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Chapter 2: Problem 45

Which of the following processes is an irreversible reaction? (A) \(\mathrm{CH}_{4}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)\) (B) \(\mathrm{HCN}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightarrow \mathrm{CN}^{-}(a q)+\mathrm{H}_{3} \mathrm{O}^{+}(a q)\) (C) \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(s) \rightarrow \mathrm{Al}^{3+}(a q)+3 \mathrm{NO}_{3}^{-}(a q)\) (D) \(2 \mathrm{Ag}^{+}(a q)+\mathrm{Ti}(s) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Ti}^{2+}(a q)\)

Short Answer

Expert verified
The irreversible reaction is Option (A).

Step by step solution

01

Analyze the First Reaction

The first reaction is \(\mathrm{CH}_{4}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)\), a combustion reaction. Combustion reactions are generally irreversible under standard conditions as they release energy, and the products do not react to form the original reactants.
02

Analyze the Other Reactions

In the other reactions, they seem to be a form of disassociation or a redox reaction, which could technically run in the reverse direction under the right conditions (providing the opposite ions to the solution, or changing the redox potentials by altering the concentrations). Thus, these reactions could be reversible.
03

Conclusion

Considering all the reactions, the first reaction is the most likely to be completely irreversible under the given conditions, as combustion reactions often are.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Combustion Reactions
When you think about burning something, you're thinking about a combustion reaction. These reactions occur when a substance, typically a hydrocarbon, reacts with oxygen to form carbon dioxide and water.
  • They usually release significant amounts of energy in the form of heat and light.
  • This release of energy makes them exothermic reactions.
Combustion reactions are generally irreversible. This is because once the energy is released, the products (like carbon dioxide and water) don't spontaneously turn back into the original fuel and oxygen under standard conditions. Combustion of methane (CH4) is a classic example, and this is precisely what you've seen in Reaction A of the exercise.
Redox Reactions
Redox reactions are all about the transfer of electrons between species. In these reactions, one species loses electrons (it's oxidized), while another gains electrons (it's reduced).
  • An easy way to remember: OIL RIG鈥擮xidation Is Loss, Reduction Is Gain.
  • Redox reactions can be as simple as combining substances or as complex as involving multiple steps.
The reaction seen in Reaction D from the exercise is a redox reaction. Here, silver ions are reduced to solid silver, while titanium is oxidized to form titanium ions. Depending on the components and conditions, redox reactions can sometimes be reversed, unlike combustion reactions.
Chemical Equilibrium
Chemical equilibrium is a crucial concept, where the forward and reverse reactions occur at the same rate, so the concentrations of reactants and products remain constant over time.
  • It's not that the reactions stop; rather, they proceed at the same rate in both directions.
  • This state of balance doesn't mean equal amounts of reactants and products, but rather a constant ratio between them.
Equilibrium conditions are often seen in reactions that can easily reverse, such as dissociation reactions. In Reaction B from the exercise, the dissociation of HCN in water is reversible, hence governed by equilibrium dynamics, unlike the irreversible nature of combustion.
Reaction Types
Understanding the various types of reactions is key to mastering chemistry. Here are a few:
  • Combustion: Involves burning a substance in oxygen.
  • Redox: Based on the transfer of electrons among substances.
  • Dissociation: Typically involves an ionic compound dissociating into its ions.
In the given exercise, combustion is highlighted by the burning of methane with oxygen. Redox is well-exemplified by the reaction involving silver and titanium ions. Meanwhile, dissociation is demonstrated by the breakdown of Al(NO3)3 into aluminum and nitrate ions in Reaction C. Recognizing these types helps identify reaction properties like reversibility, energy change, and the nature of reactants and products.

