The discovery of subatomic particles paved the way for numerous later breakthroughs.

An element’s atomic structure describes its nucleus and electron arrangement. Protons, electrons, and neutrons make up the atomic structure of matter.

The atom’s nucleus is made up of protons and neutrons, surrounded by electrons. An element’s atomic number describes the number of protons in its nucleus.

Protons and electrons are equal in neutral atoms. However, atoms can receive or lose electrons to strengthen their stability, resulting in an ion.

The number of protons and electrons in an atom determines its atomic structure. This is why various elements have varied properties.

Atomic Models

Many scientists used atomic models to explain atomic structure in the 18th and 19th centuries. These models, each with their own advantages and disadvantages, helped shape the present atomic model. Scientists including John Dalton, J.J. Thomson, Ernest Rutherford, and Niels Bohr made major contributions to the discipline. This section discusses their theories on atomic structure.

Dalton’s Atomic Theory

The English chemist John Dalton proposed that atoms were indivisible and indestructible. He also noted that while all atoms of a given element are identical, their size and mass vary.

Chemical reactions require rearranging of atoms to generate products. Dalton’s postulates stated that atoms were the smallest particles responsible for chemical processes.

His postulates are as follows:

• Atoms make up all matter
• It is indivisible
• Some elements include only one sort of atom
• The mass of each atom varies from element to element
• A chemical process rearranges atoms
• Atoms cannot be generated nor destroyed, but can be altered
• In addition to the conservation of mass, the atomic theory explained the Laws of Multiple Proportions and Reverse Proportions

Demerits of Dalton’s Atomic Theory

• The theory couldn’t explain isotopes.
• Nothing concerning atomic structure was well described.
• Scientists later identified atom-divisible particles inside the atom.

The discovery of subatomic particles inside atoms improved our understanding of chemical species. The following subatomic particles were discovered:

Thomson Atomic Model

Sir Joseph John Thomson, an English chemist, proposed his atomic structure model in the early 1900s.

He won the Nobel Prize for discovering “electrons”. His work uses a cathode ray experiment. The experiment’s working structure is as follows:

Cathode Ray Experiment

It has a glass tube with two holes, one for the vacuum pump and one for the gas input.

The vacuum pump keeps a “partial vacuum” inside the glass chamber. The glass tube has electrodes (cathode and anode) coupled to a high voltage power supply.

Conclusions:

Thomson defined the atomic structure as a positively charged sphere with negatively charged electrons embedded in it.

The “plum pudding model” is named after a plum pudding dish in which the pudding represents the positively charged atom and the plum pieces represent the electrons.

Thomson’s atomic structure defined atoms as being electrically neutral, with equal positive and negative charges.

Limitations of Thomson’s Atomic Model: 

Limitations of Thomson’s Atomic Model:  Also, new subatomic particle findings couldn’t fit into his atomic model.

Rutherford Atomic Theory

Alpha Ray Scattering Experiment

Conclusions:

• Rutherford determined that most of the atom’s space is vacant because most rays passed through
• This is due to the positive charge repulsion with another positive charge inside the atom
• A very strong positive charge at the atom’s core deflected 1/1000th of rays. He named it the “nucleus”
• He argued the nucleus holds most of the atom’s charge and mass
• Rutherford suggested the following atomic structure based on the above findings and conclusions

• The nucleus is the core of an atom, containing most of the charge and mass
• Atomic structure is round
• Electrons orbit the nucleus in a circular orbit, just like planets do

Rutherford Atomic Model Limits

The electrons will lose all their energy if they had to rotate around the nucleus, thereby explaining the atom’s stability.

A continuous spectrum is expected if electrons constantly circle the nucleus. What we see is a line spectrum.

Rutherford’s Model, J.J. Thomson Model

Particles atomic

Protons

• Protons are positively charged atoms. A proton’s charge is 1e, or 1.602 x 10-19
• A proton weighs around 1.672 x 10-24 grams
• Protons weigh 1800 times more than e–  
• The total number of protons in an element’s atoms equals its atomic number

Neutrons

• A neutron has roughly the same mass as a proton, 1.674 x 10-24
• Neutrons have no charge and are electrically neutral
• The number of protons in an element’s nucleus is the same, while the number of neutrons varies

Electrons

• An electron’s charge is -1e, or -1.602 x 10-19 C
• An electron’s mass is 9.1 x 10-31 Kg
• Because electrons have a low mass, they are neglected when determining atom mass

Isotope Atomic Structure

Atomic nucleons are made up of nucleons. A nucleon is a proton or neutron. Each element has a unique atomic number that describes the number of protons in it. However, an element’s atomic structure might have different nucleon counts.

