_{\mathrm{f}}^{\circ}\right)\) for nitroglycerin.

Determine the mass of CO2 produced by burning enough of each of the following fuels to produce 1.00 x 102 kJ of heat. Which fuel contributes least to global warming per kJ of heat produced?

Methanol (CH3OH) has been suggested as a fuel to replace gasoline. Write a balanced equation for the combustion of methanol, find , and determine the mass of carbon dioxide emitted per kJ of heat produced. Use the information from the previous exercise to calculate the same quantity for octane, C8H18. How does methanol compare to octane with respect to global warming?

The citizens of the world burn the fossil fuel equivalent of 7 x 1012 kg of petroleum per year. Assume that all of this petroleum is in the form of octane (C8H18) and calculate how much CO2 (in kg) is produced by world fossil fuel combustion per year. (Hint: Begin by writing a balanced equation for the combustion of octane.) If the atmosphere currently contains approximately 3 x 1015 kg of CO2, how long will it take for the world’s fossil fuel combustion to double the amount of atmospheric carbon dioxide?

In a sunny location, sunlight has a power density of about 1 kW/m2. Photovoltaic solar cells can convert this power into electricity with 15% efficiency. If a typical home uses 385 kWh of electricity per month, how many square meters of solar cells would be required to meet its energy requirements? Assume that electricity can be generated from the sunlight for 8 hours per day.

The kinetic energy of a rolling billiard ball is given by \(\mathrm{KE}=1 / 2 \mathrm{mv}^2\). Suppose a \(0.17 \mathrm{~kg}\) billiard ball is rolling down a pool table with an initial speed of \(4.5 \mathrm{~m} / \mathrm{s}\). As it travels, it loses some of its energy as heat. The ball slows down to \(3.8 \mathrm{~m} / \mathrm{s}\) and then collides head-on with a second billiard ball of equal mass. The first billiard ball completely stops and the second one rolls away with a velocity of \(3.8 \mathrm{~m} / \mathrm{s}\). Assume the first billiard ball is the system and calculate w, q, and \(\Delta E\) for the process.

A 100 W lightbulb is placed in a cylinder equipped with a moveable piston. The lightbulb is turned on for 0.015 hour, and the assembly expands from an initial volume of 0.85 L to a final volume of 5.88 L against an external pressure of 1.0 atm. Use the wattage of the lightbulb and the time it is on to calculate \(\Delta E\) in joules (assume that the cylinder and lightbulb assembly is the system and assume two significant figures). Calculate w and q.

Evaporating sweat cools the body because evaporation is an endothermic process: Estimate the mass of water that must evaporate from the skin to cool the body by 0.50 °C. Assume a body mass of 95 kg and assume that the specific heat capacity of the body is 4.0 J/g °C.

LP gas burns according to the exothermic reaction: What mass of LP gas is necessary to heat 1.5 L of water from room temperature (25.0 °C) to boiling (100.0 °C)? Assume that during heating, 15% of the heat emitted by the LP gas combustion goes to heat the water. The rest is lost as heat to the surroundings.

Use standard enthalpies of formation to calculate the standard change in enthalpy for the melting of ice. (The for H2O(s) is –291.8 kJ/mol.) Use this value to calculate the mass of ice required to cool 355 mL of a beverage from room temperature (25.0 °C) to 0.0 ????C. Assume that the specific heat capacity and density of the beverage are the same as those of water.

Dry ice is solid carbon dioxide. Instead of melting, solid carbon dioxide sublimes according to the equation: When carbon dioxide sublimes, the gaseous CO2 is cold enough to cause water vapor in the air to condense, forming fog. When dry ice is added to warm water, heat from the water causes the dry ice to sublime more quickly. The evaporating carbon dioxide produces a dense fog often used to create special effects. In a simple dry ice fog machine, dry ice is added to warm water in a Styrofoam cooler. The dry ice produces fog until it evaporates away, or until the water gets too cold to sublime the dry ice quickly enough. Suppose that a small Styrofoam cooler holds 15.0 liters of water heated to 85 °C. Use standard enthalpies of formation to calculate the change in enthalpy for dry ice sublimation, and calculate the mass of dry ice that should be added to the water so that the dry ice completely sublimes away when the water reaches 25 °C. Assume no heat loss to the surroundings. [The for CO2(s) is –427.4 kJ/mol.]

A 25.5 g aluminum block is warmed to 65.4 °C and plunged into an insulated beaker containing 55.2 g water initially at 22.2 °C. The aluminum and the water are allowed to come to thermal equilibrium. Assuming that no heat is lost, what is the final temperature of the water and aluminum?

If 50.0 mL of ethanol ( density = 0.789 g/mL ) initially at \(7.0 \ ^\circ C\) is mixed with 50.0 mL of water ( density = 1.0 g/mL ) initially at \(28.4 \ ^\circ C\) in an insulated beaker, and assuming that no heat is lost, what is the final temperature of the mixture?

Palmitic acid (C16H32O2) is a dietary fat found in beef and butter. The caloric content of palmitic acid is typical of fats in general. Write a balanced equation for the complete combustion of palmitic acid and calculate the standard enthalpy of combustion. What is the caloric content of palmitic acid in Cal/g? Do the same calculation for table sugar (sucrose, C12H22O11). Which dietary substance (sugar or fat) contains more Calories per gram? The standard enthalpy of formation of palmitic acid is -208 kJ/mol and that of sucrose is -2226.1 kJ/mol. [Use H2O () in the balanced chemical equations because the metabolism of these compounds produces liquid water.]

Hydrogen and methanol have both been proposed as alternatives to hydrocarbon fuels. Write balanced reactions for the complete combustion of hydrogen and methanol and use standard enthalpies of formation to calculate the amount of heat released per kilogram of the fuel. Which fuel contains the most energy in the least mass? How does the energy of these fuels compare to that of octane (C8H18)?

Problem 107E Derive a relationship between ?H and ?E for a process in which the temperature of a fixed amount of an ideal gas changes.

Under certain nonstandard conditions, oxidation by O2(g) of 1 mol of SO2(g) to SO3(g) absorbs 89.5 kJ. The enthalpy of formation of SO3(g) is -204.2 kJ under these conditions. Find the enthalpy of formation of SO2(g).

