Relative Atomic Mass Of Oxygen

Every atom has its own unique relative atomic mass (RAM) based on a standard comparison or relative scale e.g. It has been based on hydrogen H = 1 amu and oxygen O = 16 amu in the past (amu = relative atomic mass unit).; The relative atomic mass scale is now based on an isotope of carbon, namely, carbon-12, nuclide symbol, which is given the value of 12.0000 amu. Density is the mass of a substance that would fill 1 cm 3 at room temperature. Relative atomic mass The mass of an atom relative to that of carbon-12. This is approximately the sum of the number of protons and neutrons in the nucleus. Where more than one isotope exists, the value given is the abundance weighted average.

Naturally occurring oxygen is composed of three stable isotopes, 16 O, 17 O, and 18 O, with 16 O being the most abundant (99.762% natural abundance). Depending on the terrestrial source, the standard atomic weight varies within the range of 15.999 03, 15.999 77 (the conventional value is 15.999). Oxygen has a molar mass of 15.9994 g/mol and nitrogen has a molar mass of 14.0067 g/mol. Since both of these elements are diatomic in air - O 2 and N 2, the molar mass of oxygen gas is aprox. 32 g/mol and the molar mass of nitrogen gas is aprox.

Taken from http://www.sizes.com/units/atomic_mass_unit.htm

History of the atomic mass unit

Stanislao Cannizzaro (1826–1910), the pioneer in this field, adopted the hydrogen atom as a standard of mass and set its atomic weight at 2. Others accepted the idea of using a specific atom as a standard of mass, but preferred a more massive standard in order to reduce experimental error.

As early as 1850, chemists used a unit of atomic weight based on saying the atomic weight of oxygen was 16. Oxygen was chosen because it forms chemical compounds with many other elements, simplifying determination of their atomic weights. Sixteen was chosen because it was the lowest whole number that could be assigned to oxygen and still have an atomic weight for hydrogen that was not less than 1.

The 0=16 scale was formalized when a committee appointed by the Deutsche Chemische Gesellschaft called for the formation of an international commission on atomic weights in March 1899. A commission of 57 members was formed. Since the commission carried on its business by correspondence, the size proved unwieldy, and the Gesellschaft suggested a smaller committee be elected. A 3-member International Committee of Atomic Weights was duly elected, and in 1903 issued its first report, using the 0=16 scale.⁵ Openttd server.

Taking isotopes into account

The discovery of isotopes complicated the picture. In nature, pure oxygen is composed of a mixture of isotopes: some oxygen atoms are more massive than others.

This was no problem for the chemists’ calculations as long as the relative abundance of the isotopes in their reagents remained constant, though it confirmed that oxygen’s atomic weight was the only one that in principle would be a whole number (hydrogen’s, for example, was 1.000 8).

Physicists, however, dealing with atoms and not reagents, required a unit that distinguished between isotopes. At least as early as 1927⁶ Latest onenote version for mac. physicists were using an atomic mass unit defined as equal to one-sixteenth of the mass of the oxygen-16 atom (the isotope of oxygen containing a total of 16 protons and neutrons).

In 1919, isotopes of oxygen with mass 17 and 18 were discovered.⁷ Thus the two amu’s clearly diverged: one based on one-sixteenth of the average mass of the oxygen atoms in the chemist’s laboratory, and the other based on one-sixteenth of the mass of an atom of a particular isotope of oxygen.

In 1956, Alfred Nier (at the bar in the Hotel Krasnapolski in Amsterdam) and independently A. Ölander⁸, both members of the Commission on Atomic Masses of the IUPAP, suggested to Josef Mattauch that the atomic weight scale be based on carbon-12. That would be okay with physicists, since carbon-12 was already used as a standard in mass spectroscopy. The chemists resisted making the amu one-sixteenth the mass of an oxygen-16 atom; it would change their atomic weights by about 275 parts per million. Making the amu one-twelfth the mass of a carbon-12 nucleus, however, would lead to only a 42 parts per million change, which seemed within reason.

Mattauch set to work enthusiastically proselytizing the physicists, while E. Wichers lobbied the chemists.⁹ In the years 1959–1961 the chemists and physicists resolved to use the isotope carbon-12 as the standard, setting its atomic mass at 12.

