Atomic Models - Comparision Of Size Of Atom & Closest Approach

Atomic Models

Comparision of Size of Atom & Closest approach Next Slide:

Thompson’s Model of the Atom

Proposed by J.J. Thompson in late 19th century
Atom is made up of a positively charged solid sphere
Electrons are embedded in the sphere
Particles are evenly distributed Next Slide:

Rutherford’s Model of the Atom

Proposed by Ernest Rutherford in early 20th century
Atom consists of a small, dense, positively charged nucleus
Electrons revolve around the nucleus in orbits
Most of the atom’s mass is concentrated in the nucleus
Nucleus is positively charged Next Slide:

Bohr’s Model of the Atom

Proposed by Niels Bohr in 1913
Electrons revolve around the nucleus in specific energy levels or orbits
Electrons can jump from one energy level to another by absorbing or emitting energy
Each energy level can hold a specific number of electrons
Bohr’s model explained the line spectrum of elements Next Slide:

Quantum Mechanical Model of the Atom

Proposed in the 1920s
Based on the principles of quantum mechanics
Describes the behavior of electrons in terms of probabilities
Electrons are not in fixed orbits but occupy orbitals around the nucleus
Orbitals are regions of space with the highest probability of finding an electron Next Slide:

Comparing Sizes of Atoms

The size of an atom is determined by the distance between the outermost electron and the nucleus
According to Thompson’s model, atoms have a uniform size due to evenly distributed electrons
Rutherford’s model brought the concept of a concentrated positive nucleus, resulting in smaller atomic sizes
Bohr’s model introduced the idea of specific energy levels, affecting the atomic size Next Slide:

Closest Approach of Atoms

Closest approach of atoms occurs when their electron clouds overlap as they come close to each other
Thompson’s model suggests no repulsion or attraction between atoms as they are uniformly distributed
Rutherford’s model predicts repulsion due to concentrated positive charges in the nucleus
Bohr’s model also predicts repulsion but to a lesser extent compared to Rutherford’s model Next Slide:

Electron Shielding

Electron shielding refers to the effect of inner electrons on outer electrons in an atom
Due to electron shielding, outer electrons experience reduced effective nuclear attraction
In Thompson’s model, electrons are evenly distributed, so electron shielding is not significant
In Rutherford’s model, electron shielding is not considered
Bohr’s and Quantum Mechanical models take electron shielding into account Next Slide:

Atomic Radii Trend

Atomic radii generally increase down a group (vertical column) on the periodic table
Atomic radii generally decrease across a period (horizontal row) on the periodic table
The trend can be explained by the increasing number of energy levels and effective nuclear charge Next Slide:

Summary

Atomic models such as Thompson’s, Rutherford’s, Bohr’s, and Quantum Mechanical have evolved over time
Each model provides a different understanding of the atom’s structure and behavior
The size of an atom and closest approach are influenced by the atomic model
Electron shielding affects the interaction between atoms
Atomic radii display trends on the periodic table due to energy levels and effective nuclear charge

Atomic Structure

Atoms are composed of three subatomic particles: protons, neutrons, and electrons
Protons have a positive charge, neutrons are neutral, and electrons have a negative charge
The number of protons determines the element and is called the atomic number
The sum of protons and neutrons gives the atomic mass

Protons and Neutrons

Protons and neutrons are located in the nucleus of an atom
Protons have a mass of approximately 1 atomic mass unit (u)
Neutrons also have a mass of approximately 1 atomic mass unit (u)
Protons and neutrons are collectively known as nucleons

Electrons

Electrons are negatively charged particles
Electrons orbit around the nucleus in specific energy levels
The energy levels are represented by the principal quantum number (n)
Each energy level can hold a maximum number of electrons

Valence Electrons

Valence electrons are the outermost electrons of an atom
Valence electrons are involved in chemical reactions and bonding
The number of valence electrons influences an element’s chemical properties
Valence electrons are located in the outermost energy level of an atom

Example: Electron Configuration of Oxygen (O)

Oxygen has an atomic number of 8
Its electron configuration is 1s² 2s² 2p⁴
There are 2 electrons in the 1s orbital, 2 in the 2s orbital, and 4 in the 2p orbital