One tablespoon of peanut butter has a mass of 16 g. It is combusted in a calorimeter whose heat capacity is \(120.0\mathrm{\ kJ}/^{\circ}\mathrm{C}\). The temperature of the calorimeter rises from \(22.2\ ^{\circ}\mathrm{C}\) to \(25.4\ ^{\circ}\mathrm{C}\). Find the food caloric content of peanut butter.

A mixture of 2.0 mol of \(H_2 (g)\) and 1.0 mol of \(O_2 (g)\) is placed in a sealed evacuated container made of a perfect insulating material at \(25 \ ^\circ C\). The mixture is ignited with a spark and reacts to form liquid water. Determine the temperature of the water.

Problem 111E A 20.0 L volume of an ideal gas in a cylinder with a piston is at a pressure of 3.0 atm. Enough weight is suddenly removed from the piston to lower the external pressure to 1.5 atm. The gas then expands at constant temperature until its pressure is 1.5 atm. Find ?E , ?H , q , and w for this change in state.

When 10.00 g of phosphorus is burned in O2(g) to form P4O10(s) , enough heat is generated to raise the temperature of 2950 g of water from 18.0 °C to 38.0 °C. Calculate the enthalpy of formation of P4O10(s) under these conditions.

The for the oxidation of S in the gas phase to SO3 is -204 kJ/mol and for the oxidation of SO2 to SO3 is 89.5 kJ/mol. Find the enthalpy of formation of SO2 under these conditions.

The \(\Delta H_{\mathrm{f}}^{\mathrm{o}}\) of \(\mathrm{TiI}_{3}(s)\) is –328 kJ/mol and the for the reaction \(2 \mathrm{Ti}(s)+3 \mathrm{I}_{2}(g) \ \longrightarrow \ 2 \mathrm{TiI}_{3}(s)\) is –839 kJ. Calculate the \(\Delta H\) of sublimation of \(\mathrm{I}_{2}(s)\), which is a solid at \(25^{\circ} \mathrm{C}\).

A gaseous fuel mixture contains 25.3% methane (CH4), 38.2% ethane (C2H6), and the rest propane (C2H8) by volume. When the fuel mixture contained in a 1.55 L tank, stored at 755 mmHg and 298 K, undergoes complete combustion, how much heat is emitted? (Assume that the water produced by the combustion is in the gaseous state.)

A copper cube measuring 1.55 cm on edge and an aluminum cube measuring 1.62 cm on edge are both heated to 55.0 °C and submerged in 100.0 mL of water at 22.2 °C. What is the final temperature of the water when equilibrium is reached? (Assume a density of 0.998 g/mL for water.)

A gaseous fuel mixture stored at 745 mmHg and 298 K contains only methane (CH4) and propane (C3H8). When 11.7 L of this fuel mixture is burned, it produces 769 kJ of heat. What is the mole fraction of methane in the mixture? (Assume that the water produced by the combustion is in the gaseous state.)

A pure gold ring and a pure silver ring have a total mass of 14.9 g. The two rings are heated to 62.0 °C and dropped into 15.0 mL of water at 23.5 °C. When equilibrium is reached, the temperature of the water is 25.0 °C. What is the mass of each ring? (Assume a density of 0.998 g/mL for water.)

A typical frostless refrigerator uses 655 kWh of energy per year in the form of electricity. Suppose that all of this electricity is generated at a power plant that burns coal containing 3.2% sulfur by mass and that all of the sulfur is emitted as SO2 when the coal is burned. If all of the SO2 goes on to react with rainwater to form H2SO4, what mass of H2SO4 does the annual operation of the refrigerator produce? (Hint: Assume that the remaining percentage of the coal is carbon and begin by calculating for the combustion of carbon.)

A large sport utility vehicle has a mass of 2.5 x 103 kg . Calculate the mass of CO2 emitted into the atmosphere upon accelerating the SUV from 0.0 mph to 65.0 mph. Assume that the required energy comes from the combustion of octane with 30% efficiency. (Hint: Use KE = 1/2 mv2 to calculate the kinetic energy required for the acceleration.)

Combustion of natural gas (primarily methane) occurs in most household heaters. The heat given off in this reaction is used to raise the temperature of the air in the house. Assuming that all the energy given off in the reaction goes to heating up only the air in the house, determine the mass of methane required to heat the air in a house by \(10.0^{\circ} \mathrm{C}\). Assume that the house dimensions are \(30.0 \ \mathrm{m} \times 30.0 \ \mathrm{m} \times 3.0 \ \mathrm{m}\), specific heat capacity of air is \(30 \ \mathrm{J} / \mathrm{K} \cdot \mathrm{mol}\), and 1.00 mol of air occupies 22.4 L for all temperatures concerned.

When backpacking in the wilderness, hikers often boil water to sterilize it for drinking. Suppose that you are planning a backpacking trip and will need to boil 35 L of water for your group. What volume of fuel should you bring? Assume that the fuel has an average formula of C7H16, 15% of the heat generated from combustion goes to heat the water (the rest is lost to the surroundings), the density of the fuel is 0.78 g/mL, the initial temperature of the water is 25.0 °C, and the standard enthalpy of formation of C7H16 is –224.4 kJ/mol.

An ice cube of mass 9.0 g is added to a cup of coffee. The coffee’s initial temperature is 90.0 °C and the cup contains 120.0 g of liquid. Assume the specific heat capacity of the coffee is the same as that of water. The heat of fusion of ice (the heat associated with ice melting) is 6.0 kJ/mol. Find the temperature of the coffee after the ice melts.

Find , , q , and w for the freezing of water at –10.0 °C. The specific heat capacity of ice is 2.04 J/g °C and its heat of fusion (the quantity of heat associated with melting) is –332 J/g.

Problem 125E Starting from the relationship between temperature and kinetic energy for an ideal gas, find the value of the molar heat capacity of an ideal gas when its temperature is changed at constant volume. Find its molar heat capacity when its temperature is changed at constant pressure.

Problem 127E The heat of vaporization of water at 373 K is 40.7 kJ/mol. Find q , w , ?E , and ?H for the evaporation of 454 g of water at this temperature at 1 atm.

Problem 126E An amount of an ideal gas expands from 12.0 L to 24.0 L at a constant pressure of 1.0 atm. Then the gas is cooled at a constant volume of 24.0 L back to its original temperature. Then it contracts back to its original volume. Find the total heat flow for the entire process.

