Atomic Structure
Discovery of Atom:
John Dalton's discovery and theory of the atom marked a foundational moment in the field of chemistry
and physics. Here’s a concise summary of his contribution in four key points:
1. Atomic Theory Proposition (Early 1800s): John Dalton proposed that all matter is made up of
small, indivisible particles called atoms. This idea was not entirely new, as ancient philosophers
like Democritus had speculated about the existence of atoms. However, Dalton provided the first
scientific basis for the concept.
2. Combination of Atoms: Dalton theorized that atoms combine in fixed ratios to form compounds.
This was based on his observations and measurements of the weights in which elements combine
to form compounds, leading to the law of multiple proportions.
3. Atoms of Each Element are Identical: Dalton suggested that all atoms of a given element are
identical in mass and properties, but different from the atoms of any other element. This was a
crucial step in distinguishing between different elements and in explaining why elements react in
consistent ways.
4. Chemical Reactions: According to Dalton, chemical reactions involve the rearrangement of
atoms. Atoms are neither created nor destroyed in chemical reactions, a principle that underpins
the law of conservation of mass in chemical processes.
Dalton's atomic theory laid the groundwork for modern chemistry by providing a clear, quantitative
framework for understanding how materials combine and react at the atomic level.
Discovery of the electron:
The electron was the first fundamental particle. It was discovered by J.J. Thomson in 1897 by utilizing
Faraday's study of electrical discharge in partially evacuated tubes, known as cathode ray tubes.
The production and properties of cathode rays were critical in the discovery of the electron. Here's an
easy-to-understand description of how cathode rays are produced and their key properties:
Production of Cathode Rays
1. Vacuum Tube: Cathode rays are generated in a vacuum tube, which is essentially a glass tube
from which most of the air has been removed to create a near-vacuum condition.
2. Electrodes: Inside the vacuum tube, there are two metal plates: one called the cathode and the
other called the anode. The cathode is negatively charged, and the anode is positively charged.
3. High Voltage: A high voltage is applied across these electrodes. This voltage causes the cathode
to emit electrons, though this was not known at the time of their discovery.
4. Emission: These emitted particles (later identified as electrons) are attracted to the positively
charged anode, creating a beam that travels from the cathode to the anode. This beam is what we
call cathode rays.
Figure: Observations in the cathode ray tube experiment
Properties of Cathode Rays
1. Travel in Straight Lines: Cathode rays travel in straight lines from the cathode to the anode within
the vacuum tube.
2. Shadow Formation: If an object is placed in the path of these rays inside the vacuum tube, it casts
a shadow on the glass tube, showing that the rays travel in straight lines.
3. Magnetic and Electric Fields: Cathode rays are deflected by magnetic and electric fields,
indicating that they have charge. The direction of deflection was used to deduce that the charge
was negative.
4. Fluorescence: When cathode rays strike certain materials coated on the end of the tube, they cause
these materials to glow or fluoresce.
5. X-ray production: They produce X-rays when they strike against the surface of hard metals like
tungsten, molybdenum, etc.
Through these observations, particularly the deflection in magnetic and electric fields, J.J. Thomson
concluded in 1897 that cathode rays consisted of previously unknown negatively charged particles, which
he named electrons. This was a groundbreaking discovery that significantly advanced the understanding
of atomic structure.
Discovery of the proton:
The production and properties of anode rays, also known as canal rays, were instrumental in the discovery
of the proton. Here's an easy-to-understand overview of how anode rays are produced and their
characteristics:
Production of Anode Rays
1. Modified Vacuum Tube: Anode rays are produced in a specialized vacuum tube equipped with a
perforated cathode (a cathode with holes). This tube is similar to the one used for generating
cathode rays but with this crucial difference.
2. Gas Inside the Tube: Unlike the cathode ray tube, which is near-vacuum, the anode ray tube
contains a low pressure of gas (like hydrogen or helium).
3. High Voltage: When a high voltage is applied between the cathode and the anode, electrons
(cathode rays) are emitted from the cathode. These electrons pass through the holes in the cathode
and collide with the gas atoms.