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons
Isotopes have the same number of protons and electrons, but different atomic masses
Isotopes are denoted by the element’s name followed by the atomic mass (e.g., Carbon-12, Carbon-13)

Radioactive Decay

Radioactive decay is the process by which unstable nuclei undergo spontaneous transformation
Three common types of radioactive decay are alpha decay, beta decay, and gamma decay
Alpha decay involves the emission of an alpha particle (2 protons and 2 neutrons)
Beta decay involves the emission of a beta particle (an electron or positron)
Gamma decay involves the emission of gamma rays (high-energy photons)

Example: Alpha Decay of Uranium-238

Uranium-238 undergoes alpha decay
It emits an alpha particle and transforms into Thorium-234
The alpha particle consists of 2 protons and 2 neutrons

Half-Life

The half-life of a radioactive substance is the time it takes for half of the initial quantity to decay
Half-life is a constant characteristic of each radioactive isotope
Half-life can be used to determine the age of fossils, archaeological artifacts, and geological samples

Example: Carbon-14 Dating

Carbon-14 is a radioactive isotope used for carbon dating
The half-life of Carbon-14 is approximately 5730 years
By measuring the ratio of Carbon-14 to Carbon-12 in a sample, the age of the sample can be estimated

Atomic Spectra

Atomic spectra are the unique patterns of light emitted or absorbed by atoms
Each element has a distinct atomic spectrum
Atomic spectra can be used to identify the presence of specific elements
Atomic spectra are obtained through spectroscopy Next Slide:

Emission Spectra

Emission spectra are produced when atoms emit light
When electrons transition from higher energy levels to lower energy levels, they release energy in the form of photons
Each electron transition corresponds to a specific wavelength of light
Emission spectra consist of bright lines at specific wavelengths Next Slide:

Absorption Spectra

Absorption spectra are produced when atoms absorb light
When atoms are exposed to light of specific wavelengths, electrons absorb photons and transition to higher energy levels
The absorbed wavelengths are missing in the resulting spectrum, creating dark lines on a continuous spectrum
Each element has a unique absorption spectrum Next Slide:

The Balmer Series

The Balmer series is a set of spectral lines in the emission spectrum of hydrogen
The Balmer series is in the visible region of the electromagnetic spectrum
The series is described by the Balmer formula: Lambda = R_H(1/2^2 - 1/n^2) where Lambda is the wavelength, R_H is the Rydberg constant, and n is the principal quantum number Next Slide:

Bohr’s Formula for Energy Levels

According to Bohr’s model, the energy of an electron in a particular energy level is given by: E = (-13.6 eV) / n^2 where E is the energy, -13.6 eV is the ionization energy of hydrogen, and n is the principal quantum number Next Slide:

Photoelectric Effect

The photoelectric effect refers to the emission of electrons when light is incident on a material surface
The effect can only be explained by considering light as composed of particles called photons
The energy of a photon is given by E = hf, where E is the energy, h is Planck’s constant, and f is the frequency Next Slide:

Wave-Particle Duality

Wave-particle duality is the concept that light and matter have both wave-like and particle-like properties
Experiments such as the double-slit experiment support this duality
The wave-particle duality is central to quantum mechanics Next Slide:

The Uncertainty Principle

The uncertainty principle states that the position and momentum of a particle cannot both be precisely determined
The more accurately we know the position of a particle, the less accurately we know its momentum, and vice versa
The uncertainty principle is a fundamental principle in quantum mechanics Next Slide:

Quantum Numbers

Quantum numbers describe the properties and characteristics of electrons in an atom
There are four types of quantum numbers: principal quantum number (n), azimuthal quantum number (l), magnetic quantum number (m_l), and spin quantum number (m_s)
Quantum numbers determine the energy, shape, orientation, and spin of an electron Next Slide:

Heisenberg’s Uncertainty Principle

Heisenberg’s uncertainty principle states that it is impossible to simultaneously know the precise position and momentum of a subatomic particle
The principle arises from the wave-particle duality of particles
The uncertainty principle places limits on our ability to measure certain properties of particles Next Slide: