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Chapter 11: Problem 18

Hydrogen reacts explosively with oxygen. However, a mixture of \(\mathrm{H}_{2}\) and \(\mathrm{O}_{2}\) can exist indefinitely at room temperature. Explain why \(\mathrm{H}_{2}\) and \(\mathrm{O}_{2}\) do not react under these conditions.

Short Answer

Expert verified
Hydrogen and oxygen do not react explosively at room temperature because the average kinetic energy of the molecules is not sufficient to overcome the activation energy barrier required for the reaction to occur. In the absence of a catalyst or a higher temperature, hydrogen and oxygen can coexist indefinitely without undergoing an explosive reaction.

Step by step solution

01

Recall characteristics of chemical reactions

In a chemical reaction, reactants need to collide with each other at an appropriate orientation and sufficient energy for the reaction to take place. The minimum energy required for a successful collision is called the activation energy.
02

Understand activation energy and reaction rates

Activation energy is the threshold energy that molecules need to attain in order for a chemical reaction to occur. If the molecules do not possess the required energy, they will not successfully react even if they collide with the proper orientation. The rate of a reaction depends on factors such as temperature, concentration, and presence of a catalyst.
03

Consider room temperature conditions

At room temperature, the average kinetic energy of the hydrogen and oxygen molecules is relatively low. Consequently, most collisions between the molecules do not possess sufficient energy to overcome the activation energy barrier for the explosive reaction to occur.
04

Role of a catalyst

A catalyst can lower the activation energy of a reaction, allowing it to occur more easily at lower temperatures. However, in this scenario, there is no catalyst present to initiate the explosive reaction between hydrogen and oxygen at room temperature.
05

Conclusion

Hydrogen and oxygen do not react explosively at room temperature because the average kinetic energy of the molecules at this temperature is not high enough to overcome the activation energy barrier. In the absence of a catalyst or a higher temperature, these two gases can coexist indefinitely without undergoing a reaction.

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Most popular questions from this chapter

Hydrogen peroxide and the iodide ion react in acidic solution as follows: $$\mathrm{H}_{2} \mathrm{O}_{2}(a q)+3 \mathrm{I}^{-}(a q)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{I}_{3}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)$$ The kinetics of this reaction were studied by following the decay of the concentration of \(\mathrm{H}_{2} \mathrm{O}_{2}\) and constructing plots of \(\ln \left[\mathrm{H}_{2} \mathrm{O}_{2}\right]\) versus time. All the plots were linear and all solutions had \(\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]_{0}=8.0 \times 10^{-4} \mathrm{mol} / \mathrm{L} .\) The slopes of these straight lines depended on the initial concentrations of \(\mathrm{I}^{-}\) and \(\mathrm{H}^{+} .\) The results follow: The rate law for this reaction has the form $$\text { Rate }=\frac{-\Delta\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]}{\Delta t}=\left(k_{1}+k_{2}\left[\mathrm{H}^{+}\right]\right)\left[\mathrm{I}^{-}\right]^{m}\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]^{n}$$ a. Specify the order of this reaction with respect to \(\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]\) and \(\left[\mathrm{I}^{-}\right]\) b. Calculate the values of the rate constants, \(k_{1}\) and \(k_{2}\) c. What reason could there be for the two-term dependence of the rate on \(\left[\mathrm{H}^{+}\right] ?\)

A first-order reaction has rate constants of \(4.6 \times 10^{-2} \mathrm{s}^{-1}\) and \(8.1 \times 10^{-2} \mathrm{s}^{-1}\) at \(0^{\circ} \mathrm{C}\) and \(20 .^{\circ} \mathrm{C},\) respectively. What is the value of the activation energy?

Consider the reaction $$3 \mathrm{A}+\mathrm{B}+\mathrm{C} \longrightarrow \mathrm{D}+\mathrm{E}$$ where the rate law is defined as $$-\frac{\Delta[\mathrm{A}]}{\Delta t}=k[\mathrm{A}]^{2}[\mathrm{B}][\mathrm{C}]$$ An experiment is carried out where \([\mathrm{B}]_{0}=[\mathrm{C}]_{0}=1.00 \space M\) and \([\mathrm{A}]_{0}=1.00 \times 10^{-4} \mathrm{M}\) a. If after \(3.00 \min ,[A]=3.26 \times 10^{-5} M,\) calculate the value of \(k\) b. Calculate the half-life for this experiment. c. Calculate the concentration of \(\mathrm{B}\) and the concentration of A after 10.0 min.

The activation energy for a reaction is changed from \(184 \space\mathrm{kJ} /\) mol to \(59.0 \space\mathrm{kJ} / \mathrm{mol}\) at \(600 .\) K by the introduction of a catalyst. If the uncatalyzed reaction takes about 2400 years to occur, about how long will the catalyzed reaction take? Assume the frequency factor \(A\) is constant, and assume the initial concentrations are the same.

One of the concerns about the use of Freons is that they will migrate to the upper atmosphere, where chlorine atoms can be generated by the following reaction: $$\mathrm{CCl}_{2} \mathrm{F}_{2}(g) \stackrel{h v}{\longrightarrow} \mathrm{CF}_{2} \mathrm{Cl}(g)+\mathrm{Cl}(g)$$ Chlorine atoms can act as a catalyst for the destruction of ozone. The activation energy for the reaction $$\mathrm{Cl}(g)+\mathrm{O}_{3}(g) \longrightarrow \mathrm{ClO}(g)+\mathrm{O}_{2}(g)$$is \(2.1 \mathrm{kJ} / \mathrm{mol} .\) Which is the more effective catalyst for the destruction of ozone, Cl or NO? (See Exercise 75.)
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