4. Ionization: The collision of electrons with gas atoms ionizes these atoms, stripping them of one
or more electrons and producing positively charged ions.
5. Acceleration: These positive ions are then accelerated towards the cathode due to their positive
charge, moving in the opposite direction to the cathode rays. These moving positive ions constitute
the anode rays.
Figure: Observations in the cathode ray tube experiment
Properties of Anode Rays
1. Direction: Anode rays travel in straight lines from the anode towards the cathode, opposite to the
direction of the cathode rays.
2. Dependence on Gas: The nature of the anode rays (specifically, the mass of the particles) depends
on the type of gas present in the tube. Each type of gas generates ions characteristic of its atomic
structure.
3. Electric and Magnetic Fields: Anode rays are deflected by magnetic and electric fields, which
indicates they are charged. The direction of deflection helps determine the positive charge of the
particles.
4. Mass Observation: The deflection patterns and mass measurements led to the realization that
anode rays consisted of positive particles. Each particle's mass was specific to the gas used,
corresponding to either atomic or molecular ions.
Through these experiments, notably those conducted by Eugen Goldstein in 1886 who first observed canal
rays, and later by Ernest Rutherford, scientists concluded that anode rays contained positively charged
components of atoms. Rutherford's later experiments, which involved the deflection of these particles by
thin metal foils, ultimately led to the discovery of the proton in 1917. This work fundamentally deepened
the understanding of atomic structure, identifying the proton as a core component of the atom's nucleus.
J.J. Thomson’s model:
• J.J. Thomson proposed the "plum pudding" model of the atom after he discovered the electron in
1897, which depicted the atom as a sphere of positive charge with negatively charged electrons
embedded within it.
• This model suggested that the positive charge was uniformly distributed across the atom.
• The electrons were scattered throughout to maintain electrical neutrality.
Figure: J.J Thomson’s atomic model
However, Thomson's model was eventually disproven by Ernest Rutherford's gold foil experiment in
1909, which demonstrated that atoms have a dense central nucleus, leading to the development of the
nuclear model of the atom.
Rutherford’s Nuclear Atomic Model--Discovery of Nucleus:
The alpha scattering experiment, conducted by Ernest Rutherford, is a famous experiment that led to the
discovery of the atomic nucleus. Here's an easy explanation of how it worked:
Setup:
• Alpha Particles: Rutherford used alpha particles, which are tiny, positively charged particles.
• Gold Foil: He directed these alpha particles at an extremely thin sheet of gold foil of a thickness
of 0.00004 cm.
• Detection Screen: Around the foil, he placed a screen coated with a material (ZnS) that glows
when hit by alpha particles. This helped him see where the particles ended up after hitting the foil.
Figure: Pictorial representation of Rutherford’s experiment
What Happened:
• Straight Through: Most alpha particles went straight through the foil without changing direction.
This suggested that most of the space inside the atom is empty.
• Deflected Particles: Some alpha particles were deflected (bounced off at angles). This was
unexpected because, according to earlier models of the atom, the particles should have passed
through with minimal deflection.
• Bounced Back: A few alpha particles (1 in 20000) bounced back toward the source. This was
surprising and crucial because it indicated that something very dense and heavy within the atom
was repelling the alpha particles.
Figure: Different types of deflection of alpha-particles by s thin metal foil are shown more
simply.
Conclusion:
• Discovery of the Nucleus: The deflections and bounce-backs led Rutherford to conclude that all
the positive charge and most of the mass of the atom were concentrated in a small area in the center
of the atom, which he called the nucleus.
• Mostly Empty Space: The fact that most particles went straight through without any deflection
showed that atoms are mostly space, with the electrons orbiting the dense nucleus.
This experiment was crucial because it disproved the earlier "plum pudding" model of the atom and led
to the development of the nuclear model of the atom, where the atom is mostly space with a small, dense
nucleus at the center surrounded by orbiting electrons.
