Why wasn't fluorine used as a standard for atomic mass?

In summary, the concept of relative atomic mass was first introduced by scientists in the early 1800s, with the lightest element, hydrogen, being used as the standard with an atomic weight of 1.00. However, Prout's hypothesis that all atomic masses would be exact multiples of hydrogen was disproved by Jöns Jakob Berzelius, who showed that some elements had atomic masses that were not whole numbers. This was later explained by the discovery of isotopes and the neutron in the early 1900s. This led to the development of two different atomic-mass scales, one used by chemists and the other by physicists. The chemists' scale was based on a natural mixture of oxygen isotopes,
  • #1
PainterGuy
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Hi,

I found the two sections in red a little confusing. I'd appreciate it if you could help me with it.

The first scientists to determine relative atomic masses were John Dalton and Thomas Thomson between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Relative atomic mass (Atomic weight) was originally defined relative to that of the lightest element, hydrogen, which was taken as 1.00, and in the 1820s, Prout's hypothesis stated that atomic masses of all elements would prove to be exact multiples of that of hydrogen. Berzelius, however, soon proved that this was not even approximately true, and for some elements, such as chlorine, relative atomic mass, at about 35.5, falls almost exactly halfway between two integral multiples of that of hydrogen. Still later, this was shown to be largely due to a mix of isotopes, and that the atomic masses of pure isotopes, or nuclides, are multiples of the hydrogen mass, to within about 1%.

In the 1860s, Stanislao Cannizzaro refined relative atomic masses by applying Avogadro's law (notably at the Karlsruhe Congress of 1860). He formulated a law to determine relative atomic masses of elements: the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight and determined relative atomic masses and molecular masses by comparing the vapor density of a collection of gases with molecules containing one or more of the chemical element in question.[4]

In the 20th century, until the 1960s, chemists and physicists used two different atomic-mass scales. The chemists used a "atomic mass unit" (amu) scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to only the atomic mass of the most common oxygen isotope (16_O, containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to two different tables of atomic mass. The unified scale based on carbon-12, 12C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the chemists' scale. This was adopted as the 'unified atomic mass unit'. The current International System of Units (SI) primary recommendation for the name of this unit is the dalton and symbol 'Da'. The name 'unified atomic mass unit' and symbol 'u' are recognized names and symbols for the same unit.
Source: https://en.wikipedia.org/wiki/Atomic_mass#History

The existence of isotopes was first suggested around 1913[reference]. The first periodic table by Mendeleev in 1869 used non-integer atomic masses for some elements [reference].

The following two questions are about the first section in red.

Question 1:
It says, "this was shown to be largely due to a mix of isotopes". Shouldn't it be completely or wholly due to a mix of isotopes?

Question 2:
What kind of explanation the chemists of 19th century had for fractional atomic masses; such as for chlorine 35.5?

I found the following explanation. Different isotopes exists as a result of different number of neutrons in a nucleus therefore, in my opinion, saying "by the discovery of isotopes and the neutron" is little bit redundant; but, yes, the discovery of neutron around 1930 added more to the explanation.

Why "almost 1%"?

The discrepancy between Prout's hypothesis and the known variation of some atomic weights to values far from integral multiples of hydrogen, was explained between 1913 and 1932 by the discovery of isotopes and the neutron. According to the whole number rule of Francis Aston, Prout's hypothesis is correct for atomic masses of individual isotopes, with an error of at most 1%.
Source: https://en.wikipedia.org/wiki/Prout's_hypothesisThe following question is about the second section in red.

Question 3:
I don't understand why it "led to two different tables of atomic mass". Could you please help me with it?
 
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  • #2
Q1: Not wholly, it is also in small part due to atomic masses being not exact multiples of hydrogen.
Q2: I don't know much about the discovery of isotopes, but I don't think the phrase is redundant. They could have discovered isotopes (atoms of the same element with different masses) before the neutron was discovered. We know that isotopes have different numbers of neutrons in the nucleus, but they could have discovered isotopes experimentally before knowing the explanation for them.
Q3: The relative atomic mass of natural oxygen (average over all isotopes) is slightly different from that of pure 16O, so values for other elements will differ depending which is chosen as the standard for u = 16.000. On the 12C scale, 16O has an atomic weight of 15.995, while natural oxygen is 15.999.
 
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  • #3
Thank you!

mjc123 said:
Q1: Not wholly, it is also in small part due to atomic masses being not exact multiples of hydrogen.

