Chapter 6 Thermodynamics

What are standard conditions

298K and 1atm

heat of vaporization of water in kj/mol

40.7 kj/mol

heat of fusion of water

6.02 kJ/mol

A reaction occurs that results in a set of products with more stable bonds and more orderly arrangement than were present in the reactants. Which of the following is true of this reaction? A. The enthalpy and entropy changes are negative. B. The enthalpy change is positive, and the entropy change is negative. C. The enthalpy change is negative, and the entropy change is positive. D. The enthalpy and entropy changes are positive.

A

A substance has a standard heat of formation of 0. This is best explained by which of the following statements? A. The substance is an element in its standard state. B. At low temperatures, entropy makes virtually no contribution to changes in Gibbs free energy. C. ΔH is a state function. D. The bond energies within the substance are very negative.

A

According to the laws of thermodynamics, an endothermic reaction is most likely to be spontaneous if the: A. change in entropy is positive and the temperature is high. B. change in entropy is negligible and the temperature is low. C. change in entropy is positive and the temperature is low. D. change in entropy is negative and the temperature is high.

A

Consider the following reaction: NH3(g) + HCl(g) → NH4Cl(g) ΔG° = -91.1 kJ Which of the following is a true statement? A. None of the other answers are correct. B. The entropy change of the reaction is positive. C. The reaction is always spontaneous. D. The reaction is never spontaneous.

A

Given the enthalpies of formation for NO and NO2 are 90 kJ/mol and 33 kJ/mol, respectively, what is the reaction enthalpy for 2 NO(g) + O2(g) ? 2 NO2(g)? A. -114 kJ/mol B. -57 kJ/mol C. -24 kJ/mol Your Answer D. Inadequate information provided to answer the question

A

The formation of CO2(g) from the reaction of CO(g) and O2(g) is spontaneous at room temperature. What must also be true for this reaction? A. It is exothermic. B. It is endergonic. C. It is endothermic. D. It is isentropic.

A

Which of the following lists hydrogen halides in terms of increasing standard heats of formation? A. HF < HCl < HBr < HI B. HI < HBr < HCl < HF C. HBr < HF < HCl < HI D. HF < HI < HCl < HBr

A

onsider the following reactions: Which reaction proceeds at a faster rate? Reaction 1: S(s) + O2(g) → SO2(g) ΔG° = -300.1 kJ Reaction 2: Cu2S(s) → 2 Cu(s) + SO2(g) ΔG° = -217.3 kJ A. The reaction rates cannot be determined with the information provided. B. Reaction 2 C. Reaction 1 D. Reaction 1 and Reaction 2 proceed at the same rate.

A

The combustion of methanol is given by the reaction: 2 CH3OH (g) + 3O2 (g) —> 2 CO2 (g) + 4 H2O (g) Delta H = -1352 kJ A. How much heat is produced when 16 grams of O2 reacts with excess methanol? B. Is the reaction exothermic or endothermic? C. How many moles of carbon dioxide are produced when 676 kJ of heat is produced?

A: 225 KJ B: Exothermic C. 1 mole

For the endothermic reaction: 2 CO2 (g) —> 2 CO (g) + O2 (g), which of the following is true? A. Delta H is positive, and delta S is positive B. Delta H is positive, and delta S is negative C. Delta H is negative, and delta S is positive D. Delta H is negative, and delta S is negative

A; Reaction is endothermic, which makes delta H positive. This eliminates C and D. There are two moles of gases on the left and 3 moles of gas on the right; this represents and increase in entropy so entropy is positive. Answer A is correct

Given the following enthalpies of formation, what is the reaction enthalpy for C3H4(g) + 4 O2(g) → 3 CO2(g) + 2 H2O(l)? ΔH°f for C3H4(g) 185 kJ/mol ΔH°f for CO2(g) -286 kJ/mol ΔH°f for H2O(l) -394 kJ/mol A. -1831 kJ/mol B. -1461 kJ/mol C. -495 kJ/mol D. 865 kJ/mol

A; The reaction enthalpy can be calculated by taking the sum of the enthalpies of formation of the products and subtracting from it the sum of the enthalpies of formation of the reactants after multiplying each by their stoichiometric coefficients. The enthalpy of formation for elements in their standard state is defined as zero (hence we will not need it for our solution).