Find , , q , and w for the change in state of 1.0 mol H2O(l) at 80 °C to H2O(g) at 110 °C. The heat capacity of H2O(l) = 75.3 J/mol K , heat capacity of H2O(g) = 25.0 J/mol K , and the heat of vaporization of H2O is 40.7 x 103 J/mol at 100 °C.

The heat of combustion of liquid octane (C8H18) to carbon dioxide and liquid water at 298 K is –1303 kJ/mol. Find for this reaction.

Find for the combustion of ethanol (C2H6O) to carbon dioxide and liquid water from the following data. The heat capacity of the bomb calorimeter is 34.65 kJ/K and the combustion of 1.765 g of ethanol raises the temperature of the calorimeter from 294.33 K to 295.84 K.

Which statement is true of the internal energy of the system and its surroundings following a process in which \(\Delta E_{\mathrm{sys}}=+65\mathrm{\ kJ}\)? Explain. a. The system and the surroundings both lose 65 kJ of energy. b. The system and the surroundings both gain 65 kJ of energy. c. The system loses 65 kJ of energy and the surroundings gain 65 kJ of energy. d. The system gains 65 kJ of energy and the surroundings lose 65 kJ of energy.

Problem 132E The internal energy of an ideal gas depends only on its temperature. Which statement is true of an isothermal (constanttemperature) expansion of an ideal gas against a constant external pressure? Explain. a. ?E is positive b. w is positive c. q is positive d. ?E is negative

Problem 133E Which expression describes the heat evolved in a chemical reaction when the reaction is carried out at constant pressure? Explain. a. ?E - w b. ?E c. ?E - q

Problem 134E Two identical refrigerators are plugged in for the first time. Refrigerator A is empty (except for air) and refrigerator B is filled with jugs of water. The compressors of both refrigerators immediately turn on and begin cooling the interiors of the refrigerators. After 2 hours, the compressor of refrigerator A turns off while the compressor of refrigerator B continues to run. The next day, the compressor of refrigerator A can be heard turning on and off every few minutes, while the compressor of refrigerator B turns off and on every hour or so (and stays on longer each time). Explain these observations.

Problem 135E A 1 kg cylinder of aluminum and 1 kg jug of water, both at room temperature, are put into a refrigerator. After 1 hour, the temperature of each object is measured. One of the objects is much cooler than the other. Which one is cooler and why?

Problem 136E Two substances A and B, initially at different temperatures, are thermally isolated from their surroundings and allowed to come into thermal contact. The mass of substance A is twice the mass of substance B, but the specific heat capacity of substance B is four times the specific heat capacity of substance A. Which substance will undergo a larger change in temperature?

Problem 137E When 1 mol of a gas burns at constant pressure, it produces 2418 J of heat and does 5 J of work. Determine ?E , ?H , q , and w for the process.

Problem 138E In an exothermic reaction, the reactants lose energy and the reaction feels hot to the touch. Explain why the reaction feels hot even though the reactants are losing energy. Where does the energy come from?

Problem 139E Which statement is true of a reaction in which ?V is positive? Explain. a. ?H = ?E b. ?H > ?E c. ?H <?E

What is the meaning of ekklesia (para. 43)? What does King mean when he invokes the true ekklesia?

How does King balance the twin appeals to religion and patriotism throughout Letter from Birmingham Jail? Do you think he puts more emphasis on religion or patriotism? Why do you think he makes that choice?

In the later 1960s, Alice Walker wrote an essay titled The Civil Rights Movement: What Good Was It? How would you answer her question today? What good do you believe has resulted from the civil rights movement?

In the long sentence in paragraph 14 (beginning with But when you have seen), why does King arrange the when clauses in the order that he does? Try repositioning them, and then discuss the difference in effect.

What rhetorical strategies are used in paragraph 25? Identify at least four

What are the chief rhetorical strategies used in paragraph 31? Identify at least five

What is internal energy? Is internal energy a state function?

Trace one of the following patterns of figurative language throughout Kings letter: darkness and light, high and low, sickness and health.

King uses repetition of single words or phrases, of sentence structures, and of sounds. Focusing on a passage of one or more paragraphs, discuss the effect of this use of repetition.

Considering the final three paragraphs as Kings conclusion, discuss whether you believe it is rhetorically effective.

If you were asked to develop a series of suggestions to the area where you live for a healthier environment based on Putnams research, what would you recommend?

Identify one or two of the ways people establish social connectedness for example, by belonging to an organ ized religion, volunteering, participating in associations, spending time with friends. What role(s) do you think social media might play in increasing social capital?

What is heat capacity? Explain the difference between heat capacity and specific heat capacity.

Explain how the high specific heat capacity of water can affect the weather in coastal regions.

If two objects, A and B, of different temperature come into direct contact, what is the relationship between the heat lost by one object and the heat gained by the other? What is the relationship between the temperature changes of the two objects? (Assume that the two objects do not lose any heat to anything else.)

What is pressurevolume work? How is it calculated?

What is calorimetry? Explain the difference between a coffeecup calorimeter and a bomb calorimeter. What is each designed to measure?

What is the change in enthalpy 1H2 for a chemical reaction? How is H different from E? 2

Explain the difference between an exothermic and an endothermic reaction. Give the sign of H for each type of reaction.

From a molecular viewpoint, where does the energy emitted in an exothermic chemical reaction come from? Why does the reaction mixture undergo an increase in temperature even though energy is emitted?

From a molecular viewpoint, where does the energy absorbed in an endothermic chemical reaction go? Why does the reaction mixture undergo a decrease in temperature even though energy is absorbed?

Is the change in enthalpy for a reaction an extensive property? Explain the relationship between H for a reaction and the amounts of reactants and products that undergo reaction.

Explain how the value of H for a reaction changes upon each operation: a. multiplying the reaction by a factor. b. reversing the reaction. Why do these relationships hold?

What is Hesss law? Why is it useful?

What is a standard state? What is the standard enthalpy change for a reaction?

What is the standard enthalpy of formation for a compound? For a pure element in its standard state?

How do you calculate Ho rxn from tabulated standard enthalpies of formation?

What are the main sources of the energy consumed in the United States?

What are the main environmental problems associated with fossil fuel use?

Explain global climate change. What causes global warming? What is the evidence that global warming is occurring?

