Question: Where does an atomic number come from?
About the article "Subatomic particle", section "Dividing an atom".

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The atomic mass and relative isotopic mass are sometimes confused, or incorrectly used, as synonyms of relative atomic mass (also known as atomic weight) and the standard atomic weight (a particular variety of atomic weight).
However, as noted in the introduction, atomic weight and standard atomic weight represent terms for averages of isotopic abundances, not for single nuclides.
As such, atomic weight and standard atomic weight often differ numerically from relative isotopic mass, and they can also have different units than atomic mass when this quantity is not expressed in unified atomic mass units (see the linked article for atomic weight).
The atomic mass (relative isotopic mass) is defined as the mass of a single atom, which can only be one isotope (nuclide) at a time, and is not an abundance-weighted average, as in the case of relative atomic mass/atomic weight.
The atomic mass or relative isotopic mass of each isotope and nuclide of a chemical element is therefore a number that can in principle be measured to a very great precision, since every specimen of such a nuclide is expected to be exactly identical to every other specimen, as all atoms of a given type in the same energy state, and every specimen of a particular nuclide, are expected to be exactly identical in mass to every other specimen of that nuclide.
For example, every atom of oxygen-16 is expected to have exactly the same atomic mass (relative atomic mass) as every other atom of oxygen-16.
In the case of many elements that have one naturally occurring isotope (mononuclidic elements) or one dominant isotope, the actual numerical similarity/difference between the atomic mass of the most common isotope, and the relative atomic mass or (standard) atomic weight can be small or even nil, and does affect most bulk calculations.
However, such an error can exist and even be important when considering individual atoms for elements that are not mononuclidic.
For non-mononuclidic elements that have more than one common isotope, the numerical difference in relative atomic mass (atomic weight) from even the most common relative isotopic mass, can be half a mass unit or more (e.g.
see the case of chlorine where atomic weight and standard atomic weight are about 35.45).
The atomic mass (relative isotopic mass) of an uncommon isotope can differ from the relative atomic mass, atomic weight, or standard atomic weight, by several mass units.
Atomic masses expressed in unified atomic mass units (i.e.
relative isotopic masses) are always close to whole-number values, but never (except in the case of carbon-12) exactly a whole number, for two reasons: protons and neutrons have different masses, and different nuclides have different ratios of protons and neutrons.
atomic masses are reduced, to different extents, by their binding energies.
The ratio of atomic mass to mass number varies from about 0.99884 for 56Fe to 1.00782505 for 1H.
Any mass defect due to nuclear binding energy is experimentally a small fraction (less than 1%) of the mass of a nucleon, and is an even smaller percentage compared to the average mass per nucleon in carbon-12, which is moderately strongly-bound compared with other atoms.
Since protons and neutrons differ from each other in mass by an even smaller fraction (about 0.0014 u), rounding the atomic mass of any given nuclide to the nearest whole number always gives the nucleon count, or mass number.
The neutron count (neutron number) may then be derived by subtracting the number of protons (atomic number). 
The amount that the ratio of atomic masses to mass number deviates from 1 is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number.
This corresponds to the fact that nuclear fission in an element heavier than zirconium produces energy, and fission in any element lighter than niobium requires energy.
On the other hand, nuclear fusion reactions: fusion of two atoms of an element lighter than scandium produces energy, whereas fusion in elements heavier than calcium requires energy.
The difference mass number minus atomic mass is not maximized at iron-56 (where it is around 0.06 atomic mass units), but at heavier elements where it reaches about 0.10 atomic mass units.
Here are some values of the ratio of atomic mass to mass number: 
 Relative atomic mass (not to be confused with relative isotopic mass) is a synonym for atomic weight, and in some circumstances may even be synonymous with "standard" atomic weight (depending on the sample, see below).
It is an average atomic mass, or the weighted mean of the atomic masses of all the atoms of a particular chemical element found in a particular sample, which is then standarized by comparison to carbon-12.
Relative atomic mass is frequently used as a synonym for the "standard" atomic weight and it is correct to do so if the relative atomic mass used is that for an element from Earth under defined conditions.
However, relative atomic mass covers more than standard atomic weights, and is a less specific term that may more broadly refer to non-terrestrial environments and highly specific terrestrial environments that deviate from Earth-average or have different certainties (number of significant figures) than do the standard atomic weights.
"Standard" atomic weight refers to the relative atomic mass or atomic weight of an element "in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances."
Because these standard atomic weights are an average (mean) of relative isotopic masses for a given element from different sources (places on Earth), standard atomic weights are subject to natural variation.
An uncertainty in brackets is therefore included in sources of standard atomic weights, which often reflects natural variability in isotopic distribution for an element, rather than uncertainty in measurement.
Although there is an attempt to cover the range of variability on Earth with standard atomic weight figures, there are known cases of mineral samples which contain elements with atomic weights that are outliers from the standard atomic weight range.
Lithium represents a unique case where the natural abundances of the isotopes have in some cases been found to have been perturbed by human isotopic separation activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources, such as rivers.
For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning.
Therefore, for synthetic elements the total nucleon count of the most stable isotope (i.e., the isotope with the longest half-life) is listed in brackets, in place of the standard atomic weight.
When the term "atomic weight" is used in chemistry, usually it is the more specific standard atomic weight that is implied.
It is standard atomic weights that are used in periodic tables and many standard references in ordinary terrestrial chemistry.   
All consideration of nuclear electrons ended with James Chadwick's discovery of the neutron in 1932.
An atom of gold now was seen as containing 118 neutrons rather than 118 nuclear electrons, and its positive charge now was realized to come entirely from a content of 79 protons.
After 1932, therefore, an element's atomic number Z was also realized to be identical to the proton number of its nuclei. 
Internal conversion is a radioactive decay process where an excited nucleus interacts electromagnetically with an electron in one of the lower atomic orbitals, causing the electron to be emitted (ejected) from the atom.
Thus, in an internal conversion process, a high-energy electron is emitted from the radioactive atom, but not from a nucleon in the nucleus.
Instead, the electron is ejected as a result of an interaction between the entire nucleus and an outside electron that interacts with it.
For this reason, the high-speed electrons from internal conversion are not beta particles, since the latter come from beta decay, where they are newly created in the process.
Since no beta decay takes place during internal conversion, the element atomic number does not change, and thus (as is the case with gamma decay) no transmutation of one element to another is seen.
However, since an electron is lost, an otherwise neutral atom becomes ionized.
Also, no neutrino is emitted during internal conversion.
Internally converted electrons do not have the characteristic energetically spread spectrum of beta particles, which results from varying amounts of decay energy being carried off by the neutrino (or antineutrino) in beta decay.
Internally converted electrons, which carry a fixed fraction of the characteristic decay energy, have a well-specified discrete energy.
The energy spectrum of a beta particle is thus a broad hump, extending to a maximum decay energy value, while the spectrum of internally converted electrons has a sharp peak.