van 't Hoff equation

From Wikipedia, the free encyclopedia

The van 't Hoff equation relates the change in the equilibrium constant, K_eq, of a chemical reaction to the change in temperature, T, given the standard enthalpy change, ΔH^⊖, for the process. It was proposed by Dutch chemist Jacobus Henricus van 't Hoff in 1884 in his book Études de dynamique chimique (Studies in Dynamic Chemistry).^[1] This equation is sometimes also referred to as the Vukančić–Vuković equation.^[2][3][4]

The van 't Hoff equation has been widely utilized to explore the changes in state functions in a thermodynamic system. The van 't Hoff plot, which is derived from this equation, is especially effective in estimating the change in enthalpy, or total energy, and entropy, or amount of disorder, of a chemical reaction

Equation

Under standard conditions

Under standard conditions, the van 't Hoff equation is^[5][6]

${\displaystyle {\frac {d}{dT}}\ln K_{\mathrm {eq} }={\frac {\Delta H^{\ominus }}{RT^{2}}}}$

where R is the ideal gas constant. This equation is exact at any one temperature. In practice, the equation is often integrated between two temperatures under the assumption that the reaction enthalpy ΔH is constant. Since in reality ΔH and ΔS do vary with temperature for most processes,^[7] the integrated equation is only approximate.

A major use of the integrated equation is to estimate a new equilibrium constant at a new absolute temperature assuming a constant standard enthalpy change over the temperature range.

To obtain the integrated equation, it is convenient to first rewrite the van 't Hoff equation as^[5]

${\displaystyle {\frac {d\ln K_{\mathrm {eq} }}{d{\frac {1}{T}}}}=-{\frac {\Delta H^{\ominus }}{R}}.}$

The definite integral between temperatures T₁ and T₂ is then

${\displaystyle \ln {\frac {K_{2}}{K_{1}}}={\frac {-\Delta H^{\ominus }}{R}}\left({\frac {1}{T_{2}}}-{\frac {1}{T_{1}}}\right).}$

In this equation K₁ is the equilibrium constant at absolute temperature T₁, and K₂ is the equilibrium constant at absolute temperature T₂.

Development from thermodynamics

Combining the definition of Gibbs free energy:

${\displaystyle \Delta G^{\ominus }=\Delta H^{\ominus }-T\Delta S^{\ominus },}$

where S is the entropy of the system, and the Gibbs free energy isotherm equation:^[8]

${\displaystyle \Delta G^{\ominus }=-RT\ln K_{\mathrm {eq} },}$

we obtain

${\displaystyle \ln K_{\mathrm {eq} }=-{\frac {\Delta H^{\ominus }}{RT}}+{\frac {\Delta S^{\ominus }}{R}}.}$

Differentiation of this expression with respect to the variable 1/T yields the van 't Hoff equation.

Provided that ΔH^⊖ and ΔS^⊖ are constant, the preceding equation gives ln K as a linear function of 1/T and hence is known as the linear form of the van 't Hoff equation. Therefore, when the range in temperature is small enough that the standard enthalpy and entropy changes are essentially constant, a plot of the natural logarithm of the equilibrium constant versus the reciprocal temperature gives a straight line. The slope of the line may be multiplied by the gas constant R to obtain the standard enthalpy change of the reaction, and the intercept may be multiplied by R to obtain the standard entropy change.

van 't Hoff isotherm

The Gibbs free energy changes with the temperature and pressure of the thermodynamic system. The van 't Hoff isotherm can be used to determine the Gibbs free energy for non-standard state reactions at a constant temperature:^[9]

${\displaystyle \left({\frac {dG}{d\xi }}\right)_{T,p}=\Delta _{\mathrm {r} }G+RT\ln Q_{\mathrm {r} },}$

where Δ_rG is the Gibbs free energy for the reaction, and Q_r is the reaction quotient. When a reaction is at equilibrium, Q_r = K_eq. The van 't Hoff isotherm can help estimate the equilibrium reaction shift. When Δ_rG < 0, the reaction moves in the forward direction. When Δ_rG > 0, the reaction moves in the backwards directions.

van 't Hoff plot

For a reversible reaction, the equilibrium constant can be measured at a variety of temperatures. This data can be plotted on a graph with ln K_eq on the y-axis and 1/T on the x axis. The data should have a linear relationship, the equation for which can be found by fitting the data using the linear form of the van 't Hoff equation

${\displaystyle \ln K_{\mathrm {eq} }=-{\frac {\Delta H^{\ominus }}{RT}}+{\frac {\Delta S^{\ominus }}{R}}.}$

This graph is called the van 't Hoff plot and is widely used to estimate the enthalpy and entropy of a chemical reaction. From this plot, −ΔH/R is the slope, and ΔS/R is the intercept of the linear fit.

