Dalton's Atomic Theory

• 'ATOM' is the smallest particle of an element.
• 'Atom' originates from the Greek word meaning 'Not to be cut'.

Dalton's Atomic Theory (1803)

• Atom is considered a smallest, dense, hard, indivisible particle.
• Each element is made of a specific type of atoms.
• Elements' properties vary due to the different atoms they contain.
• The theory provides a basis for the law of chemical combination.

Dalton's Theory Limitations

• Fails to explain different atoms' valency and mass differences.
• Discovery of isobars and isotopes showed atoms of the same element can have different atomic masses and vice versa.

Discovery of Fundamental Particles

• Late 19th-century discoveries of subatomic particles such as protons, electrons, X-rays led to the idea that atoms could be divided into smaller parts.
• Atoms consist of subatomic particles like neutron, proton, electron, neutrino, positron etc.
• Electrons, protons, and neutrons are fundamental subatomic particles

J.J. Thomson's Electron Discovery (1897)

• Electrons are negatively charged, discovered by J.J. Thomson.
• William Crookes studied electrical discharge in partially evacuated tubes known as cathode ray discharge tubes in 1879.

Cathode Ray Properties

• Cathode rays travel in straight lines.
• They produce a mechanical effect, are negatively charged and travel at high speed.
• Cathode rays can cause fluorescence.
• Cathode rays heat objects they strike due to kinetic energy transfer.
• X-rays are produced when cathode rays hit heavy metals.
• Cathode rays can ionize gases and produce scintillation on photographic plates.
• They can penetrate thin metallic sheets.
• Using variations in electric and magnetic fields, the charge to mass ratio of the electron was found is 1.758820 × 10¹¹ C/kg.
• R.S. Millikan measured the charge on each electron as -1.602 × 10⁻¹⁹ C via oil drop experiment.

Goldstein's Proton Discovery

• Protons are positively charged particles discovered by Goldstein
• Eugen Goldstein's Canal Ray experiment in 1886 led to the discovery of the proton.
• Using a discharge tube with a perforated cathode, Goldstein applied high voltage and saw a faint luminous ray.
• Anode rays' e/m value depends on the gas in the tube, and is maximized when the gas is hydrogen.

Chadwick's Neutron Discovery (1932)

• In 1932, Chadwick bombarded Beryllium with alpha particles and found neutral penetrating radiations and named the neutrons.
• When this radiation was directed towards paraffin wax (containing H₂), it ejected protons which were detected via an ionization chamber.
• Neutrons are relatively massive but neutral, therefore, can penetrate the nucleus of any element.
• Neutrons are fundamental particles in all atomic nuclei, except hydrogen (protium).

Properties of Anode Rays

• Travel in straight lines.
• Anode rays are material particles that are positively charged.
• Anode rays can be affected by external magnetic fields and photographic plates.
• The e/m ratio of anode rays is less than electrons.
• They produce flashes of light on a ZnS screen.

Composition of Matter by Mass and Charge

• Electron: Mass - 9.1096 × 10⁻³¹ kg, Charge - -1.602 × 10⁻¹⁹ Coulombs
• Proton: Mass - 1.6726 × 10⁻²⁷ kg, Charge - +1.602 × 10⁻¹⁹ Coulombs
• Neutron: Mass - 1.6749 × 10⁻²⁷ kg, Charge - Neutral

Thomson's Atomic Model

• J.J. Thomson proposed the first detailed atomic model.
• Thomson's model proposed the atom is a uniform sphere of positive charge with electrons at places within.
• This model is known as the Plum-Pudding model because electrons are distributed within its positively charged region.

Thomson's Model Limitations

• The mass is considered evenly spread and does not reflect the movement of electrons.

Rutherford's α-Scattering Experiment

• Rutherford bombarded a thin gold foil with high-speed alpha particles.
• Angular deflections of scattered alpha-particles were studied.
• Most alpha-particles passed through without deflection.
• Some alpha-particles deflected at small angles.
• Very few alpha-particles deflected at large angles, some deflected by 180°.
• Most of the alpha particles went straight implied large empty space within the atom.
• Deflection of a few alpha particles suggested positive charge and most of the mass is concentrated in a small space within the atom.
• Alpha-particles that come close to this point experience repulsion.

Core Tenets of Rutherford's Model

• Positively charged heavy mass occupies a small volume known as the nucleus, present at the center of the atom.
• Few alpha-particles strongly scattered or returned, suggesting a rigid nucleus and recoil due to direct collision.
• Number of protons increases with increasing atomic number, increasing repulsion.

Rutherford's Definition of Atomic Size

• Size of nucleus is 10⁻¹³ cm, equated to 1 fermi, which is 10⁻¹⁵ m
• Size of atom = 10⁻⁸ cm equating to 1 Å, which is 10⁻¹⁰ m
• The density of the nucleus is approximately +10¹⁷ kg/m³.

Limitations of Rutherford's Model

• The model does not follow Maxwell's theory, where a charged particle emits energy when revolving in an attractive field, reduces its loop, and collides with the nucleus.

Definitions related to Nucleus

• Mass Number (A) = Number of Protons (Z) + Number of Neutrons (n).
• Number of Neutrons (n) = Mass Number (A) – Number of Protons (Z)

Spectrum of Light

• Occurs when a continuous amount of energy is emitted.
• In reality discontinuous spectrum occur.

Atomic Number (Z)

• It equals the number of protons in the atom's nucleus and represents electrons in the neutral atom.

