Atomic Structure: Sub-Atomic Particles

• Chemical behavior diversity in elements arises from differences in atoms' internal structure.
• Early Indian and Greek philosophers proposed atoms as fundamental matter building blocks around 400 B.C.
• The word "atom" comes from the Greek "a-tomio," meaning "uncut-able" or "non-divisible."
• These early atomic ideas were speculative without experimental basis.
• John Dalton proposed the first atomic theory of matter with a firm scientific basis, in 1808.
• Dalton's atomic theory successfully explained laws of conservation of mass, constant composition, and multiple proportions.
• Dalton atomic theory could not explain experiments showing electrically charged substances when rubbed.
• Late 19th and early 20th-century experiments revealed atoms consist of sub-atomic particles (electrons, protons, neutrons), challenging Dalton's theory.

Discovery of Sub-Atomic Particles

• Insights came from experiments involving electrical discharge through gases.
• Basic principle: like charges repel, unlike charges attract each other.

Discovery of Electron

• In 1830, Michael Faraday showed electricity passing through electrolyte solutions causes chemical reactions at electrodes.
• Faraday's work suggested electricity has a particulate nature.
• Mid-1850s: scientists study electrical discharge in partially evacuated tubes (cathode ray discharge tubes).
• Cathode ray tubes are made of sealed glass with two metal electrodes.
• Electrical discharge visible at low pressures, high voltages.
• High voltage application initiates current flow, with particles moving from negative (cathode) to positive electrode (anode).
• Current flow confirmed by making a hole in the anode and coating the tube with zinc sulfide (phosphorescent material).

Cathode Ray Results

• Cathode rays originate at the cathode and move toward the anode.
• Rays are not visible, but their behavior can be observed using fluorescent or phosphorescent materials.
• In absence of electric or magnetic fields, the rays travel in straight lines.
• Electric or magnetic fields deflect cathode rays in a manner consistent with negatively charged particles.
• Conclusion: cathode rays consist of negatively charged particles called electrons.
• Cathode ray (electron) characteristics do not depend on electrode material or gas in the tube.
• Electrons are a fundamental constituent of all atoms.

Charge to Mass Ratio of Electron

• In 1897, J.J. Thomson measured the ratio of electron charge (e) to mass (me).
• Thomson used a cathode ray tube with electric and magnetic fields perpendicular to each other and the electron path.
• Applying only electric field: electrons deviate and hit point A.
• Applying only magnetic field: electrons hit point C.
• Balancing electrical and magnetic field strength returns the electron to its original path (point B).
• Thomson argued that deviation amount depends on the particle's negative charge magnitude.
• Stronger charge results in stronger interaction and deflection.
• Deviation depends on the particle mass; lighter particles deflect more.
• Deflection grows with increasing voltage or magnetic field strength.
• Thomson determined e/me = 1.758820 × 10¹¹ C kg⁻¹.
• me is electron mass in kg, and e is electron charge magnitude in coulombs (C).
• Electrons are negatively charged and denoted as -e.

Charge on the Electron

• R.A. Millikan (1868-1953) used the oil drop experiment (1906-14) to find the electron charge
• The process involves observing the motion of charged oil droplets in an electric field to determine the charge
• Millikan found the electron charge to be -1.6 x 10^-19C
• The present accepted value of electrical charge is -1.602176 x 10^-19C
• Combining the charge results with Thomson's e/m ratio gives the electron mass:
  • me = (1.602176 x 10^-19C) / (1.758820 x 10^11 C kg-1) = 9.1094 x 10^-31 kg

Discovery of Protons and Neutrons

• Modified cathode ray tubes lead to the discovery of canal rays or positively charged particles.
• Positively charged particles differ from cathode rays.
• Unlike cathode rays, mass of positively charged particles depends on the gas present in the tube.
• Charge to mass ratio relies on gas origin.
• Some carry a multiple of the fundamental electrical charge unit.
• Behavior in magnetic/electrical fields is opposite to electrons.
• The smallest, lightest positive ion from hydrogen called a proton, was characterized in 1919.
• Chadwick (1932) discovered neutrons by bombarding beryllium with alpha particles.
• Neutrons are electrically neutral particles with slightly more mass than protons.

Atomic Models

• Experiments suggested Dalton's indivisible atom is composed of sub-atomic particles with positive and negative charges.
• Major problems for scientists:
  • To account for atomic stability
  • To explain elemental behavior in physical and chemical properties.
  • To explain molecule formation via atom combination
  • To understand electromagnetic radiation

Millikan's Oil Drop Method

• Oil droplets in the form of mist, produced by the atomiser, were allowed to enter through a tiny hole in the upper plate of electrical condenser
• Downward motion of these droplets viewed through telescope with micrometer eye piece
• Millikan measured oil droplet mass and the air inside the chamber was ionized by passing a beam on X-rays through it
• Electrical charge acquired through collision with gaseous ions
• By measuring the effects of electrical forces applied to the plate's voltage, Millikan concluded the magnitude of the charge on the droplets on oil is an integral multiple of the electrical charge
• q=ne, where n=1, 2, 3