Convert between energy units: a. 534 kWh to J b. 215 kJ to Cal c. 567 Cal to J d. 2.85 * 103 J to cal

Convert between energy units: a. 231 cal to kJ b. 132 * 104 kJ to kcal c. 4.99 * 103 kJ to kWh d. 2.88 * 104 J to Cal

Suppose that a person eats 2387 Calories per day. Convert this amount of energy into each unit: a. J b. kJ c. kWh

A particular frost-free refrigerator uses about 745 kWh of electrical energy per year. Express this amount of energy in each unit: a. J b. kJ c. Cal

Which statement is true of the internal energy of a system and its surroundings during an energy exchange with a negative Esys? a. The internal energy of the system increases and the internal energy of the surroundings decreases. b. The internal energy of both the system and the surroundings increases. c. The internal energy of both the system and the surroundings decreases. d. The internal energy of the system decreases and the internal energy of the surroundings increases.

Which statement is true of the internal energy of a system and its surroundings during an energy exchange with a negative Esys? a. The internal energy of the system increases and the internal energy of the surroundings decreases. b. The internal energy of both the system and the surroundings increases. c. The internal energy of both the system and the surroundings decreases. d. The internal energy of the system decreases and the internal energy of the surroundings increases.

Identify each energy exchange as primarily heat or work and determine whether the sign of E is positive or negative for the system. a. Sweat evaporates from skin, cooling the skin. (The evaporating sweat is the system.) b. A balloon expands against an external pressure. (The contents of the balloon is the system.) c. An aqueous chemical reaction mixture is warmed with an external fl ame. (The reaction mixture is the system.)

Identify each energy exchange as primarily heat or work and determine whether the sign of E is positive or negative for the system. a. A rolling billiard ball collides with another billiard ball. The fi rst billiard ball (defi ned as the system) stops rolling after the collision. b. A book is dropped to the fl oor. (The book is the system). c. A father pushes his daughter on a swing. (The daughter and the swing are the system).

A system releases 622 kJ of heat and does 105 kJ of work on the surroundings. What is the change in internal energy of the system?

A system absorbs 196 kJ of heat and the surroundings do 117 kJ of work on the system. What is the change in internal energy of the system?

The gas in a piston (defined as the system) warms and absorbs 655 J of heat. The expansion performs 344 J of work on the surroundings. What is the change in internal energy for the system?

The air in an inflated balloon (defined as the system) warms over a toaster and absorbs 115 J of heat. As it expands, it does 77 kJ of work. What is the change in internal energy for the system?

We pack two identical coolers for a picnic, placing twenty-four 12-ounce soft drinks and 5 pounds of ice in each. However, the drinks that we put into cooler A were refrigerated for several hours before they were packed in the cooler, while the drinks that we put into cooler B were at room temperature. When we open the two coolers 3 hours later, most of the ice in cooler A is still present, while nearly all of the ice in cooler B has melted. Explain this difference.

A kilogram of aluminum metal and a kilogram of water are each warmed to \(75^{\ \circ}\mathrm{C}\) and placed in two identical insulated containers. One hour later, the two containers are opened and the temperature of each substance is measured. The aluminum has cooled to \(35^{\ \circ}\mathrm{C}\) while the water has cooled only to \(66^{\ \circ}\mathrm{C}\). Explain this difference.

How much heat is required to warm 1.50 L of water from 25.0 C to 100.0 C? (Assume a density of 1.0 g/mL for the water.)

How much heat is required to warm 1.50 kg of sand from 25.0 C to 100.0 C?

Suppose that 25 g of each substance is initially at 27.0 C. What is the final temperature of each substance upon absorbing 2.35 kJ of heat? a. gold b. silver c. aluminum d. water

An unknown mass of each substance, initially at 23.0 C, absorbs 1.95 * 103 J of heat. The final temperature is recorded as indicated. Find the mass of each substance. a. Pyrex glass 1Tf = 55.4 C2 b. sand 1Tf = 62.1 C2 c. ethanol 1Tf = 44.2 C2 d. water 1Tf = 32.4 C2 51.

How much work (in J) is required to expand the volume of a pump from 0.0 L to 2.5 L against an external pressure of 1.1 atm?

The average human lung expands by about 0.50 L during each breath. If this expansion occurs against an external pressure of 1.0 atm, how much work (in J) is done during the expansion?

The air within a piston equipped with a cylinder absorbs 565 J of heat and expands from an initial volume of 0.10 L to a final volume of 0.85 L against an external pressure of 1.0 atm. What is the change in internal energy of the air within the piston?

A gas is compressed from an initial volume of 5.55 L to a final volume of 1.22 L by an external pressure of 1.00 atm. During the compression the gas releases 124 J of heat. What is the change in internal energy of the gas?

When 1 mol of a fuel burns at constant pressure, it produces 3452 kJ of heat and does 11 kJ of work. What are E and H for the combustion of the fuel?

The change in internal energy for the combustion of 1.0 mol of octane at a pressure of 1.0 atm is 5084.3 kJ. If the change in enthalpy is 5074.1 kJ, how much work is done during the combustion?

Determine whether each process is exothermic or endothermic and indicate the sign of H . a. natural gas burning on a stove b. isopropyl alcohol evaporating from skin c. water condensing from steam

Determine whether each process is exothermic or endothermic and indicate the sign of H . a. dry ice evaporating b. a sparkler burning c. the reaction that occurs in a chemical cold pack used to ice athletic injuries

Consider the thermochemical equation for the combustion of acetone (C3H6O) , the main ingredient in nail polish remover. C3H6O(l) + 4 O2(g) h 3 CO2(g) + 3 H2O(g) Hrxn = -1790 kJ If a bottle of nail polish remover contains 177 mL of acetone, how much heat is released by its complete combustion? The density of acetone is 0.788 g/mL.

What mass of natural gas (CH4) must burn to emit 267 kJ of heat? CH4(g) + 2 O2(g) h CO2(g) + 2 H2O(g) Hrxn = -802.3 kJ

Nitromethane (CH 3 NO 2 ) burns in air to produce significant amounts of heat. 2 CH3NO2(l) + 3 >2 O2(g)h2 CO2(g) + 3 H2O(l) + N2(g) Hrxn = -1418 kJ How much heat is produced by the complete reaction of 5.56 kg of nitromethane?

Titanium reacts with iodine to form titanium(III) iodide, emitting heat. 2 Ti(s) + 3 I2(g) h 2 TiI3(s) Hrxn = -839 kJ Determine the masses of titanium and iodine that react if 1.55 * 10 3 kJ of heat is emitted by the reaction.