Consider the following reaction: CH4 + 2 O2 → CO2 + 2 H2O ΔH° = -860 kJ What is the best explanation for why the reaction is exothermic? A. CO2 is less stable than CH2. B. The energy required to break reactant bonds is less than the energy released when product bonds are formed. C. More bonds are broken in the reactants than are formed in the products. D. The O=O bond is stronger than the C=O bond, and the C—H bond is stronger than the H—O bond.

B

During the electrolysis of liquid water into hydrogen and oxygen gas at standard temperature and pressure, energy is: A. absorbed during the breaking of H—H bonds and the reaction is spontaneous. B. released during the formation of H—H bonds and the reaction is nonspontaneous. C. absorbed during the formation of O=O bonds and the reaction is spontaneous. D. released during the breaking of O—H bonds and the reaction is nonspontaneous.

B

Given the following bond energies, what is the enthalpy of the combustion of methane? Bond Dissociation Energy H-C 414 kJ/mol H-O 460 kJ/mol O-O 146 kJ/mol O=O 497 kJ/mol C=O 803 kJ/mol A. -917 kJ/mol B. -796 kJ/mol C. 796 kJ/mol D. 917 kJ/mol

B

If a chemical reaction is found to be spontaneous under a given set of conditions, then which of the following must be negative? A.Cp B. ΔG C. ΔH D. ΔS

B

Of the following reactions, which would have the greatest positive entropy change A. 2 NO2 (g) + O2 (g) —> 2 NO (g) B. 2 HCl (aq) + Mg(s) —> MgCl (aq) + H2 (g) C. 2 H2O (g) + Br2 (g) + SO2(g) —> 2 HBr (g) + H2SO4 (aq) D. 2 I- (aq) + Cl2 (g) —> I2 (s) + 2 Cl- (aq)

B

What could make the following nonspontanoeus endothermic reaction spontaneous? 2 H2O (l) --> 2 H2(g) + O2 (g) A. Decreasing volume B. Increasing temperature C. Decreasing temperature D. The reaction will always be nonspontaneous.

B

Which of the following processes would have a negative delta S? A. The evaporation of a liquid B. The freezing of a liquid C. The melting of a solid D. The sublimation of a solid

B

Which of the following reactions results in the greatest increase in ΔS? A. N2(g) + O2(g) → 2 NO(g) B. CO2(s) → CO2(g) C. 2 C3H8O(l) + 9 O2(g) → 6 CO2(g) + 8 H2O(l) D. NH4NO3(s) → NH4NO3(aq)

B

f it is discovered that a certain nonspontaneous reaction becomes spontaneous is the temperature is lowered, then which of the following must be true? A. Delta S is negative, and delta H is positive B. Delta S is negative, and delta H is negative C. Delta S is positive, delta H is positive D. Delta S is positive, delta H is negative

B

A chemist wishes to determine the energy of a chlorine-chlorine bond. She knows the strength of a H—Cl bond and a H—H bond are 433 kJ/mol and 436 kJ/mol, respectively. Given the following reaction, what is the strength of a chlorine-chlorine bond?H2(g) + Cl2(g) → 2 HCl(g) ΔH° = -187 kJ/mol A. 187 kJ/mol B. 243 kJ/mol C. 676 kJ/mol D. 1056 kJ/mol

B;

At 100 K, a certain reaction has the following values for ΔH and ΔS: ΔH = 25 kJ and ΔS = 50 J/K. Neglecting any variation in ΔH and ΔS, at what temperature will ΔG = 0? A. 100 K B. 500 K C. 1000 K D. 250 K

B;

Metal-organic frameworks (MOFs) are systems of metals tethered together by organic linkers forming porous solids with exposed, active metal centers. MOFs are often used to bind and hold gases such as H2, often allowing for a more dense packing of gas than allowed in the pure liquid state. If the binding of gas to metals in the MOF is spontaneous, which must be true? A. The enthalpy of binding is positive. B. The enthalpy of binding is negative. C. The change in free energy associated with binding is positive. D. The entropy of reaction is positive.

