STRUCTURE OF ATOM - PART 1

 DISCOVERY OF ELECTRONS

1. Discovery of Electrolytic Effects (1830s):
• Michael Faraday demonstrated that passing electricity through an electrolyte solution led to chemical reactions at the electrodes.
• These reactions resulted in the liberation and deposition of matter at the electrodes.
• Faraday formulated laws related to electrolytic effects, which are studied in Class XII.
• These observations hinted at the particulate nature of electricity, suggesting that it consists of discrete particles.
2. Cathode Ray Discharge Tube Experiments (Mid-1850s):
• Scientists, particularly Faraday, focused on studying electrical discharge in cathode ray discharge tubes (partially evacuated tubes).
• A cathode ray tube is a glass tube containing two thin metal electrodes, cathode (negative electrode) and anode (positive electrode), sealed inside.
• The experiments were conducted at very low pressures and high voltages, achieved by evacuating the gas from the glass tube.
• Applying a high voltage led to the flow of current through a stream of particles moving from the cathode to the anode.
• These particles were termed "cathode rays" or "cathode ray particles."
• To visualize these rays, a hole was made in the anode, and the back of the tube behind the anode was coated with phosphorescent material (zinc sulphide).
• When the cathode rays passed through the anode and struck the zinc sulphide coating, it produced a bright spot, indicating their path.
3. Characteristics of Cathode Rays (Cathode Rays are Electrons):
• Cathode rays were observed to originate from the cathode (negative electrode) and move toward the anode (positive electrode).
• These rays themselves were invisible, but their behavior was evident when they interacted with certain materials that fluoresced or phosphoresced upon contact.
• Television picture tubes are examples of cathode ray tubes, where images are produced through the fluorescence of specific materials on the screen.
• In the absence of electrical or magnetic fields, cathode rays traveled in straight lines, suggesting their uncharged nature.
• However, when subjected to electrical or magnetic fields, the behavior of cathode rays resembled that of negatively charged particles, indicating the presence of negatively charged constituents.
• It was concluded that cathode rays were composed of negatively charged particles known as "electrons."
• Notably, the properties of cathode rays (electrons) remained consistent regardless of the material used for the electrodes or the type of gas within the cathode ray tube.
• This led to the significant conclusion that electrons are fundamental components of all atoms.

In summary, the experiments with cathode ray discharge tubes provided crucial evidence for the existence of electrons as discrete, negatively charged particles. This discovery played a pivotal role in understanding the atomic and subatomic structure of matter and laid the foundation for modern physics.

 

1. Measurement of Charge-to-Mass Ratio (1897):
• British physicist J.J. Thomson conducted experiments in 1897 to measure the ratio of the electrical charge (e) to the mass of the electron (me).
• Thomson used a cathode ray tube setup and applied perpendicular electric and magnetic fields to the path of the electrons.
• When only an electric field was applied, electrons were deflected from their original path, hitting the cathode ray tube at point A.
• Similarly, with only a magnetic field, electrons struck the cathode ray tube at point C.
• By carefully adjusting the strengths of the electric and magnetic fields, Thomson was able to bring the electrons back to their original path, hitting the screen at point B.
• Thomson's goal was to find the balance of field strengths that would nullify the electron's deflection.
2. Factors Affecting Particle Deflection:
• Thomson deduced that the amount of deflection experienced by particles due to electric or magnetic fields depended on several factors:
• (i) Magnitude of the Negative Charge: Greater magnitude of charge on a particle resulted in stronger interaction with electric or magnetic fields, leading to more significant deflection.
• (ii) Mass of the Particle: Lighter particles experienced greater deflection when subjected to the same fields.
• (iii) Strength of the Electric or Magnetic Field: Increasing the voltage across the electrodes or the strength of the magnetic field resulted in increased deflection of electrons from their original path.
3. Calculation of Charge-to-Mass Ratio (e/m_e):
• Thomson carried out precise measurements of the deflections experienced by electrons in the presence of electric and magnetic fields.
• Based on these measurements, Thomson determined the value of the charge-to-mass ratio (e/m_e) to be:
• e/m_e = 1.758820 × 10¹¹ C/ kg
• In this equation, m_e represents the mass of the electron in kilograms, and e represents the magnitude of the charge on the electron in coulombs (C).
4. Nature of Electron Charge:
• Since electrons are negatively charged, the charge on an electron is denoted as -e.

In summary, J.J. Thomson's experiments using the cathode ray tube and the application of perpendicular electric and magnetic fields allowed him to accurately determine the charge-to-mass ratio of electrons. His work provided valuable insights into the fundamental properties of electrons and their behavior in electric and magnetic fields.

 

1. 
   Millikan's Oil Drop Experiment (1906-1914):
• R.A. Millikan conducted an experiment called the oil drop experiment between 1906 and 1914.
• The purpose of the experiment was to determine the fundamental charge carried by individual electrons.
2. Determination of Electron Charge:
• Millikan's oil drop experiment involved suspending tiny oil droplets in a chamber using an upward air flow.
• By carefully controlling the air flow and observing the motion of the droplets, Millikan was able to measure the terminal velocity of the droplets.
• By applying electrical fields to the chamber, Millikan could manipulate the motion of the droplets, counteracting gravity and adjusting their terminal velocities.
• By balancing gravitational force and electrical force, he calculated the charge carried by individual oil droplets, which could be related to the charge of an electron.
• Millikan found that the charge on the electron was approximately -1.6 × 10^-19 Coulombs (C).
3. Comparison with Accepted Charge Value:
• The currently accepted value of the elementary charge (the charge of an electron) is -1.602176 × 10^-19 C.
• Millikan's experiment provided a charge value that was remarkably close to the accepted value, further confirming the quantized nature of electric charge.
4. Mass of the Electron Determination:
• Millikan's determination of the charge on the electron, combined with J.J. Thomson's earlier measurement of the charge-to-mass ratio (e/m_e), allowed for the calculation of the mass of the electron.
• Using Thomson's value of e/m_e and Millikan's charge value, the mass of the electron was determined as approximately 9.1094 × 10^-31 kilograms (kg).

In summary, R.A. Millikan's oil drop experiment played a crucial role in accurately determining the charge of an electron. This experiment, along with J.J. Thomson's earlier work, provided essential data that allowed scientists to calculate the mass of the electron and contribute to the understanding of the fundamental properties of subatomic particles.

 

Discovery of Canal Rays and Positive Particles:

1. Electrical Discharge and Canal Rays:
• Modifying the cathode ray tube setup led to the discovery of canal rays, which were associated with positively charged particles.
• Unlike cathode rays (electrons), canal rays carried positively charged particles.
2. Characteristics of Positively Charged Particles:
• Mass Dependence on Gas Nature: The mass of positively charged particles depended on the type of gas present in the cathode ray tube. These particles were essentially positively charged gaseous ions.
• Charge-to-Mass Ratio Variation: The charge-to-mass ratio of these particles varied based on the specific gas from which they originated.
• Charge Multiples: Some of the positively charged particles carried a multiple of the fundamental unit of electrical charge.
• Behavior in Fields: The behavior of these particles in electric or magnetic fields was opposite to that observed for electrons or cathode rays.

Discovery of Protons and Neutrons: 3. Discovery of Protons (1919):

• The lightest and smallest positive ion was obtained from hydrogen gas.
• This positively charged particle was characterized in 1919 and named the "proton."
• Protons are essential components of the nucleus of atoms.

4. Discovery of Neutrons (1932):
• James Chadwick discovered electrically neutral particles known as neutrons in 1932.
• Chadwick performed experiments involving bombarding a thin sheet of beryllium with alpha particles.
• As a result of this bombardment, electrically neutral particles were emitted from the beryllium nuclei.
• These neutral particles had a mass slightly greater than that of protons.
• Chadwick named these neutral particles "neutrons."