By measuring the equilibrium constant, K_eq, at different temperatures, the van 't Hoff plot can be used to assess a reaction when temperature changes.^[10][11] Knowing the slope and intercept from the van 't Hoff plot, the enthalpy and entropy of a reaction can be easily obtained using

${\displaystyle {\begin{aligned}\Delta H&=-R\times {\text{slope}},\\\Delta S&=R\times {\text{intercept}}.\end{aligned}}}$

The van 't Hoff plot can be used to quickly determine the enthalpy of a chemical reaction both qualitatively and quantitatively. Change in enthalpy can be positive or negative, leading to two major forms of the van 't Hoff plot.

Endothermic reactions

van 't Hoff plot for an endothermic reaction

For an endothermic reaction, heat is absorbed, making the net enthalpy change positive. Thus, according to the definition of the slope:

${\displaystyle {\text{slope}}=-{\frac {\Delta H}{R}},}$

for an endothermic reaction, ΔH > 0 (and the gas constant R > 0), so

${\displaystyle {\text{slope}}=-{\frac {\Delta H}{R}}<0 .="" annotation="">$

Thus, for an endothermic reaction, the van 't Hoff plot should always have a negative slope.

Exothermic reactions

van 't Hoff plot for an exothermic reaction

For an exothermic reaction, heat is released, making the net enthalpy change negative. Thus, according to the definition of the slope:

${\displaystyle {\text{slope}}=-{\frac {\Delta H}{R}},}$

fom an exothermic reaction, ΔH < 0, so

${\displaystyle {\text{slope}}=-{\frac {\Delta H}{R}}>0.}$

Thus, for an exothermic reaction, the van 't Hoff plot should always have a positive slope.

Error propagation

Using the fact that ΔG^⊖ = −RT ln K = ΔH^⊖ − TΔS^⊖, it would appear that two measurements of K would suffice to be able to obtain a value of ΔH^⊖:

${\displaystyle \Delta H^{\ominus }=R{\frac {\ln K_{1}-\ln K_{2}}{{\frac {1}{T_{2}}}-{\frac {1}{T_{1}}}}},}$

where K₁ and K₂ are the equilibrium constant values obtained at temperatures T₁ and T₂ respectively. The precision of ΔH^⊖ values obtained in this way is highly dependent on the precision of the equilibrium constant values. A typical pair of temperatures might be 25 and 35 °C (298 and 308 K). For these temperatures

${\displaystyle {\frac {1}{T_{2}}}-{\frac {1}{T_{1}}}={\frac {1}{298}}-{\frac {1}{308}}=1.1\times 10^{-4}{\text{ K}}.}$

Inserting this value into the expression for ΔH^⊖:

${\displaystyle \Delta H^{\ominus }\approx 76\left(\ln K_{1}-\ln K_{2}\right){\text{ kJ/mol}}.}$

Now, error propagation shows that the error on ΔH^⊖ will be about 76 kJ/mol times the error on (ln K₁ − ln K₂), or about 110 kJ/mol times the error on the ln K values. Assume for example that the error on each ln K is σ ≈ 0.05, a small but reasonable value. The error on ΔH^⊖ will be about 5 kJ/mol. So, even though the individual stability constants were determined with good precision, the enthalpy calculated in this way is subject to a significant error.

The entropy will then be obtained from ΔS^⊖ = 1/T(ΔH^⊖ + RT ln K). In this expression the error on the second term is negligible compared to the error on the first term. The magnifying factor is then 76 kJ/mol ÷ 298 K, so for an error of 0.05 in the logarithms the error on ΔS^⊖ will be of the order of 17 J/(K mol).

When equilibrium constants are measured at three or more temperatures, values of ΔH^⊖ will be obtained by straight-line fitting. In this case the error on the standard enthalpy will be magnified to a somewhat lesser, but still substantial, extent.

Applications of the van 't Hoff plot

van 't Hoff analysis

van 't Hoff analysis

In biological research, the van 't Hoff plot is also called van 't Hoff analysis.^[12] It is most effective in determining the favored product in a reaction.

Assume two products B and C form in a reaction:

a A + d D → b B,

a A + d D → c C.

In this case, K_eq can be defined as ratio of B to C rather than the equilibrium constant.

When B/C > 1, B is the favored product, and the data on the van 't Hoff plot will be in the positive region.

When B/C < 1, C is the favored product, and the data on the van 't Hoff plot will be in the negative region.

Using this information, a van 't Hoff analysis can help determine the most suitable temperature for a favored product.

Recently, a van 't Hoff analysis was used to determine whether water preferentially forms a hydrogen bond with the C-terminus or the N-terminus of the amino acid proline.^[13] The equilibrium constant for each reaction was found at a variety of temperatures, and a van 't Hoff plot was created. This analysis showed that enthalpically, the water preferred to hydrogen bond to the C-terminus, but entropically it was more favorable to hydrogen bond with the N-terminus. Specifically, they found that C-terminus hydrogen bonding was favored by 4.2–6.4 kJ/mol. The N-terminus hydrogen bonding was favored by 31–43 J/(K mol).

This data alone could not conclude which site water will preferentially hydrogen-bond to, so additional experiments were used. It was determined that at lower temperatures, the enthalpically favored species, the water hydrogen-bonded to the C-terminus, was preferred. At higher temperatures, the entropically favored species, the water hydrogen-bonded to the N-terminus, was preferred.

Mechanistic studies

van 't Hoff plot in mechanism study

A chemical reaction may undergo different reaction mechanisms at different temperatures.^[14]

In this case, a van 't Hoff plot with two or more linear fits may be exploited. Each linear fit has a different slope and intercept, which indicates different changes in enthalpy and entropy for each distinct mechanisms. The van 't Hoff plot can be used to find the enthalpy and entropy change for each mechanism and the favored mechanism under different temperatures.

${\displaystyle {\begin{aligned}\Delta H_{1}&=-R\times {\text{slope}}_{1},&\Delta S_{1}&=R\times {\text{intercept}}_{1};\\[5pt]\Delta H_{2}&=-R\times {\text{slope}}_{2},&\Delta S_{2}&=R\times {\text{intercept}}_{2}.\end{aligned}}}$

In the example figure, the reaction undergoes mechanism 1 at high temperature and mechanism 2 at low temperature.

Temperature dependence

Temperature-dependent van 't Hoff plot

The van 't Hoff plot is linear based on the assumption that the enthalpy and entropy are constant with temperature changes. However, in some cases the enthalpy and entropy do change dramatically with temperature. A first-order approximation is to assume that the two different reaction products have different heat capacities. Incorporating this assumption yields an additional term c/T² in the expression for the equilibrium constant as a function of temperature. A polynomial fit can then be used to analyze data that exhibits a non-constant standard enthalpy of reaction:^[15]

${\displaystyle \ln K_{\mathrm {eq} }=a+{\frac {b}{T}}+{\frac {c}{T^{2}}},}$

where

${\displaystyle {\begin{aligned}\Delta H&=-R\left(b+{\frac {2c}{T}}\right),\\\Delta S&=R\left(a-{\frac {c}{T^{2}}}\right).\end{aligned}}}$

Thus, the enthalpy and entropy of a reaction can still be determined at specific temperatures even when a temperature dependence exists.

Surfactant self-assembly

The van 't Hoff relation is particularly useful for the determination of the micellization enthalpy ΔH^⊖
_m of surfactants from the temperature dependence of the critical micelle concentration (CMC):

${\displaystyle {\frac {d}{dT}}\ln \mathrm {CMC} ={\frac {\Delta H_{\mathrm {m} }^{\ominus }}{RT^{2}}}.}$

However, the relation loses its validity when the aggregation number is also temperature-dependent, and the following relation should be used instead:^[16]

${\displaystyle RT^{2}\left({\frac {\partial }{\partial T}}\ln \mathrm {CMC} \right)_{P}=-\Delta H_{\mathrm {m} }^{\ominus }(N)+T\left({\frac {\partial }{\partial N}}\left(G_{N+1}-G_{N}\right)\right)_{T,P}\left({\frac {\partial N}{\partial T}}\right)_{P},}$

with G_{N + 1} and G_N being the free energies of the surfactant in a micelle with aggregation number N + 1 and N respectively. This effect is particularly relevant for nonionic ethoxylated surfactants^[17] or polyoxypropylene–polyoxyethylene block copolymers (Poloxamers, Pluronics, Synperonics).^[18] The extended equation can be exploited for the extraction of aggregation numbers of self-assembled micelles from differential scanning calorimetric thermograms.^[19]