Thomson Model of Atom

• Also called plum pudding, raisin pudding or watermelon
• J.J Thomson proposed atom possesses spherical shape with radius approx 10^-10m, positive charge distributed uniformly
• Electrons embedded to give stable electrostatic arrangement
• visualized as pudding or watermelon with positive charge, plums, seeds, electron embedded
• Important feature: mass of atom assumed to be uniformly distributed
• Model explained overall neutrality of atom but contradicted later results
• Thomson awarded Nobel Prize in Physics in 1906 for electricity conduction in gases investigations.
• Röntgen (1845-1923) discovered rays that can cause fluorescence in fluorescent materials, named them X-rays, produced effectively when electrons strike the dense metal anode, called targets

Radioactivity

• Becquerel (1852-1908) noticed certain elements emit radiation spontaneously.
• This phenomenon is called radioactivity, elements radioactive elements.
• Marie Curie, Pierre Curie, Rutherford, and Frederick Soddy developed the field.
• Rutherford proved a-rays are high-energy particles with two positive charge units and four atomic mass units.
• α particles = helium nuclei (combined with two electrons = helium gas).
• β rays are negatively charged, like electrons.
• γ rays are high-energy radiations such as X-rays, are neutral in nature and do not consist of particles
• Penetrating power (least to greatest): alpha, beta (100x alpha), gamma (1000xalpha).

Rutherford's Nuclear Model of Atom

• Rutherford and students (Geiger and Marsden) bombarded thin gold foil with a-particles.
• Experiment: high-energy a-particles directed at thin gold foil (~100 nm), with a circular fluorescent zinc sulfide screen.
• Expected results: a-particles would pass through with slight slowing/deflection(even distribution of mass)
• Actual results were unexpected.
  • Most a-particles passed through undeflected.
  • Small fraction deflected at small angles.
  • Very few bounced back (deflected nearly 180 degrees).
• Rutherford's conclusions:
  • Most of the space in atom is empty
  • Positive charge and mass concentrated in a small volume (nucleus)- enormous repulsive force
  • Nucleus volume is small compared to total atomic volume.
• Radius of atom ~10⁻¹⁰ m, radius of nucleus 10⁻¹⁵ m. If a cricket ball is the nucleus, then the atom's radius would be about 5 km.
• Nuclear Model of the Atom with Postulations:
  • Most of the mass and positive charge found in the nucleus
  • Nucleus surrounded by very fast electrons in orbits
  • Electrons and nucleus electrostatic forces of attraction

Atomic Number and Mass Number

• Positive charge on nucleus due to protons.
• The charge on proton equals the electron charge magnitude (but opposite sign).
• Number of protons defines the atomic number (Z)
  • Hydrogen has 1 proton, sodium has 11 (Z=1, Z=11) for example
• In order to keep the electrical neutrality, the number of electrons in an atom is equal to the number of protons (atomic number, Z)
• atomic number (Z) = # of protons in the nucleus = # of electrons in a neutral atom

Nucleons

• Protons and neutrons are nucleons
• Mass number (A) = total number of nucleons.
• Mass number (A) = number of protons (Z) + number of neutrons (n).

Isobars and Isotopes

• Composition noted by showing element symbol(X) mass number(A as superscript to left) and atomic number(Z as subscript to left).

• Isobars have same mass number but different atomic number (example ¹⁴C and ¹⁴N).

• Isotopes have same atomic number but different mass numbers.

• Isotopes differ due to containing different number of neutrons

• 99.985% of hydrogen atom has 1 proton= protium(¹H)

• Some hydrogen contains isotopes: containing 1 proton+1 neutron = deuterium(²D), possessing 1 proton+2 neutrons= tritium(³T)

• Chemical properties controlled by number of electrons which is influenced by the number of protons in the nucleus

• Neutrons have little effect on chemical qualities of an element = Isotopes have same chemical behavior.

Problems

• Bromine has atomic number (Z) = 35, mass number (A) = 80. It is neutral. Protons = electrons = 35, Neutrons = 80 - 35 = 45.
• Species example is neutral with each having 18 electrons, 16 protons and 16 neutrons. Atomic number = 16(sulfur/S), mass number= 16+16=32. Must be an anion with two negative charges(18-16=2)
• Symbol= ³²S²⁻ _Always find if species is a neutral atom, cation or anion or if it is neutral atom, equation (2.3) is valid. ie #protons = #electrons= atomic number; If the species is an ion, determine whether the number of protons are larger (cation, positive ion) is (anion, negative ion)