In summary, the modified cathode ray tube experiments led to the discovery of canal rays carrying positively charged particles. These particles exhibited distinct characteristics, such as varying mass depending on the gas, charge-to-mass ratio variability, and some carrying multiples of the fundamental charge unit. The discovery of protons, the smallest positive ions, and their characteristics was significant in understanding atomic structure. Subsequently, the need for electrically neutral particles was addressed by James Chadwick, who discovered neutrons through experiments involving beryllium and alpha particles. These discoveries contributed to the development of the modern atomic model and deepened our understanding of the constituents of the atom.

 

Transition from Dalton's Model and Challenges:

1. Observations and Sub-Atomic Particles:
• Experiments outlined in previous sections suggested that Dalton's idea of an indivisible atom was no longer tenable.
• It became evident that atoms were composed of sub-atomic particles carrying both positive and negative charges.
2. Challenges for Scientists:
• After the discovery of sub-atomic particles, scientists faced significant challenges:
• Stability of Atoms: Explaining the stability of atoms, given the presence of charged particles.
• Comparison of Elements: Understanding the behavior of elements based on physical and chemical properties.
• Molecular Formation: Explaining how different kinds of molecules formed through the combination of various atoms.
• Electromagnetic Radiation: Understanding the origin and nature of electromagnetic radiation absorbed or emitted by atoms.

Proposed Atomic Models:

       3. J.J. Thomson's Model:

• J.J. Thomson proposed an atomic model based on the presence of negatively charged electrons embedded in a positively charged "pudding-like" matrix.
• This model was known as the "plum pudding" model.
• However, this model had difficulty explaining the stability of atoms and their behavior.

4. Ernest Rutherford's Model:
• Ernest Rutherford conducted the famous gold foil experiment.
• His model proposed that the atom has a small, dense, positively charged nucleus at its center.
• Electrons orbit the nucleus at a distance, much like planets orbiting the sun.
• The vast majority of the atom's mass is concentrated in the nucleus.
• This model explained the results of Rutherford's experiment and provided a better understanding of the atom's structure.

The experiments involving sub-atomic particles led to the realization that Dalton's concept of indivisible atoms was no longer accurate. This posed challenges related to atom stability, element behavior, molecular formation, and electromagnetic radiation. Different atomic models were proposed to address these challenges. J.J. Thomson's "plum pudding" model and Ernest Rutherford's nuclear model were two notable attempts. Rutherford's model, with a dense nucleus and orbiting electrons, proved to be more successful in explaining experimental observations and laid the foundation for modern atomic theory.

 

Millikan's Oil Drop Method:

1. Experimental Setup:
• Millikan's oil drop experiment involved using oil droplets in mist form, created by an atomizer.
• The oil droplets were allowed to enter an electrical condenser through a small hole in the upper plate.
• The motion of these droplets as they fell downward was observed using a telescope equipped with a micrometer eyepiece.
• By measuring the rate at which the oil droplets fell, Millikan was able to determine their mass.
2. Ionization of Air:
• The chamber containing the oil droplets was filled with air, which was ionized by passing a beam of X-rays through it.
• The ionization process created gaseous ions within the chamber.
3. Charge Acquisition by Droplets:
• The oil droplets acquired electrical charge by colliding with the gaseous ions created through air ionization.
4. Effect of Electric Fields:
• By applying voltage to the plates of the electrical condenser, an electric field was created within the chamber.
• Depending on the charge on the droplets and the polarity and strength of the applied voltage, the motion of the charged oil droplets could be manipulated.
5. Observations and Conclusions:
• Millikan observed the behavior of the charged oil droplets under the influence of the electric field.
• By carefully measuring the effects of the electrical field strength on the droplets' motion, Millikan reached a significant conclusion.
• He deduced that the magnitude of the electrical charge (q) carried by the oil droplets was always a whole number multiple (n) of the fundamental electrical charge (e).
• Mathematically, this relationship can be expressed as: q = n * e, where n can be any positive integer (1, 2, 3...).

In summary, Millikan's Oil Drop Method involved observing the motion of charged oil droplets in the presence of an electric field to determine the fundamental charge carried by these droplets. By analyzing the effects of electric fields on the droplets' behavior, Millikan found that the charge on the droplets was quantized, meaning it existed in multiples of the elementary charge (e). This experiment played a crucial role in confirming the quantized nature of electric charge and provided essential data for understanding sub-atomic particles.