Mass Number (A)

• Is the sum of neutrons and protons, collectively known as nucleons.

Atomic Weight

• It is the average of the weights of all isotopes of an element.
• For two isotopes, Z₁ and Z₂, with weights W₁ and W₂, respectively, with % abundance (X₁ and X₂), the average atomic weight =(W₁X₁ + W₂X₂)/(X₁ + X₂).

Isotopes

• Soddy discovered isotopes which indicate the same element has different atomic forms with same atomic numbers, similar chemical properties, and varying physical characteristics.

Isobars

• Aston discovered Isobars are different atoms of different elements that share a similar mass number, but differ in their atomic number. Meaning they do not have the same chemical properties

Isotones

• These are atoms of different elements with an identical number of neutrons,

Isodiaphers

• These are atoms of different elements with the same difference between the number of neutrons and protons.

Isosters

• Are molecules sharing the same number of atoms and electrons.

Isoelectronic

• Atoms, ions, molecules possessing a similar number of electrons.

Relative Abundance

• Isotopes of an element occur in different percentages naturally.
• Average atomic mass can be calculated using relative abundance.

Wave Motion

• Transfers energy from one point to another point.
• Waves are characterized by wavelength, frequency, time period, wave number, amplitude, and velocity.

Wavelength (λ)

• The distance between two nearest crests or troughs measured in Ångstrom (Å), picometer (pm), nanometer (nm), centimeter (cm), or meter (m).

Frequency (v)

• The number of waves passing a point in 1 second, measured in Hertz (Hz), sec⁻¹, or cycle per second (cps).

Electromagnetic Waves (EMW)

• Contain electric and magnetic fields and transfer energy at the speed of light without needing a medium and have perpendicular propagating directions.
• Types include cosmic rays, gamma rays, X-rays, UV, visible light, IR, microwaves, and radio waves.

Electromagnetic Spectrum

• Arrangement of various electromagnetic radiations based on increasing wavelengths or frequencies.

Maxwell's Theory of Electromagnetic Waves

• Radiations have wave nature for interference and diffraction, consist of oscillating electric and magnetic fields perpendicular to each other and the propagation direction.
• All travel at the speed of light and is directly proportional to amplitude.

Spectrum

• When light from a source is scattered with a prism the different wave lengths deviate through different angles.

Types of Spectrum

• Continuous is when sunlight is passed through a prism, and gets dispersed into different colors.
• Line is a series of thin bright lines derived from an excitation analyzed via a spectroscope for a series of color.
• Absorption spectrum occurs when white light passes through a substance, which absorbs particular wavelengths.

Planck's Quantum Theory

• Diffraction and interference phenomena are described by wave nature of electromagnetic radiation.
• Hot bodies emitting radiation, electron ejection from a metal surface when radiation strikes it.
• The theory states atoms/molecules emit or absorb energy in discrete quantities or small packets known as quantum and not arbitrary amounnts.

Black Body Radiation

• Radiation is produced when a black body emits and absorbs radiations all frequencies

Photoelectric Effect

• Photoelectric effect occurs when light of a certain minimum frequency strikes a metal surface, ejecting electrons.
• The minimum frequency is the Threshold frequency that causes photo electric emission
• Increasing the frequency or decreading the wavelength of the light ray resultins in emission of more photons
• Maxiumums in kinetic energy occur during photoelectric effect

Bohr's Atomic Model

• Based on quantum theory, explaining the stability of the atom and sharp spectral lines.
• Atom is composed of a nucleus where the protons and neutrons are present.
• The Negatively charged electron are revolving around the nucleus.
• Quantized energy levels, explaining atomic spectra.
• Absorption or emission of radiation depend on jumps between an atom takes from moving in on stationary orbit to another.
• Emitted radiation occurs as a single quantum, and has fixed frequencies.

Important Equations for Bohr's Model

• Angular momentum = mvr = nh/2π
• Radius of orbit = r = 0.529 × (n²/Z) Å
• Velocity of electron = v = 2.188 × 10⁶ × (Z/n) m/sec
• Energy = E = -13.6 × (Z²/n²) eV / atom

Bohr's Theory Limitations

• The theory is successful for hydrogen atoms and failed in multi electron species in predicting and accounting the line spectra of atoms.
• Also, theory does not explain the presence multiple spectral lines, splitting of spectral lines or state an uncertainty principle.

Bohr-Sommerfeld Theory

• Introduced path of an electron around nucleus is an ellipse with nucleus at one of the foci. Circular orbit is special case of the ellipse.
• Velocity of electron moving in an elliptical orbit varies at different parts of the orbit. This causes relavitistic variation in mass of moving electron. Therefore, he took into account relativistic variation of mass of electron with velocity.

Quantum Numbers and Quantum Theory

• To describe the detailed picture the electron, the quantum are the identification numbers are required

Principle Quantum number

• Represents the orbit and the distance of the election
• Expressed as n
• Limited values between 1 to 7

Azimuthal Quantum Number

• Also know as orbital number to state shapes of an orbital.
• This variable is expressed as the shape to determine where electrons are located

Hund's Rule

• Orbitals are evenly filled with spin before added electrons

Pauli's exclusion Principle

• No two electrons are similar
• The quantum helps explain the wide range of physical phenomena known today
• Different stack to be at different spins

Symmetrical Distribution

• Symmetrical shaped electrons result the most attractions and high stability in orbitals