Structure of Atom – Class 9 Science Best Notes

Objective:

By the end of this lesson, readers will:

• Introduction of Atom
  • Dalton’s Atomic Theory
  • Subatomic Particles
• Discovery of Subatomic Particles
  • Electron
  • Proton
  • Neutron
• Thomson’s Model of Atom
• Rutherford’s Model of Atom
• Bohr’s Model of Atom
• Atomic Number and Mass Number
• Isotopes and Isobars
• Valency

Introduction of Atom:

An atom is the smallest unit of an element that retains the chemical properties of that element. Early scientists hypothesized about the structure of atoms, leading to several theories and models.

1. Dalton’s Atomic Theory:

An atom is a tiny, indivisible particle that makes up matter.

• All atoms of a given element are identical in mass and properties, but atoms of different elements differ.
• Atoms are in a fixed ratio to form a compound.
• Atoms can neither be created nor destroyed by any chemical process.

This is an early model proposed to explain the concept of the atom. “John Dalton” gave this theory in 1808, in which Dalton explained some important points about the basic structure and properties of the atom.

A brief history of this theory:

Almost at the end of the 18th century, attempts were made to understand how different things are made up and what the smallest unit of any matter is. Then, John Dalton, an English scientist, studied the behavior of matter and chemical reactions, based on which Dalton proposed the “atomic theory.”

Main postulates:

• The smallest unit of any matter is an atom, which cannot be divided further. That is, an atom is indivisible.
• All the atoms of an element are identical. That is, the size, mass, and chemical properties of the atoms are the same.
• All the atoms of different elements are different. That means the size, mass, and chemical properties of the atoms are also different. This is the reason why there are differences in the behavior of atoms of different elements.
• When atoms come together to form compounds, they combine in fixed ratios. That is, the proportion of elements in a compound is always the same. For example, in water (H₂O) there is always a ratio of 2 hydrogen and 1 oxygen.
• In a chemical reaction, the total mass of the matter always remains constant because in chemical reactions atoms are neither created nor destroyed, they just get re-arranged.

 Importance:

Dalton’s atomic theory was the initial and important step in explaining the concept of atom. This theory is very important because:

Understood the basic concept of matter:

• This theory gave the concept of the basic building blocks of matter and explained that everything is made up of atoms.

Explained about the Law of Conservation of Mass:

• Dalton’s theory explains the concept that mass is conservative in chemical reactions, that is, there is no spontaneous creation or destruction.

Limitations:

There was no concept of subatomic particles:

• Dalton said that atoms are indivisible, but later the discovery of electrons, protons and neutrons showed that atoms are autonomously divisible.

Concept of isotopes not explained:

• According to Dalton’s theory, the mass of the atoms of an element should be the same, but the discovery of isotopes showed that the mass of atoms of the same element can be different.

Chemical bonding not explained:

• Dalton’s theory does not explain the concept of chemical bonding or the formation of bonds between atoms.

2. Subatomic Particles:

In earlier time we knew that an atom was the smallest unit but in the present time atom is further divided into 3 subatomic particles:

• Electron: The negatively charged particle present in the outer region of an atom is called an electron.
• Proton: The positively charged particle present in the nucleus of an atom is called a proton.
• Neutron: The neutral particle also present in the nucleus of an atom is called a neutron, which has no charge.

Discovery of Subatomic Particles

1. Electron:

• Electron was discovered by J.J. Thomson.
• He proved through cathode ray tube experiment that there are negatively charged particles inside the atom.

2. Proton:

• E. Goldstein proved through anode rays (canal rays) that there are positively charged particles inside the atom which are called protons.

3. Neutron:

• Neutron was discovered by James Chadwick in 1932.
• Neutron has no charge and it is present in the nucleus.

3. Thomson’s Model of Atom:

J.J. Thomson said that the atom is a sphere of positive charge with negatively charged electrons embedded in it. This model is also known as the “plum pudding model” or “watermelon model”.

A brief history of this theory:

There is a concept that was the first attempt to explain the structure of the atom. This model was proposed by J.J. Thomson in 1897 when he discovered the electron.

J.J. Thomson performed a cathode ray tube experiment in which he observed that some negatively charged particles move from the cathode towards the anode. These particles had very light mass and carried a negative charge. He named these particles as “electrons”.

After this discovery, the question arose that if there are negatively charged particles inside the atom, then where is the positive charge? Thomson proposed his model to answer this question.

2. Thomson’s Model:

J.J. Thomson said that the atom is a sphere of positive charge with negatively charged electrons embedded in it. This model is also known as the “plum pudding model” or “watermelon model”.

• Positive Sphere: He thought that the main structure of the atom is a large sphere of positive charge that keeps the whole atom stable.
• Electrons as Embedded Particles: This positive sphere has negatively charged electrons embedded in it, like seeds in a watermelon or raisins in a pudding.