Elements with differing nucleon numbers (also known as mass numbers) are called isotopes. So, an element’s isotopes have the same number of protons but differ in neutrons.

The atomic structure of an isotope is characterised by its chemical symbol, atomic number, and mass number. For example, protium, deuterium, and tritium are all known naturally occurring hydrogen isotopes. These hydrogen isotopes’ atomic structures are shown below.

Elemental isotopes differ in stability. Isotopes’ half-lives vary as well. However, because they share electrical structures, they exhibit comparable chemical behaviour.

Some Atomic Structures

The total number of protons, electrons, and neutrons in an atom represents its structure. Several elements’ atomic structures are shown below

Hydrogen

Protium is the most prevalent hydrogen isotope on Earth. This isotope’s atomic and mass numbers are 1 and 1

It has one proton, one electron, and no neutrons ( total neutrons = mass number – atomic number)

Carbon

C12 and C13 are stable isotopes of carbon. 12C is the most abundant isotope at 98.9%. 6 protons, 6 electrons, 6 neutrons

The electrons are spread into two shells, the outermost (valence shell) having four electrons. Carbon’s tetravalency allows it to establish chemical connections with various elements.

Oxygen

Oxygen has three stable isotopes: 18O, 17O, and 16O. But oxygen-16 is the most common.

Oxygen atom structure: With an atomic number of 8 and a mass number of 16, this isotope has 8 protons and 8 neutrons. An oxygen atom has 6 valence electrons and 8 total.

Bohr’s  Atomic Theory

In 1915, Neils Bohr proposed his atomic model. This is the most extensively used atomic model based on Planck’s theory of quantization.

1. Atomic electrons are arranged in distinct “stationary orbits”
2. Quantum numbers can represent these shells’ energy levels
3. Electrons can move between energy levels by absorbing or shedding energy
4. An electron cannot absorb or emit energy if it is stationary
5. Electrons only orbit the nucleus in these fixed orbits
6. Stationary orbit energy is quantized

Bohr’s Atomic Theory’s Limits

1. Bohr’s atomic structure only works for single electron species like H, He+, Li2+, Be3+,…..
2. In a more precise spectrometer, each line spectrum of hydrogen was revealed to be a collection of smaller distinct lines
3. The Stark and Zeeman effects defy Bohr’s theory
4. A conjugate physical quantity cannot be measured with 100% accuracy, according to Heisenberg’s uncertainty principle. There will always be some measurement error

Drawback:

Bohr successfully measured position and momentum, two conjugate quantities (theoretically)

Stark effect: The Stark effect is the deflection of electrons in an electric field

Zeeman effect: The Zeeman effect is the deflection of electrons in a magnetic field

Matter’s Duality

The photoelectric effect indicates that electrons, which were viewed as particles, are also 

waves.Thomas Young established this with his double slit experiment.

So, since nature is symmetrical, so should light or any other matter wave be.

Quantum Numbers

• Principal Quantum number (n): It describes the orbital number or shell number of electron.
• Azimuthal Quantum numbers (l): It describes the orbital (sub-orbit) of the electron.
• Magnetic Quantum number: It describes the number of energy states in each orbit.
• Spin Quantum number(s): It describes the direction of spin, S = -½ = Anticlockwise and ½ = Clockwise.

Electronic Configuration of an Atom

The electrons have to be filled in the s, p, d, f in accordance with the following rule.

1. Aufbau’s principle: The filling of electrons should take place in according with the ascending order of energy of orbitals:

• Lower energy orbital should be filled first and higher energy levels.
• The energy of orbital α(p + l) value it two orbitals have same (n + l) value, E α n
• Ascending order of energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, . . .

2. Pauli’s exclusion principle: No two electrons can have the same four quantum numbers, and if they must be in the same energy state, they must be in opposite spies.
3. Hund’s maximum multiplicity rule: To couple degenerate (identical energy) orbitals, they must first be filled singly.

Conclusion

They are electrons, neutrons, and protons. Neutral electrons and positive protons Protons and electrons attract. Electrons are substantially smaller than protons and neutrons. The nucleus is made up of neutrons and protons. Electrons migrate in rings around the nucleus. The atom’s particles are kept together by strong forces.