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Final Answers
© 2000-2017   Gérard P. Michon, Ph.D.

Physics of Gases and Fluids

 Jan Baptista van Helmont 
 1577-1644  Otto von Guericke 
 1602-1686  Blaise Pascal 
 1623-1662  Robert Boyle 
 1627-1691  Amedeo Avogadro 
 1776-1856 Natura abhorret a vacuo.
Nature abhors a vacuum.

Articles previously on this page:

  • Raising the Titanic, with (a lot of) hydrogen.
    The above article has moved...  Click for the new location.

Related articles on this site:

Related Links (Outside this Site)

Jacques Alexandre César Charles  by  Karine Larochelle.
Van der Waals equation of state   |   Calculator  by  Dieter Bingemann.
Classical Kinetic Theory of Gases:  A Crash Review.  by  Kenneth Boyd.
Wikipedia: Equation of state.

Basic Physics of Gases and Fluids

 Blaise Pascal (2017-11-20)     Vacuum and the origin of atmospheric pressure.
In still air, the vertical gradient of pressure is the weight density.

At the urging of  Blaise Pascal (1623-1662), on 19 September 1648,  Florin Périer  (Pascal's brother-in-law)  led a party of several Clermont-Ferrand notables on an ascension of nearby Puy-de-Dôme  to record the variation of air pressure with altitude, using a mercury barometer  (invented by Torricelli in 1643).

The variation they found was so large that Pascal realized he could observe the effect locally in Paris,  using the modest variation in altitude at the bottom and at the top of the freshly rebuilt  Saint-Jacques Tower  (the tallest building in Paris at the time).

A dozen years later, on 27 April 1661,  Richard Towneley (1629-1707)  and  Henri Power (1623-1668; FRS 1663)  farmously performed a similar experiment to measure the pressure of air at different altitudes on  Pendle Hill in Lancashire.

 Come back later, we're
 still working on this one...

 Otto von Guericke 
 1602-1686 (2006-01-14)   The Magdeburg Hemispheres :  Mighty Pressure.
A famous demonstration by the inventor of the vacuum pump.

The inventor of the air pump (1650) was the German physicist Otto von Guericke (1602-1686).  Guericke was also Mayor of  Magdeburg.  To demonstrate the might of atmospheric pressure, he built two matching copper bowls  (the Magdeburg hemispheres)  14 inches in diameter, which could be placed in airtight contact to form a hollow sphere from which air could be pumped out.

In 1654, the device was demonstrated dramatically before Emperor Ferdinand III and the Imperial Diet at Regensburg:  After Guericke had pumped (most of) the air out of the "Magdeburg sphere", two teams of horses were not enough to pull apart those hemispheres which only "thin" surrounding air was holding together.

The force holding evacuated 14" Magdeburg hemispheres together is the product of the atmospheric pressure  (about 101325 Pa)  into the cross-sectional area  (a disk 14" across has an area of about 0.1 m). ; ton of thrust...

 Clapeyron (2003-05-21)     p V   =   n R T
Who formulated the  ideal gas law ?

Emile Clapeyron (1799-1864; X1816)  gave the law its modern form in 1834.  Louis-Joseph Gay-Lussac (1778-1850)  had first put all aspects of this relation together around 1802.  Although Gay-Lussac only considered a fixed amount of gas  (constant n)  at that time, he would also pave the way to Avogadro's law with his 1809 observation that chemical gaseous reactions always entail simple ratios in the volumes of the reactants and products.

The  ideal gas law  (French:  Loi des gaz parfaits)  is deceptively simple:

p V   =   n R T

This states that the product of the pressure (p) and volume (V) of an ideal gas is proportional to its absolute temperature (T) and to the number of molecules in it;  "n" is actually the number of moles, which is the number of molecules divided by a numerical constant called Avogadro's number, whereas "R" is a constant of proportionality equal to Boltzmann's constant (k) multiplied into Avogadro's number:

R   =   k NA   =   8.3144621(75) J/K/mol.

The law was instrumental in the clarification of all the concepts involved, with the sole exception of "volume", of course...  The law can be broken down into four statements which were historically studied separately (some, or all, of these may be invalid in more elaborate models of actual fluids).  When the full law was finally formulated as above, it became possible to quantify the discrepancies between the "ideal" behavior it described and the actual properties of real gases...

Avogadro's coat of arms

Avogadro's Law (1811)

At the same temperature and [low] pressure, equal volumes
of different gases contain the same number of molecules.

This was formulated by  Amedeo Avogadro (1776-1856)  in 1811.  Like other  ideal gas laws,  it's only a rough  approximation  under most actual conditions.  However, it's so accurate at very low pressures that it can be presented as a limiting case for real gases at constant temperature:

n R T   =   lim  p V(p)
p ® 0

Boyle's coat of arms

Boyle-Mariotte Law :  Isothermal  (1662, 1676)

At a given temperature, the volume of a gas
is inversely proportional to its pressure.

The law was first formulated in print by Robert Boyle (1627-1691) in 1662.  Boyle omitted the important  isothermal  requirement which was added in 1676 by Edme Mariotte (1620-1684) when he rediscovered the whole thing independently...

A fluid which obeys the Boyle-Mariotte law is called a  Mariotte gas.  In such a gas, the product  pV  depends only on the absolute temperature (T) :

p V   =   n R  f (T)

Historically, it's essentially the simplification entailed by putting  f (T) = T  in the (idealized) above equation which helped  define  the concept of "absolute" temperature on which other aspects of the ideal law depend quantitatively:

Charles' Law :  Isobaric (1787)

At a given pressure, the volume of a gas is
proportional to its absolute temperature.

f (p)   =   n R T

The isobaric law has been named after Jacques [Alexandre César] Charles (1746-1823) to acknowledge the unpublished work he performed in 1787  (which was quoted by Gay-Lussac in 1802).