Chlorofluorocarbon

From Wikipedia, the free encyclopedia

Chlorofluorocarbons (CFCs) are fully halogenated paraffin hydrocarbons that contain only carbon (С), chlorine (Cl), and fluorine (F), produced as volatile derivative of methane, ethane, and propane. They are also commonly known by the DuPont brand name Freon. The most common representative is dichlorodifluoromethane (R-12 or Freon-12). Many CFCs have been widely used as refrigerants, propellants (in aerosol applications), and solvents. Because CFCs contribute to ozone depletion in the upper atmosphere, the manufacture of such compounds has been phased out under the Montreal Protocol, and they are being replaced with other products such as hydrofluorocarbons (HFCs)^[1] (e.g., R-410A) and R-134a.^[2][3]

Structure, properties and production

As in simpler alkanes, carbon in the CFCs bonds with tetrahedral symmetry. Because the fluorine and chlorine atoms differ greatly in size and effective charge from hydrogen and from each other, the methane-derived CFCs deviate from perfect tetrahedral symmetry.^[4]

The physical properties of CFCs and HCFCs are tunable by changes in the number and identity of the halogen atoms. In general, they are volatile but less so than their parent alkanes. The decreased volatility is attributed to the molecular polarity induced by the halides, which induces intermolecular interactions. Thus, methane boils at −161 °C whereas the fluoromethanes boil between −51.7 (CF₂H₂) and −128 °C (CF₄). The CFCs have still higher boiling points because the chloride is even more polarizable than fluoride. Because of their polarity, the CFCs are useful solvents, and their boiling points make them suitable as refrigerants. The CFCs are far less flammable than methane, in part because they contain fewer C-H bonds and in part because, in the case of the chlorides and bromides, the released halides quench the free radicals that sustain flames.

The densities of CFCs are higher than their corresponding alkanes. In general, the density of these compounds correlates with the number of chlorides.

CFCs and HCFCs are usually produced by halogen exchange starting from chlorinated methanes and ethanes. Illustrative is the synthesis of chlorodifluoromethane from chloroform:

HCCl₃ + 2 HF → HCF₂Cl + 2 HCl

The brominated derivatives are generated by free-radical reactions of the chlorofluorocarbons, replacing C-H bonds with C-Br bonds. The production of the anesthetic 2-bromo-2-chloro-1,1,1-trifluoroethane ("halothane") is illustrative:

CF₃CH₂Cl + Br₂ → CF₃CHBrCl + HBr

Reactions

The most important reaction^{[citation needed]} of the CFCs is the photo-induced scission of a C-Cl bond:

CCl₃F → CCl₂F^. + Cl^.

The chlorine atom, written often as Cl^., behaves very differently from the chlorine molecule (Cl₂). The radical Cl^. is long-lived in the upper atmosphere, where it catalyzes the conversion of ozone into O₂. Ozone absorbs UV-B radiation, so its depletion allows more of this high energy radiation to reach the Earth's surface. Bromine atoms are even more efficient catalysts; hence brominated CFCs are also regulated.

Applications

CFCs and HCFCs are used in a variety of applications because of their low toxicity, reactivity and flammability. Every permutation of fluorine, chlorine and hydrogen based on methane and ethane has been examined and most have been commercialized. Furthermore, many examples are known for higher numbers of carbon as well as related compounds containing bromine. Uses include refrigerants, blowing agents, propellants in medicinal applications and degreasing solvents.

Billions of kilograms of chlorodifluoromethane are produced annually as precursor to tetrafluoroethylene, the monomer that is converted into Teflon.^[5]

Classes of compounds, nomenclature

• Chlorofluorocarbons (CFCs): when derived from methane and ethane these compounds have the formulae CCl_mF_4−m and C₂Cl_mF_6−m, where m is nonzero.
• Hydro-chlorofluorocarbons (HCFCs): when derived from methane and ethane these compounds have the formula CCl_mF_nH_4−m−n and C₂Cl_xF_yH_6−x−y, where m, n, x, and y are nonzero.
• and bromofluorocarbons have formulae similar to the CFCs and HCFCs but also include bromine.
• Hydrofluorocarbons (HFCs): when derived from methane, ethane, propane, and butane, these compounds have the respective formulae CF_mH_4−m, C₂F_mH_6−m, C₃F_mH_8−m, and C₄F_mH_10−m, where m is nonzero.

