Taken from http://www.sizes.com/units/atomic_mass_unit.htm
- An atomic mass unit (symbolized AMU or amu) is defined as precisely 1/12 the mass of an atom of carbon-12. The carbon-12 (C-12) atom has six protons and six neutrons in its nucleus. In imprecise terms, one AMU is the average of the proton rest mass and the neutron rest mass.
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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.
The original standard of atomic weight, established in the 19th century, was hydrogen, with a value of 1. From about 1900 until 1961, oxygen was used as the reference standard, with an assigned value of 16. The unit of atomic mass was thereby defined as 1/16 the mass of an oxygen atom. Atomic Mass of Oxygen Atomic mass of Oxygen is 15.9994 u.
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
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.
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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 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.
How do you calculate atomic mass?
To calculate the atomic mass of an element, we have to calculate how much each isotope contributes to the mass of the atom. To accomplish this, we usually use an approach called the weighted average. The weighted average takes into account the mass and percentage abundance of each isotope.
What’s percentage abundance?
It is the proportion of atoms of an isotope in a sample of an element taken from the natural world. Percentage abundance is always reported as a percentage, and it is calculated as: (number of atoms of an isotope) divided by (the total number of atoms of all isotopes of that element) multiplied by 100. Fortnite remote play. Percentage abundance usually can be divided by 100 to get fractional abundance.
How do you use weighted average to calculate atomic mass?
To use weighted average, we must take into account the mass and percentage abundance of each isotope. Let’s use the data in the following table to show how weighted average is used to calculate the atomic mass of oxygen.
Oxygen Isotopes Amu
Solution
To calculate the atomic mass of oxygen using the data in the above table, we must first
- multiply the mass of each isotope by its corresponding natural abundance (percentage abundance). But, since the abundance is in %, you must also divide each abundance value by 100.
And second,
- Sum the result to get the atomic mass of the element
Thus,
Atomic mass of oxygen = 15.995 amu (99.76/100) + 16.999 amu (.04/100) + 17.999 amu (.2/100)
= 15.956612 amu + 0.0067996 amu + 0.035998 amu
Oxygen #of Protons
= 15.9994096 amu
= 16.00 amu
Note that the abundance in percent always add up to 100 %. Meaning if the previous question had left out the percentage abundance value for oxygen-18, you could have gotten it by subtracting the sum of the percentage abundance for oxygen-16 and oxygen-17 from 100% to get the percentage abundance for oxygen-18.
That is: 100% – (99.76% + .04%) = 0.2% for oxygen-18
Generally, you can apply this approach to figure out missing percentage abundance when you know the percentage abundance values for all, but one isotope.
If you examine the atomic masses on the periodic table, you will notice that they are fractional. Why are they fractional? They are fractional because atoms exist as isotopes. And isotopes do not have the same mass and abundance in nature. As a result, to calculate the atomic mass of an element, we have to calculate how much each isotope contributes to the mass of the atom.
Why are there no units of grams attached to the atomic mass of oxygen?
There are no units attached to atomic masses because atomic masses are relative atomic masses. Relative in this sense means one thing is compared to another. So, relative atomic mass means the mass of one atom is compared to the mass of another atom.
Download doc ppt pdf search engine free software. The atom to which other atoms are compared to is usually called the standard. At present, an isotope of carbon called carbon-12 (C-12) is selected as the standard and assigned an atomic mass of exactly 12 amu, where amu stands for atomic mass units. Therefore, the mass of every other atom on the periodic table is determined by how light or heavy it is when compared to the mass of C-12.
What instrument is used to measure the relative masses of atoms?
Oxygen Mass Amu
At present, mass spectrometry is the technique used to measure the relative masses of atoms and their percentage abundance in nature. In a mass spectrometer, atoms interact with a magnetic field and separate according to their mass to charge ratio. As they separate according to this ratio, their percentage abundance and relative atomic masses can be calculated. Let’s use the following example to illustrate how the relative mass of an atom is calculated using carbon-12 as the standard.
Oxygen Amu Weight
Example
If a chemist measured a sample in a mass spectrometer and determined that the mass ratio of 28Si (silicon-28) to 12C (carbon-12) is: 2.33, calculate the mass of 28Si?
Solution
Since the mass of carbon-12 is is assigned a value of 12, then it follows that the mass of 28Si = mass ratio of (28Si/12C) multiplied by mass of 12C
Thus, mass of 28Si = 2.33 x 12 amu = 27.98 amu
To learn more about atomic mass and Avogadro’s number, click here.