Gay-Lussac's Law :  Isochoric (1802)  Signature of Louis-Joseph 
 Gay-Lussac (1778-1850)

When confined to a given volume, the pressure of
a gas is proportional to its absolute temperature.

f (V)   =   n R T

This is the law of perfect gases under isochoric (or isovolumic) conditions.

(2006-09-18)   Joule's Law   (Joule's first law)
The internal energy (U) of an ideal gas depends only on its temperature.

This law goes beyond the descriptive aspects of a gas (volume, pressure and temperature) by introducing energetic aspects:

The internal energy  (U)  of a gas is  [ideally]
a function of its absolute temperature only.

Skip the following "advanced" discussion on first reading.

Using the relevant expression of dU, this "law" can be stated as follows:

0   =    (  U  )T    =       [ T (  p  )V   - p ]
vinculum vinculum

Thus, at constant volume,  Joule's first law  translates into a  differential equation  which etablishes that  p  and  T  are proportional.  It is equivalent to the isochoric law of Gay-Lussac,  which is true if and only if the molar  equation of state  is of the form:

f (V)   =   R T

For a perfect gas, the internal energy is actually proportional to the absolute temperature, under the classical assumptions of the  kinetic theory of gases, which defines temperature as proportional to the average (translational) kinetic energy of molecules and postulates equipartition of energy between the  j  active  mechanical degrees of freedom of each molecule:  Along with the 3 translational degrees, the number of rotational degrees of freedom is:

  • 0  for a spherical molecule (monoatomic gases: helium, argon, etc.). 
  • 2  for a molecule with axial symmetry  (carbon dioxide  and all diatomic gases, like hydrogen, nitrogen and oxygen). 
  • 3  otherwise  (steam, ammonia, methane, etc.).

The reason why a quantum object can't rotate about an axis of perfect symmetry is that such a rotation would be "logically" impossible to observe, as it does not entail any change in the quantum state.

At high temperature, molecules can no longer be considered perfectly rigid.  Technically, vibrational modes exist in molecules of two atoms or more.  Those behave as additional degrees of freedom that are "dormant" at low temperatures and become :active" at higher temperatures  (for reasons explained by statistical mechanics).

If the assumption is made (which is part of the ideal gas model) that there's no energy of interaction between molecules  (which thus interact only via relatively rare collisions)  the above tells the whole story:  U = (j/2) nRT

  • U  =  3/2 nRT   for monoatomic gases  (e.g., He, Ar).
  • U  =  5/2 nRT   for diatomic  or  linear  molecules  (e.g., H2, N2, CO2).
  • U  =  3  nRT   for all  other gases  (e.g., H2O, NH3, CH4).

Molecular interactions in  real  gases entail some violation of Joule's law.  Even  nonpolar  neighboring molecules induce distortions in each other's electronic clouds, causing attractive forces called  Van der Waals forces.

Joule-Thomson law  (Joule's second law):

Joule's second law states that the  enthalpy  (H = U+pV)  of a  Joule-Thomson gas  depends only on its temperature.  The molar equation of state of a Joule-Thomson gas is of the following form, which is equivalent to the isobaric law of Charles:

f (p)   =   R T

A gas which obeys  both  of Joule's laws is necessarily a  perfect gas.

(2010-12-13)   Deflating a Tire
What happens when a pressurized gas is released into the atmosphere?

When air is let out of a tire, the temperature of the tire does not change.

This non-intuitive fact holds for ideal gases  but not necessarily for real gases.  The reason is that each molecule contributes equally to the total pressure, independently of the other molecules which may or may not be there.  Each molecule that escapes thus reduces the pressure (p) and the number of moles inside the container (n) in the same proportion.  This does not entail any change in the temperature  T = (p/n) V/R.

Adiabatic cooling vs. Joule-Thomson effect   |   Letting Air Out of Tires

(2006-09-17)   Realistic Equations of State   [molar]
Better approximations for actual gases than the  ideal gas law.

The equations below are substitutes for the molar  (n = 1)  ideal gas law.  They were designed to better describe the actual behavior of real gases.  In particular, the volume of a gas under extreme pressure does not vanish but tends to a finite limit, called  covolume.  The molar covolume is traditionally denoted by the symbol  "b".  The simplest (molar) equation which accounts for this important feature of real gases has been known as the  Clausius equation of state :

p  ( V-b )   =   RT

The Clausius equation for  n  moles of gas would be:  p  ( V-nb )  =  nRT.  A  Clausius gas  thus obeys the isochoric law of Gay-Lussac and does not really contradict Avogadro's law  (since, at very low pressure, the covolume of any gas is a negligible fraction of its volume).  On the other hand, a Clausius gas  does violate the isothermal law of Boyle-Mariotte and the isobaric law of Charles.

 Johannes Diderik van der Waals 
 (1837-1923) earned a Nobel prize in 1910

Van der Waals Fluid   (1873)

A generalization of the Clausius equation appeared in the doctoral dissertation of  Johannes Diderik van der Waals  (1837-1923; Nobel 1910) :

( p + a / V2 )  ( V-b )   =   RT

 Charles Cagniard de la Tour 
 1777-1859, X1794 In the main, this equation accounts for interactions between molecules which are able to trigger a transition between liquid and gaseous phases below a certain temperature, which was apparently first identified as  critical  by the French physicist Charles Cagniard (1777-1859; X1794) in 1822.

The Irish physicist Thomas Andrews (1813-1895) is often credited for the principle of the "continuity of states" in a fluid:  The gaseous and liquid forms can be transformed into each other without undergoing a phase transition  (by traveling above the critical point).  Andrews showed this in 1869 for carbon dioxide  (which doesn't exist as a liquid below its surprisingly high triple-point pressure of  5.11 atm).

