Your Journey To The Basics Of Quantum Realm Volume II: Your Journey to The Basics Of Quantum Realm, #2

[BASICS OF QUANTUM REALM]

VOL-2

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WRITTEN BY

PRAJJWAL JHA

COMPLIED BY

ANIL THAPA

Preface

The world of particles is fascinating and equally complicated. This book deals with the simplest way to learn about the basics of mentioned world. The very first section in the book deals with the complete takedown on atom, its concepts, models and their basic characteristics leading us to harness the knowledge about it for the further uses. It also describes the basics of the atomic structures taking the dig at minuest level and about the radiations that could come out from them.

This piece of work is actually the bridge between the known classical concepts of quantum world to the more modern approach towards the particles. This connects the vedic inferences of the quantum world to the more modern concepts that will be discussed in the volume 3.

Readers are advised to have simple knowledge on the matter constituents, henceforth, it will be useful for all levels readers who are interested in knowing the constructional framework of the substances.

Authors.

Acknowledgement

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Your Journey to the Basics of Quantum Realm Volume-2 is the continuation of volume 1 which we are certain that reader is already acquainted with. Similar to the previous volume, this is written with deep researches focusing on the detailed understanding about the subject matter to the reader. In this case, various minds have worked and alot of inputs from various living and non living sources were considered.

Initiating with my family, especially my parents who have always been supportive for the completion of the book. My siblings, especially my brother Er. Ujjwal Jha has been the role model upon whom I have looked onto and my sister Ms. Sweta Jha, who has significant contribution in this piece of work. Without their support this would never have been possible in reality.

I cannot neglect the fact that, I have been strongly influenced by the teaching methodologies of my physics mentor and Assistant Professor Mr. Lok Nath Sharma, Tribhuwan University, Nepal , who has always been my doubt solver and have helped me with various resources. This piece of information strongly depicts the knowledge gaining pattern that I have received from him in my high school days. I am indebted to Biochemistry Assistant Professor Dr. Avnish Tarwey, BRSSABVM Medical College, Rajnandgaon, for his wonderful suggestions regarding the content of the book. I cannot unsee the assistance provided by the Anatomy Department of my College, for their suggestions.

Various other minds have posthumously worked for the completion of this book from both Nepal and India. Starting with Mr. Anil Thapa, Mr. Anup Dhital, Mr. Gaurav Yadav, Mr. Bishesh Kattel, Mr. Dipesh Thapa, Mr Asmit pandey , Mr Nayan Dhital ,Ms. Krisha Sapkota from Nepal who have helped directly or indirectly for the completion of book.

Similarly from India, the strong effort with which I was blessed, came from Mr. Devprayag

Sharma, Mr. Ambuj Tiwari, Mr. Ayush Kumar, Mr. Pratham Gajendra Tripathi, Ms. Sneha and all my batchmates for their assistance.

- Prajjwal Jha

Content

Atom

Introduction

CONDUCTORS AND INSULATOR

ATOMIC THEORY AND MODELAS

MODELS OF ATOMIC STRUCTURE

THE BEGININGS OF MODERN ATOMIC THEORY

STUDIES OF THE PROPERTIES OF ATOMIC SIZE OF ATOM

Discovery of electrons

Structure of the Nucleus

Discovery of Radioactivity 10.Schrödinger Wave Equation

Uncertainty in Quantum World

Beauty of Uncertainty

The Semiclassical Uncertainty Principle

Exclusion in Quantum

1. Pauli’s Exclusion Principle

2. The Philosophical Meaning of the Pauli’s Exclusion Principle

Photoelectric Effect

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INTRODUCTION

Details

Classical explanations of Photoelectric Effect

X-  Rays

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Physical Interpretation

Production, Detection and Uses

Radiation Detection and Measurements

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INTRODUCTION

APPLICATIONS OF RADIATION INTERACTIONS IN DETECTORS

QUARK

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INTRODUCTION

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JOURNEY TO THE FERMI

INTRODUCTION

Reality doesn’t exist when you are not looking at it this means that the universe may not exist if there was no observe it.

( ...For Dad )

ATOM

Introduction

Atom, smallest unit into which matter can be divided without the release of electrically charged particles. It also is the smallest unit of matter that has the characteristic properties of a chemical element. As such, the atom is the basic building block of chemistry.

Most of the atom is empty space. The rest consists of a positively charged nucleus of protons and neutrons surrounded by a cloud of negatively charged electrons. The nucleus is small and dense compared with the electrons, which are the lightest charged particles in nature. Electrons are attracted to any positive charge by their electric force; in an atom, electric forces bind the electrons to the nucleus.