The propane fuel (C3H8) used in gas barbeques burns according to the thermochemical equation: If a pork roast must absorb 1.6 * 103 kJ to fully cook, and if only 10% of the heat produced by the barbeque is actually absorbed by the roast, what mass of CO2 is emitted into the atmosphere during the grilling of the pork roast?

Charcoal is primarily carbon. Determine the mass of CO2 produced by burning enough carbon (in the form of charcoal) to produce 5.00 * 102 kJ of heat. C(s) + O2(g) h CO2(g) Hrxn = -393.5 kJ

A silver block, initially at 58.5 C, is submerged into 100.0 g of water at 24.8 C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 26.2 C. What is the mass of the silver block? 6

A 32.5 g iron rod, initially at 22.7 C, is submerged into an unknown mass of water at 63.2 C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 59.5 C. What is the mass of the water? 6

A 31.1 g wafer of pure gold, initially at 69.3 C, is submerged into 64.2 g of water at 27.8 C in an insulated container. What is the final temperature of both substances at thermal equilibrium?

A 2.85 g lead weight, initially at 10.3 C, is submerged in 7.55 g of water at 52.3 C in an insulated container. What is the final temperature of both substances at thermal equilibrium?

Two substances, A and B, initially at different temperatures, come into contact and reach thermal equilibrium. The mass of substance A is 6.15 g and its initial temperature is 20.5 C. The mass of substance B is 25.2 g and its initial temperature is 52.7 C. The final temperature of both substances at thermal equilibrium is 46.7 C. If the specific heat capacity of substance B is 1.17 J/g # C , what is the specific heat capacity of substance A? 7

A 2.74 g sample of a substance suspected of being pure gold is warmed to \(72.1 \ ^\circ C\) and submerged into 15.2 g of water initially at \(24.7 \ ^\circ C\). The final temperature of the mixture is \(26.3 \ ^\circ C\). What is the heat capacity of the unknown substance? Could the substance be pure gold?

Exactly 1.5 g of a fuel burns under conditions of constant pressure and then again under conditions of constant volume. In measurement A the reaction produces 25.9 kJ of heat, and in measurement B the reaction produces 23.3 kJ of heat. Which measurement (A or B) corresponds to conditions of constant pressure? Which one corresponds to conditions of constant volume? Explain.

In order to obtain the largest possible amount of heat from a chemical reaction in which there is a large increase in the number of moles of gas, should you carry out the reaction under conditions of constant volume or constant pressure? Explain.

When 0.514 g of biphenyl (C12H10) undergoes combustion in a bomb calorimeter, the temperature rises from 25.8 C to 29.4 C. Find Erxn for the combustion of biphenyl in kJ/mol biphenyl. The heat capacity of the bomb calorimeter, determined in a separate experiment, is 5.86 kJ/C.

Mothballs are composed primarily of the hydrocarbon naphthalene (C10H8) . When 1.025 g of naphthalene burns in a bomb calorimeter, the temperature rises from 24.25 C to 32.33 C. Find Erxn for the combustion of naphthalene. The heat capacity of the calorimeter, determined in a separate experiment, is 5.11 kJ/C. 75

Zinc metal reacts with hydrochloric acid according to the balanced equation: Zn(s) + 2 HCl(aq) h ZnCl2(aq) + H2(g) When 0.103 g of Zn( s ) is combined with enough HCl to make 50.0 mL of solution in a coffee-cup calorimeter, all of the zinc reacts, raising the temperature of the solution from 22.5 C to 23.7 C. Find Hrxn for this reaction as written. (Use 1.0 g/mL for the density of the solution and 4.18 J/g # C as the specific heat capacity.) 76

Instant cold packs used to ice athletic injuries on the field contain ammonium nitrate and water separated by a thin plastic divider. When the divider is broken, the ammonium nitrate dissolves according to the endothermic reaction: NH4NO3(s) h NH4 +(aq) + NO3 -(aq) In order to measure the enthalpy change for this reaction, 1.25 g of NH4NO3 is dissolved in enough water to make 25.0 mL of solution. The initial temperature is 25.8 C and the final temperature (after the solid dissolves) is 21.9 C. Calculate the change in enthalpy for the reaction in kJ. (Use 1.0 g/mL as the density of the solution and 4.18 J/g # C as the specific heat capacity.) Qu

For each generic reaction, determine the value of H2 in terms of H1. a. A + B h 2 C H1 2 C h A + B H2 = ? b. A + 1 >2B h C H1 2 A + B h 2 C H2 = ? c. A h B + 2 C H1 1 >2 B + C h 1 >2 A H2 = ? 78. Co

Consider the generic reaction: A + 2 B h C + 3 D H = 155 kJ Determine the value of H for each related reaction: a. 3 A + 6 B h 3 C + 9 D b. C + 3 D h A + 2 B c. 1 >2C + 3 >2 D h 1 >2A + B

Calculate Hrxn for the reaction: Fe2O3(s) + 3 CO(g) h 2 Fe(s) + 3 CO2(g) Use the following reactions and given H s. 2 Fe(s) + 3 >2 O2(g) h Fe2O3(s) H = -824.2 kJ CO(g) + 1 >2O2(g) h CO2(g) H = -282.7 kJ 80

Calculate Hrxn for the reaction: CaO(s) + CO2(g) h CaCO3(s) Use the following reactions and given H s. Ca(s) + CO2(g) + 1 >2 O2(g) h CaCO3(s) H = -812.8 kJ 2 Ca(s) + O2(g) h 2 CaO(s) H = -1269.8 kJ 81

Calculate \(\Delta H_{\mathrm{rxn}}\) for the reaction: \(5 \mathrm{C}(s)+6 \mathrm{\ H}_{2}(\mathrm{g}) \longrightarrow \mathrm{C}_{5} \mathrm{H}_{12}(l)\) Use the following reactions and given \(\Delta H \text { 's }\). \(\mathrm{C}_5\mathrm{H}_{12}(l)+8\mathrm{\ O}_2(g)\longrightarrow5\mathrm{\ CO}_2(g)+6\mathrm{\ H}_2\mathrm{O}(g)\) \(\Delta H=-3244.8 \mathrm{\ kJ}\) \(\mathrm{C}(s)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)\) \(\Delta H=-393.5 \mathrm{\ kJ}\) \(2\ \mathrm{H}_2(\mathrm{g})+\mathrm{O}_2(\mathrm{g})\longrightarrow2\mathrm{\ H}_2\mathrm{O}(\mathrm{g})\) \(\Delta H=-483.5 \mathrm{\ kJ}\)