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Most popular questions from this chapter

Which of the following pairs of ions would make the best buffer with a pH between 6 and 7? \(K_{\mathrm{a}}\) for \(\mathrm{HC}_{3} \mathrm{H}_{2} \mathrm{O}_{2}=1.75 \times 10^{-5}\) \(K_{\mathrm{a}}\) for \(\mathrm{HPO}_{4}^{2-}=4.8 \times 10^{-13}\) (A) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}\) (B) \(\mathrm{HPO}_{4}^{2-}\) and \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) (C) \(\mathrm{HC}_{3} \mathrm{H}_{2} \mathrm{O}_{2}\) and \(\mathrm{NaC}_{3} \mathrm{H}_{2} \mathrm{O}_{2}\) (D) \(\mathrm{NaOH}\) and \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\)

A sample of \(\mathrm{H}_{2} \mathrm{S}\) gas is placed in an evacuated, sealed container and heated until the following decomposition reaction occurs at \(1000 \mathrm{K} :\) $$2 \mathrm{H}_{2} \mathrm{S}(g) \rightarrow 2 \mathrm{H}_{2}(g)+\mathrm{S}_{2}(g) \qquad K_{\mathrm{c}}=1.0 \times 10^{-6}$$ If, at a given point in the reaction, the value for the reaction quotient \(Q\) is determined to be \(2.5 \times 10^{-8},\) which of the following is occurring? (A) The concentration of the reactant is decreasing while the concentration of the products is increasing. (B) The concentration of the reactant is increasing while the concentration of the products is decreasing. (C) The system has passed the equilibrium point, and the concentration of all species involved in the reaction will remain constant. (D) The concentrations of all species involved are changing at the same rate.

20.0 \(\mathrm{mL}\) of 1.0 \(\mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3}\) is placed in a beaker and titrated with a solution of \(1.0 \mathrm{M} \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2},\) resulting in the creation of a precipitate. How much \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\) must be added to reach the equivalence point? (A) 10.0 mL (B) 20.0 mL (C) 30.0 mL (D) 40.0 mL

The enthalpy values for several reactions are as follows: (I) \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2}(g) \rightarrow \mathrm{C}(s)+\mathrm{H}_{2} \mathrm{O}(g)\) \(\quad \Delta H=-131 \mathrm{kJ} / \mathrm{mol}_{\mathrm{rxn}}\) (II) \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow 3 \mathrm{H}_{2}(g)+\mathrm{CO}(g)\) \(\quad \Delta H=206 \mathrm{kJ} / \mathrm{mol}_{\mathrm{rxn}}\) (III) \(\mathrm{CO}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2}(g)\) \(\quad \Delta H=-41 \mathrm{kJ} / \mathrm{mol}_{\mathrm{rxn}}\) (IV) \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l)\) \(\quad \Delta H=-890 \mathrm{kJ} / \mathrm{mol}_{\mathrm{rxn}}\) In which of the reactions does the amount of energy released by the formation of bonds in the products exceed the amount of energy necessary to break the bonds of the reactants by the greatest amount? (A) Reaction I (B) Reaction II (C) Reaction III (D) Reaction IV

A sample of liquid butane \(\left(\mathrm{C}_{\mathrm{L}} \mathrm{H}_{10}\right)\) in a pressurized lighter is set up directly beneath an aluminum can, as show in the diagram above. The can contains 100.0 \(\mathrm{mL}\) of water, and when the butane is ignited the temperature of the water inside the can increases from \(25.0^{\circ} \mathrm{C}\) to \(82.3^{\circ} \mathrm{C}\) . The total mass of butane ignited is found to be 0.51 \(\mathrm{g}\) , the specific heat of water is \(4.18 \mathrm{J} / \mathrm{g} \cdot^{\circ} \mathrm{C},\) and the density of water is \(1.00 \mathrm{g} / \mathrm{mL} .\) (a) Write the balanced chemical equation for the combustion of one mole of butane in air. (i) How much heat did the water gain? (ii) What is the experimentally determined heat of combustion for (ii) Whane based on this experiment? Your answer should be in \(\mathrm{kJ} / \mathrm{mol}\) . (c) Given butane's density of 0.573 \(\mathrm{g} / \mathrm{mL}\) at \(25^{\circ} \mathrm{C},\) calculate how much heat would be emitted if 5.00 \(\mathrm{mL}\) of it were combusted at that temperature. (d) The overall combustion of butane is an exothermic reaction. Explain why this is, in terms of bond energies. (e) One of the major sources of error in this experiment comes from the heat that is aboorbed by the air. Why, then, might it not be a good ide to perform this experiment inside a sealed container to prevent the heat from leaving the system?
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