According to this model, the overall charge in the atom is neutral because the positive sphere and negative electrons balance each other.

Plum Pudding Model or Watermelon Model:

Plum Pudding Model: Thomson likewise called his model a plum pudding model. In this, the positive charge was viewed as the principal part of the pudding and electrons were envisioned as little raisins which are dispersed in the positive pudding.

Watermelon Model: Another model given is that of watermelon. The red piece of the watermelon addresses the positive charge and the seeds resemble electrons that are dissipated in the watermelon.

Limitation:

Thomson’s model was the initial step to make sense of the construction of the molecule, however, it had a few weaknesses:

• Stability not explained: Thomson’s model couldn’t make sense of how positive and negative charges are steady in a similar iota.
• Missing the concept of Nucleus: There was no understanding of the nucleus of the particle in this model, while Rutherford later found the nucleus.
• Position and motion of electrons not explained: Thomson’s model didn’t make sense of how electrons move regardless of whether they have a particular circle.

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Positive Sphere: 

He thought that the main structure of the atom is a large sphere of positive charge that keeps the whole atom stable.

Electrons as Embedded Particles: 

This positive sphere has negatively charged electrons embedded in it, like seeds in a watermelon or raisins in a pudding.

According to this model, the overall charge in the atom is neutral because the positive sphere and negative electrons balance each other.

3. Plum Pudding Model or Watermelon Model (Example of Tarbooz and Plum Pudding)**

Plum Pudding: 

Thomson also called his model as plum pudding model. In this, positive charge was considered to be the main part of the pudding and electrons were imagined as small raisins which are scattered in the positive pudding.

Watermelon Model: 

Another example given is that of watermelon. The red part of the watermelon represents the positive charge and the seeds are like electrons which are scattered in the watermelon.

4. Limitations of Thomson’s Model:

Thomson’s model was the first step to explain the structure of the atom, but it had some shortcomings:

Stability was not explained: 

Thomson’s model could not explain how positive and negative charges are stable in the same atom.

Concept of Nucleus Missing:

 There was no concept of the nucleus of the atom in this model, while Rutherford later discovered the nucleus.

The position and motion of electrons were not explained: 

Thomson’s model did not explain how electrons move or whether they have any specific orbit or not.

5. Conclusion

Thomson’s model made some important points about the atom, which included the existence of electrons and the way of thinking about their placement. But this model could not fully explain the structure of the atom and later scientists like Rutherford and Bohr proposed new models which were more accurate.

4. Rutherford’s Model of Atom:

Rutherford’s Model of the atom gave a new and exact method for grasping the design of the particle. This model is additionally called the “Nuclear Model of the Atom** because in this the idea of the core was perceived interestingly.

A brief history of this theory:

Rutherford proposed his model in 1911, which depended on his renowned Gold Foil Test. In this trial, he made light emission particles strike a sheet of gold foil and noticed their example.

-Alpha Particles: 

These are decidedly accused particles that move with extremely high energy.

Observation of Gold Foil Experiment:

• The greatest alpha particles went straight through the gold foil with no obstacle.

• A few alpha particles redirected marginally, meaning their way different.

• Not many particles skipped straight back.

These perceptions were surprising and couldn’t be made sense of by Thomson’s model. So Rutherford proposed his new model.

2. Rutherford’s Model of the Atom (Rutherford ka Model)

The model given Rutherford’s perceptions incorporates a few central issues:

– Concept of Nucleus: 

Rutherford said that all the positive charge and mass of the molecule is gathered in a little focal piece of the particle, which is known as the **nucleus**. This core is tiny and thick.

– Position of Electrons: 

Electrons move outside the core, a ways off. These electrons spin in “roundabout circles” around the core, very much like planets rotate around the sun.

– Mostly Empty Space of Atom: 

As indicated by Rutherford’s investigation there is for the most part void space inside the molecule. Because of this, a large portion of the alpha particles went through the foil with practically no deterrent.

3. Structure of Rutherford’s Model (Structure of Rutherford’s Model)

– Nucleus (Center): 

In the focal point of the molecule, there is the core which contains a positive charge and the most extreme mass of the particle. This core is tiny and thick.

– Electrons in Orbits (Movement of Electrons): 

Electrons move around the core in fixed circles and they convey a negative charge.

– Mostly Empty Space: 

A large portion of the molecule is vacant, because of this unfilled space the vast majority of the wire particles went through the gold foil.

4. Limitations of Rutherford’s Model (Shortcomings of Rutherford’s Model)

Rutherford’s model made sense of a great deal yet it likewise had a few inadequacies:

– Electron Strength Issue (Falling of Electrons):

Electrons that spin around the core will lose charge and energy. As per this hypothesis, they ought to fall inside the core, however this doesn’t occur. Rutherford’s model couldn’t make sense of this thing.

