Question: What is the charge of alpha particles?
About the article "Atomic number", section "The proton and the idea of nuclear electrons".

Candidate passages:
In the years 1899 and 1900, physicists Ernest Rutherford (working in McGill University in Montreal, Canada) and Paul Villard (working in Paris) separated radiation into three types: eventually named alpha, beta, and gamma by Rutherford, based on penetration of objects and deflection by a magnetic field.
Alpha rays were defined by Rutherford as those having the lowest penetration of ordinary objects.
Rutherford's work also included measurements of the ratio of an alpha particle's mass to its charge, which led him to the hypothesis that alpha particles were doubly charged helium ions (later shown to be bare helium nuclei).
In 1907, Ernest Rutherford and Thomas Royds finally proved that alpha particles were indeed helium ions.
To do this they allowed alpha particles to penetrate a very thin glass wall of an evacuated tube, thus capturing a large number of the hypothesized helium ions inside the tube.
They then caused an electric spark inside the tube, which provided a shower of electrons that were taken up by the ions to form neutral atoms of a gas.
Subsequent study of the spectra of the resulting gas showed that it was helium and that the alpha particles were indeed the hypothesized helium ions.
Because alpha particles occur naturally, but can have energy high enough to participate in a nuclear reaction, study of them led to much early knowledge of nuclear physics.
Rutherford used alpha particles emitted by radium bromide to infer that J. J. Thomson's Plum pudding model of the atom was fundamentally flawed.
In Rutherford's gold foil experiment conducted by his students Hans Geiger and Ernest Marsden, a narrow beam of alpha particles was established, passing through very thin (a few hundred atoms thick) gold foil.
The alpha particles were detected by a zinc sulfide screen, which emits a flash of light upon an alpha particle collision.
Rutherford hypothesized that, assuming the "plum pudding" model of the atom was correct, the positively charged alpha particles would be only slightly deflected, if at all, by the dispersed positive charge predicted.
It was found that some of the alpha particles were deflected at much larger angles than expected (at a suggestion by Rutherford to check it) and some even bounced almost directly back.
Although most of the alpha particles went straight through as expected, Rutherford commented that the few particles that were deflected was akin to shooting a fifteen-inch shell at tissue paper only to have it bounce off, again assuming the "plum pudding" theory was correct.
It was determined that the atom's positive charge was concentrated in a small area in its center, making the positive charge dense enough to deflect any positively charged alpha particles that came close to what was later termed the nucleus.
Note: Prior to this discovery, it was not known that alpha particles are themselves atomic nuclei, nor was the existence of protons or neutrons known.
After this discovery J.J. Thomson's "plum pudding" model was abandoned, and Rutherford's experiment led to the Bohr model (named for Niels Bohr) and later the modern wave-mechanical model of the atom.
Rutherford went on to use alpha particles to accidentally produce what he later understood as a directed nuclear transmutation of one element to another, in 1917.
Transmutation of elements from one to another had been understood since 1901 as a result of natural radioactive decay, but when Rutherford projected alpha particles from alpha decay into air, he discovered this produced a new type of radiation which proved to be hydrogen nuclei (Rutherford named these protons).
Further experimentation showed the protons to be coming from the nitrogen component of air, and the reaction was deduced to be a transmutation of nitrogen into oxygen in the reaction 14N + α → 17O + p  This was the first-discovered nuclear reaction.
To the pictures on the right: According to the energy-loss curve by Bragg it is recognizable that the alpha particle loses indeed more energy on the end of the trace. 
The energy of the alpha emitted in alpha decay is mildly dependent on the half-life for the emission process, with many orders of magnitude differences in half-life being associated with energy changes of less than 50% (see alpha decay).
The energy of alpha particles emitted varies, with higher energy alpha particles being emitted from larger nuclei, but most alpha particles have energies of between 3 and 7 MeV (mega-electron-volts), corresponding to extremely long and extremely short half-lives of alpha-emitting nuclides, respectively.
This energy is a substantial amount of energy for a single particle, but their high mass means alpha particles have a lower speed (with a typical kinetic energy of 5 MeV; the speed is 15,000 km/s, which is 5% of the speed of light) than any other common type of radiation (β particles, neutrons, etc.)
Because of their charge and large mass, alpha particles are easily absorbed by materials, and they can travel only a few centimetres in air.
They can be absorbed by tissue paper or the outer layers of human skin (about 40 micrometres, equivalent to a few cells deep). 
The nucleus was discovered in 1911, as a result of Ernest Rutherford's efforts to test Thomson's "plum pudding model" of the atom.
The electron had already been discovered earlier by J.J. Thomson himself, and knowing that atoms are neutral, Thomson postulated that there must be a positive charge as well.
In his plum pudding model, Thomson stated that an atom consisted of negative electrons randomly scattered within a sphere of positive charge.
Ernest Rutherford later devised an experiment that involved the deflection of alpha particles at a thin sheet of metal foil.
He reasoned that if Thomson's model were correct, the immense alpha particles would easily pass through the foil with very little deviation in their paths.
To his surprise, many of the particles were deflected at very large angles.
Because the mass of alpha particles is about 8000 times that of an electron, it became apparent that a very strong force was present that allowed the particles to be deflected.
He realized that the plum pudding model could not be accurate and that the deflections of the alpha particles could only be caused by a center of concentrated charge that contained most of the atom's mass.
Thus, the idea of a nuclear atom—an atom with a dense center of positive charge—became justified. 
  The atom's electron cloud does not influence alpha particle scattering.
Much of an atom's charge (specifically, its positive charge) is concentrated in a relatively tiny volume at the center of the atom, known today as the nucleus.
The magnitude of this charge is proportional to (up to a charge number that can be approximately half of) the atom's atomic mass - the remaining mass is now known to be mostly attributed to neutrons.
This concentrated central mass and charge is responsible for deflecting both alpha and beta particles.
The mass of heavy atoms such as gold is mostly concentrated in the central charge region, since calculations show it is not deflected or moved by the high speed alpha particles, which have very high momentum in comparison to electrons, but not with regard to a heavy atom as a whole.  
Though the word "atom" originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles.
The constituent particles of an atom are the electron, the proton and the neutron; all three are fermions.
However, the hydrogen-1 atom has no neutrons and the hydron ion has no electrons.
The electron is by far the least massive of these particles at , with a negative electrical charge and a size that is too small to be measured using available techniques.
It is the lightest particle with a positive rest mass measured.
Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges.
If an atom have more or fewer electrons than its atomic number, then is become respectively negatively or positively charged as a whole; a charged atom is called ion.
Electrons are known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details.
Protons have a positive charge and a mass 1,836 times that of the electron, at .
The number of protons in an atom is called its atomic number.
Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei.
By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton.
Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or , the heaviest of the three constituent particles, but it can be reduced by the nuclear binding energy.
Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of —although the 'surface' of these particles is not sharply defined.
The neutron was discovered in 1932 by the English physicist James Chadwick.
In the Standard Model of physics, electrons are truly elementary particles with no internal structure.
However, both protons and neutrons are composite particles composed of elementary particles called quarks.
There are two types of quarks in atoms, each having a fractional electric charge.
Protons are composed of two up quarks (each with charge +) and one down quark (with a charge of −).
Neutrons consist of one up quark and two down quarks.
This distinction accounts for the difference in mass and charge between the two particles.
The quarks are held together by the strong interaction (or strong force), which is mediated by gluons.
The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more).
The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.