B; MOFs are often used to bind and hold gases such as H2, often allowing for a more dense packing of gas than allowed in the pure liquid state. If the binding of gas to metals in the MOF is spontaneous, the enthalpy of binding is negative. As we know the reaction is spontaneous (i.e., the change in free energy is negative), and the change in entropy is certainly negative (as gas molecules are transferred to the solid state), the enthalpy of the reaction must be negative to satisfy ΔG = ΔH - TΔS.

A graduate student estimates the chemical energy in a solution before and after a reaction. The chemical potential energy following the reaction is significantly reduced and the student hypothesizes that the unaccounted energy transitioned to thermal energy. Which of the following laws of thermodynamics best explains this observation? A. Zeroth law B. First law C. Second law D. Third law

B; The first law of thermodynamics stipulates that energy is always conserved and in this example, it is simply changing form from chemical energy to thermal energy (choice B is correct). The zeroth law states that two systems at equilibrium with a third are in equilibrium with each other (choice A is incorrect). The second law states that all processes proceed toward disorder (choice C is incorrect) and the third law states that a perfect crystal possesses no entropy at absolute zero (choice D is incorrect).

Given the following reactions, what is the enthalpy for the reaction above? C2H2(g) + 2 H2(g) → C2H6(g) ΔH° = -94.5 kJ/mol 2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g) ΔH° = -566 kJ/mol 2 H2O(g) → 2 H2(g) + O2(g) ΔH° = 142.4 kJ/mol A. -519 kJ/mol B. 470 kJ/mol C. 519 kJ/mol D. 1036 kJ/mol

B; This question can be solved utilizing Hess's Law, which allows for determination of the reaction enthalpy by summing the enthalpies of a series of other reactions. In order for the given reactions to be added to give the desired equation, we must reverse all three equations and multiply the first and third equations by two. This results in ΔH° values for the three equations of 189 kJ/mol, 566 kJ/mol, and -284.8 kJ/mol, respectively. These can be summed to give the reaction enthalpy of 470 kJ/mol (choice B is correct).

A spontaneous reaction can be made non-spontaneous if: A. ΔH > 0, ΔS > 0, and the temperature is raised. B. ΔH < 0, ΔS < 0, and the temperature is raised. C. ΔH > 0, ΔS < 0, and the temperature is lowered. D. ΔH < 0, ΔS > 0, and the temperature is lowered.

B; ΔG is the indicator of spontaneity. If ΔG is negative, the reaction is spontaneous, while if it is positive, the reaction is non-spontaneous. The equation that describes this is: ΔG = ΔH - TΔS A negative ΔH and positive ΔS will always be spontaneous at any temperature, while a positive ΔH and negative ΔS will always be non-spontaneous at any temperature (eliminate choices C and D). A positive ΔHand positive ΔS will be non-spontaneous at low temperatures, and will become spontaneous as temperature increases (i.e. increasing the contribution of the entropy term), so choice A can be eliminated. A negative ΔH and negative ΔS will be spontaneous at low temperatures, and will become non-spontaneous as temperature increases, thus making choice B the best answer.

Calculate the standard enthalpy change for the reaction CaCO3(s) CaO(s) + CO2(g) given the standard heats of formation: CaCO3(s) = -1209.6 kj/mol CaO(s) = -635.5 kj/mol CO2(g) = -393.5 A. 570 kJ/mol B. 178 kJ/mol C. -570 kJ/mol D. -178 kJ/mol

B; ΔHrxn is the difference between the total ΔHf of the products and the total ΔHf of the reactants. Thus, ΔHrxn = Σ(ΔHf, products) - Σ(ΔHf, reactants) = (-635.5 kJ/mol - 393.5 kJ/mol) - (-1206.9 kJ/mol) = 178 kJ/mol

A 36 gram sample of water requires 93.4 kJ to sublime. What are the heats of fusion and vaporization for water? A. ∆Hfus = -20 kJ/mol, DHvap = 66.7 kJ/mol B. ∆Hfus = 40.7 kJ/mol, DHvap = 6.0 kJ/mol C. ∆Hfus = 6.0 kJ/mol, DHvap = 40.7 kJ/mol D. ∆Hfus = 12.0 kJ/mol, DHvap = 81.4 kJ/mol

C

At 298 K, ΔH° = 436 kJ and ΔS° = 100 J/K for the reaction H2(g) → 2 H(g). Therefore, this reaction: A. is at equilibrium at 298 K. B. is spontaneous at 298 K. C. is not spontaneous at 298 K but will be spontaneous at high enough temperatures D. will never be spontaneous, regardless of the temperature.

