Henry's law | Wikipedia audio article

This is an audio version of the Wikipedia Article:


00:01:08 1 Fundamental types and variants of Henry's law constants
00:03:44 2 Values of Henry's law constants
00:03:57 3 Temperature dependence
00:05:21 4 Effective Henry's law constants spaniH/isubeff/sub
00:05:27 5 Dependence on ionic strength (Sechenov equation)
00:07:22 6 Non-ideal solutions
00:09:51 6.1 Solvent mixtures
00:14:46 7 Miscellaneous
00:15:11 7.1 In geochemistry
00:16:12 7.2 Comparison to Raoult's law
00:16:34 8 See also
00:16:59 9 References
00:18:20 10 External links
00:18:35 The Henry volatility defined via concentration (
00:19:42 The Henry volatility defined via aqueous-phase mixing ratio (
00:20:43 The dimensionless Henry volatility
00:22:05 Values of Henry's law constants
00:22:24 Temperature dependence
00:25:29 Effective Henry's law constants Heff
00:29:27 Dependence on ionic strength (Sechenov equation)
00:31:41 Non-ideal solutions
00:35:14 Solvent mixtures
00:36:24 Miscellaneous
00:36:33 In geochemistry
00:37:40 Comparison to Raoult's law
00:40:08 See also



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SUMMARY
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In chemistry, Henry's law is a gas law that states that the amount of dissolved gas in a liquid is proportional to its partial pressure above the liquid. The proportionality factor is called Henry's law constant. It was formulated by the English chemist William Henry, who studied the topic in the early 19th century. In his publication about the number of gases absorbed by water,[1] he described the results of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed under the common pressure of the atmosphere."

An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater divers that changes during decompression, leading to decompression sickness. An everyday example is given by one's experience with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is almost pure carbon dioxide, at a pressure higher than atmospheric pressure. After the bottle is opened, this gas escapes, moving the partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes out of solution.