Numbering system

A special numbering system is used for fluorinated alkanes, prefixed with Freon-, R-, CFC- and HCFC-, where the rightmost value indicates the number of fluorine atoms, the next value to the left is the number of hydrogen atoms plus 1, and the next value to the left is the number of carbon atoms less one (zeroes are not stated), and the remaining atoms are chlorine.

Freon-12, for example, indicates a methane derivative (only two numbers) containing two fluorine atoms (the second 2) and no hydrogen (1-1=0). It is therefore CCl₂F₂.

Another, easier equation that can be applied to get the correct molecular formula of the CFC/R/Freon class compounds is this to take the numbering and add 90 to it. The resulting value will give the number of carbons as the first numeral, the second numeral gives the number of hydrogen atoms, and the third numeral gives the number of fluorine atoms. The rest of the unaccounted carbon bonds are occupied by chlorine atoms. The value of this equation is always a three figure number. An easy example is that of CFC-12, which gives: 90+12=102 -> 1 carbon, 0 hydrogens, 2 fluorine atoms, and hence 2 chlorine atoms resulting in CCl₂F₂. The main advantage of this method of deducing the molecular composition in comparison with the method described in the paragraph above is that it gives the number of carbon atoms of the molecule.

History

Carbon tetrachloride (CCl₄) was used in fire extinguishers and glass "anti-fire grenades" from the late nineteenth century until around the end of World War II. Experimentation with chloroalkanes for fire suppression on military aircraft began at least as early as the 1920s. Freon is a trade name for a group of CFCs which are used primarily as refrigerants, but also have uses in fire-fighting and as propellants in aerosol cans. Bromomethane is widely used as a fumigant. Dichloromethane is a versatile industrial solvent.

The Belgian scientist Frédéric Swarts pioneered the synthesis of CFCs in the 1890s. He developed an effective exchange agent to replace chloride in carbon tetrachloride with fluoride to synthesize CFC-11 (CCl₃F) and CFC-12 (CCl₂F₂).

In the late 1920s, Thomas Midgley, Jr. improved the process of synthesis and led the effort to use CFC as refrigerant to replace ammonia (NH₃), chloromethane (CH₃Cl), and sulfur dioxide (SO₂), which are toxic but were in common use. In searching for a new refrigerant, requirements for the compound were: low boiling point, low toxicity, and to be generally non-reactive. In a demonstration for the American Chemical Society, Midgley flamboyantly demonstrated all these properties by inhaling a breath of the gas and using it to blow out a candle^[6] in 1930.^[7][8]

Commercial development and use

During World War II, various chloroalkanes were in standard use in military aircraft, although these early halons suffered from excessive toxicity. Nevertheless, after the war they slowly became more common in civil aviation as well. In the 1960s, fluoroalkanes and bromofluoroalkanes became available and were quickly recognized as being highly effective fire-fighting materials. Much early research with Halon 1301 was conducted under the auspices of the US Armed Forces, while Halon 1211 was, initially, mainly developed in the UK. By the late 1960s they were standard in many applications where water and dry-powder extinguishers posed a threat of damage to the protected property, including computer rooms, telecommunications switches, laboratories, museums and art collections. Beginning with warships, in the 1970s, bromofluoroalkanes also progressively came to be associated with rapid knockdown of severe fires in confined spaces with minimal risk to personnel.

By the early 1980s, bromofluoroalkanes were in common use on aircraft, ships, and large vehicles as well as in computer facilities and galleries. However, concern was beginning to be expressed about the impact of chloroalkanes and bromoalkanes on the ozone layer. The Vienna Convention for the Protection of the Ozone Layer did not cover bromofluoroalkanes as it was thought, at the time, that emergency discharge of extinguishing systems was too small in volume to produce a significant impact, and too important to human safety for restriction.