You are right. One of the reasons might be that hydrogen (i.e. hydrogen-1, protium or light hydrogen) is one proton and one electron atom. Other atoms contain neutrons and more number of electrons. A neutron is a bit heavier than proton.

mjc123 said:
Q2: I don't know much about the discovery of isotopes, but I don't think the phrase is redundant. They could have discovered isotopes (atoms of the same element with different masses) before the neutron was discovered. We know that isotopes have different numbers of neutrons in the nucleus, but they could have discovered isotopes experimentally before knowing the explanation for them.

Actually this is what happened. They discovered isotopes experimentally around 1913 and neutron was found quite later around 1930.

mjc123 said:
Q3: The relative atomic mass of natural oxygen (average over all isotopes) is slightly different from that of pure 16O, so values for other elements will differ depending which is chosen as the standard for u = 16.000. On the 12C scale, 16O has an atomic weight of 15.995, while natural oxygen is 15.999.

In the 20th century, until the 1960s, chemists and physicists used two different atomic-mass scales. The chemists used a "atomic mass unit" (amu) scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to only the atomic mass of the most common oxygen isotope (16_O, containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to two different tables of atomic mass.
Source: https://en.wikipedia.org/wiki/Atomic_mass#History

I think I get it now. Personally, I would say the choice of physicists was more sensible and practical. They chose pure O-16 as a standard and weighted all other elements against this standard. On the other hand, in my opinion, the choice of chemists was somewhat confusing. They were using natural mixture of oxygen isotopes as mass standard. It requires one to explain where that mixture has been taken from. I don't think natural mixture of oxygen isotopes comes in same proportion everywhere around the world. The proportion might be different in mixtures where one is taken in London and the other in Washington.

Thanks a lot!
 
  • #4
PainterGuy said:
Personally, I would say the choice of physicists was more sensible and practical. They chose pure O-16 as a standard and weighted all other elements against this standard. On the other hand, in my opinion, the choice of chemists was somewhat confusing. They were using natural mixture of oxygen isotopes as mass standard. It requires one to explain where that mixture has been taken from. I don't think natural mixture of oxygen isotopes comes in same proportion everywhere around the world. The proportion might be different in mixtures where one is taken in London and the other in Washington.

1. Oxygen was suggested by chemists in 1898, before isotopes were discovered. It was initially chosen as it easily reacts with most elements, so it was quite convenient to use as a practical reference.

2. For most elements composition of the mixture is quite constant and doesn't depend on the source of the element. B is a notable exception, as well as some actinides (difference in atomic masses/isotope composition was what lead to the discovery of Oklo nuclear reactors). Plus, this is knowledge we have now, at the time of the first attempts to codify atomic masses neither isotopes nor nuclear reactions (which are a basic reason of discrepancies) were known.
 
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  • #5
PainterGuy said:
A neutron is a bit heavier than proton.
While true, this suggests that the mass of a nucleus is the sum of the masses of the neutrons and protons that make it up. That is not correct. There is a discrepancy -- a "mass defect".

If you put a neutron and a proton together (i.e. as a deuterium atom) then the mass of the two together is a bit less than the sum of the masses of the two separately.

Neutron: 1.00866491588 u
Proton: 1.007276466621 u
Sum: 2.01594138209 u
Deuterium: 2.0141017926 u
 
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  • #6
jbriggs444 said:
While true, this suggests that the mass of a nucleus is the sum of the masses of the neutrons and protons that make it up. That is not correct. There is a discrepancy -- a "mass defect".

If you put a neutron and a proton together (i.e. as a deuterium atom) then the mass of the two together is a bit less than the sum of the masses of the two separately.

Neutron: 1.00866491588 u
Proton: 1.007276466621 u
Sum: 2.01594138209 u
Deuterium: 2.0141017926 u

Thank you for bringing this up. I didn't know about it.

Although somewhat unrelated to the original question, it wouldn't be fair if this is not addressed to complete the discussion.