At the  critical temperature,  the isothermal curve which gives pressure as a function of volume has an inflexion point with an horizontal tangent:

 (  p  )T   =   0   (  2p  )T,T   =   0
vinculum vinculum
V V2

Differentiating the equation of state of a Van der Waals fluid,  we obtain:

( V-b ) dp  +  ( p  -  a / V2  +  2 a b / V3 ) dV   =   R dT

Along an isothermal curve  (dT = 0)  this yields:

 (  p  )T   =   -   p  -  a / V2  +  2 a b / V3
vinculum vinculum
V V - b

At the critical point, that quantity vanishes  (the tangent to the isothermal is horizontal).  Since  p  is thereby stationary at that point, the isothermal derivative of this expression is simply equal to its derivative with respect to  V.  The vanishing of that  second derivative  at the critical point  (which is an inflexion point of the critical isothermal)  thus boils down to the following:

0   =   d/dV ( -  a / V2  +  2 a b / V3 )     Therefore:  V = 3b

With this, we obtain  p  from the previous vanishing expression and and  T  from the equation of state.  The critical point is thus characterized by:

Vc = 3b ,       pc = a / 27 b2 ,       Tc = 8a / 27 R b

Using those  critical values  as units of volume, pressure and temperature, the above  Van der Waals equation  takes on the following  reduced form :

( p + 3/V 2 )  ( V - 1/3 )   =   8T/3

The principle of corresponding states says that different substances at the same reduced pressure and temperature will have the same reduced volume.  This fails for real fluids (especially at high pressure) but it's true of any 3-parameter model similar to the Van der Waals model  (see next).

Entropy of a Van der Waals fluid   |   Speed of sound in a Van der Waals fluid

Dieterici Fluid  (1899) :

The Dieterici equation of state is another 3-parameter model which predicts a liquid-vapor transition and an associated critical state:

p (V-b)   =   RT exp(-a / RTV)

The innovative introduction of an exponential term was justified theoretically by Conrad Dieterici  (1858-1929).  The above predicts a critical compressibility factor  (Z = pV/RT)  of 0.27067, which is more realistic than the  Van der Waals  value of  0.375.  In spite of this and other advantages, the Dieterici equation remains far less popular than the Van der Waals equation.

Carnahan-Starling equation  (1969)  and beyond :

In 1969,  Norman F. Carnahan (b.1942)  and  K.E. Starling proposed the following 2-parameter equation of state,  expressed in terms of  y = b/4V :

p V   =   RT  ( 1 + y + y2 - y3 ) / (1 - y) 3

The CS equation is based on a model of rigid nonattracting spheres and directly improves upon the aforementioned  Clausius equation of state.

To account for the interaction between molecules,  the following 3-parameter equations have been proposed as refinements of the  Van der Waals  and  Dieterici  approaches, respectively:

( p + a / V2 )  V   =   RT  ( 1 + y + y2 - y3 ) / (1 - y) 3
p V   =   RT exp(-a / RTV)  ( 1 + y + y2 - y3 ) / (1 - y) 3

See:   The Dieterici alternative to the van der Waals approach...   by  Richard J. Sadus  (2002).

All equations of state featuring three parameters  (a,b,R)  are identical for all substances, if stated in units of the critical temperature, pressure and volume.

A few characteristics of some actual gases :
SubstanceCritical StateInversion Boiling

(1 atm)
Tcpcmolar  Vc Ti
WaterH2O 647.10 K22.07  MPa55.95 mL  373.15 K
MethaneCH4 190.56 K4.59  MPa98.60 mL  111.55 K
CFC-12 CCl2F2 385 K4.12  MPa209 mL  243 K
NitrogenN2 126.21 K3.39  MPa90.08 mL 621 K77.35 K
OxygenO2 154.58 K5.036 MPa73.39 mL 764 K90.19 K
 CO2 304.19 K7.38  MPa91.90 mL  Sublimation
Air 603 K 
HydrogenH2 32.98 K1.315 MPa66.93 mL 202 K20.28 K
HeliumHe 5.19 K0.227 MPa57.48 mL 40 K 4.216 K
NeonNe 44.40 K2.76  MPa41.70 mL 231 K27.10 K
ArgonAr 150.87 K4.898 MPa75.23 mL 723 K87.29 K
KryptonKr 209.45 K5.50  MPa91.19 mL  119.93 K
XenonXe 289.73 K5.80  MPa118.28 mL  165.05 K
RadonRa 378.15 K6.28  MPa   211.30 K
g2hn (Yahoo! 2007-07-07)   How do I obtain the (molar) Van der Waals constants  (a, b, R) for ethane,  knowing the critical density is  0.21 g/cc  at  tc = 32.1°C  and  pc = 48.8 atm?

The mass of a mole of ethane  (C2H6  CAS 74-84-0)  being  30.07 g  the critical molar volume is  Vc = 30.07 / 0.21 = 143.19 cc.

Therefore,  b = Vc / 3 = 47.73 mL = 0.00004773 m3.

Since  pc = 48.8 atm = 4944660 Pa, we have   a = 27 b2 pc = 0.30415 J.m3.

Tc = 273.15 + 32.1 = 305.25 K  gives  R = 8a / (27 b T) = 6.18534 J/K.

In this context,  R  is just another adjustable parameter but it's useful to check that its value is not too far from the  ideal gas constant :  8.3145 J/K.

Negative pressure in liquids:   Trees are Freaking Awesome!  by  Derek A. Muller  (Veritassium, 2012-10-30).

(2006-11-11)  Virial  equation of state   [molar]
pV   =   RT  [ 1 + B(T) / V + C(T) / V2 + D(T) / V3 + ... ]

This general expansion was proposed in 1885 by  Max Thiesen  (1849-1936).

In 1901,  H. K. Onnes  (1853-1926; Nobel 1913)  coined the name  second virial coefficient;  for  B(T).  C(T)  is the third virial coefficient, etc.

We're following the usual academic practice to equate the first virial coefficient to unity.  For a "best fit", this would assume that the constant R may be adjusted at will...  In practice (especially when comparing several actual gases) it can be better to use the standard value of the gas constant for R and allow values of the first virial coefficient which may differ (slightly) from unity.

Under the proper mathematical assumptions, the ith virial coefficient is ascribed to those mutual interactions within a set of  i  molecules which do not reduce to interactions within a smaller set  (in a way loosely reminiscent of inclusion-exclusion enumeration).