Calculate Hrxn for the reaction: CH4(g) + 4 Cl2(g) h CCl4(g) + 4 HCl(g) Use the following reactions and given H s. C(s) + 2 H2(g) h CH4(g) H = -74.6 kJ C(s) + 2 Cl2(g) h CCl4(g) H = -95.7 kJ H2(g) + Cl2(g) h 2 HCl(g) H = -92.3 kJ Ent

Write an equation for the formation of each compound from its elements in their standard states , and find Hf for each from Appendix IIB . a. NH3(g) b. CO2(g) c. Fe2O3(s) d. CH4(g)

Write an equation for the formation of each compound from its elements in their standard states , and find Hrxn for each from Appendix IIB . a. NO2(g) b. MgCO3(s) c. C2H4(g) d. CH3OH(l)

Hydrazine (N2H4) is a fuel used by some spacecraft. It is normally oxidized by N2O4 according to the equation: N2H4(l) + N2O4(g) h 2 N2O(g) + 2 H2O(g) Calculate Hrxn for this reaction using standard enthalpies of formation.

Pentane \(\left(\mathrm{C}_5 \mathrm{H}_{12}\right)\) is a component of gasoline that burns according to the following balanced equation: \(\mathrm{C}_5 \mathrm{H}_{12}(l)+8 \mathrm{O}_2(g) \longrightarrow 5 \mathrm{CO}_2(g)+6 \mathrm{H}_2 \mathrm{O}(g)\) Calculate \(\Delta H_{\mathrm{rxn}}^{\circ}\) for this reaction using standard enthalpies of formation. (The standard enthalpy of formation of liquid pentane is \(-146.8 \mathrm{~kJ} / \mathrm{mol}\).)

Use standard enthalpies of formation to calculate Hrxn for each reaction: a. C2H4(g) + H2(g) h C2H6(g) b. CO(g) + H2O(g) h H2(g) + CO2(g) c. 3 NO2(g) + H2O(l) h 2 HNO3(aq) + NO(g) d. Cr2O3(s) + 3 CO(g) h 2 Cr(s) + 3 CO2(g)

Use standard enthalpies of formation to calculate Hrxn for each reaction: a. 2 H2S(g) + 3 O2(g) h 2 H2O(l) + 2 SO2(g) b. SO2(g) + 1 >2 O2(g) h SO3(g) c. C(s) + H2O(g) h CO(g) + H2(g) d. N2O4(g) + 4 H2(g) h N2(g) + 4 H2O(g)

During photosynthesis, plants use energy from sunlight to form glucose (C6H12O6) and oxygen from carbon dioxide and water. Write a balanced equation for photosynthesis and calculate Hrxn .

Ethanol (C2H5OH) can be made from the fermentation of crops and has been used as a fuel additive to gasoline. Write a balanced equation for the combustion of ethanol and calculate Hrxn .

Top fuel dragsters and funny cars burn nitromethane as fuel according to the balanced combustion equation: 2 CH3NO2(l) + 3 >2O2(g) h 2 CO2(g) + 3 H2O(l) + N2(g) Hrxn = -1418 kJ The enthalpy of combustion for nitromethane is -709.2 kJ/mol. Calculate the standard enthalpy of formation 1Hf2 for nitromethane. 92

The explosive nitroglycerin 1C3H5N3O92 decomposes rapidly upon ignition or sudden impact according to the balanced equation: 4 C3H5N3O9(l) h 12 CO2(g) + 10 H2O(g) + 6 N2(g) + O2(g) Hrxn = -5678 kJ Calculate the standard enthalpy of formation 1Hf2 for nitroglycerin. Ene

Determine the mass of CO2 produced by burning enough of each of the following fuels to produce 1.00 * 102 kJ of heat. Which fuel contributes least to global warming per kJ of heat produced? a. CH4(g) + 2 O2(g) h CO2(g) + 2 H2O(g) Hrxn = -802.3 kJ b. C3H8(g) + 5 O2(g) h 3 CO2(g) + 4 H2O(g) Hrxn = -2043 kJ c. C8H18(l) + 25>2 O2(g) h 8 CO2(g) + 9 H2O(g) Hrxn = -5074.1 kJ 94.

Methanol (CH3OH) has been suggested as a fuel to replace gasoline. Write a balanced equation for the combustion of methanol, find Hrxn , and determine the mass of carbon dioxide emitted per kJ of heat produced. Use the information from the previous exercise to calculate the same quantity for octane, C8H18 . How does methanol compare to octane with respect to global warming?

The citizens of the world burn the fossil fuel equivalent of 7 * 1012 kg of petroleum per year. Assume that all of this petroleum is in the form of octane (C8H18) and calculate how much CO2 (in kg) is produced by world fossil fuel combustion per year. (Hint: Begin by writing a balanced equation for the combustion of octane.) If the atmosphere currently contains approximately 3 * 1015 kg of CO2 , how long will it take for the worlds fossil fuel combustion to double the amount of atmospheric carbon dioxide?

In a sunny location, sunlight has a power density of about \(1\ \mathrm{kW}/\mathrm{m}^2\). Photovoltaic solar cells can convert this power into electricity with 15% efficiency. If a typical home uses 385 kWh of electricity per month, how many square meters of solar cells would be required to meet its energy requirements? Assume that electricity can be generated from the sunlight for 8 hours per day.

The kinetic energy of a rolling billiard ball is given by KE = 1 >2mv2 . Suppose a 0.17 kg billiard ball is rolling down a pool table with an initial speed of 4.5 m/s. As it travels, it loses some of its energy as heat. The ball slows down to 3.8 m/s and then collides head-on with a second billiard ball of equal mass. The first billiard ball completely stops and the second one rolls away with a velocity of 3.8 m/s. Assume the first billiard ball is the system and calculate w , q , and E for the process.

A 100 W lightbulb is placed in a cylinder equipped with a moveable piston. The lightbulb is turned on for 0.015 hour, and the assembly expands from an initial volume of 0.85 L to a final volume of 5.88 L against an external pressure of 1.0 atm. Use the wattage of the lightbulb and the time it is on to calculate E in joules (assume that the cylinder and lightbulb assembly is the system and assume two significant figures). Calculate w and q .