Because of the nature of quantum mechanics, no single image has been entirely satisfactory at visualizing the atom’s various characteristics, which thus forces physicists to use complementary pictures of the atom to explain different properties. In some respects, the electrons in an atom behave like particles orbiting the nucleus. In others, the electrons behave like waves frozen in position around the nucleus. Such wave patterns, called orbitals, describe the distribution of individual electrons. The behaviour of an atom is strongly influenced by these orbital properties, and its chemical properties are determined by orbital groupings known as shells.

This chapter opens with a broad overview of the fundamental properties of the atom and its constituent particles and forces. Following this overview is a historical survey of the most influential concepts about the atom that have been formulated through the centuries.

Atomic Model

Most matter consists of an agglomeration of molecules, which can be separated relatively easily. Molecules, in turn, are composed of atoms joined by chemical bonds that are more difficult to break. Each individual atom consists of smaller particles—namely, electrons and nuclei. These particles are electrically charged, and the electric forces on the charge are responsible for holding the atom together. Attempts to separate these smaller constituent particles require

ever-increasing amounts of energy and result in the creation of new subatomic particles, many of which are charged.

As noted in the introduction to this chapter, an atom consists largely of empty space. The nucleus is the positively charged centre of an atom and contains most of its mass. It is composed of protons, which have a positive charge, and neutrons, which have no charge. Protons, neutrons, and the electrons surrounding them are long-lived particles present in all ordinary, naturally occurring atoms. Other subatomic particles may be found in association with these three types of particles. They can be created only with the addition of enormous amounts of energy, however, and are very short-lived.

All atoms are roughly the same size, whether they have 3 or 90 electrons. Approximately 50 million atoms of solid matter lined up in a row would measure 1 cm (0.4 inch). A convenient unit of length for measuring atomic sizes is the angstrom (Å), defined as 10−10 metre. The radius of an atom measures 1–2 Å. Compared with the overall size of the atom, the nucleus is even more minute. It is in the same proportion to the atom as a marble is to a football field. In volume the

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nucleus takes up only 10−14 metres of the space in the atom—i.e., 1 part in 100,000. A convenient unit of length for measuring nuclear sizes is the femtometre (fm), which equals 10−15 metre. The diameter of a nucleus depends on the number of particles it contains and ranges from about 4 fm for a light nucleus such as carbon to 15 fm for a heavy nucleus such as lead. In spite of the small size of the nucleus, virtually all the mass of the atom is concentrated there. The protons are massive, positively charged particles, whereas the neutrons have no charge and are slightly more massive than the protons. The fact that nuclei can have anywhere from 1 to nearly 300 protons and neutrons accounts for their wide variation in mass. The lightest nucleus, that of hydrogen, is 1,836 times more massive than an electron, while heavy nuclei are nearly 500,000 times more massive.

Basic properties

Atomic number

The single most important characteristic of an atom is its atomic number (usually denoted by the letter Z), which is defined as the number of units of positive charge (protons) in the nucleus. For example, if an atom has a Z of 6, it is carbon, while a Z of 92 corresponds to uranium. A neutral atom has an equal number of protons and electrons so that the positive and negative charges exactly balance. Since it is the electrons that determine how one atom interacts with another, in the end it is the number of protons in the nucleus that determines the chemical properties of an atom.

Atomic mass and isotopes

The number of neutrons in a nucleus affects the mass of the atom but not its chemical properties. Thus, a nucleus with six protons and six neutrons will have the same chemical properties as a nucleus with six protons and eight neutrons, although the two masses will be different. Nuclei with the same number of protons but different numbers of neutrons are said to be isotopes of each other. All chemical elements have many isotopes.

It is usual to characterize different isotopes by giving the sum of the number of protons and neutrons in the nucleus—a quantity called the atomic mass number. In the above example, the first atom would be called carbon-12 or 12C (because it has six protons and six neutrons), while the second would be carbon-14 or 14C.

The mass of atoms is measured in terms of the atomic mass unit, which is defined to be 1/12 of the mass of an atom of carbon-12, or 1.660538921 × 10−24 gram. The mass of an atom consists of the mass of the nucleus plus that of the electrons, so the atomic mass unit is not exactly the same as the mass of the proton or neutron.

The electron

Charge, mass, and spin

Scientists have known since the late 19th century that the electron has a negative electric charge. The value of this charge was first measured by the American physicist Robert Millikan between 1909 and 1910. In Millikan’s oil-drop experiment, he suspended tiny oil drops in a

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chamber containing an oil mist. By measuring the rate of fall of the oil drops, he was able to determine their weight. Oil drops that had an electric charge (acquired, for example, by friction when moving through the air) could then be slowed down or stopped by applying an electric force. By comparing applied electric force with changes in motion, Millikan was able to determine the electric charge on each drop. After he had measured many drops, he found that the charges on all of them were simple multiples of a single number. This basic unit of charge was the charge on the electron, and the different charges on the oil drops corresponded to those having 2, 3, 4,... extra electrons on them. The charge on the electron is now accepted to be 1.602176565 × 10−19 coulomb. For this work Millikan was awarded the Nobel Prize for Physics in 1923.

