STRUCTURE OF THE ATOM-Notes

We have already learned that atoms and molecules are the basic building blocks of matter. The presence of different kinds of matter is because different atoms make them up. Now, we may wonder: what makes the atom of one element different from that of another, and are atoms truly indivisible as Dalton proposed, or do they contain smaller parts? In this chapter, we will explore these questions. We will study subatomic particles and the models that explain how they are arranged within the atom. By the end of the 19th century, scientists faced the challenge of uncovering the structure of the atom and explaining its properties. Through various experiments, especially those involving electricity and conduction, it became clear that atoms are not indivisible but contain internal components.

To understand how matter behaves when charged, simple experiments can help reveal the presence of electrical forces.

When we perform simple activities like combing dry hair or rubbing a glass rod with silk, we notice that objects can attract small pieces of paper or balloons.
This happens because rubbing causes objects to become electrically charged.
These observations show that atoms are not indivisible; they contain smaller charged particles.

Around the late 19th and early 20th centuries, scientists began uncovering the hidden structure of the atom through experiments and observations.

By 1900, scientists had discovered that atoms included at least one tiny negatively charged particle called the electron, discovered by J.J. Thomson.
Earlier, in 1886, E. Goldstein had found positively charged canal rays, leading to the discovery of another particle, the proton.
The proton carries a positive charge equal in size but opposite to that of the electron, and its mass is about 2000 times greater.

Once the main subatomic particles had been identified, scientists turned to the question of how these particles were arranged inside the atom.

Electrons are represented by \(e^-\) and protons by \(p^+\). The proton’s mass is taken as one unit with a charge of +1, while the electron’s mass is negligible and its charge is –1.
Atoms were thus thought to consist of protons and electrons whose charges balance each other. Since electrons can be easily removed, they are located on the outer part of the atom, while protons remain inside.
The major question that followed was how these particles are arranged within the atom, a topic explored in the next section.

In Chapter 3, we learned about Dalton’s atomic theory, which explained that atoms are indivisible and indestructible. However, when we found out about the discovery of electrons and protons inside the atom, we realized that Dalton’s theory wasn’t completely correct. This made us curious about how these particles are arranged within an atom. To explain this, many scientists came up with different atomic models. J.J. Thomson was the first to propose a model for the structure of the atom.

Thomson’s model of an atom, often called the “plum pudding model,” was one of the earliest ideas about how atoms are structured. According to this model, the atom is like a sphere of positive charge with tiny, negatively charged electrons embedded in it—similar to the way plums are spread through a pudding.

We can imagine the atom as a soft, positively charged ball, where electrons are scattered throughout, balancing the overall charge. Thomson believed that the positive and negative charges were mixed together in an even way, which made the atom electrically neutral as a whole. This model helped scientists understand that electrons exist inside atoms and that atoms can be broken down further.

However, as new experiments were done, this model was found to have some limitations. Even so, Thomson’s “plum pudding” idea was an important step in our journey to unravel the mysterious structure of atoms, opening the door for future discoveries and more accurate models.

Thomson proposed that:

• An atom consists of a positively charged sphere, and the electrons are embedded in it.
• The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral.

Rutherford’s model of an atom marked a major turning point in our understanding of atomic structure. Through a famous experiment—the gold foil experiment—Ernest Rutherford and his team bombarded a thin sheet of gold with alpha particles. They were surprised to observe that while most particles passed straight through, a few were deflected at large angles, and some even bounced straight back.

Rutherford’s Expectations:

• Alpha particles would pass straight through the gold foil without any deflection.
• The positive charge in the atom was thought to be spread out evenly, so no significant obstacles would stop or deflect the alpha particles.
• Only minor deviations from the straight path were expected due to the “plum pudding” model.

Rutherford’s Actual Outcomes:

• Most alpha particles did indeed pass straight through the gold foil.
• Some alpha particles were deflected at small angles.
• A few alpha particles bounced back or were deflected at very large angles.

Rutherford’s Conclusion:

• Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected.
• Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.
• A very small fraction of α-particles were deflected by \(180^{\circ}\), indicating that all the positive charge and mass of the gold atoms were concentrated in a very small volume within the atom.

Rutherford’s nuclear model of an atom:

• There is a positively charged centre in an atom called a nucleus. Nearly all the mass of an atom resides in the nucleus.
• The electrons revolve around the nucleus in circular paths.
• The size of the nucleus is very small as compared to the size of the atom.

Drawbacks of Rutherford’s model of the atom

• It couldn’t explain atomic stability:
  Classical physics said electrons moving in circles around the nucleus should lose energy and collapse into the nucleus, but atoms are actually stable.
• It did not account for electron energy levels:
  Rutherford’s model didn’t explain why atoms emit specific colours of light or why electrons can only occupy certain energy states.
• It failed to describe chemical properties:
  The model didn’t show how the arrangement of electrons affects how atoms bond with each other, which is essential for chemistry.
• No explanation for electron movement:
  Electrons were said to orbit the nucleus, but the model didn’t explain how they stay at fixed distances or avoid losing energy.
• It ignored quantum theory:
  Rutherford’s model only used classical physics, missing quantum mechanics, which is crucial for understanding atoms.

Bohr’s model of the atom brought a fresh perspective to atomic science and answered many questions left unresolved by previous models.
Niels Bohr imagined the atom as a miniature solar system, but with a key difference that electrons did not just move randomly around the nucleus.
Instead, he proposed that electrons revolve around the nucleus in specific circular paths known as energy levels or shells.

Neils Bohr postulation

• Only certain special orbits known as discrete orbits of electrons are allowed inside the atom.
• While revolving in discrete orbits, the electrons do not radiate energy.

Energy levels are the fixed orbits around the nucleus where electrons reside in an atom, as described by Bohr’s model. Each energy level corresponds to a specific amount of energy an electron holds. Electrons can move between these levels by absorbing or releasing energy, often as light.

The energy level closest to the nucleus is called the ground state, while levels further away are known as excited states. These levels are often labelled as K, L, M, N, and so on, starting from the nucleus outward.

Understanding energy levels helps explain why atoms are stable and why different elements give off unique colours of light. In summary, energy levels are key to how electrons behave and how atoms display their distinct properties.

Neutrons are particles found in the nucleus of an atom, along with protons. Unlike protons, which have a positive charge, neutrons have no charge; they are neutral. Discovered by James Chadwick in 1932, neutrons add to the mass of an atom but do not affect its charge. The number of neutrons in an atom can vary, creating different isotopes of the same element. Neutrons help keep the nucleus stable and are a key part of understanding atomic structure.

Electrons in an atom are not scattered randomly; rather, they occupy specific orbits or shells around the nucleus, much like seats arranged around a table. These shells are given names like K, L, M, and N, starting from the one closest to the nucleus and moving outward. Each shell has a certain maximum capacity for holding electrons, and electrons fill the lower-energy shells first before moving to higher ones.

The rules for filling these shells were explained by Bohr and Bury. According to their scheme, the maximum number of electrons a shell can hold is given by the formula \[2n^2\] where \(n\) is the number of the shell (starting from 1 for K, 2 for L, and so on). For example, the K shell can have up to 2 electrons, the L shell up to 8, the M shell up to 18, and so forth.

Electrons fill these shells in order—the innermost shell first, then the next one, and so on—until all the electrons in the atom have been placed. This arrangement helps keep atoms stable and determines how they interact with other atoms in chemical reactions.
By understanding how electrons are distributed in shells, we can predict the properties and behaviour of an element much more easily.

Valency is the number of electrons an atom needs to gain, lose, or share to fill its outermost shell and become stable. The electrons in the outermost shell, called valence electrons, decide how an atom will bond and react with others.

For example, hydrogen has one electron in its outermost shell and needs one more to be stable, so its valency is 1.

Oxygen has six electrons in its outermost shell and needs two more for stability, giving it a valency of 2.

Understanding valency and outermost electrons helps us know how atoms join together to form molecules.

Name of Element	Symbol	Atomic Number	Number of Protons	Number of Neutrons	Number of Electrons	Valency	Electron Distribution (K L M N)
Hydrogen	H	1	1	0	1	1	1 – – –
Helium	He	2	2	2	2	0	2 – – –
Lithium	Li	3	3	4	3	1	2 1 – –
Beryllium	Be	4	4	5	4	2	2 2 – –
Boron	B	5	5	6	5	3	2 3 – –
Carbon	C	6	6	6	6	4	2 4 – –
Nitrogen	N	7	7	7	7	3	2 5 – –
Oxygen	O	8	8	8	8	2	2 6 – –
Fluorine	F	9	9	10	9	1	2 7 – –
Neon	Ne	10	10	10	10	0	2 8 – –
Sodium	Na	11	11	12	11	1	2 8 1 –
Magnesium	Mg	12	12	12	12	2	2 8 2 –
Aluminium	Al	13	13	14	13	3	2 8 3 –
Silicon	Si	14	14	14	14	4	2 8 4 –
Phosphorus	P	15	15	16	15	3,5	2 8 5 –
Sulphur	S	16	16	16	16	2	2 8 6 –
Chlorine	Cl	17	17	18	17	1	2 8 7 –
Argon	Ar	18	18	22	18	0	2 8 8 0

Atomic number is one of the most important ways to identify and distinguish elements. Every atom has a small, dense nucleus made up of positively charged protons. The atomic number is simply the count of these protons present in the nucleus. This number is always unique for each element and is given the symbol ‘Z’ in chemistry.

All atoms of the same element have exactly the same atomic number, and that’s what defines the element itself.
For instance, if an atom has only one proton, we call it hydrogen and assign it atomic number Z = 1. If an atom contains six protons, it is known as carbon, and its atomic number is Z = 6. In other words, the atomic number is the total number of protons in the nucleus, and it determines an atom’s identity as well as its place on the periodic table.

Mass number is a straightforward but essential concept for understanding what makes one atom different from another, besides its atomic number. It is the total count of two types of particles inside the atom’s nucleus: protons and neutrons. While protons define the atomic number, neutrons add to the atom’s mass without changing its identity.

Every atom’s mass number is calculated by adding the number of protons and neutrons together. This sum is always a whole number and is symbolized by ‘A’. For example, if an atom of carbon has 6 protons and 6 neutrons, its mass number will be \[A=12\] If a nitrogen atom has 7 protons and 7 neutrons, then \[A=14\] In the notation for an atom, the atomic number, mass number and symbol of the element are to be written as: \[^A_Z\mathrm{X}\] For example, nitrogen is written as \[^{14}_{7}\mathrm{N}\]

Although atoms of the same element share the same atomic number, their mass numbers may differ due to varying numbers of neutrons. This is why some elements exist in different forms, called isotopes. In summary, mass number tells us the total number of protons and neutrons in the atom’s nucleus, which helps us understand both the variety and stability of elements.

Isotopes are a fascinating feature of elements that help explain the hidden diversity within atoms. All atoms of a particular element share the same number of protons in their nucleus, which determines their atomic number and chemical identity. However, atoms of the same element can have different numbers of neutrons, and this creates isotopes.

Isotopes are simply atoms of the same element that have the same atomic number but different mass numbers, because their neutron counts differ. For example, all hydrogen atoms have one proton, but some hydrogen atoms have no neutrons (protium), some have one neutron (deuterium), and some have two neutrons (tritium). These versions of hydrogen are its isotopes.

Even though isotopes of an element behave the same way in chemical reactions, their physical properties like mass and stability can vary. This special feature allows isotopes to be used in medicine, research, and industry in unique ways, such as in medical imaging or as tracers. In summary, isotopes show us that atoms with the same identity can have subtle differences, making the study of elements much more interesting and useful.

Applications

• An isotope of uranium is used as a fuel in nuclear reactors.
• An isotope of cobalt is used in the treatment of cancer.
• An isotope of iodine is used in the treatment of goitre.

Isobars are another interesting concept in chemistry that helps us understand the variety of atoms found in nature. Unlike isotopes, isobars are atoms of different elements that share the same mass number but have different atomic numbers. This means they have the same total number of protons and neutrons in their nuclei, but different numbers of protons and neutrons individually.

For example, consider sodium \(^{11}_{23}\mathrm{Na}\) and magnesium \(^{12}_{23}\mathrm{Mg}\): both atoms have a mass number of 23, but sodium has 11 protons and magnesium has 12 protons. Because their neutron count adjusts accordingly, these elements end up with the same mass number.

Isobars do not have the same chemical properties since their atomic numbers, and therefore their identities as elements, are different. However, their similar mass gives them some related physical properties. Understanding isobars helps us see that the mass number alone does not define an element—it’s the combination of protons and neutrons that matters. It’s a concept that brings more depth and richness to our study of atomic structure.

• Credit for the discovery of electron and proton goes to J.J. Thomson and E.Goldstein, respectively.
• J.J. Thomson proposed that electrons are embedded in a positive sphere.
• Rutherford’s alpha-particle scattering experiment led to the discovery of the atomic nucleus.
• Rutherford’s model of the atom proposed that a very tiny nucleus is present inside the atom and electrons revolve around this nucleus. The stability of the atom could not be explained by this model.
• Neils Bohr’s model of the atom was more successful. He proposed that electrons are distributed in different shells with discrete energy around the nucleus. If the atomic shells are complete, then the atom will be stable and less reactive.
• J. Chadwick discovered presence of neutrons in the nucleus of an atom. So, the three sub-atomic particles of an atom are: (i) electrons, (ii) protons and (iii) neutrons. Electrons are negatively charged, protons are positively charged and neutrons have no charges. The mass of an electron is about \(frac{1}{2000}\) times the mass of an hydrogen atom. The mass of a proton and a neutron is taken as one unit each.
• Shells of an atom are designated as K,L,M,N,….
• Valency is the combining capacity of an atom.
• The atomic number of an element is the same as the number of protons in the nucleus of its atom.
• The mass number of an atom is equal to the number of nucleons in its nucleus.
• Isotopes are atoms of the same element, which have different mass numbers.
• Isobars are atoms having the same mass number but different atomic numbers.
• Elements are defined by the number of protons they possess.