The  virial expansion  of the  (molar)  Van der Waals equation of state  is:

pV   =   RT  [ 1 + (b-a/RT) / V + b2/V2 + b3/V3 + b4/V4 + ... ]

An equation of state where all virial coefficients would vanish beyond the second one would  not  be good enough to predict a liquid-vapor phase transition.

A fluid so described could not have a critical state  (where the first and second isothermal derivatives of pressure over volume vanish together).  Proving this is an interesting little exercise, which is left to the reader...

Boyle Temperature :

Real gases obey the Boyle-Mariotte law only  approximately.  At low pressures (i.e., large molar volumes) the approximation is tightest for a temperature which makes the second virial coefficient  B(T)  vanish.  Such a temperature is a characteristic of a given gas called its  Boyle temperature.

The Boyle temperature for a  Van der Waals gas  is thus:  TB = a / Rb  which is  27/8  = 3.375 times the aforementioned critical temperature.

Experimentally however, we're told that the ratio of the Boyle temperature to the critical temperature is consistently around 2.75.  The Dieterici  equation of state  gives a ratio of about 2.52 instead.

When starting from zero pressure at a constant temperature below the  Boyle temperature,  the product pV first decreases with pressure to reach a minimum, then it increases with pressure.  On the other hand, above the  Boyle temperature,  the product pV always increases with pressure.

What's viscosity?

A viscous deformation is a lateral deformation which occurs at a rate that's proportional to shear stress.  Such a deformation is most readily observed with thick paste, but it occurs much faster in not-so-thick liquids (low viscosity).  Technically, tar pitch is actually a thick liquid of extremely high viscosity.  On the other hand, contrary to urban legends, glasses are amorphous solids at room temperature.  They do not behave like "supercooled liquids" (unless heated above a certain transition point) but they may "creep" like crystalline solids do.

For crystalline solids and/or rocks, what occurs [over geological periods of time] is a so-called "plastic" deformation.  The rate of plastic flow is not directly proportional to stress but rather to some increasing function of stress with a zero derivative at the origin (so that there's no plastic deformation at all when the stress is low).  For example, in what's called "Power Law Creep", the strain rate is proportional to some n-th power of stress (n>1) and to a temperature factor exp(-Q/RT), which indicates how easily crystal dislocations may move at an absolute temperature T.  For typical rocks, the exponent n is around 3 (it ranges from slightly below 2 up to 8 or so) and the "activation energy" Q is roughly 200 kJ/mol (it may range from below 100 kJ/mol up to about 500 kJ/mol).  Viscous flow is a special case (n = 1 and Q = 0) of Power Law Creep...

Quantitatively, viscosity is the ratio of shear stress to shear strain rate.  (These two are proportional in the case of what's called a Newtonian fluid.) Units of viscosity are thus homogeneous to the product of a stress (or pressure) by a time.  Therefore, the SI unit of viscosity is the pascal-second (Pa×s), also called poiseuille (Pl), which is 10 times as large as the corresponding [deprecated] CGS unit, the poise (P) of 1 dyne-second per square centimeter.  The centipoise (cP) is a surviving synonym of the millipascal-second (mPa×s), which happens to be the viscosity of water at a temperature of about 20.2°C.

  • At 20°C, the viscosity of water is about 1.005 mPa×s and decreases with temperature:  At  t°C, the viscosity in millipascal-second is within 0.2% of
    1.005 exp(-0.024475 (t-20) (1 - 0.8787 (t-20) / (t+98.5) )).
  • At 20°C, the viscosity of air is about 0.01808 mPa×s and increases with temperature; it's about 0.01709 mPa×s at 0°C.
Dynamic Viscosity of Liquids  (in  mPa×s)
mPa×s = cP0°C20°C25°C100°C
Tar Pitch  2 1011   
Glycerin 120701485 95014.8
Castor Oil 986   
Heavy Oil 450   
Olive Oil 81   
SAE 20W   5.7
SAE 10W   4.2
SAE 5W   3.9
Sea Water1.8241.070 0.279
Methanol 0.590.544 
Acetone 0.3260.306 
Diethyl Ether0.2830.2330.224 
Dynamic Viscosity of Gases
Argon 22.17 
Nitrogen 16.6017.5020.90
CO213.80 18.50
Steam  12.60

Data compiled from multiple sources.
The dynamic viscosity of  gases  is roughly proportional to the square root of the absolute temperature (T).  It changes very little with pressure.

 Human blood  at 37°C  is about 0.45 mPa×s

The above is called absolute or dynamic viscosity, as opposed to the relative or kinematic viscosity, which is defined as the quotient of the dynamic viscosity by the mass density, expressed in units that are homogeneous to the ratio of an area and a time (m2/s). 

(Dynamic viscosity)(mPa×s  or  cP)
=   (Kinematic viscosity)(cSt) * (Density)(kg / L)

The [deprecated] CGS unit of kinematic viscosity is the stokes (St), which is 10000 times smaller than the SI unit (m2/s).  The  centistokes  (cSt)  is more commonly used in practice.  A kinematic viscosity is routinely obtained directly with a [Saybolt or Engler]  viscometer,  by timing the passage of a calibrated volume of liquid under its own weight through a small aperture.  The reciprocal of [either flavor of] viscosity is called fluidity.

Laminar Flow   |   Fluid Characteristics Chart  (Engineer's Edge)
Viscosity gradings  (dynamic, kinematic, Saybolt, SAE, ISO)  by Roy Beardmore

Michel Couvreux (TransMineral USA.  2003-04-29)       Permeability :
Mortar made from our natural hydraulic lime  (NHL 5, with twice its volume of sand)  has vapor-exchange properties listed as 0.55 gram of air per hour, per square meter, per mmHg [for one-coat renders].
What's its permeability?

Short Answer:   10 perm-inches (US)   or   14.5  ng / s.m.Pa.  Read on...

