 In chemistry, '''Henry's law''' is one of the gas laws, formulated by William Henry in 1803. It states that: :''At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid.'' An equivalent way of stating the law is that the solubility of a gas in a liquid at a particular temperature is proportional to the pressure ''of that gas'' above the liquid. Henry's law has since been shown to apply for a wide range of dilute solutions, not merely those of gases.  An everyday example of Henry's law is given by carbonated soft drinks. Before the bottle or can is opened, the gas above the drink is almost pure carbon dioxide at a pressure slightly higher than atmospheric pressure. The drink itself contains dissolved carbon dioxide. When the bottle or can is opened, some of this gas escapes, giving the characteristic hiss (or &quot;pop&quot; in the case of a champagne bottle). Because the pressure above the liquid is now lower, some of the dissolved carbon dioxide comes out of solution as bubbles. If a glass of the drink is left in the open, the concentration of carbon dioxide in solution will come into equilibrium with the carbon dioxide in the air, and the drink will go &quot;flat&quot;.  ==Formula and the Henry's law constant== Henry's law can be put into mathematical terms (at constant temperature) as :p = k_{\rm H}\, c where ''p'' is the partial pressure of the solute in the gas above the solution, ''c'' is the concentration of the solute  and ''k''H is a constant with the dimensions of pressure divided by concentration. The constant, known as the Henry's law constant, depends on the solute, the solvent and the temperature.  Some values for ''k''H for gases dissolved in water at 298&amp;nbsp;K include:  :oxygen (O2) : 769.2 L·atm/mol :carbon dioxide (CO2) : 29.4 L·atm/mol :hydrogen (H2) : 1282.1 L·atm/mol  There are other forms of Henry's Law, each of which defines the constant ''k''H differently and requires different dimensional units. In particular, the &quot;concentration&quot; of the solute in solution may also be expressed as a mole fraction or as a molality.  ===Other forms of Henry's law=== There are various other forms of Henry's Law which are discussed in the technical literature.[ University of Arizona chemistry class notes][ An extensive list of Henry's law constants, and a conversion tool]  {| class=&quot;wikitable&quot; |+ '''Table 1: Some forms of Henry's law and constants (gases in water at 298 K), derived from  ! equation: || k_{\mathrm{H,pc}} = \frac{p}{c} || k_{\mathrm{H,cp}} = \frac{c}{p} ||  k_{\mathrm{H,px}} = \frac{p}{x}  ||  k_{\mathrm{H,cc}} = \frac{c_{\mathrm{aq}}}{c_{\mathrm{gas}}}  |- ! units: || \frac{\mathrm{L} \cdot \mathrm{atm}}{\mathrm{mol}} ||  \frac{\mathrm{mol}}{\mathrm{L} \cdot \mathrm{atm}} || \rm atm\, || ''dimensionless'' |- |align=center| O2 ||align=center| 769.23||align=center| 1.3 ||align=center| 4.259 ||align=center| 3.180 |- |align=center| H2 ||align=center| 1282.05 ||align=center| 7.8 ||align=center| 7.099 ||align=center| 1.907 |- |align=center| CO2 ||align=center| 29.41 ||align=center| 3.4 ||align=center| 0.163 ||align=center| 0.8317 |- |align=center| N2 ||align=center| 1639.34  ||align=center| 6.1 ||align=center| 9.077 ||align=center| 1.492 |- |align=center| He ||align=center| 2702.7 ||align=center| 3.7||align=center| 14.97 ||align=center| 9.051 |- |align=center| Ne ||align=center| 2222.22 ||align=center| 4.5 ||align=center| 12.30 ||align=center| 1.101 |- |align=center| Ar ||align=center| 714.28 ||align=center| 1.4 ||align=center| 3.955 ||align=center| 3.425 |- |align=center| CO ||align=center| 1052.63  ||align=center| 9.5 ||align=center| 5.828 ||align=center| 2.324 |}  where: :''c'' = amount concentration of gas in solution (in mol/L) :''p'' = partial pressure of gas above the solution (in atm) :''x'' = mole fraction of gas in solution (dimensionless)  As can be seen by comparing the equations in the above table, the Henry's law constant ''k''H,pc is simply the inverse of the constant ''k''H,cp. Since all ''k''H may be referred to as Henry's law constants, readers of the technical literature must be quite careful to note which version of the Henry's Law equation is being used.  It should also be noted the Henry's Law is a limiting law that only applies for ''sufficiently dilute'' solutions. The range of concentrations in which it applies becomes narrower the more the system diverges from ideal behavior. Roughly speaking, that is the more chemically ''different'' the solute is from the solvent.   