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1-a brief history of nuclear physics

If we want to understand nuclear physics, we should first know its history.

Nuclear physics began with Henri Becquerel's discovery of radioactivity while investigating the phosphorescence of uranium salts in 1896. The following year, J.J. Thomson's discovery of the electron led to the realization that atoms had an inner structure, and Thomson proposed the plum pudding model of the atom, although this model was not entirely correct and we will discuss this later. In the 20th century, physicists discovered alpha, beta, and gamma rays emitted from the atom. In 1911, Otto Hahn and James Chadwick's experiments showed that beta radiation was more continuous according to the separation of the decay spectrum.

Perhaps the most important event in nuclear physics was Rutherford's discovery of the nucleus, and we will discuss this in more detail when we talk about scattering experiments. In 1932, James Chadwick discovered the neutron. Einstein's research on special relativity also led to the discovery of radioactivity in the same year that Thomson discovered the electron, leading to the realization of the internal structure of the atom, and Thomson proposed the plum pudding model of the atom, although this model was not entirely correct, and we will discuss this later.

2) Limits of nuclear physics:

The first section was just a brief history of nuclear physics. Of course, we will talk about all of these in more detail later, but before that, let's talk about the part that separates nuclear physics from quantum mechanics and atomic physics.

First of all, nuclear physics means nucleus, which means it deals with the nucleus of the atom.

It deals with events that concern the nucleus, such as binding energy between nucleons in the nucleus, nuclear radius, and nuclear forces. If the atom is considered as a whole, it would be the subject of atomic physics.

In addition, nuclear physics deals with topics such as fission, fusion, nuclear reactors, and atomic bomb structure.

3)Some basic concepts:

For those who have not previously studied some basic concepts in high school or university, let's briefly discuss some of them.

Atomic number (Z) = The atomic number is the true identity of an atom and is equivalent to the number of protons. In a neutral atom, the number of protons and electrons are equal, so the atomic number is also equal to the number of electrons.

We said that the periodic table is designed according to atomic number, which is why it is an element's identity.

Mass number (A) = Most of an atom's mass consists of protons and neutrons, and electrons are usually neglected because they are very small compared to protons and neutrons. Therefore, the mass number is taken as the sum of protons and neutrons, and the masses of protons and neutrons are usually assumed to be equal to each other because they are very close.

Isotope = Isotope atoms have the same number of protons but different numbers of neutrons, but the most important thing we need to know here is that only the physical property of neutrons changes. There will be no chemical change, so the atom will still be a member of the same element.

Isobar = Isobar atoms are those with the same mass number, but there is no similarity between their numbers of protons or neutrons; they are just the same by chance. The reason why the concept of isobar is important is that we will frequently use it in later nuclear physics articles when talking about concepts such as the condition for a nucleus to be stable and the stability valley.

4-ATOM MODELS: Physicists have created an almost perfect atom model through years of experience and work to better understand the atom, but before this modern atom model, many atom models were created, most of which you may have learned in high school. Nevertheless, we will continue to explain them. a) Dalton Atom Model: Model: Dalton's atomic model assumptions: Elements are made up of extremely small, indivisible particles called atoms. All atoms of a particular element are identical in size, mass, and chemical properties. However, atoms of one element are different from those of all other elements. Compounds are composed of atoms of more than one element. The number of atoms of each element in a given compound is always in the same simple ratio (Law of definite proportions). Chemical reactions only involve the separation, combination, or rearrangement of atoms; they do not result in the creation or destruction of atoms (Law of conservation of mass). Relevant laws: John Dalton's atomic theory allows for a better definition of chemical change topics: Law of conservation of mass: The total mass of the reactants in a chemical reaction is equal to the total mass of the products. Law of multiple proportions: If two elements form more than one compound, then the ratio of the masses of the second element that combine with a fixed mass of the first element will always be in a small whole number ratio. For example, in water (H2O), 2 grams of hydrogen combines with 16 grams of oxygen, while in hydrogen peroxide (H2O2), 2 grams of hydrogen combines with 32 grams of oxygen. Shortcomings and errors in Dalton's Atomic Theory: Not all atoms of an element are the same. At that time, the existence of isotopes could not be detected as neutron particles had not yet been discovered. All atoms of an element must have the same number of protons and electrons, although the number of neutrons may differ, which would result in a different atom of the same element. Atoms are not solid, but instead, they have a hollow structure. The smallest known particle is not an atom. Nowadays, it is claimed that there are more than seventy types of particles that make up the nucleus of an atom. Not all molecules of a compound are the same, just as not all atoms of an element are the same. The idea that atoms cannot be broken down or reassembled is contradicted by the fact that atoms can emit alpha radiation.