C

At STP, a certain liquid will spontaneously vaporize, even though the reaction is endothermic. The most likely explanation for this is that: A. change in enthalpy has no effect on reaction spontaneity. B. temperature has no effect on reaction spontaneity. C. this reaction results in an increase in entropy. D. endothermic reactions are always spontaneous.

C

Compared to uncatalyzed reactions, the Ea - ΔHrxn for catalyzed reactions is: A. higher because catalysts increase Ea. B. higher because catalysts decrease ΔHrxn. C. lower because catalysts decrease Ea. D. lower because catalysts increase ΔHrxn.

C

If H denotes enthalpy and S denotes entropy, then a chemical system will always react spontaneously when: A. ΔH > 0 and ΔS < 0. B. ΔH < 0 and ΔS < 0. C. ΔH < 0 and ΔS > 0. D. ΔH > 0 and ΔS > 0.

C

What must be true about spontaneous, endothermic reaction? A. Delta H is negative B. Delta G is positive C. Delta S is positive D. Delta S is negative

C

Determine ΔH for the reaction CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l) given the following information: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) ΔH = -802 kJ/mol 2 H2O(g) 2 H2O(l) ΔH = -88 kJ/mol A. 714 kJ/mol B. -714 kJ/mol C. -890 kJ/mol D. 890 kJ/mol

C; -890 kJ/mol is ΔH for the reaction CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l) given the following information: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) ΔH = -802 kJ/mol 2 H2O(g) 2 H2O(l) ΔH = -88 kJ/mol The overall reaction is just the sum of the two individual reactions. Therefore the overall ΔHrxn is (-802 kJ/mol) + (-88 kJ/mol) = -890 kJ/mol.

Which of the following substances does NOT have a heat of formation equal to zero at standard conditions? A. F2 (g) B. Cl (2) C. Br2 (g) D. I2 (s)

C; Br2 is a liquid under standard conditions

A researcher investigates an endothermic reaction found to be non-spontaneous in a temperature-independent manner. Which of the following is the reaction being studied? A. 2 H2O(g) → 2 H2(g) + O2(g) B. CO2(s) → CO2(l) C. N2(g) + 2 O2(g) → 2 NO2(g) D. N2O4(g) → 2 NO2(g)

C; The reaction described is non-spontaneous in a temperature-independent manner. As it is endothermic, this indicates that the reaction must have a negative ΔS given ΔG = ΔH - TΔS. The only reaction displaying a negative change in entropy is N2(g) + 2 O2(g) → 2 NO2(g) because the total number of gaseous moles decreases (choice C is correct, and choices A and D are not). In addition, choice B can be eliminated since as a substance changes phase from solid to liquid to gas, the particles are more disordered, so entropy increases.

Given the following reactions, what is the reaction enthalpy for 3 C(graphite) + 4 H2(g) → C3H8(g)? C3H8(g) + 2 O2 → 3 C(graphite) + 4 H2O(l) ΔH° = -1036 kJ/mol H2(g) + ½ O2(g) → H2O(l) ΔH° = -285 kJ/mol A. -1321 kJ/mol B. -751 kJ/mol C. -104 kJ/mol D. 75 kJ/mol

C; This question can be solved utilizing Hess's Law, which allows for determination of the reaction enthalpy by summing the enthalpies of a series of other reactions. In order for the given reactions to be added to give the desired equation, we must reverse the first reaction and multiply the second reaction by four. This results in ΔH° values of 1036 kJ/mol and -1140 kJ/mol, respectively. These can then be summed to give the reaction enthalpy of -104 kJ/mol (choice C is correct).

