Chlorofluoromethanes (CFMs) are carbon compounds of chlorine and fluorine and are also known as Freons. Examples are Freon-11 \(\left(\mathrm{CFCl}_{3}\right)\) and Freon-12 \(\left(\mathrm{CF}_{2} \mathrm{Cl}_{2}\right),\) which were used as aerosol propellants. Freons have also been used in refrigeration and air-conditioning systems. In 1995 Mario Molina, F. Sherwood Rowland, and Paul Crutzen were awarded the Nobel Prize mainly for demonstrating how these and other CFMs contribute to the "ozone hole" that develops at the end of the Antarctic winter. In other parts of the world, reactions such as those shown below occur in the upper atmosphere where ozone protects the earth's inhabitants from harmful ultraviolet radiation. In the stratosphere CFMs absorb high-energy radiation from the sun and split off chlorine atoms that hasten the decomposition of ozone, \(\mathrm{O}_{3}\). Possible reactions are $$ \begin{array}{ll} \mathrm{O}_{3}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{O}_{2}(g)+\mathrm{ClO}(g) & \Delta H^{\circ}=-126 \mathrm{~kJ} \\ \mathrm{ClO}(g)+\mathrm{O}(g) \longrightarrow \mathrm{Cl}(g)+\mathrm{O}_{2}(g) & \Delta H^{\circ}=-268 \mathrm{~kJ} \\ \mathrm{O}_{3}(g)+\mathrm{O}(g) \longrightarrow 2 \mathrm{O}_{2}(g) & \Delta H^{\circ}=? \end{array} $$ The \(\mathrm{O}\) atoms in the second equation come from the breaking apart of \(\mathrm{O}_{2}\) molecules caused by ultraviolet radiation from the sun. Use the first two equations to calculate the value of \(\Delta H^{\circ}\) (in kilojoules) for the last equation, the net reaction for the removal of \(\mathrm{O}_{3}\) from the atmosphere.

Step by step solution

01

Write down the given reactions and their enthalpies of reaction

First, we list the two given reactions with their respective \( \Delta H^{\circ} \) values:1. \( \mathrm{O}_{3}(g) + \mathrm{Cl}(g) \longrightarrow \mathrm{O}_{2}(g) + \mathrm{ClO}(g) \) with \( \Delta H^{\circ} = -126 \mathrm{~kJ} \) 2. \( \mathrm{ClO}(g) + \mathrm{O}(g) \longrightarrow \mathrm{Cl}(g) + \mathrm{O}_{2}(g) \) with \( \Delta H^{\circ} = -268 \mathrm{~kJ} \) We want to find the \( \Delta H^{\circ} \) for the reaction:3. \( \mathrm{O}_{3}(g) + \mathrm{O}(g) \longrightarrow 2 \mathrm{O}_{2}(g) \)

02

Manipulate the given reactions to obtain the target reaction

To find the enthalpy change for the target reaction, we need to add or subtract the given reactions in such a way that the final equation only shows the reactants and products of the target reaction. We need to cancel out the \( \mathrm{Cl} \) and \( \mathrm{ClO} \) from the right side of the first reaction and the left side of the second reaction.When we add the two reactions, we get:\[\begin{align*}&\mathrm{O}_{3}(g) + \mathrm{Cl}(g) \longrightarrow \mathrm{O}_{2}(g) + \mathrm{ClO}(g) &\Delta H^{\circ} = -126 \mathrm{~kJ} \&\mathrm{ClO}(g) + \mathrm{O}(g) \longrightarrow \mathrm{Cl}(g) + \mathrm{O}_{2}(g) &\Delta H^{\circ} = -268 \mathrm{~kJ} \end{align*}\]The \( \mathrm{Cl} \) and \( \mathrm{ClO} \) cancel out, and we are left with:\[\mathrm{O}_{3}(g) + \mathrm{O}(g) \longrightarrow 2 \mathrm{O}_{2}(g)\]