 

J.J. Thomson's Model of the Atom:

1. Thomson's Proposal (1898):
• In 1898, J.J. Thomson introduced his atomic model, suggesting that an atom possesses a spherical shape with a radius of approximately 10^-10 meters.
• According to this model, the positive charge is evenly spread throughout the atom.
2. Electron Arrangement:
• Electrons are embedded within the positively charged sphere in a manner that establishes the most stable electrostatic arrangement.
• This arrangement aims to achieve electrostatic equilibrium within the atom.
3. Variety of Names:
• Thomson's model has been referred to by various names, including plum pudding, raisin pudding, and watermelon model.
• It is often visualized as a positively charged pudding or watermelon containing plums or seeds (representing electrons) embedded within it.
4. Uniform Mass Distribution:
• A significant characteristic of this model is the assumption that the atom's mass is uniformly distributed throughout its volume.
• This assumption implies that the positive charge and mass are distributed uniformly within the atom.
5. Explanation of Neutrality:
• Although this model was successful in explaining the overall neutrality of the atom (the equal number of positive and negative charges), it had limitations when compared to later experimental results.
6. Limitations and Later Experiments:
• Despite explaining neutrality, Thomson's model did not align with the outcomes of subsequent experiments.
• As more research and experiments were conducted, new insights emerged that required a more accurate depiction of atomic structure.
7. Recognition and Nobel Prize:
• J.J. Thomson's contributions to the field of physics, including his theoretical and experimental investigations on the conduction of electricity by gases, earned him the Nobel Prize in Physics in 1906.

In summary, J.J. Thomson's atomic model proposed a spherical atom with a uniform positive charge distribution and embedded electrons. Although the model explained atomic neutrality, it faced inconsistencies with later experimental findings. Thomson's significant work in the realm of electricity conduction in gases led to his Nobel Prize recognition in 1906.

 

Rutherford's Nuclear Model of the Atom:

Rutherford's Experiment:

1. Alpha Particle Scattering Experiment:
• Rutherford, along with his students Hans Geiger and Ernest Marsden, conducted the alpha particle scattering experiment.
• High-energy alpha particles were directed at a very thin gold foil (approximately 100 nm thick) in the presence of a fluorescent zinc sulfide screen.
2. Unexpected Results:
• The results of the experiment were unexpected and contradicted Thomson's model of the atom.
• Thomson's model predicted that alpha particles would pass through a uniform distribution of mass in the gold atoms without significant deflection.
3. Observed Outcomes:
• (i) Most alpha particles passed through the gold foil undeflected.
• (ii) A small fraction of alpha particles were deflected by small angles.
• (iii) A very few alpha particles (approximately 1 in 20,000) were deflected nearly 180 degrees, bouncing back.

Conclusions and Rutherford's Model: 4. Key Conclusions:

• (i) Most of the atom is empty space, as evidenced by the majority of undeflected alpha particles passing through.
• (ii) Deflected alpha particles indicated the presence of a concentrated positive charge, contrary to Thomson's model.
• (iii) Rutherford's calculations showed that the volume occupied by the nucleus is extremely small compared to the total atom's volume. The nucleus is incredibly dense.

5. Rutherford's Nuclear Model:
• Rutherford proposed the nuclear model of the atom based on his observations and conclusions.
• (i) The atom contains a small, dense, positively charged nucleus, where most of the positive charge and mass are concentrated.
• (ii) Electrons revolve around the nucleus in circular orbits with high speeds, akin to planets orbiting the sun in the solar system.
• (iii) The attractive electrostatic forces between the positively charged nucleus and the negatively charged electrons keep the atom stable.

Rutherford's alpha particle scattering experiment yielded unexpected results that led to the development of the nuclear model of the atom. This model depicted the atom as having a concentrated, positively charged nucleus surrounded by electrons in circular orbits. The analogy to the solar system helped visualize this atomic structure. Rutherford's model laid the foundation for modern atomic theory and contributed significantly to our understanding of the atom's structure.