Evaporating sweat cools the body because evaporation is an endothermic process: H2O(l) h H2O(g) Ho rxn = +44.01 kJ Estimate the mass of water that must evaporate from the skin to cool the body by 0.50 C. Assume a body mass of 95 kg and assume that the specific heat capacity of the body is 4.0 J/g # C . 1

LP gas burns according to the exothermic reaction: C3H8(g) + 5 O2(g) h 3 CO2(g) + 4 H2O(g) Hrxn = -2044 kJ What mass of LP gas is necessary to heat 1.5 L of water from room temperature (25.0 C) to boiling (100.0 C)? Assume that during heating, 15% of the heat emitted by the LP gas combustion goes to heat the water. The rest is lost as heat to the surroundings.

Use standard enthalpies of formation to calculate the standard change in enthalpy for the melting of ice. (The Hf for H2O(s) is 291.8 kJ/mol.) Use this value to calculate the mass of ice required to cool 355 mL of a beverage from room temperature (25.0 C) to 0.0 C. Assume that the specific heat capacity and density of the beverage are the same as those of water. 1

Dry ice is solid carbon dioxide. Instead of melting, solid carbon dioxide sublimes according to the equation: CO2(s) h CO2(g) When dry ice is added to warm water, heat from the water causes the dry ice to sublime more quickly. The evaporating carbon dioxide produces a dense fog often used to create special effects. In a simple dry ice fog machine, dry ice is added to warm water in a Styrofoam cooler. The dry ice produces fog until it evaporates away, or until the water gets too cold to sublime the dry ice quickly enough. Suppose that a small Styrofoam cooler holds 15.0 liters of water heated to 85 C. Use standard enthalpies of formation to calculate the change in enthalpy for dry ice sublimation, and calculate the mass of dry ice that should be added to the water so that the dry ice completely sublimes away when the water reaches 25 C. Assume no heat loss to the surroundings. [The Ho f for CO2(s) is 427.4 kJ/mol.] 1

A 25.5 g aluminum block is warmed to 65.4 C and plunged into an insulated beaker containing 55.2 g water initially at 22.2 C. The aluminum and the water are allowed to come to thermal equilibrium. Assuming that no heat is lost, what is the final temperature of the water and aluminum?

If 50.0 mL of ethanol ( density = 0.789 g>mL ) initially at 7.0 C is mixed with 50.0 mL of water ( density = 1.0 g>mL ) initially at 28.4 C in an insulated beaker, and assuming that no heat is lost, what is the final temperature of the mixture?

Palmitic acid 1C16H32O22 is a dietary fat found in beef and butter. The caloric content of palmitic acid is typical of fats in general. Write a balanced equation for the complete combustion of palmitic acid and calculate the standard enthalpy of combustion. What is the caloric content of palmitic acid in Cal/g? Do the same calculation for table sugar (sucrose, C12H22O11 ). Which dietary substance (sugar or fat) contains more Calories per gram? The standard enthalpy of formation of palmitic acid is -208 kJ/mol and that of sucrose is -2226.1 kJ/mol. [Use H2O1l2 in the balanced chemical equations because the metabolism of these compounds produces liquid water.]

Hydrogen and methanol have both been proposed as alternatives to hydrocarbon fuels. Write balanced reactions for the complete combustion of hydrogen and methanol and use standard enthalpies of formation to calculate the amount of heat released per kilogram of the fuel. Which fuel contains the most energy in the least mass? How does the energy of these fuels compare to that of octane (\(C_8 H_18\))?

Derive a relationship between H and E for a process in which the temperature of a fixed amount of an ideal gas changes.

Under certain nonstandard conditions, oxidation by O2(g) of 1 mol of SO2 (g) to SO3(g) absorbs 89.5 kJ. The enthalpy of formation of SO3(g) is -204.2 kJ under these conditions. Find the enthalpy of formation of SO2(g) .

One tablespoon of peanut butter has a mass of 16 g. It is combusted in a calorimeter whose heat capacity is 120.0 kJ>C. The temperature of the calorimeter rises from 22.2 C to 25.4 C. Find the food caloric content of peanut butter. 1

A mixture of 2.0 mol of H2(g) and 1.0 mol of O2(g) is placed in a sealed evacuated container made of a perfect insulating material at 25 C. The mixture is ignited with a spark and reacts to form liquid water. Determine the temperature of the water.

A 20.0 L volume of an ideal gas in a cylinder with a piston is at a pressure of 3.0 atm. Enough weight is suddenly removed from the piston to lower the external pressure to 1.5 atm. The gas then expands at constant temperature until its pressure is 1.5 atm. Find AE, AH, q and w for this change in state.

When 10.00 g of phosphorus is burned in O2(g) to form P4O10(s) , enough heat is generated to raise the temperature of 2950 g of water from 18.0 C to 38.0 C. Calculate the enthalpy of formation of P4O10(s) under these conditions.

The \(\Delta H\) for the oxidation of S in the gas phase to \(\mathrm{SO}_{3}\) is -204 kJ/mol and for the oxidation of \(\mathrm{SO}_{2}\) to \(\mathrm{SO}_{3}\) is 89.5 kJ/mol. Find the enthalpy of formation of \(\mathrm{SO}_{2}\) under these conditions.

The Ho f of TiI3(s) is 328 kJ/mol and the H for the reaction 2 Ti(s) + 3 I2(g) h 2 TiI3(s) is 839 kJ. Calculate the H of sublimation of I2(s) , which is a solid at 25 C. 11

A gaseous fuel mixture contains 25.3% methane (CH4) , 38.2% ethane (C2H6) , and the rest propane (C3H8) by volume. When the fuel mixture contained in a 1.55 L tank, stored at 755 mmHg and 298 K, undergoes complete combustion, how much heat is emitted? (Assume that the water produced by the combustion is in the gaseous state.)

A gaseous fuel mixture stored at 745 mmHg and 298 K contains only methane (CH4) and propane (C3H8) . When 11.7 L of this fuel mixture is burned, it produces 769 kJ of heat. What is the mole fraction of methane in the mixture? (Assume that the water produced by the combustion is in the gaseous state.)

A copper cube measuring 1.55 cm on edge and an aluminum cube measuring 1.62 cm on edge are both heated to 55.0 C and submerged in 100.0 mL of water at 22.2 C. What is the final temperature of the water when equilibrium is reached? (Assume a density of 0.998 g/mL for water.)