Regulation

Since the late 1970s, the use of CFCs has been heavily regulated because of their destructive effects on the ozone layer. After the development of his electron capture detector, James Lovelock was the first to detect the widespread presence of CFCs in the air, finding a mole fraction of 60 ppt of CFC-11 over Ireland. In a self-funded research expedition ending in 1973, Lovelock went on to measure CFC-11 in both the Arctic and Antarctic, finding the presence of the gas in each of 50 air samples collected, and concluding that CFCs are not hazardous to the environment. The experiment did however provide the first useful data on the presence of CFCs in the atmosphere. The damage caused by CFCs was discovered by Sherry Rowland and Mario Molina who, after hearing a lecture on the subject of Lovelock's work, embarked on research resulting in the first publication suggesting the connection in 1974. It turns out that one of CFCs' most attractive features—their low reactivity— is key to their most destructive effects. CFCs' lack of reactivity gives them a lifespan that can exceed 100 years, giving them time to diffuse into the upper stratosphere.^[9] Once in the stratosphere, the sun's ultraviolet radiation is strong enough to cause the homolytic cleavage of the C-Cl bond.

NASA projection of stratospheric ozone, in Dobson units, if chlorofluorocarbons had not been banned. Animated version.

By 1987, in response to a dramatic seasonal depletion of the ozone layer over Antarctica, diplomats in Montreal forged a treaty, the Montreal Protocol, which called for drastic reductions in the production of CFCs. On 2 March 1989, 12 European Community nations agreed to ban the production of all CFCs by the end of the century. In 1990, diplomats met in London and voted to significantly strengthen the Montreal Protocol by calling for a complete elimination of CFCs by the year 2000. By the year 2010, CFCs should have been completely eliminated from developing countries as well.

Ozone-depleting gas trends

Because the only CFCs available to countries adhering to the treaty is from recycling, their prices have increased considerably. A worldwide end to production should also terminate the smuggling of this material. However, there are current CFC smuggling issues, as recognized by the United Nations Environmental Programme (UNEP) in a 2006 report titled "Illegal Trade in Ozone Depleting Substances". UNEP estimates that between 16,000–38,000 tonnes of CFCs passed through the black market in the mid-1990s. The report estimated between 7,000 and 14,000 tonnes of CFCs are smuggled annually into developing countries. Asian countries are those with the most smuggling; as of 2007, China, India and South Korea were found to account for around 70% of global CFC production,^[10] South Korea later to ban CFC production in 2010.^[11] Possible reasons for continued CFC smuggling were also examined: the report noted that many banned CFC producing products have long lifespans and continue to operate. The cost of replacing the equipment of these items is sometimes cheaper than outfitting them with a more ozone-friendly appliance. Additionally, CFC smuggling is not considered a significant issue, so the perceived penalties for smuggling are low. In 2018 public attention was drawn to the issue, that at an unknown place in east Asia an estimated amount of 13.000 metric tons anually of CFCs have been produced since about 2012 in violation of the protocol.^[12][13] While the eventual phaseout of CFCs is likely, efforts are being taken to stem these current non-compliance problems.

By the time of the Montreal Protocol, it was realised that deliberate and accidental discharges during system tests and maintenance accounted for substantially larger volumes than emergency discharges, and consequently halons were brought into the treaty, albeit with many exceptions.

Regulatory gap

While the production and consumption of CFCs are regulated under the Montreal Protocol, emissions from existing banks of CFCs are not regulated under the agreement. In 2002, there were an estimated 5,791 kilotons of CFCs in existing products such as refrigerators, air conditioners, aerosol cans and others.^[14] Approximately one-third of these CFCs are projected to be emitted over the next decade if action is not taken, posing a threat to both the ozone layer and the climate.^[15] A proportion of these CFCs can be safely captured and destroyed.

Regulation and DuPont

In 1978 the United States banned the use of CFCs such as Freon in aerosol cans, the beginning of a long series of regulatory actions against their use. The critical DuPont manufacturing patent for Freon ("Process for Fluorinating Halohydrocarbons", U.S. Patent #3258500) was set to expire in 1979. In conjunction with other industrial peers DuPont formed a lobbying group, the "Alliance for Responsible CFC Policy," to combat regulations of ozone-depleting compounds.^[16] In 1986 DuPont, with new patents in hand, reversed its previous stance and publicly condemned CFCs.^[17] DuPont representatives appeared before the Montreal Protocol urging that CFCs be banned worldwide and stated that their new HCFCs would meet the worldwide demand for refrigerants.^[17]

Phasing-out of CFCs

Use of certain chloroalkanes as solvents for large scale application, such as dry cleaning, have been phased out, for example, by the IPPC directive on greenhouse gases in 1994 and by the volatile organic compounds (VOC) directive of the EU in 1997. Permitted chlorofluoroalkane uses are medicinal only.