Note:   Besides the obvious problem with units, this question raises the issue of converting a rating about the diffusion of air into the kind of water-vapor rating which is clearly expected in the construction business.  We must also dispose of the fact that the stated rating is about a thickness of the material which is not explicitely specified:
      In the US, the thickness of "one-coat renders" is traditionally quoted as 3/8"  (there are normally 3 coats, including a thin outer coat, totalling 7/8" in thickness: 3/8+3/8+1/8).  However, the link quoted in the question makes explicit reference to a "French standard" (which we've been unable to locate).  This seems to imply that "one-coat renders" is shorthand for a metric reference thickness of 1 cm (1 cm is only 5% more than 3/8").  We'll be assuming this much in the following article.  If you know better, please let us know...

There are two kinds of permeabilityGeologists consider the flow of water or oil through rocks and deal with hydraulic permeability (which we'll examine last) based on "Darcy's Law".  On the other hand, the diffusion permeability used in the construction industry refers to "Fick's Law", the statement that the flux of a diffusing substance is proportional to its concentration gradient.  For gases and vapors, partial pressures (more precisely,  fugacities) are used instead of "concentrations"...  Air and vapor ratings are different.  A layer of given thickness is rated according to permeance (= permeability divided by thickness):

Permeance, for "Water Vapor Transmission" (WVT):

The permeance of a vapor barrier [more precisely, a "vapor diffusion retarder" or VDR] is a measure of how fast it lets water vapor through, when its two sides have different partial pressures of water vapor.  Permeance may be stated either in perms [US perms] or in metric perms (which are about 52% larger than US perms).  These units are defined as follows:

  • 1 US perm  =  1 grain (gn) of water vapor per hour, per ft2, per inHg.
  • 1 metric perm  =  1 gram (g) of water vapor per day, per m2, per mmHg.

A perm is exactly  32399455 / 49161192, or about 0.659 of a metric perm.

Note that this "metric perm" unit is not metric at all, since it's based on nonmetric units like the day for time, and the millimeter of mercury (mmHg) for pressure.  The strict metric equivalents of both of the above units are:

  • 57.213494670945101394-  ng / s.m2.Pa   =   1 US perm.
  • 86.812682389543557170+  ng / s.m2.Pa   =   1 metric perm.

The above SI unit (ng / s.m2.Pa) is homogeneous to a time divided by a length and actually boils down to a  "picosecond per meter"  (ps / m)  but it's never expressed this way.  Unfortunately, this very metric unit has [wrongly] been dubbed a "metric perm" by some sources...

The value of a perm in SI units does not depend on temperature !  Please ignore all those well-meaning Internet sources [ 1 | 2 | 3 | 4 ] which state otherwise by listing the above equivalences as valid at 0°C, while giving slightly higher values (+0.418 %) for 23°C.  This nonsense comes from a poor understanding of units of pressure:  The millimeter of mercury (mmHg) and the inch of mercury (inHg) are units based on the conventional density of mercury, which is exactly 13595.1 g/L...  The fact that this is very close to the actual density of mercury at 0°C is not directly relevant to unit conversions...  Even less relevant is the density at 23°C (about 13538.5 g/L) which happens to be 0.418 % less [here we go].  This only means that, if we were measuring pressure with an actual column of mercury at 23°C, we would have to divide the readings in millimeters by about 1.00418 to obtain actual pressures expressed in so-called mmHg.  Got it?  (If you didn't, you're in good company, since even standardization organizations have been known to introduce such silly "units" of pressures, whose values vary with temperature, just like the actual densities of mercury and water do.)

Permeance is occasionally given in g/m2/day (DIN 53122) where a "standard" vapor-pressure difference is understood, which corresponds to a difference in relative humidity of 85% at 23°C between both sides (for example, 0%||85% or 15%||100%).  As the saturation vapor-pressure over liquid water (100% r.h.) at 23°C is about 21.080 mmHg, the standard vapor-pressure difference is 17.918 mmHg, making the above unit equivalent to about (1/17.918) of a metric perm, or 0.084683 of a US perm  (the conversion factor is about 11.8).

"Water-Vapor" Permeability:

The above permeance is equal to the [diffusion] permeability of the material divided by its thickness.  Permeability is a characteristic of the material itself, which may be expressed in "perm-inches":  If a material has a permeability of one perm-inch, a one-inch layer of it has a permeance of one perm, a two-inch layer has a permeance of  ½ perm, etc.  (For the record, units of permeability are homogeneous to a time, and a perm-inch is about 1.45322 picoseconds.)

WVT Properties of Some Construction Materials   (perm = "US" perm)
MaterialPermeance for common thicknessWater-Vapor
Still Air (ideally)1 cm 3053600120
Gypsum Board1/2" 7993340
Phenolic Foam1/2" 5261426
Tyvek HomeWrap ®  50590 
NHL5 Mortar (1:2)1 cm 2530010
NHL5 Mortar (1:2)7/8" 1113510
OSB 7/16" 2240.86
Elastosil ® SG2.2 mm 1.3150.11
Styrofoam (tm)1"1121
(*) DIN 53122 rating in g/m2/day is for 18 mmHg  (85% difference in relative humidity at 23°C)
Note:  Data was compiled from various sources and may have been adjusted for self-consistency.

The mass diffusivity of water vapor in air (D) is known to be 2.42 10-5 m2/s...  This is the coefficient to use in Fick's equation  Jx = -D (dC/dx), which relates the mass flux (J) to the gradient of mass concentration (C).  The partial pressure (p) of a gas is obtained from C by introducing the molar mass (M) into the ideal gas law:  p = RTC/M.  Thus we have:  J = -(DM/RT) (dp/dx).  In other words, the diffusion permeability of ideal water vapor is DM/RT, where R is the molar gas constant: 8.314 472(15) J/K/mol.  In the case of water vapor (M = 0.018 kg) diffusing in air at T = 300 K, this gives a permeability of about 175 ng/s.m.Pa, or 120 perm-inches, as shown in the first line of the above table.