When a sample of solid NH4NO3 is dissolved in water, the reaction flask becomes cold to the touch. What can be concluded about the given thermodynamic values for the solvation process? A. ΔG < 0, ΔH < 0, ΔS > 0 B. ΔG > 0, ΔH < 0, ΔS > 0 C. ΔG < 0, ΔH > 0, ΔS > 0 D. ΔG < 0, ΔH > 0, ΔS < 0

C; When a sample of solid N4NO3 is dissolved in water, the reaction flask becomes cold to the touch. In can be concluded that ΔG < 0, ΔH > 0, ΔS > 0 for the solvation process. Since the solvation process is observed to occur, it is spontaneous, so ΔG < 0. If the reaction flask is cold to the touch, then the enthalpy change of solvation is positive (ΔH > 0) as this is an endothermic process. Finally, dissolving a solid in a liquid tremendously increases the disorder of the two components as they mix, so ΔS > 0.

A gas is observed to undergo condensation. Which of the following is true about the process? A. Delta H is positive, delta S is positive B. Delta H is positive, and delta S is negative C. Delta H is negative, and delta S is positive D. Delta H is negative, and delta S is negative

D

In electrochemistry, which of the following best demonstrates the first law of thermodynamics? A. A galvanic cell proceeding spontaneously B. The overall increase in entropy observed following a reaction C. The chemical reaction proceeding to a more probable state D. Conversion of equal quantities of chemical energy into electrical energy

D

N2(g) + 3 H2(g) → 2 NH3(g), ΔH = -22 kcal/mol If ΔG is negative, then the reaction as written is: A. spontaneous and endothermic. B. spontaneous with an increase in entropy. C. nonspontaneous due to the decrease in entropy. D. spontaneous and exothermic

D

Which of the following must result in a negative free energy change for a reaction? A. The enthalpy change is negative, and the entropy change is negative. B. The entropy change is positive. C. The enthalpy change is negative. D. The enthalpy change is negative, and the entropy change is positive.

D

Which of the following phase changes involves the greatest increase in entropy? A. Vaporization B. Freezing C. Deposition D. Sublimation

D

Which of the following processes does NOT contribute to the change in enthalpy, Delta H RXN, of a chemical reaction. A. Phase Change B. Formation of stronger intermolecular forces C. Breaking covalent bonds D. The presence of a heterogenous catalyst.

D

Which of the following reactions has the most positive ∆Srxn? A. C(diamond) C(graphite) B. 6 H2O(g) 6 H2O(l) C. 6 CO2(g) + 6 H2O(g) C6H12O6(s) D. CO2(s) CO2(g)

D

Which of the following should have the highest enthalpy of vaporization? A. N2 B. Br2 C. Hg D. Al

D

Which of the following would have the biggest decrease in entropy for a gas in a piston cylinder? A. Compression B. Doing work on the gas C. Cooling D. Deposition

D

Perspiration is important in maintaining normal body temperature. Compared to water at its boiling point, which of the following is true about water at normal human body temperature? A. Intermolecular forces are weaker B. Average kinetic energy is greater C. More energy is required for gas expansion D. The heat required for vaporization is higher

D; Perspiration is important in maintaining normal body temperature. Compared to water at its boiling point, it is true the heat required for vaporization of water is higher at normal human body temperature. As water on the skin vaporizes it absorbs energy and cools the body. Temperature is a measure of average kinetic energy, so water has less average kinetic energy at 37°C compared to 100°C. As temperature increases, water requires more and more energy to expand the volume of its gaseous phase. Although the elimination of "more energy is required for gas expansion" might not be obvious, "the heat required for vaporization is higher" is a better answer. Water is much more likely to vaporize at its boiling point than at body temperature. It takes additional energy to go from the liquid to gas phase at lower temperatures. Therefore the heat required to vaporize is higher when water is at lower temperatures.

When a bond is being formed, what is the enthalpy?

Delta H < 0 (exothermic)

When bonds are broken, what is the enthalpy?

Delta H > 0 (endothermic)

What is delta G on a reaction diagram?