A pure gold ring and a pure silver ring have a total mass of 14.9 g. The two rings are heated to 62.0 C and dropped into 15.0 mL of water at 23.5 C. When equilibrium is reached, the temperature of the water is 25.0 C. What is the mass of each ring? (Assume a density of 0.998 g/mL for water.) Ch

A typical frostless refrigerator uses 655 kWh of energy per year in the form of electricity. Suppose that all of this electricity is generated at a power plant that burns coal containing 3.2% sulfur by mass and that all of the sulfur is emitted as SO2 when the coal is burned. If all of the SO2 goes on to react with rainwater to form H2SO4 , what mass of H2SO4 does the annual operation of the refrigerator produce? (Hint: Assume that the remaining percentage of the coal is carbon and begin by calculating Ho rxn for the combustion of carbon.)

A large sport utility vehicle has a mass of 2.5 * 103 kg . Calculate the mass of CO2 emitted into the atmosphere upon accelerating the SUV from 0.0 mph to 65.0 mph. Assume that the required energy comes from the combustion of octane with 30% efficiency. (Hint: Use KE = 1 >2 mv2 to calculate the kinetic energy required for the acceleration.)

Combustion of natural gas (primarily methane) occurs in most household heaters. The heat given off in this reaction is used to raise the temperature of the air in the house. Assuming that all the energy given off in the reaction goes to heating up only the air in the house, determine the mass of methane required to heat the air in a house by \(10.0 \ ^\circ C\). Assume that the house dimensions are \(30.0 m \times 30.0 m \times 3.0 m\), specific heat capacity of air is \(30 \ J/K \cdot mol\) , and 1.00 mol of air occupies 22.4 L for all temperatures concerned.

When backpacking in the wilderness, hikers often boil water to sterilize it for drinking. Suppose that you are planning a backpacking trip and will need to boil 35 L of water for your group. What volume of fuel should you bring? Assume that the fuel has an average formula of C7H16 , 15% of the heat generated from combustion goes to heat the water (the rest is lost to the surroundings), the density of the fuel is 0.78 g/mL, the initial temperature of the water is 25.0 C, and the standard enthalpy of formation of C7H16 is 224.4 kJ/mol.

An ice cube of mass 9.0 g is added to a cup of coffee. The coffees initial temperature is 90.0 C and the cup contains 120.0 g of liquid. Assume the specific heat capacity of the coffee is the same as that of water. The heat of fusion of ice (the heat associated with ice melting) is 6.0 kJ/mol. Find the temperature of the coffee after the ice melts.

Find H , E , q , and w for the freezing of water at 10.0 C. The specific heat capacity of ice is 2.04 J/g # C and its heat of fusion (the quantity of heat associated with melting) is 332 J/g. 12

Starting from the relationship between temperature and kinetic energy for an ideal gas, find the value of the molar heat capacity of an ideal gas when its temperature is changed at constant volume. Find its molar heat capacity when its temperature is changed at constant pressure.

An amount of an ideal gas expands from 12.0 L to 24.0 L at a constant pressure of 1.0 atm. Then the gas is cooled at a constant volume of 24.0 L back to its original temperature. Then it contracts back to its original volume. Find the total heat flow for the entire process.

The heat of vaporization of water at 373 K is 40.7 kJ/mol. Find q , w , E , and H for the evaporation of 454 g of water at this temperature at 1 atm.

Find E , H , q , and w for the change in state of 1.0 mol H2O(l) at 80 C to H2O(g) at 110 C. The heat capacity of H2O(l) = 75.3 J/mol K , heat capacity of H2O(g) = 25.0 J/mol K , and the heat of vaporization of H2O is 40.7 * 103 J/mol at 100 C. 12

The heat of combustion of liquid octane \(\left(\mathrm{C}_{8} \mathrm{H}_{18}\right)\) to carbon dioxide and liquid water at 298 K is -1303 kJ/mol. Find \(\Delta E\) for this reaction.

Find H for the combustion of ethanol (C2H6O) to carbon dioxide and liquid water from the following data. The heat capacity of the bomb calorimeter is 34.65 kJ/K and the combustion of 1.765 g of ethanol raises the temperature of the calorimeter from 294.33 K to 295.84 K.

Which statement is true of the internal energy of the system and its surroundings following a process in which Esys = +65 kJ? Explain. a. The system and the surroundings both lose 65 kJ of energy. b. The system and the surroundings both gain 65 kJ of energy. c. The system loses 65 kJ of energy and the surroundings gain 65 kJ of energy. d. The system gains 65 kJ of energy and the surroundings lose 65 kJ of energy.

The internal energy of an ideal gas depends only on its temperature. Which statement is true of an isothermal (constanttemperature) expansion of an ideal gas against a constant external pressure? Explain. a. E is positive b. w is positive c. q is positive d. E is negative

Which expression describes the heat evolved in a chemical reaction when the reaction is carried out at constant pressure? Explain. a. E - w b. E c. E - q 1

Two identical refrigerators are plugged in for the first time. Refrigerator A is empty (except for air) and refrigerator B is filled with jugs of water. The compressors of both refrigerators immediately turn on and begin cooling the interiors of the refrigerators. After 2 hours, the compressor of refrigerator A turns off while the compressor of refrigerator B continues to run. The next day, the compressor of refrigerator A can be heard turning on and off every few minutes, while the compressor of refrigerator B turns off and on every hour or so (and stays on longer each time). Explain these observations.

A 1 kg cylinder of aluminum and 1 kg jug of water, both at room temperature, are put into a refrigerator. After 1 hour, the temperature of each object is measured. One of the objects is much cooler than the other. Which one is cooler and why?

Two substances A and B, initially at different temperatures, are thermally isolated from their surroundings and allowed to come into thermal contact. The mass of substance A is twice the mass of substance B, but the specific heat capacity of substance B is four times the specific heat capacity of substance A. Which substance will undergo a larger change in temperature?

When 1 mol of a gas burns at constant pressure, it produces 2418 J of heat and does 5 J of work. Determine E , H , q , and w for the process.

In an exothermic reaction, the reactants lose energy and the reaction feels hot to the touch. Explain why the reaction feels hot even though the reactants are losing energy. Where does the energy come from?

Which statement is true of a reaction in which V is positive? Explain. a. H = E b. H 7 E c. H 6 E Answers