Concerning NHL mortar, the unit given in Michel's question seems equivalent to 24 metric perms (since the time basis is an hour instead of a day), so that 0.55 of such a unit would be exactly 13.2 metric perms (about 20 US perms) if it was a true WVT rating.  Unfortunately it's explicitely specified as an air rating instead, so we must somehow convert that into a water-vapor rating by considering the different diffusion properties of air and water vapor...

If we assume that most pores in the mortar are much larger than the dimensions of the molecules of the gases involved, we have microscopic spaces in which the molecules bounce in thermal equilibrium, in much the same way they would in an open space.  Under this simplification, each molecule of an ideal gas of molar mass M would essentially diffuse at a rate proportional to its average speed, which is itself proportional to Ö(T/M) at temperature T.

  • Water vapor (H2O) has a molar mass of 18.0153 g.
  • Dry air is essentially 79% N2 (28.0135 g) and 21% O2 (31.9988 g).

Ignoring interesting details, like the nitrogen enrichment of air diffusing through mortar, we'll consider that air is roughly a simple gas of molar mass 28.96456 g.  Such a gas would diffuse about 1.268 times slower than water vapor (this is the square root of the molar mass ratio).  Accordingly, the permeance is roughly 1.268 times larger for water vapor than for air, so the WVT permeance of a 1 cm layer of NHL mortar is about 25.40 US perms (16.74 metric perms);  the permeability of 1:2 NHL5 mortar is thus very nearly equal to 10 perm-inches.

Hydraulic Permeability

Darcy's Law (of which there's a more general tensor form) states that a porous material lets a fluid through at a rate inversely proportional to its viscosity and directly proportional to both the cross-sectional area and the gradient of the applied pressure (also called hydraulic gradient).  The material's [intrinsic] permeability is simply the coefficient of proportionality in this law:

(flow rate)   =   (permeability) (area) (pressure gradient)  /  (viscosity)

The pressure gradient between two parallel sides is their pressure difference divided by the thickness of the material between them.

The permeability coefficient happens to be homogeneous to a surface area.

 Come back later, we're
 still working on this one...

ASTM E 96 Standard Test   |   ASTM E 96 Testing
J.C. Alvarez Master's Thesis: Evaluation of moisture diffusion theories in porous materials
Fick's Law   |   Laws of Gas Transport   |   Plastics and Elastomers

(Mark Barnes, UK. 2000-12-05)
What are the resonant frequencies of a given volume of enclosed air?

The amplitude U of a wave propagating at celerity V obeys the  wave equation :

 Pierre-Simon Laplace 
   2 U     =     2 U   +   2 U   +   2 U  
Vinculum Vinculum Vinculum Vinculum Vinculum
V 2 t 2 x 2 y 2 z 2
 =DU [D is the Laplacian operator]

As used above, the term "celerity" is best reserved for the "phase speed" which appears in this equation.  Hereafter, V is the  speed of sound  in air.

The elementary solutions of this wave equation whose amplitudes are zero on the walls of a rectangular box are obtained as the products of single-dimensional solutions for a segment (i.e., sine waves whose half-wavelengths are a submultiple of the segment's length, so they are zero at both ends).  Such an elementary solution of frequency n has therefore the following form for a rectangular box of dimensions A, B and C, when the origin of coordinates is one corner of the box (in this, a, b and c are integers):

U(x,y,z,t) = sin(apx/A) sin(bpy/B) sin(cpz/C) exp(2ip n t)

The Laplacian DU is thus   -[(a/A)2+(b/B)2+(c/C)2]p2U   whereas the second time derivative of U is simply   -[4n2]p2U  . As the above wave equation tells us, the ratio of those two square brackets is equal to the square of the speed of sound  V.  This gives us the following formula for a resonant frequency (n) of air in a rectangular box (where a, b and c are any positive integers):

n   =   ½V Ö a2/A2 + b2/B2 + c2/C2

For a cylindrical box, say, you would solve a 2-dimensional equation for the disk (using Bessel functions) and combine things as above into another formula for n.

The lowest resonant frequency is, of course, obtained for a = b = c = 1.  The next frequencies are not multiples of this one...  The general solution of the wave equation for the cavity is a superposition of the above vibrations (à la Fourier).

We assumed perfectly rigid walls.  The relative correction for vibrating walls will typically be as small as the ratio of the vibration amplitude (probably less than a millimeter) to the wavelength in air (about a meter at 343 Hz, for air at 20°C)

The notion that resonant frequencies are directly related to volume is a myth, but we may investigate what happens to the lowest resonant frequency (a = b = c = 1) for a constant volume L3 = ABC  (A = xL, B = yL, C = L/xy):

n   =   no Ö  ( 1/x2+1/y2+x2y2 ) / 3

This happens to have a minimum at n = no in the case of the cube (x = y = 1). For a cavity that's not too different from a cube (which may not help much with guitar design) n is thus pretty close to ½Ö3 V/L, where L is the cubic root of the volume, and V is the speed of sound...

To conclude, we may examine temperature dependency:  As the thermal dilatation of the instrument itself is very minute, the frequencies are essentially proportional to the  speed of sound V, which is itself proportional to the square root of the absolute temperature T (for an ideal gas).  At room temperature (T = 295 K), this means that a variation of one kelvin (1°C) will induce a frequency change of about 0.17%.  This translates musically into (take your pick):  0.00244 octave, 0.0293 semitones, 0.735 savart, 1.465 centitone, 2.93 cents or 2.44 millioctaves.  That's not much:  It takes a difference of about 10°C for wind instruments to be off by 30 cents (30% of a semitone). The shift is thus about one semitone between a warm interior and freezing weather outdoors.  In an orchestra, all wind instruments are affected the same way, and string instruments could adjust and/or retune, so everybody could still be in tune.  Pianos are a problem, keep them indoors...