• Every atom has its own unique relative atomic mass (RAM) based on a standard comparison or relative scale e.g. it has been based on hydrogen H = 1 amu and oxygen O = 16 amu in the past (amu = relative atomic mass unit).
• The relative atomic mass scale is now based on an isotope of carbon, namely, carbon-12, nuclide symbol , which is given the value of 12.0000 amu.
  • The unit 'amu' is now being replaced by a lower case u, where u is the symbol for the unified atomic mass unit.
    • Therefore one atom of carbon, isotopic mass 12, equals 12 u, or,
    • 1 u = ¹/₁₂th the mass of one atom of the carbon-12 isotope.
  • Note that for the standard nuclide notation, , the top left number is the mass number (12) and the bottom left number is the atomic/proton number (6).
• In other words the relative atomic mass of an element is now based on the arbitrary value of the carbon-12 isotope being assigned a mass of 12.0000 by international agreement!
  • Examples are shown in the Periodic Table diagram above.
  • Note
    • (i) Because of the presence of neutrons in the nucleus, the relative atomic mass is usually at least double the atomic/proton number because there are usually more neutrons than protons in the nucleus (mass proton = 1, neutron = 1). Just scan the periodic table above and examine the pairs of numbers.
      • You should also notice that generally speaking the numerical difference between the atomic/proton number and the relative atomic mass tends to increase with increasing atomic number. This has consequences for nuclear stability.
    • (ii) For many calculation purposes, relative atomic masses are usually quoted and used at this academic level to zero or one decimal place eg.
      • e.g. hydrogen H = 1.0 or ~1, calcium Ca= 40.0 or ~40, chlorine Cl = 35.5, copper Cu = 63.6 or ~64, silver Ag 107.9 or ~108 etc.
    • At A level, values of relative atomic masses may be quoted to one or two decimal places.
      • Many atomic masses are known to an accuracy of four decimal places, but for some elements, isotopic composition varies depending on the mineralogical source, so four decimal places isn't necessarily more accurate!
• In using the symbol A_r for RAM, you should bear in mind that the letter A on its own usually means the mass number of a particular isotope and amu is the acronym shorthand for atomic mass units.
• However there are complications due to isotopes and so very accurate atomic masses are never whole integer numbers.
• Isotopes are atoms of the same element with different masses due to different numbers of neutrons. The very accurate relative atomic mass scale is based on a specific isotope of carbon, carbon-12, ¹²C = 12.0000 units exactly, for most purposes C = 12 is used for simplicity.
• For example hydrogen-1, hydrogen-2, and hydrogen-3, are the nuclide notation for the three isotopes of hydrogen, though the vast majority of hydrogen atoms have a mass of 1. When their accurate isotopic masses, and their % abundance are taken into account the average accurate relative mass for hydrogen = 1.008, but for most purposes H = 1 is good enough!

What Is The Relative Atomic Mass Of Oxygen

The Relative Atomic Mass Of Oxygen Is 16 Explain Its Meaning

• The strict definition of relative atomic mass (A_r) is that it equals the average mass of all the isotopic atoms present in the element compared to ¹/₁₂th the mass of a carbon-12 atom (relative isotopic mass of 12.0000).
  • So, in calculating relative atomic mass you must take into account the different isotopic masses of the same elements, but also their % abundance in the element.
  • Therefore you need to know the percentage (%) of each isotope of an element in order to accurately calculate the element's relative atomic mass.
  • For approximate calculations of relative atomic mass you can just use the mass numbers of the isotopes, which are obviously all integers ('whole numbers'!) e.g. in the two calculations below.
  • To the nearest whole number, isotopic mass = mass number for a specific isotope.

• Example 1.1 Calculating the relative atomic mass of bromine and
  • bromine consists of two isotopes, 50% ⁷⁹Br and 50% ⁸¹Br, calculate the A_r of bromine from the mass numbers (top left numbers).
  • A_r = [ (50 x 79) + (50 x 81) ] /100 = 80
  • So the relative atomic mass of bromine is 80 or RAM or A_r(Br) = 80
  • Note the full working shown. Yes, ok, you can do it in your head BUT many students ignore the %'s and just average all the isotopic masses (mass numbers) given, in this case bromine-79 and bromine-81.
  • The element bromine is the only case I know where averaging the isotopic masses actually works! so beware!

Relative Atomic Mass Of Oxygen Gas

• chlorine consists of two isotopes, 75% chlorine-35 and 25% chlorine-37, so using these two mass numbers ..
• .. think of the data based on 100 atoms, so 75 have a mass of 35 and 25 atoms have a mass of 37.
• The average mass = [ (75 x 35) + (25 x 37) ] / 100 = 35.5
• So the relative atomic mass of chlorine is 35.5 or RAM or A_r(Cl) = 35.5
• Note: ³⁵Cl and ³⁷Cl are the most common isotopes of chlorine, but, there are tiny percentages of other chlorine isotopes which are usually ignore
• 
  
  
  
  How to calculate relative atomic mass with accurate relative isotopic masses
  Using data from modern very accurate mass spectrometers
  (a) Accurate calculation of relative atomic mass (need to know and define what relative isotopic mass is)
  Relative isotopic mass is defined as the accurate mass of a single isotope of an element compared to ¹/₁₂th the mass of a carbon-12 atom e.g. the accurate relative isotopic mass of the cobalt-5 is 58.9332
  

  This definition of relative isotopic mass is a completely different from the definition of relative atomic mass, except both are based on the same international standard of atomic mass i.e. 1 unit (1 u) = 1/12th the mass of a carbon-12 isotope (¹²C).
  

  If we were to redo the calculation of the relative atomic mass of chlorine (example 1.1 above), which is quite adequate for GCSE purposes (and maybe A level too), but more accurately at A level, we might do ..
  

  chlorine is 75.77% ³⁵Cl of isotopic mass 34.9689 and 24.23% ³⁷Cl of isotopic mass 36.9658
  so A_r(Cl) = [(75.77 x 34.9689) + (24.23 x 36.9658)] / 100
  = 35.4527 (but 35.5 is usually ok in calculations pre-university!)
  

  
  Atomic Structure Notes, with further RAM calculations.
  (b) Calculations of % composition of isotopes
  
  It is possible to do the reverse of a relative atomic mass calculation if you know the A_r and which isotopes are present.
  The A_r of boron is 10.81 and consists of only two isotopes, boron-10 and boron-11
  The relative atomic mass of boron was obtained accurately in the past from chemical analysis of reacting masses but now mass spectrometers can sort out all of the isotopes present and their relative abundance.

  If you let X = % of boron 10, then 100-X is equal to % of boron-11

  Therefore A_r(B) = (X x 10) + [(100-X) x 11)] / 100 = 10.81

  -X + 1100 = 1081, 1100 - 1081 = X (change sides change sign!)

  so naturally occurring boron consists of 19% ¹⁰B and 81% ¹¹B