J. J. Thomson created a cathode ray tube in his laboratory, and as he expected, the rays emitted from the cathode were directed towards the anode. Thomson wanted to examine these rays in more detail, so he placed a fluorescent screen in front of the anode, with a small hole cut out of it. The cathode rays that hit the fluorescent screen caused small dots to light up on the screen. This led him to realize that the rays were composed of particles.

To determine whether these particles had an electric charge, he placed two parallel metal plates in their path, which were charged with opposite electric charges using a second battery. This created an electric field between the plates, and if the particles that made up the cathode rays had an electric charge, their path should have been deflected. When he performed the experiment, he observed that the path of the cathode rays was indeed deflected, and the deflection occurred in the direction of the positively charged plate. Since opposite charges attract each other, it was concluded that the particles that make up the cathode rays were negatively charged.

Thomson needed more information to determine the characteristic properties of the particle and its velocity, which could be useful for his work. He created a magnetic field that would affect the path of the particle coming out of the cathode and make it move in a straight line, as if it were not under any influence. By using the magnitudes of the electric and magnetic forces, which are opposite in direction and thus do not deflect the particle, he was able to obtain the velocity information through the energy equivalence. Later on, using the equivalence of forces, he would also be able to determine the charge-to-mass ratio of the particle. His calculation and reasoning were entirely correct, and the value he found was very close to the actual value. He repeated the experiment under different conditions, especially by changing the cathode material and the gas inside the tube, but the result never changed. The particle with a negative charge, regardless of the material, was a fundamental particle, and Thomson deemed it appropriate to name it the "electron."

Thomson's experiment and subsequent fundamental physics calculation are seen as an important step in the development of the atomic theory because they led to the creation of a new atomic model. Thomson discovered the electron, which dealt a heavy blow to Dalton's indivisible atoms. Regardless of the material used in the experiment, the results did not change, indicating that each element's atoms should not be completely different from one another, as Dalton had proposed. In each atom, the electron that he discovered should be able to find its place, and since this electron could leave the atom and move around inside the tube, the idea of the atom's indivisibility had to be abandoned. On the other hand, the electron was a negatively charged particle, while atoms were neutral. Therefore, there had to be positively charged particles inside the atom to balance this charge. Another observation was that the electron's charge-to-mass ratio was very high, indicating that the electron had a very small mass. In Thomson's experiment, electrons moved at low (~0.1c) speeds, so calculations of the m/e ratio could be made without the need for special relativity.[1] Based on all this information, Thomson created a new atomic model, which stated that the atom was made up of positively charged matter and embedded electrons with a negative charge:

1:The atom is made up of positively charged matter.

2:Electrons are embedded in this positive matter and do not move.

3: The mass of the electrons is very small, so all the mass of the atom is made up of the positive matter.

4:The atom is spherical in shape.

Thomson did not propose in the Silliman lectures given at Yale University in 1903 that electrons should be embedded in a maddening continuum of positive charge, motionless and dispersed in forms like raisins in a cake. At about the same time, a Japanese physicist in Tokyo named Hantaro Nagaoka proposed a "Saturnian model". According to this model, electrons were orbiting around the central positively charged matter, just like the rings around Saturn or the planets around the Sun. It is known today that this model is closer under the supervision of Nagaoka. This model of Thomson was corrected by Ernest Rutherford in 1911 with the planet or core model.[1]

c) Rutherford Atomic Model

Perhaps this is the atomic model that we should talk about the most, but we will talk about it in more detail later when we talk about it in scattering experiments.