For example, a helium atom containing four nucleons has a mass about 0.8% less than the total mass of four hydrogen nuclei (which contain one nucleon each). The helium nucleus has four nucleons bound together, and the binding energy which holds them together is, in effect, the missing 0.8% of mass.
[...]
The latter scenario is the case with nuclei such as helium: to break them up into protons and neutrons, one must inject energy. On the other hand, if a process existed going in the opposite direction, by which hydrogen atoms could be combined to form helium, then energy would be released. The energy can be computed using E = Δm c2 for each nucleus, where Δm is the difference between the mass of the helium nucleus and the mass of four protons (plus two electrons, absorbed to create the neutrons of helium).
Source: https://en.wikipedia.org/wiki/Nuclear_binding_energy#Mass_defect

1599203123074.png


Nuclear binding energy is the minimum energy that would be required to disassemble the nucleus of an atom into its component parts. These component parts are neutrons and protons, which are collectively called nucleons. The binding energy is always a positive number, as we need to spend energy in moving these nucleons, attracted to each other by the strong nuclear force, away from each other. The mass of an atomic nucleus is less than the sum of the individual masses of the free constituent protons and neutrons, according to Einstein's equation E=mc2. This 'missing mass' is known as the mass defect, and represents the energy that was released when the nucleus was formed.
Source: https://en.wikipedia.org/wiki/Nuclear_binding_energy

Nuclear binding energy is the energy required to disassemble a nucleus into the free, unbound neutrons and protons it is composed of. It is the energy equivalent of the mass defect, the difference between the mass number of a nucleus and its measured mass.[5][6] Nuclear binding energy derives from the nuclear force or residual strong force, which is mediated by three types of mesons.
Source: https://en.wikipedia.org/wiki/Binding_energy#Types_of_binding_energy

This article is about the force that holds nucleons together in a nucleus. For the force that holds quarks together in a nucleon, see Strong interaction. Not to be confused with weak nuclear force.

The nuclear force (or nucleon–nucleon interaction or residual strong force) is a force that acts between the protons and neutrons of atoms. Neutrons and protons, both nucleons, are affected by the nuclear force almost identically. Since protons have charge +1 e, they experience an electric force that tends to push them apart, but at short range the attractive nuclear force is strong enough to overcome the electromagnetic force. The nuclear force binds nucleons into atomic nuclei.

[...]

The nuclear force as a residual of the strong force
The nuclear force is a residual effect of the more fundamental strong force, or strong interaction. The strong interaction is the attractive force that binds the elementary particles called quarks together to form the nucleons (protons and neutrons) themselves. This more powerful force, one of the fundamental forces of nature, is mediated by particles called gluons. Gluons hold quarks together through color charge which is analogous to electric charge, but far stronger. Quarks, gluons, and their dynamics are mostly confined within nucleons, but residual influences extend slightly beyond nucleon boundaries to give rise to the nuclear force.
Source: https://en.wikipedia.org/wiki/Nuclear_force#The_nuclear_force_as_a_residual_of_the_strong_forceRelevant links:
https://physics.stackexchange.com/questions/360576/relation-between-strong-forces-and-binding-energy
https://physics.stackexchange.com/questions/562319/what-exactly-does-the-weak-force-do?
https://physics.stackexchange.com/questions/77196/weak-force-attractive-or-repulsive
 
  • #7
PainterGuy said:
[snip] Personally, I would say the choice of physicists was more sensible and practical. They chose pure O-16 as a standard and weighted all other elements against this standard. On the other hand, in my opinion, the choice of chemists was somewhat confusing. They were using natural mixture of oxygen isotopes as mass standard. It requires one to explain where that mixture has been taken from. I don't think natural mixture of oxygen isotopes comes in same proportion everywhere around the world. The proportion might be different in mixtures where one is taken in London and the other in Washington.

Thanks a lot!
It would have made/would make most sense to use flourine, but they didn't know that at the time.

Fluorine has only one stable isotope, flourine 19, and both flourine 18 and 20 are very short lived (seconds or minutes) so do not persist if formed by cosmic radiation. A measurement of natural flourine would therefore be essentially a precise multiple of Dalton units (notwithstanding the slight reduction in the mass of nucleons up to iron/nickel).

This is the case also for indium and holmium, but no other elements (in the seconds/minutes range of decay).
 
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Related to Why wasn't fluorine used as a standard for atomic mass?

What is atomic mass?

Atomic mass is the average mass of an atom of an element, taking into account all of its naturally occurring isotopes.

What are isotopes?

Isotopes are atoms of the same element that have different numbers of neutrons in their nucleus. This results in slightly different masses for each isotope.

How is atomic mass calculated?

Atomic mass is calculated by multiplying the mass of each isotope by its relative abundance (the percentage of that isotope found in nature), and then adding together the results.

Why is atomic mass listed as a decimal on the periodic table?

Atomic mass is listed as a decimal on the periodic table because it is an average of all the naturally occurring isotopes of an element, which have slightly different masses. The decimal reflects this average.

How does atomic mass affect an element's properties?

The atomic mass of an element does not directly affect its properties. However, different isotopes of an element may have slightly different chemical and physical properties due to their different masses.

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