(J. L. of Canada. 2000-11-23)   The Whole Atmosphere
The atmosphere is made of molecules, piled up on top of each other.  It starts at earth's surface and becomes indistinguishable from the vacuum of space a few hundred kilometers up.  Now, consider a 1 cm by 1 cm square sitting at sea level and all the air that's sitting on top of that square, from the surface all the way up to outer space.  What is the approximate mass of the air in that one square-centimeter column?

The short answer is that pressure comes from the weight of the column of air above the surface. The standard atmospheric pressure of 101325 Pa exerts a force of 10.1325 N on a surface of a square centimeter.  The problem is to estimate the mass from the weight. As we shall see, Earth's gravity does not vary much for most of the atmosphere, so we may use the standard value of gravity (9.80665 m/s) and come up with a mass of approximately 10.1325 / 9.80665, or roughly 1.033 kg over each square centimeter at sea level.

Indeed, Earth's gravity is approximately constant throughout the atmosphere:  At an altitude equal to 0.5% of the Earth radius (about 32km), it is only 1% less than at sea level, and most of the atmosphere is below that point (more than 99% of it is said to be below an altitude of about 70 km, where gravity is only down 2%).

Consider a column of air at rest (it's essentially part of a slim cone whose apex is at the center of the Earth).  The weight of all the molecules of air ultimately results in a single force pS exerted by the pressure p at sea-level on the surface area S at sea-level. Now the weight (that's force, not mass) of each molecule of mass m is slightly less than the weight at sea level mg (g being gravity at sea-level) so that the mass is slightly more than pS/g, but not much more, as previously observed.

Taking the standard values g=9.80665 m/s2 and p=101325 Pa, a surface S=1 cm2= 0.0001 m2 would have approximately 1.03323 kg of air above it. As we've seen, this is actually a slight underestimate of the true value, so we may state that a cm2 at sea-level has slightly more than a kilogram of air above it.

The interesting thing is to estimate the total mass of the atmosphere. The surface of the Earth's is very close to 5.1 1014 m2 (that's the surface of the reference ellipsoid and that's also very close to the surface area of a sphere whose radius is the "conventional" radius of the Earth, namely 6371 km).  Therefore, the above gives a total mass of about 5.27 1018 kg.

Now there's a problem: the mass of the atmosphere is listed as 5.136 1018 kg in my copy of the 1995 CRC Handbook of Chemistry and Physics (page 14-7).  We should not have been troubled at all about the small (2.6%) discrepancy from our crude estimate, except that it turns out to be in the wrong direction: If the variation of gravity with altitude was the only factor to correct, the true number should be higher, not lower.  You may remark that more air is packed at the poles (the temperature's lower, the density's higher) with a larger surface gravity (9.832186 m/s2 at the poles).  This makes a correction in the right direction, but it cannot exceed 0.26% (that's how much the gravity at the pole exceeds the standard value of g) whereas our discrepancy is 10 times as large.  What else could be wrong?  Well, our value of the "standard" atmospheric pressure was precisely designed to make this particular computation work out right and it is not at fault...

The reason for the discrepancy is that we overlooked land masses in the computation! Whatever volume is occupied by land is not available for air and does decrease the total mass of the atmosphere.  To get a 2.6% difference in the mass of the atmosphere, we would need the average land (30% of the surface of the globe) to decrease by about 9% the mass of air above it, as compared to sea level. That looks like a pretty reasonable figure, and so does the CRC value of 5.136 1018 kg for the mass of the entire Earth's atmosphere.

The molar mass of air is about 28.966 g/mol, so there are about 1.773 1020 moles in the atmosphere, that's a grand total of about 1.068 1044 individual molecules. The square root of that is 1.033 1022, which is 0.01716 moles, or the number of molecules in 0.497 g of air. This occupies a volume of about 400cc at 11°C under 1 atm. In the folklore of physics, this volume was once known as "Caesar's last breath".  This volume is about equal to a human breath, so each time we inhale, we take in about one of the molecules that Julius Caesar last exhaled more than 2000 years ago (the 2000-year delay implies that air has been thoroughly mixed since then).  This is so because the whole atmosphere is to a human breath (400cc) what a human breath is to a single molecule.

We could spoil everybody's fun by pointing out that quantum mechanics tells us that identical molecules cannot be distinguished, even in principle, so that the whole concern is really fallacious (you just can't paint a molecule red)...

(2016-07-14)   Composition of dry air at sea-level
Surveyors  assume the concentration of carbon dioxide to be  450 ppm.

Calculating the average "molar mass of air"  (28.966 g/mol)  from its main constituents at sea-level is fairly delicate.  The value used on  Numericana  is the sum of the last column in the following table, properly rounded:

Simplified composition of dry air at sea-level :
Gasmol/molair g/molg/molair
  N2  0.78076  28.0134  21.8717
  O2  0.20944  31.9988  6.70183
  Ar  0.009322  39.9481  0.372396
  CO2    0.000450  44.0095  0.0198043
  Ne  0.00001818  21.1797  0.00036687
  He  0.00000524  4.026  0.0000209736
  CH4  0.00000222  16.04  0.0000356088
  Kr  0.00000114  83.798  0.0000955297
  H2  0.00000055  2.01588  0.00000110873
  N2O  0.00000031  44.013  0.000013644
  CO  0.00000020  28.01  0.000005602
  Xe  0.000000087 131.293  0.0000114225
  O3  0.000000040  47.9982  0.00000191993
  NO2  0.000000020  46.0055  0.00000092011
  SO2  0.000000010  64.0638  0.000000640638
  NH3  0.000000003  17.0305  0.0000000510916
  Air  1.000000000   28.966  28.9663256

The above is a simplified version of the actual composition of the atmosphere  (the complex equlibrium between nitrogen oxides and ozone is summarized by two species of fixed concentrations).  It has been adjusted to make the rounded fractions have an exact sum of 1,  while matching the standard molar cocentration of  450 ppm  for  CO2  quoted by geodetic surveyors.  The above yields a value of 28.9663 g/mol for air.  (another "standard" using slightly different values, including 397 ppm for carbon dioxide, yields 28.9657 g/mol for air).