It also only applies simply for solutions where the solvent does not react chemically with the gas being dissolved. A common example of a gas that does react with the solvent is carbon dioxide, which forms carbonic acid (H2CO3) to a certain degree with water.  ===Temperature dependence of the Henry constant=== When the temperature of a system changes, the Henry constant will also change. This is why some people prefer to name it Henry coefficient. There are multiple equations assessing the effect of temperature on the constant. These forms of the van 't Hoff equation are examples:  : k_{\rm H,pc}(T) = k_{\rm H,pc}(T^\ominus)\, \exp{ \left[ -C \, \left( \frac{1}{T}-\frac{1}{T^\ominus}\right)\right]}\,   : k_{\rm H,cp}(T) = k_{\rm H,cp}(T^\ominus)\, \exp{ \left[ C \, \left( \frac{1}{T}-\frac{1}{T^\ominus}\right)\right]}\,   where :''k''H for a given temperature is the Henry's Law constant (as defined in the first section of this article). Notice that the correct sign of C depends on whether ''k''H,pc or ''k''H,cp is used. :''T'' is the thermodynamic temperature, :''T'' o refers to the standard temperature (298 K).  This equation is only an approximation, and should be used only when no better, experimentally-derived formula is known for a given gas.  The following table lists some values for constant ''C'' (in kelvins) in the equation above:  {|class=&quot;wikitable&quot; |+ '''Table 2: Values of ''C''''' |'''Gas''' || align=center| O2 || align=center| H2 ||  align=center| CO2 || align=center| N2 || align=center| He || align=center| Ne || align=center| Ar || align=center| CO |- | ''C''(K) || align=center| 1700  ||align=center| 500 || align=center| 2400 || align=center| 1300|| align=center| 230 || align=center| 490|| align=center|  1300 || align=center| 1300 |} Because solubility of permanent gases usually decreases with increasing temperature at around the room temperature, the partial pressure a given gas concentration has in liquid must increase. While heating water (saturated with nitrogen) from 25 °C to 95 °C the solubility will decrease to about 43% of its initial value. This can be verified when heating water in a pot: small bubbles evolve and rise, long before the water reaches boiling temperature. Similarly, carbon dioxide from a carbonated drink escapes much faster when the drink is not cooled because of the increased partial pressure of CO2 in higher temperatures. Partial pressure of CO2 in the gas phase in equilibrium with seawater doubles with every 16 K increase in temperature.Takahashi, T; Sutherland, SC; Sweeney, C; Poisson, A; Metzl, N; Tilbrook, B; Bates, N; Wanninkhof, R; Feely, RA; Sabine, C; Olafsson, J; Nojiri, Y  &quot;''Global sea-air CO2 flux based on climatological surface ocean pCO2 and seasonal biological and temperature effects''&quot;  Deep-Sea Research (Part II, Topical Studies in Oceanography) [Deep-Sea Research (II Top. Stud. Oceanogr.)] '''49''', 9-10, pp. 1601-1622, 2002  The constant ''C'' may be regarded as:  : C = -\frac{\Delta_{\rm solv}H}{R} = -\frac{p^\star} for a volatile solute; ''c''o&amp;nbsp;= 1&amp;nbsp;mol/L. For non-ideal solutions, the activity coefficient ''γc'' depends on the concentration and must be determined at the concentration of interest. The activity coefficient can also be obtained for non-volatile solutes, where the vapor pressure of the pure substance is negligible, by using the Gibbs–Duhem relation: :\sum_i n_i\, {\rm d}\mu_i = 0 By measuring the change in vapor pressure (and hence chemical potential) of the solvent, the chemical potential of the solute can be deduced.  The standard state for a dilute solution is also defined in terms of infinite-dilution behavior. Although the standard concentration ''c''o is taken to be 1&amp;nbsp;mol/L by convention, the standard state is a hypothetical solution of 1&amp;nbsp;mol/L in which the solute has its limiting infinite-dilution properties. This has the effect that all non-ideal behavior is described by the activity coefficient: the activity coefficient at 1&amp;nbsp;mol/L is not necessarily unity (and is frequently quite different from unity).  All the relations above can also be expressed in terms of molalities rather than concentrations, e.g.: :\mu = \mu_m^\ominus + RT\ln{\left( \frac{\gamma_m m}{m^\ominus}\right)}\,, where \gamma_m = \frac{k_{p^\star} for a volatile solutes; ''m''o&amp;nbsp;= 1&amp;nbsp;mol/kg. The standard chemical potential ''μm''o, the activity coefficient ''γm'' and the Henry's law constant ''k''H,''m'' all have different numerical values when molalities are used in place of concentrations.  