Bromofluoroalkanes have been largely phased out and the possession of equipment for their use is prohibited in some countries like the Netherlands and Belgium, from 1 January 2004, based on the Montreal Protocol and guidelines of the European Union.

Production of new stocks ceased in most (probably all) countries in 1994.^{[citation needed]} However many countries still require aircraft to be fitted with halon fire suppression systems because no safe and completely satisfactory alternative has been discovered for this application. There are also a few other, highly specialized uses. These programs recycle halon through "halon banks" coordinated by the Halon Recycling Corporation^[18] to ensure that discharge to the atmosphere occurs only in a genuine emergency and to conserve remaining stocks.

The interim replacements for CFCs are hydrochlorofluorocarbons (HCFCs), which deplete stratospheric ozone, but to a much lesser extent than CFCs.^[19] Ultimately, hydrofluorocarbons (HFCs) will replace HCFCs. Unlike CFCs and HCFCs, HFCs have an ozone depletion potential (ODP) of 0.^[20] DuPont began producing hydrofluorocarbons as alternatives to Freon in the 1980s. These included Suva refrigerants and Dymel propellants.^[21] Natural refrigerants are climate friendly solutions that are enjoying increasing support from large companies and governments interested in reducing global warming emissions from refrigeration and air conditioning. Hydrofluorocarbons are included in the Kyoto Protocol because of their very high Global Warming Potential and are facing calls to be regulated under the Montreal Protocol^{[dubious – discuss][22]} due to the recognition of halocarbon contributions to climate change.^[23]

On 21 September 2007, approximately 200 countries agreed to accelerate the elimination of hydrochlorofluorocarbons entirely by 2020 in a United Nations-sponsored Montreal summit. Developing nations were given until 2030. Many nations, such as the United States and China, who had previously resisted such efforts, agreed with the accelerated phase out schedule.^[24]

Development of alternatives for CFCs

Work on alternatives for chlorofluorocarbons in refrigerants began in the late 1970s after the first warnings of damage to stratospheric ozone were published.

The hydrochlorofluorocarbons (HCFCs) are less stable in the lower atmosphere, enabling them to break down before reaching the ozone layer. Nevertheless, a significant fraction of the HCFCs do break down in the stratosphere and they have contributed to more chlorine buildup there than originally predicted. Later alternatives lacking the chlorine, the hydrofluorocarbons (HFCs) have an even shorter lifetimes in the lower atmosphere.^[19] One of these compounds, HFC-134a, is now used in place of CFC-12 in automobile air conditioners. Hydrocarbon refrigerants (a propane/isobutane blend) are also used extensively in mobile air conditioning systems in Australia, the USA and many other countries, as they have excellent thermodynamic properties and perform particularly well in high ambient temperatures.

One of the natural refrigerants (along with ammonia and carbon dioxide), hydrocarbons have negligible environmental impacts and are also used worldwide in domestic and commercial refrigeration applications, and are becoming available in new split system air conditioners.^[25] Various other solvents and methods have replaced the use of CFCs in laboratory analytics.^[26]

In Metered-dose inhalers (MDI), a non-ozone effecting substitute was developed as a propellant, known as "hydrofluoroalkane."^[27]

Environmental impacts

As previously discussed, CFCs were phased out via the Montreal Protocol due to their part in ozone depletion. However, the atmospheric impacts of CFCs are not limited to its role as an active ozone reducer. This anthropogenic compound is also a greenhouse gas, with a much higher potential to enhance the greenhouse effect than CO₂.

Infrared absorption bands trap heat from escaping earth's atmosphere. In the case of CFCs, the strongest of these bands are located in the spectral region 7.8–15.3 µm ^[28] – referred to as an atmospheric window due to the relative transparency of the atmosphere within this region.^[29] The strength of CFC bands and the unique susceptibility of the atmosphere, at which the compound absorbs and emits radiation, are two factors that contribute to CFCs' "super" greenhouse effect.^[30] Another such factor is the low concentration of the compound. Because CO₂ is close to saturation with high concentrations,^{[clarification needed]} it takes more of the substance to enhance the greenhouse effect. Conversely, the low concentration of CFCs allow their effects to increase linearly with mass.^[30]

Tracer of ocean circulation

Because the time history of CFC concentrations in the atmosphere is relatively well known, they have provided an important constraint on ocean circulation. CFCs dissolve in seawater at the ocean surface and are subsequently transported into the ocean interior. Because CFCs are inert, their concentration in the ocean interior reflects simply the convolution of their atmospheric time evolution and ocean circulation and mixing.