Composition of Dry Air

(2016-08-03)   Humidity.  Moisture content of  clear  atmospheric air.
Relative humidity,  absolute humidity and specific humidity.

Moist air can be viewed as a mixture of water vapor and  dry air  (consisting itself of fixed molar fractions of all other atmospheric gases, including carbon dioxide).  The molar ratio of those two components is very nearly equal to the ratio of their respective partial pressures.  Dalton's law (1801)  states that the sum of the partial pressures is equal to the total pressure measured by an ordinary barometer.

Relative Humidity   (RH, j) :

At some given ambient temperature  T,  the partial pressure of water vapor  pv  cannot exceed the  saturation pressure  p*(T)  at which vapor would condensate into liquid droplets  (as dew, fog or clouds).  The composition of clear moist air is thus most commonly specified in terms of its  relative humidity  (RH)  defined as the following ratio, which can vary between  0  and  1  and is usually expressed as a percentage  (0 to 100%).

j   =   RH   =   pv / p*(T)

The function  p*(T)  can be obtained by integrating  Clapeyron's equation.

p*   =    ò   L    dT
Vinculum Vinculum

 Come back later, we're
 still working on this one...

Wikipedia :   Humidity   |   Equilibrium vapor pressure  =  p*(T)

 Montgolfier (2008-06-04)   The first hot-air balloon flew 225 years ago.
First public demonstration of a  Montgolfière :  June 4, 1783.

Joseph Montgolfier (1740-1810) and  Etienne Montgolfier (1745-1799) were two of the 16 children of a wealthy French paper manufacturer.  On June 4, 1783, they made the first public demonstration of an hot-air balloon which they privately called  Seraphina.  It was  11 m across and its sections were held together by more than 2000 buttons...  The balloon was not tethered and had no payload but it was intended as a prototype of an  aérostat  large enough to take people into the air.  Seraphina rose more than  1000 m  in the sky of Annonay,  before landing in a vineyard about  2 km  away from the take-off point.  That event was commemorated worlwide by Google on June 4, 2008, with the following logo:

 Google Logo for June 4, 2008

The first manned  untethered  flight, however, occurred several months later:

On November 21, 1783, a silk balloon made by  Etienne Montgolfier  took off at 13:54 pm, carrying two people aboard a wicker nacelle:  the physicist  Jean-François Pilâtre de Rozier (1754-1785)  and the marquis  François-Laurent d'Arlandes (1742-1809).  They flew from the  Jardins de la Muette  to  la Butte-aux-Cailles  in Paris  (a distance of about 9 km).  The flight lasted 28 minutes and an altitude of  960 m  was reached.  The balloon had a height of roughly  23 m,  a diameter of  15 m  and a volume of  2200 m3

A balloon of that size would contain about  2800 kg  of air.  So, by Charles' law and Archimedes' principle, 280 kg  would be lifted by raising the temperature of that air  10%  (roughly,  30°C  or  54°F  above atmospheric temperature).  At first, the Mongolfier brothers thought that burning wool and straw released an unknown light gas  (dubbed "gaz Montgolfier")   de Saussure  but the physicist Horace de Saussure (1740-1799) pointed out that the effect was purely thermal.  The Montgolfier brothers had experimented with hydrogen  (the newly discovered "flammable air")  but could not design a membrane impermeable to hydrogen.  The erroneous early reports that their first successful balloons were filled with a light gas had one happy consequence:  The physicist Jacques Charles (1746-1823) rushed to "reproduce" their success and came up with the first successfull hydrogen balloon shortly thereafter (see below).
 Academie des sciences

The official report for the French  Académie des Sciences  was written by  Benjamin FranklinAnother report was written by d'Arlandes himself.

Less than two weeks later  (on December 1st, 1783)  Jacques Charles, accompanied by Marie-Noël Robert, would fly from the  Jardin des Tuileries  with a more practical balloon  (dubbed  La Charlière)  which had been crafted by the brothers Jean and Marie-Noël Robert.  It was a much smaller balloon inflated with hydrogen and its envelope was made of silk impregnated with elastic gum.

Dirigible  balloons were conceived shortly thereafter (1784) by a brilliant mathematician and future revolutionary general, Jean-Baptiste Meusnier (1754-1793)  who is still remembered for his studies of minimal surfaces and of lines drawn on a curved surface  (Meusnier's Theorem, 1776).

The gas-balloon technology pioneered by Charles quickly supplanted hot-air balloons (which were only revived in the late 1950s).  Yet, on one glorious day of 1783, it's the  montgolfière  which captured imaginations by allowing Rozier and d'Arlandes to become the  first  humans to  fly !

Pioneering the Balloon:: 1783-1900  by  Greg Goebel

 Space-filling model of sulfur hexafluoride
(2011-07-31)   Sulfur Hexafluoride   (SF)
A very heavy gas and a good electrical insulator.

Under atmospheric pressure,  sulfur hexafluoride  is a colorless gas above 209.3 K  (-64°C).  Its molar mass of  146.055 g/mol.  makes it  5 times  heavier than air.

Being stable and non-toxic,  sulfur hexafluoride  is safe to use in popular demonstrations,  unlike the following heavier gases:

Above 64.05°C, the heaviest known gas is  uranium hexafluoride  (UF352.01933 g/mol)  which is highly toxic.

In the atmosphere,  sulfur hexafluoride  has a lifetime  (life expectancy)  of  3200 years  (that's a half-life of 3200 ln 2  =  4600 years).  It's the most potent greenhouse gas ever evaluated by the IPCC:  On time scales of 20, 100 and 500 years,  sulfur hexafluoride  is respectively 16300, 22800 and 32600 times more potent than  carbon dioxide.  In spite of its low current abundancy, the gas thus contributes to an estimated  0.2%  of the observed global warming and it's banned in Europe for all industrial uses, except in high-voltage  gas-insulated switchgear (GIS)  where it's irreplaceable.

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