Hydrogen is the chemical element with the symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table.Hydrogen is the most abundant chemical substance in the universe, constituting roughly 75% of all baryonic mass. Parameter Gaseous Normal Hydrogen Gaseous para-Hydrogen; density at 0 °C, (mol/cu cm)X10+3: 0.04460: 0.05459: compressibility factor at 0 °C: 1.00042: 1.0005. The atomic mass is useful in chemistry when it is paired with the mole concept: the atomic mass of an element, measured in amu, is the same as the mass in grams of one mole of an element. Thus, since the atomic mass of iron is 55.847 amu, one mole of iron atoms would weigh 55.847 grams. The nucleus of this isotope consists of only a single proton (atomic number = mass number = 1) and its mass is 1.007825 amu. Hydrogen is generally found as diatomic hydrogen gas H 2, or it combines with other atoms in compounds —monoatomic hydrogen is rare. The H–H bond is one of the strongest bonds in nature, with a bond dissociation.

The major stable isotope ¹H is also known as protium. The minor stable isotope ²H is known asdeuterium, with symbol D. The low concentration of ²H in normal sources of hydrogen may have delayed itsdiscovery until 1931 when hydrogen isotope fractionation was demonstrated by distillation, electrolysis, evaporation,and in environmental samples. The atomic weight of hydrogen has been based on mass-spectrometricmeasurements since 1938. In its report for 1961, CIAAW recommended A_r(H) = 1.007 97(1) basedon the average and the range of ²H concentrations measured in hydrogen extracted from fresh- and saltwaters; however, it was noted that substances other than water could have a wider range of atomicweights.

The currently accepted best measurement of the isotopic composition of hydrogen from asingle natural source was performed on VSMOW (distributed by the IAEA and NIST), the referencematerial endorsed by CIAAW as the basis of the delta scale for relative isotope-ratio measurements. According to this measurement, VSMOW has an amount fraction of x(²H) = 0.000 155 74(5), correspondingto A_r(H)_VSMOW = 1.007 981 75(5). The uncertainty of that value corresponds to a δ²H uncertaintyof 0.3 ‰, which is equal to or slightly smaller than typical uncertainties of most relative isotope-ratio measurements of H. Variations in the isotopic composition of hydrogen in chemicals and naturalterrestrial systems are known to exceed +1000 ‰, which is much larger than the uncertainty due to isotope-ratio measurements.

For water sources, the range of published δ²H values extends from −495 ‰ (A_r(H) = 1.007 9042), to +129 ‰ (A_r(H) =1.008 0020). Seawater, the largest reservoir of water near the Earth's surface, has a relatively uniformisotopic composition and atomic weight near that of VSMOW; whereas precipitation, polar ice,lakes, rivers, and groundwaters have atomic weights that range widely, generally decreasing with latitudeand elevation and increasing with evaporation. The highest δ²H value reported for a material ofnatural terrestrial origin is +180 ‰ for atmospheric H₂. An unusual anthropogenic occurrence ofwater from a H₂S well yielded δ²H values as high as +375 ‰ (A_r(H)= 1.008 0404). Hydrogen samples with low atomic weights, some of which are outside the range of the standard atomic-weight uncertainties, have been reported from some types of natural and artificialH₂ gases, hence the annotation 'g'. The naturally occurring hydrogen sample with the lowest atomicweight (δ²H = −836 ‰ and A_r(H) = 1.007 8507) is H₂ gas collectedfrom a natural gas well in Kansas, USA. That sample, and other similar ²H-depleted H₂-rich naturalgases elsewhere, may have formed by natural chemical reduction of water during low-temperaturereactions with ultramafic (Fe-Mg-silicate) rocks. Hydrogen gases produced artificially by electrolysisand as by-products of petrochemical processing commonly are depleted in ²H. Though not natural,those gases are considered to be important because they are used commonly in laboratories.Commercial tank H₂ has been reported to have δ²H as low as −813 ‰ (A_r(H) = 1.007 8543).

The radioactive isotope ³H, also known as tritium with symbol T, decays by negative beta emissionto ³He with a half-life of 12.3 years. Tritium is formed naturally in the atmosphere by cosmic-rayreactions such as ¹⁴N(n,t)¹²C and artificially in nuclear reactors. Large quantities of tritium were injectedinto the atmosphere as a by-product of thermonuclear bomb tests, mostly in the 1950s and 1960s.Tritium reacts in the atmosphere to form HTO and other compounds that are distributed with widelyvarying concentrations in the near-surface environment of the earth. Those variations, and other localanthropogenic ³H anomalies, are used commonly in environmental studies; however, concentrations oftritium in normal sources of hydrogen are too low by several orders of magnitude to have a measurable effect onthe atomic weight of hydrogen.

