A general ubraintv-jp.comistryubraintv-jp.comTextmaporganized around the textbookubraintv-jp.comistry: Principles, Patterns, and Applicationsby Bruce A. Averill

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Learning Objectives

To understand the relationship among temperature, pressure, and solubility. The understand that the solubility of a solid may increase or decrease with increasing temperature, To understand that the solubility of a gas decreases with an increase in temperature and a decrease in pressure.

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Experimentally it is found that the solubility of most compounds depends strongly on temperature and, if a gas, on pressure as well. As we shall see, the ability to manipulate the solubility by changing the temperature and pressure has several important consequences.

## Effect of Temperature on the Solubility of Solids

Figure $$\PageIndex{1}$$ shows plots of the solubilities of several organic and inorganic compounds in water as a function of temperature. Although the solubility of a solid generally increases with increasing temperature, there is no simple relationship between the structure of a substance and the temperature dependence of its solubility. Many compounds (such as glucose and $$\ce{CH_3CO_2Na}$$) exhibit a dramatic increase in solubility with increasing temperature. Others (such as $$\ce{NaCl}$$ and $$\ce{K_2SO_4}$$) exhibit little variation, and still others (such as $$\ce{Li_2SO_4}$$) become less soluble with increasing temperature.

api/deki/files/41609/67558bdc4beb64e06b29db7b4c8d74bb.jpg?revision=1&size=bestfit&width=374&height=426" />Figure $$\PageIndex{2}$$: Solubilities of Several Common Gases in Water as a Function of Temperature at Partial Pressure of 1 atm. The solubilities of all gases decrease with increasing temperature. (CC BY-SA-NC; anonymous)

The decrease in the solubilities of gases at higher temperatures has both practical and environmental implications. Anyone who routinely boils water in a teapot or electric kettle knows that a white or gray deposit builds up on the inside and must eventually be removed. The same phenomenon occurs on a much larger scale in the giant boilers used to supply hot water or steam for industrial applications, where it is called “boiler scale,” a deposit that can seriously decrease the capacity of hot water pipes (Figure $$\PageIndex{3}$$). The problem is not a uniquely modern one: aqueducts that were built by the Romans 2000 years ago to carry cold water from alpine regions to warmer, drier regions in southern France were clogged by similar deposits. The ubraintv-jp.comistry behind the formation of these deposits is moderately complex and will be described elsewhere, but the driving force is the loss of dissolved $$\ce{CO2}$$ from solution. Hard water contains dissolved $$\ce{Ca^{2+}}$$ and $$\ce{HCO3^{-}}$$ (bicarbonate) ions. Calcium bicarbonate ($$\ce{Ca(HCO3)2}$$ is rather soluble in water, but calcium carbonate ($$\ce{CaCO3}$$) is quite insoluble. A solution of bicarbonate ions can react to form carbon dioxide, carbonate ion, and water:

\<\ce{2HCO3^{-}(aq)-> CO3^{2-}(aq)+ H2O(l)+ CO2(aq)} \label{13.9}\>

Heating the solution decreases the solubility of $$\ce{CO2}$$, which escapes into the gas phase above the solution. In the presence of calcium ions, the carbonate ions precipitate as insoluble calcium carbonate, the major component of boiler scale.

There are many causes of fish kill, but oxygen depletion is the most common cause. (Public Domain;United States Fish and Wildlife Service)

A similar effect is seen in the rising temperatures of bodies of water such as the k0oi89Chesapeake Bay, the largest estuary in North America, where \lobal warming has been implicated as the cause. For each 1.5°C that the bay’s water warms, the capacity of water to dissolve oxygen decreases by about 1.1%. Many marine species that are at the southern limit of their distributions have shifted their populations farther north. In 2005, the eelgrass, which forms an important nursery habitat for fish and shellfish, disappeared from much of the bay following record high water temperatures. Presumably, decreased oxygen levels decreased populations of clams and other filter feeders, which then decreased light transmission to allow the eelsgrass to grow. The complex relationships in ecosystems such as the Chesapeake Bay are especially sensitive to temperature fluctuations that cause a deterioration of habitat quality.

## Effect of Pressure on the Solubility of Gases: Henry’s Law

External pressure has very little effect on the solubility of liquids and solids. In contrast, the solubility of gases increases as the partial pressure of the gas above a solution increases. This point is illustrated in Figure $$\PageIndex{4}$$, which shows the effect of increased pressure on the dynamic equilibrium that is established between the dissolved gas molecules in solution and the molecules in the gas phase above the solution. Because the concentration of molecules in the gas phase increases with increasing pressure, the concentration of dissolved gas molecules in the solution at equilibrium is also higher at higher pressures.

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Due to the low Henry’s law constant for $$\ce{O2}$$ in water, the levels of dissolved oxygen in water are too low to support the energy needs of multicellular organisms, including humans. To increase the $$\ce{O2}$$ concentration in internal fluids, organisms synthesize highly soluble carrier molecules that bind $$\ce{O2}$$ reversibly. For example, human red blood cells contain a protein called hemoglobin that specifically binds $$\ce{O2}$$ and facilitates its transport from the lungs to the tissues, where it is used to oxidize food molecules to provide energy. The concentration of hemoglobin in normal blood is about 2.2 mM, and each hemoglobin molecule can bind four $$\ce{O2}$$ molecules. Although the concentration of dissolved $$\ce{O2}$$ in blood serum at 37°C (normal body temperature) is only 0.010 mM, the total dissolved $$\ce{O2}$$ concentration is 8.8 mM, almost a thousand times greater than would be possible without hemoglobin. Synthetic oxygen carriers based on fluorinated alkanes have been developed for use as an emergency replacement for whole blood. Unlike donated blood, these “blood substitutes” do not require refrigeration and have a long shelf life. Their very high Henry’s law constants for $$\ce{O2}$$ result in dissolved oxygen concentrations comparable to those in normal blood.