Difference in energy between reactants and products

What is the first law of thermodynamics?

Energy cannot be created or destroyed

What is the second law of thermodynamics?

Every energy transfer or transformation increases the entropy of the universe.

What are the surroundings in thermodynamics?

Everything else

If a reaction is thermodynamically favorable, will the reaction take place rapidly

Just because a reaction is thermodynamically favorable does not automatically mean that it will be taking place rapidly.

What is ΔG when ΔH is negative and ΔS positive

Negative (spontaneous)

Is energy needed to make a bond?

No; energy is released

What is ΔG when ΔH is positive and ΔS is negative

Positive (nonspontaneous)

What is enthalpy?

The heat content of a system at constant pressure

What is Hess's Law?

The total enthalpy change of a reaction is independent of the route taken.

Is energy needed to break a bond?

Yes

What is the third law of thermodynamics?

absolute zero cannot be reached

Enthalpy is a measure of heat energy that is released or absorbed with ________ are broken or formed

bonds

When energy flows out of a system into the surroundings, the energy of the system ________ and the energy of the surroundings _______

decreases; increases

A nonspontaneous reaction is

endergonic

If the products have weaker bonds than the reactants, then more energy is put in during the breaking of reactant bonds than is released. In this case, the reaction is ________

endothermic

What is Gibbs free energy?

energy available to do work

What is activation energy?

energy needed to start a reaction

A spontaneous reaction is

exergonic

If the products have stronger bonds than the reactants, then more energy is released in the making of the products and the reaction is __________

exothermic

In what direction does heat flow?

hot to cold

What affect does the transition from solid to liquid have on entropy?

increases

When energy flows into the system, the energy of the system __________ and the energy of the surroundings _________

increases; decreases

What is a system in thermodynamics?

matter under study

What is entropy?

measure of disorder

What is the zeroth law of thermodynamics?

objects are in thermal equilibrium only when their temperatures are equal

Do particles in solution of undissolved liquids have more entopy

particles in solution

What is the standard heat of formation?

the amount of energy required to make one mole of a compound from its constituent elements in their natural or standard state

What is thermodynamics?

the study of energy transformations

Gibbs free energy equation

ΔG = ΔH - TΔS

What is ΔG when ΔH is positive and ΔS is positive

ΔG = ΔH - TΔS Spontaneous at high temps Nonspontaneous at low temps

What is ΔG when ΔH is negative and ΔS is negative

ΔG = ΔH - TΔS Spontaneous at low temps Nonspontanoeus at high temps

standard heat of formation equation

ΔH°(rxn) = (∑n x ΔH°f, products) - (∑n x ΔH°f, reactants)

change in entropy equation

ΔS = S(products) - S(reactants)

First law of thermodynamics equation

ΔU = Q - W ΔU = change in internal energy Q = heat W = work (done by the system)

Enthalpy equation

∆H = Hproducts - Hreactants

O-H = 460 BDE kj/mol O-O = 180 BDE kj/mol O=O = 498 BDE Reaction: 2 H2O2 (aq) --> 2 H2O (l) + O2 (g)

-140 kj/mol

Enthalpies of Formation Example NH3 (g) = -46 Delta Hf (kj/mol) CO2 (g) = -393 Delta Hf (kj/mol) CO (NH2)2 (s) = -333 Delta Hf (kj/mol) H2O = -242 Delta Hf (kj/mol) Reaction: 2 NH3 (g) + CO2 (g) —> CO(NH2)2 (s) + H2O What is the standard heat of formation?

-70 KJ/mol

What is the ΔH° for the following reaction under standard conditions if the ΔH°f of CH4(g) = -75 kj/mol, ΔH° of CO2(g) = -393 kj/mol, and ΔH° of H2O (l) = -286 kj/mol Reaction: CH4(g) + 2O2 (g) ---> CO2 (g) + 2H2O (l)

-895 kj/mol

What is the standard heat of formation of a diatomic element?

0

What is the standard heat of formation of an element in its standard state?

0

All solids and liquids are assumed to be pure, and solutions are considered to be at a concentration of _______ M

1