Postulates of Rutherford’s Model:
Rutherford's atomic model, developed in 1911 by Ernest Rutherford, was a groundbreaking advancement
in understanding the structure of the atom. Here’s a detailed explanation of the model:
Central Nucleus: The atom has a central core, the nucleus, which contains most of the atom’s mass and
all of its positive charge.
Nucleus is Tiny: Despite containing most of the mass, the nucleus is tiny compared to the overall size of
the atom. This was a revolutionary idea because it implied that the atom is mostly empty space.
Electron Orbits: Electrons, which are negatively charged, orbit the nucleus at relatively large distances.
This arrangement helps maintain the atom’s overall electrical neutrality, with the negative charges
(electrons) balancing out the positive charge of the nucleus.
Electron-Nucleus Interaction: The attraction between the negatively charged electrons and the positively
charged nucleus keeps the electrons in their orbits, preventing them from flying away. This was an
important consideration for maintaining the stability of the atomic structure.
Nuclear Density: The nucleus is very dense. Because, when an alpha particle, which is also positively
charged, comes close enough to the nucleus, it is strongly repelled.
Compared with the solar model: Electrons revolve around the nucleus in closed orbits with a fast speed
and hence almost all the space around the nucleus is occupied by the revolving electrons. The electrons
keep revolving in orbits around the nucleus as the planets revolve around the sun. Thus, Rutherford's
model of the atom bears a close resemblance with the solar system in which the massive sun plays the role
of the massive nucleus and the planets play the role of the revolving electrons.
Drawback of Rutherford’s Nuclear Atomic Model:
• This atomic model failed to explain the stability of atoms.
• According to Maxwell's theory, during acceleration charged particles would radiate energy.
Revolving electrons will lose energy and finally fall into the nucleus.
• This model of the atom also failed to explain the existence of definite lines in the hydrogen
spectrum.
• The model didn’t give details about how electrons are arranged, which affects chemical reactions.
• It was unclear why different elements have different atomic masses and how the mass relates to
the number of protons and neutrons.
Postulates of Bohr’s Model:
• In an atom, electrons (negatively charged) revolve around the positively charged nucleus in a
definite circular path called orbits or shells.
• Each orbit or shell has a fixed energy and these circular orbits are known as orbital shells.
• The energy levels are represented by an integer (n=1, 2, 3…) known as the quantum number. This
range of quantum numbers starts from the nucleus side with n=1 having the lowest energy level.
The orbits n=1, 2, 3, 4… are assigned as K, L, M, N…. shells and when an electron attains the
lowest energy level, it is said to be in the ground state.
• Bohr introduced the idea that the angular momentum of an electron in a stable orbit is quantized.
Electron revolves only in those circular orbits far in which the angular momentum (L) is an integral
ℎ
multiple of 2𝜋
𝑛ℎ
L= mvr = 2𝜋
Where, m=mass of electron, v= velocity of electron, r= Bohr radius, n= orbital’s number which
can be in integral number 1,2,3,4……….
• The electrons in an atom move from a lower energy level to a higher energy level by gaining the
required energy and an electron moves from a higher energy level to a lower energy level by losing
energy.
ℎ𝑐
∆𝐸 = 𝐸2 − 𝐸1 = ℎ𝜈 =
𝜆
Where, ∆𝐸 is the energy difference between the two levels.
h is Planck's constant (6.626 × 10-34 J.s)
𝜈 is the frequency of the emitted or absorbed light and 𝜆 is the wavelength.
Drawbacks of Bohr’s Model:
The following are the fundamental limitations of Bohr’s Model of the hydrogen atom.
• Bohr’s model no longer observes the Heisenberg Uncertainty Principle.
• The Neils Bohr atomic version speculation considers electrons to have each recognized function
and momentum simultaneously, which is unthinkable as indicated with the aid of using Heisenberg.
• The Bohr atomic version no longer makes an accurate prediction of large-sized atoms and furnishes
enough statistics that are simplest for smaller atoms.
• This model can’t explain those atoms that have more than one electron.
• Bohr’s model does not make clear the Zeeman effect whilst the spectrum is cut up into some
wavelengths in the sight of a magnetic field.