03

Calculate the enthalpy change for the target reaction

Now, we sum the enthalpies of reaction for the two given reactions to find the missing enthalpy for the target reaction. Since the reactions are added directly and nothing is reversed or multiplied, we simply add the \( \Delta H^{\circ} \) values:\[ \Delta H^{\circ}_{\text{total}} = \Delta H^{\circ}_{1} + \Delta H^{\circ}_{2} \]So,\[ \Delta H^{\circ}_{\text{total}} = -126 \mathrm{~kJ} + (-268 \mathrm{~kJ}) \]\[ \Delta H^{\circ}_{\text{total}} = -394 \mathrm{~kJ} \]The enthalpy change \( \Delta H^{\circ} \) for the reaction \( \mathrm{O}_{3}(g) + \mathrm{O}(g) \longrightarrow 2 \mathrm{O}_{2}(g) \) is \( -394 \mathrm{~kJ} \)

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Key Concepts

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

Chlorofluoromethanes

Chlorofluoromethanes (CFMs), widely known as Freons, are chemically stable compounds composed of carbon, chlorine, and fluorine. Initially, their applications included aerosol propellants, refrigeration, and air conditioning due to their non-toxic and non-flammable properties. However, the environmental consequences of using CFMs were not immediately recognized.

These compounds are now infamous for their role in ozone depletion. They were designed to be very stable, which was originally seen as a benefit because they wouldn't react with other substances and cause harm. Unfortunately, this stability means they persist in the atmosphere long enough to reach the stratosphere, where they then release chlorine atoms upon exposure to ultraviolet radiation. It's this chlorine that plays a key part in the breakdown of ozone (O₃).

Despite their phased out production under the Montreal Protocol due to their adverse environmental effects, CFMs remain an important subject of study in the field of environmental chemistry and are used as a classic example when teaching about anthropogenic impacts on the ozone layer.

Ozone Depletion

Ozone depletion refers to the thinning and reduction of the ozone layer in Earth's stratosphere. The ozone layer is crucial for life on Earth because it absorbs the majority of the Sun's harmful ultraviolet (UV) radiation. Ozone molecules are composed of three oxygen atoms (O₃) and are naturally formed when oxygen molecules (O₂) are split by UV light, allowing individual oxygen atoms to recombine with O₂ to form O₃.

However, certain substances can accelerate the breakdown of ozone. The most infamous culprits are chlorofluoromethanes, which upon UV radiation exposure release chlorine atoms. These chlorine atoms react with ozone, resulting in the formation of ordinary oxygen molecules and chlorine monoxide (ClO), ultimately reducing the overall ozone concentration. This process is particularly concerning in polar regions, leading to the seasonal formation of 'ozone holes'.

The study of ozone depletion not only encompasses its causes but also its global implications—including increased UV exposure that can lead to higher rates of skin cancer and cataracts, and the impact on climate and ecosystems.

Thermochemical Equations

Thermochemical equations are a type of chemical equation that includes the heat changes of the reaction, expressed through the enthalpy change (\( \textstyle \text{Δ}H \textstyle \text{°} \)). Enthalpy (\textstyle \text{H}) is the measure of the total energy of a thermodynamic system, including both internal energy and the energy required to make room for it by displacing its environment.

When writing thermochemical equations, it's crucial to specify the states of the reactants and products (for example, as gas (g), liquid (l), or solid (s)) and to balance the equation in accordance to the law of conservation of mass. The sign of \textstyle \text{Δ}H indicates whether the reaction is exothermic (-) or endothermic (+), and the magnitude gives the amount of energy released or absorbed.

In the example provided, the problem involves calculating the enthalpy change for the net reaction that removes ozone from the atmosphere. By carefully manipulating the given thermochemical equations for the reactions involving chlorine and ozone, we used the principle of Hess's Law to add the enthalpy changes of individual steps and arrive at the total enthalpy change for the desired reaction. As complex as this process sounds, it's simply a way to apply conservation of energy to chemical reactions—a concept pivotal in the study of thermodynamics and environmental chemistry.