– Range Clarification not given:

Rutherford’s model couldn’t make sense of the range of an iota, which is an extraordinary component of a component.

These deficiencies were improved by Niels Bohr in his model.

5. Conclusion:

Rutherford’s model was a “Nuclear Model” which makes sense of the places of the core and electrons and their development. This model was a vital stage in grasping our nuclear construction. It made sense that an iota has a thick core around which electrons move and the particle is generally occupied with void space.

5. Bohr’s Model of Atom:

Bohr’s Model of the Particle is a model that makes sense of the design of the molecule and makes sense of the idea of the movement of electrons and energy levels. This model was proposed by Niels Bohr in 1913, which tends to certain limits of Rutherford’s model and gives significant experiences into nuclear solidness and electron course of action.

1. Foundation – Birth of Bohr Model

Rutherford’s model made sense of the course of action of the core and electrons in the iota, yet it couldn’t make sense of electron solidness. As per Rutherford’s model, electrons move consistently in the bird brace of the core and ought to fall into the core while losing energy. To take care of this issue, Bohr gave his idea of “Quantized Energy Levels”.

2. Bohr’s Model of the Particle (Bohr ka Model)

A few primary concerns of Bohr’s model are:

-Fixed Circles:

Bohr said that electrons rotate around the core in fixed circles. These are additionally called shells or energy levels. These circles are “quantized”, that is, electrons can exist and move just in unambiguous circles.

-Quantized Energy Levels:

Each circle or energy level has a decent energy. Electrons don’t lose energy in their circle and consequently stay stable. This means that as long as the electron stays at its energy level, it won’t lose energy and won’t fall into the core.

-Energy Assimilation and Emanation:

On the off chance that the electron bounces from its lower energy circle to a higher energy circle, it ingests energy. What’s more, on the off chance that it comes from a higher circle to a lower circle, it emanates energy. This energy is consistently in unambiguous sums and is delivered as photons.

3. Construction of Bohr’s Model (Design of Bohr’s Model)

-Core (focus):

It is the focal point of the iota, which contains protons and neutrons. It contains the positive charge and significant mass of the particle.

-Electrons in Fixed Circles:

Electrons rotate around the core in fixed circles which have their own particular energy. Each circle is additionally called an energy level or shell and is named K, L, M, N…

-Naming of Energy Levels:

The principal circle is the nearest to the core and has the least energy (K shell), and the ensuing circles have dynamically more energy.

4. Significant Marks of Bohr’s Model:

Stable Circles:

Electrons stay stable in their proper circles and don’t lose energy.

Energy Ingestion or Discharge:

At the point when the electron changes circle, it retains or emanates energy, and this energy is in a proper sum.

Energy Level Portrayal:

Each circle is addressed as an energy level or shell, in which electrons can move without changing their situation.

5. Restrictions of Bohr’s Model:

Bohr’s model was extremely useful in grasping nuclear design, yet there were a few limits:

Not material for multi-electron molecules:

This model turns out just for single electron iotas like hydrogen. In multi-electron particles, electron associations likewise happen which are excluded from this model.

Not a glaringly obvious Reason for Sub-energy Levels:

Bohr’s model didn’t make sense of sub-energy levels (subshells) and electron’s perplexing way of behaving.

Thought about Fragmented After the Advancement of Quantum Mechanics:

After the disclosure of quantum mechanics, this model was thought of as deficient and misrepresented, because this model didn’t address the electron’s wave-molecule duality and vulnerability standard.

6. Conclusion

Bohr’s model was a significant stage in understanding nuclear design, which makes sense of the idea of a molecule’s solidness and electron energy levels. This model gave the idea of electron’s decent circles and quantized energy levels which later turned into areas of strength for the improvement of quantum mechanics.

6. Atomic Number and Mass Number:

Atomic Number (Z):

The number of protons in an atom is called atomic number.

Atomic number (Z) = Number of protons

Mass Number (A):

The total number of protons and neutrons in an atom is called mass number.

Mass number (A) = Number of protons + Number of neutrons

7. Isotopes and Isobars:

Isotopes:

Such atoms which have same atomic number but different mass number are called isotopes.

Example: Three isotopes of Hydrogen – protium, deuterium, and tritium.

Isobars:

Such atoms which have same mass number but different atomic number are called isobars.

Example:  Carbon-14 and Nitrogen-14.

8. Valency:

Valency is an atom’s ability to combine with other atoms, based on the number of electrons in its outermost shell.

Note: Elements with a full outer shell have a valency of zero, as they are stable (e.g., noble gases).

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