The charge on the proton is equal in magnitude to that on the electron but opposite in sign—that is, the proton has a positive charge. Because opposite electric charges attract each other, there is an attractive force between electrons and protons. This force is what keeps electrons in orbit around the nucleus, something like the way that gravity keeps Earth in orbit around the Sun.

The electron has a mass of about 9.109382911 × 10−28 gram. The mass of a proton or neutronis about 1,836 times larger. This explains why the mass of an atom is primarily determined by the mass of the protons and neutrons in the nucleus.

The electron has other intrinsic properties. One of these is called spin. The electron can be pictured as being something like Earth, spinning around an axis of rotation. In fact, most elementary particles have this property. Unlike Earth, however, they exist in the subatomic world and are governed by the laws of quantum mechanics. Therefore, these particles cannot spin in any arbitrary way, but only at certain specific rates. These rates can be 1/2, 1, 3/2, 2,... times a basic unit of rotation. Like protons and neutrons, electrons have spin 1/2.

Particles with half-integer spin are called fermions, for the Italian American physicist Enrico Fermi, who investigated their properties in the first half of the 20th century. Fermions have one important property that will help explain both the way that electrons are arranged in their orbits and the way that protons and neutrons are arranged inside the nucleus. They are subject to the Pauli exclusion principle (named for the Austrian physicist Wolfgang Pauli), which states that no two fermions can occupy the same state—for example, the two electrons in a helium atom must have different spin directions if they occupy the same orbit.

Because a spinning electron can be thought of as a moving electric charge, electrons can be thought of as tiny electromagnets. This means that, like any other magnet, an electron will respond to the presence of a magnetic field by twisting. (Think of a compass needle pointing north under the influence of Earth’s magnetic field.) This fact is usually expressed by saying that electrons have a magnetic moment. In physics, magnetic moment relates the strength of a magnetic field to the torque experienced by a magnetic object. Because of their intrinsic spin, electrons have a magnetic moment given by −9.28 × 10−24 joule per tesla.

Orbits and energy levels

Unlike planets orbiting the Sun, electrons cannot be at any arbitrary distance from the nucleus; they can exist only in certain specific locations called allowed orbits. This property, first explained by Danish physicist Niels Bohr in 1913, is another result of quantum

mechanics—specifically, the requirement that the angular momentum of an electron everything else in the quantum world, come in discrete bundles called quanta.

In the Bohr atom electrons can be found only in allowed orbits, and these allowed orbits are at different energies. The orbits are analogous to a set of stairs in which the gravitational potential energy is different for each step and in which a ball can be found on any step but never in between.

The laws of quantum mechanics describe the process by which electrons can move from one allowed orbit, or energy level, to another. As with many processes in the quantum world, this process is impossible to visualize. An electron disappears from the orbit in which it is located and reappears in its new location without ever appearing any place in between. This process is called a quantum leap or quantum jump, and it has no analog in the macroscopic world.

Because different orbits have different energies, whenever a quantum leap occurs, the energy possessed by the electron will be different after the jump. For example, if an electron jumps from a higher to a lower energy level, the lost energy will have to go somewhere and in fact will be emitted by the atom in a bundle of electromagnetic radiation. This bundle is known as a photon, and this emission of photons with a change of energy levels is the process by which atoms emit light (for instance LASER).

In the same way, if energy is added to an atom, an electron can use that energy to make a quantum leap from a lower to a higher orbit. This energy can be supplied in many ways. One common way is for the atom to absorb a photon of just the right frequency. For example, when white light is shone on an atom, it selectively absorbs those frequencies corresponding to the energy differences between allowed orbits.

Each element has a unique set of energy levels, and so the frequencies at which it absorbs and emits light act as a kind of fingerprint, identifying the particular element. This property of atoms has given rise to spectroscopy, a science devoted to identifying atoms and molecules by the kind of radiation they emit or absorb.

This picture of the atom, with electrons moving up and down between allowed orbits, accompanied by the absorption or emission of energy, contains the essential features of the Bohr atomic model, for which Bohr received the Nobel Prize for Physics in 1922. His basic model does not work well in explaining the details of the structure of atoms more complicated than hydrogen, however. This requires the introduction of quantum mechanics. In quantum mechanics each orbiting electron is represented by a mathematical expression known as a wave function—something like a vibrating guitar string laid out along the path of the electron’s orbit. These waveforms are called orbitals.

In the quantum mechanical version of the Bohr atomic model, each of the allowed electron orbits is assigned a quantum number n that runs from 1 (for the orbit closest to the nucleus)