Geiger and Geiger Marsden in Rutherford's students By sending +2 charged alpha particles ( +2) to a gold layer behind which a film was placed, the paths of the rays after they hit the plate were drawn.

Geiger Marsden calculated the diameter of the atom with a very small deviation in his experiment.

The Rutherford model of the atom is likened to the Solar System. Most of the rays sent through the sun passed directly through the plate. It has been likened to a proton-filled nucleus and the electrons revolving around it to planets.

As a result of the experiment, the following findings were reached:

There are large voids in the atom.

It has been observed that a small part of the rays are refracted and a very small part is reflected. From here, it was learned: + (positive) charges in the atom are collected in a small volume called the nucleus.

He stated that most of the atomic mass is concentrated in the nucleus.

There are as many electrons as positively charged particles in the atom, around the nucleus, and electrons occupy most of the atom's volume.

Shortcomings and Errors of the Rutherford Atomic Model

It did not find the neutron (but predicted its existence).

He could not fully explain the location and movement of electrons.

With the Rutherford atomic model, the foundations of the Modern Atomic Theory and the Bohr atomic model were laid.

d)satürn atom modeli:

The Japanese physicist Nagaoka, whom few people know, discovered the nucleus of the atom before Rutherford and suggested it, but it was not given much attention.

Nagaoka's model made two predictions:

a very large atomic center (by analogy with a massive planet)

electrons bound around the center by electrostatic forces (in analogy with the gravitationally bound ring orbiting Saturn).

Both predictions were confirmed by Ernest Rutherford in his 1911 paper proposing the atomic nucleus, which also mentions Nagaoka's model. But other details of the model were wrong. In particular, the electrically charged rings will be unstable due to the impulse interruption. Nagaoka abandoned the model he proposed in 1908.

Rutherford and Niels Bohr presented the more viable Bohr model in 1913.


Assumptions of Bohr's theory

In 1913, Niels Bohr proposed the Bohr theory using the spectrum lines of the hydrogen atom and Planck's quantum theory. In the light of this information, Bohr's postulates can be summarized as follows:

Electrons in an atom move in orbits at a certain distance from the nucleus, and the angular momentum in these orbits is integer multiples of h/2 (where h is Planck's constant and is the number pi). Every steady state has a constant energy.

At any stable energy level, the electron moves in a circular orbit. These orbitals are called energy levels or shells.

While the electron is in one of the stable states, the atom does not emit radiation. However, when it moves from a higher energy level to a lower energy level, it emits a quantum (packet) of light equal to the energy difference between the levels. Here E = Eson-Eilk relation is valid.

Stable levels at which electron movement is possible are designated by letters such as K, L, M, N, O or, with the lowest energy level being 1, each energy level + an integer and generally denoted by "n". (n: 1,2,3 ...¥)

Deficiencies and mistakes in the Bohr Atomic Model

Since electrons are very fast, not only classical physics but also relativity must be considered.

The Bohr Model of the Atom can only explain the spectra of single-electron atoms (hydrogen). It cannot explain the spectra of multi-electron atoms.

Wave-particle duality (De Broglie Hypothesis) is not considered in the Bohr Atomic Model.

According to the Heisenberg uncertainty principle, the position and velocity of the electron in an atom cannot be determined simultaneously with complete precision. Therefore, the concept of "orbit" is wrong.

There are no neutrons. Neutrons were discovered by James Chadwick in 1932.

It was insufficient to explain the formation of atomic and intermolecular bonds.

note: some of the articles are sampled from wikipedia (some atom models)


A-Radius of core,load distribution

B-binding energy

C-proton neutron leak lines

D-neutron stars

e-The nature of nuclear forces

G- The importance of Deuteron

F-scattering theory

H-shell model

K-nuclear vibrations

L-alpha decay

M-beta decay

N-nuclear reactions

I-nuclear reactors



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