CIAAW

Hydrogen
A_r(H) = [1.007 84, 1.008 11] since 2009
The name derives from the Greek hydro for 'water' and genes for 'forming' because it burned in airto form water. Hydrogen was discovered by the English physicist Henry Cavendish in 1766.

Natural variations of hydrogen isotopic composition

Isotopic reference materials of hydrogen.

So then, why isn't the atomic mass of Hydrogen exactly 1?

If you check a periodic table, you'll see that Hydrogen actually has a mass of 1.00794. If hydrogen is the lightest of all substances, then why not give it a mass of exactly 1 on our relative mass scale?

There are three reasons:

• First, atoms have isotopes, and these isotopes do not all have the same mass. The mass of the atoms in nature - what we use as the atomic mass - is a weighted average of all these different isotopes.
  

• The second reason is historical. Once upon a time, way back before 1961, there actually were two sets of atomic masses (though everybody called them atomic weights then). One scale was used by physicists; the other by chemists. Both were based on weights compared to Oxygen, rather than Hydrogen. Oxygen was used because it combines with a lot of things to form oxides. This made it a better choice as a standard because of the ease of chemical analysis. Oxygen was set to have an atomic mass of 16, which was just about 16 times as heavy as Hydrogen being 1. Unfortunately, Chemists picked naturally occurring Oxygen, which is a mixture of isotopes of Oxygen-16, Oxygen-17, and Oxygen-18. After all when you made an oxide of an element you would do so in naturally occurring oxygen. Physicists picked the pure isotope Oxygen-16, because they tended to make their measurements on the basis of mass spectrometry.
  Though the ratio of any two atom's masses was the same on either scale, it was horribly confusing, so in 1961, a compromise was reached. Instead of using either Hydrogen, or Oxygen as the standard, the isotope of Carbon with 6 protons and 6 neutrons in its nucleus (Carbon-12) was given a mass of exactly 12. It was a good choice, since it was in between the two previously used standards, and meant that nothing had to change too much.

Hydrogen Atomic Mass Unit

• The third reason is the most important of all. If a hydrogen atom has only one proton, and carbon-12 has 6 protons and 6 neutrons to make up its mass of twelve, why isn't the mass of hydrogen 1/12 of that of carbon-12?

  Mass of 1 hydrogen atom	Mass of sub-atomic particles	Mass of 1 carbon-12 atom
  1.00794	6 protons = 6 x 1.007277 6.043662 6 neutrons = 6 x 1.008665 6.051990 6 electrons = 6 x 0.000548 0.003288 Total 12.098940	12.0 exactly

  If you think about it, Hydrogen at 1.00794 is more than 1/12 of the weight of carbon-12 (as you can see from the above table, if you multiply 12 times the mass of a single hydrogen atom it comes to more than 12). The reason for this effect is nuclear binding energy. After all, the protons in the nucleus are all positive, and so the nucleus should just repel itself apart. It doesn't of course, so something must be 'binding' it together. This nuclear binding energy makes the mass of all atoms (except hydrogen-1, which only has 1 proton) slightly lighter that what you'd get by adding up the mass of the sub-atomic particles. Einstein's famous equation E = mc² shows us that we can get the necessary binding energy from the mass of the sub-atomic particles. So the mass of any multi-nucleon atom is less than the sum of the weights of its separated parts. Its this change in mass when the nucleus changes size that is the source of the enormous amount of energy in nuclear reactions.

So we could have set hydrogen to be exactly 1, but then we'd have had to really revise the atomic weight table back in 1961. If hydrogen was assigned a mass of 1 exactly, then oxygen would have become 15.87, quite a difference from the mass chemists were using. Choosing carbon-12 as the reference standard meant the least change was necessary. Still, if you do really accurate calculations based on the old and the new scale you can see some differences. For example, on the pre-1961 atomic weight scale the molecular weight of table salt, Sodium chloride NaCl would have been 58.45. On today's scale it is 58.44. Dbeaver not showing databases. The difference is just 0.02%, so for most purposes it wouldn't matter.

Hold it! You just used the term molecular weight. Isn't that wrong? Yes, of course it is, but for Sodium chloride, we shouldn't even use the term molecular mass. Instead we should use the term 'formula mass', because Sodium Chloride really isn't a molecule of NaCl.

Hydrogen Atomic Mass And Atomic Number

Copyright © 1998 - 2008 David Dice