CFC and SF₆ tracer-derived age of ocean water

Chlorofluorocarbons (CFCs) are anthropogenic compounds that have been released into the atmosphere since the 1930s in various applications such as in air-conditioning, refrigeration, blowing agents in foams, insulations and packing materials, propellants in aerosol cans, and as solvents.^[31] The entry of CFCs into the ocean makes them extremely useful as transient tracers to estimate rates and pathways of ocean circulation and mixing processes.^[32] However, due to production restrictions of CFCs in the 1980s, atmospheric concentrations of CFC-11 and CFC-12 has stopped increasing, and the CFC-11 to CFC-12 ratio in the atmosphere have been steadily decreasing, making water dating of water masses more problematic.^[32] Incidentally, production and release of sulfur hexafluoride (SF₆) have rapidly increased in the atmosphere since the 1970s.^[32] Similar to CFCs, SF₆ is also an inert gas and is not affected by oceanic chemical or biological activities.^[33] Thus, using CFCs in concert with SF₆ as a tracer resolves the water dating issues due to decreased CFC concentrations.

Using CFCs or SF₆ as a tracer of ocean circulation allows for the derivation of rates for ocean processes due to the time-dependent source function. The elapsed time since a subsurface water mass was last in contact with the atmosphere is the tracer-derived age.^[34] Estimates of age can be derived based on the partial pressure of an individual compound and the ratio of the partial pressure of CFCs to each other (or SF₆).^[34]

Partial pressure and ratio dating techniques

The age of a water parcel can be estimated by the CFC partial pressure (pCFC) age or SF₆ partial pressure (pSF₆) age. The pCFC age of a water sample is defined as:

${\displaystyle pCFC={\frac {[CFC]}{F(T,S)}}}$

where [CFC] is the measured CFC concentration (pmol kg⁻¹) and F is the solubility of CFC gas in seawater as a function of temperature and salinity.^[35] The CFC partial pressure is expressed in units of 10–12 atmospheres or parts-per-trillion (ppt).^[36] The solubility measurements of CFC-11 and CFC-12 have been previously measured by Warner and Weiss^[36] Additionally, the solubility measurement of CFC-113 was measured by Bu and Warner^[37] and SF₆ by Wanninkhof et al.^[38] and Bullister et al.^[39] Theses authors mentioned above have expressed the solubility, F, at a total pressure of 1 atm as:

${\displaystyle \ln F=a_{1}+a_{2}\left({\frac {100}{T}}\right)+a_{3}\ln \left({\frac {T}{100}}\right)+a_{4}\left({\frac {T}{100}}\right)^{2}+S\left[b_{1}+b_{2}\left({\frac {T}{100}}\right)+b_{3}\left({\frac {T}{100}}\right)\right]}$

where F = solubility expressed in either mol l⁻¹ or mol kg⁻¹ atm⁻¹, T = absolute temperature, S = salinity in parts per thousand (ppt), a₁, a₂, a₃, b₁, b₂, and b₃ are constants to be determined from the least squares fit to the solubility measurements.^[37] This equation is derived from the integrated Van 't Hoff equation and the logarithmic Setchenow salinity dependence.^[37]

It can be noted that the solubility of CFCs increase with decreasing temperature at approximately 1% per degree Celsius.^[34]

Once the partial pressure of the CFC (or SF₆) is derived, it is then compared to the atmospheric time histories for CFC-11, CFC-12, or SF₆ in which the pCFC directly corresponds to the year with the same. The difference between the corresponding date and the collection date of the seawater sample is the average age for the water parcel.^[34] The age of a parcel of water can also be calculated using the ratio of two CFC partial pressures or the ratio of the SF₆ partial pressure to a CFC partial pressure.^[34]

Safety

According to their material safety data sheets, CFCs and HCFCs are colorless, volatile, toxic liquids and gases with a faintly sweet ethereal odor. Overexposure at concentrations of 11% or more may cause dizziness, loss of concentration, central nervous system depression and/or cardiac arrhythmia. Vapors displace air and can cause asphyxiation in confined spaces. Although non-flammable, their combustion products include hydrofluoric acid, and related species.^[40] Normal occupational exposure is rated at 0.07% and does not pose any serious health risks.^[41]