Problems With The Rutherford Model

• The Rutherford's model describes atoms as small scale solar system, nucleus acting as sun, electrons act as planets, but it does not explain atoms
• Applies the principles of classical mechanics to atom as a small scale solar system, which doesn't fit since it applies mass
• Planets' well described orbits work because of gravitational forces with this expression :
• G* {(m₁ m₂)/r²} where 'G= Gravitational force' and m₁, m₂ are masses and 'r' describes distance from center of the masses
• The nucleus mathematically similar to solar system (coloumb force); F= {(kq₁q₂)/r}, "k= Constant, q₁ and q₂ represent charges, r= Separation distance of charges. However, when bodies are accelerated, there can be electromagnetic radiation
• Accelerations with accelerated speeds = Maxwell Electromagnetic Theory = Charged particles release electromagnetic radiation.
• Electrons emit radiation/radiation carries electron movement and the orbit will shrink. Shrinking= electrons should hit nucleus. However this doesn't happen
• Rutherford theory is not correct
• There can be electrostatic attraction where nucleus is stable

Problems that the Rutherford Model can explain

• Says nothing about distribution of electrons about nucleus/ electron strength

Developments Towards Bohr's Model of the Atom

• Radiation's studies with matter provided information on atom/molecules building
• Bohr used info from matter interaction w/ radiation to better Rutherford Model by:
  • Radiation having both wave and particles features(dual characteristic)
  • Atomic spectra w/ experimental results
• Experimental parts first

Wave Nature of Electromagnetic Radiation

• Physicists research absorbent of heated objects in mid-1800s - Known as thermal radiation and consisted electromagnetic waves w frequencies
• Maxwell(1870) described relations among matter of electrically charged parts/magnetic fields
• When particle charged moves under acceleration = alternating electrical/magnetic fields made and sent out.
• Electromagnetic Radiation: Transmission via waves
• Light's forms are known via speculation Light thought to be made of bits= corpusles (Newton) & Light is from oscillations of electricity+ magnetism(Maxwell) (Fig 2.6)

EM Wave Simple Properties - Oscillating

• Oscillating Electric/Magnetic Fields = perpendicular
• EM Waves don't need medium, thus can move in vacuum

EM Radiation

• Different Electromagnetic Radiation Types:Wavelength/ frequency is different
  • = EM spectrum(Fig.2.7) - Different areas are identified by name - Broadcast, Microwaves(Radar) - Infrared(Heating) - Ultra-Violet (Component of sunlight)
• Wavelengths are noted using meters
• Hertz is used to recognize frequencies(Heinrich Hertz called the units Hertz)
• 3.0*10^8 m s^-1 speeds all radiation (Speed of light)
• The relation of light, frequencies, and wave length goes from this expression: c(light speed) = v(frequency)* λ (Wavelengths) and "C= vλ"

Wave or wave number formula

• Wave number equation goes from Wave numbers = wave length^-1
• wave length= c(Speed of light) / v's Frequency"

Planks Quantum Theory

• Diffraction + Interferences = EM Radiation described by wave form
• However: Body emissions cant be by wave mechanics
• Black body radiation phenomenon:Hot bodies transmit electromagnetic radiation everywhere – High temp= radiation visible for each point in EM Spectrum. As temperature rises % of short wavelength(blue light) radiation is made
• An Ideal Body Absorbs(Black body) = Emits radiation @ all regularity.
• Emitted regularity depends @ temperature.
• Radiation attempts will predict functions of length but will not be explained correctly by wave theory
• Radiation with Black bodies = Oscillators interaction with energy to find frequency changes with interactions with electromagnetic radiation
• Planck = Radiation is energy divided discreetly
  • Atoms discharge energy limited &not continuously
  • Quantum = Energy amount that is smallest - Proportional to frequencies in (2.6) E =hv H = Planck's constant: 6.626* 10^-34 J's
• Staircase Analogy = You can stand on different steps, but cannot stand between them(energies follow this)

Photoelectric Effects

• Hertz- Electrons discharge via rays on metals = Photoelectric Effects(1887)
  • Beam touching = Electrons emitted
  • Count electrons emitted
  • Energy measured using battery - Minimum level- photoelectric effect is not seen = Threshold frequencies. Energies >v. Kinetic energy is present
  • Kinetic energies are faster as light gains energy.
• Energy/Speed can be found regardless of light's speeds
• • Brightness isn't what provides energy/speed but speed's frequency. Classical Physics says the opposite;

Theory's Flaws

• Beam of sunlight goes onto metal surface = Shooting range
• Electrons shot out by instant energy
• More frequencies = High energy, More electrons
  • Kinetic Electrons = frequency to electromagnetic radiations
• Equation: Light is able to explain photoelectric effects

Electromagnetic Radiations

• Waves or movement/particle = double nature
• Radiation intermix with matter = Particle traits while waves/diffractions = wave traits
  • Ancient thinking wasn't convinced
  • Microscope and electrons exhibit properties with wave particle

Problems of light

De Brogile Equation

• Mass and velocity calculate wavelength

Heisenberg's Uncertainty Principle

• Cannot Determine momentum or location easily of electrons simultaneously

Reasons For Failure/Limitations in The Bohr Model

• Finer details - Hydrogen molecules aren't accounted
  • Hydrogen, helium, and two electron elements unable to split
• Bonded molecules aren't able to be made.

Towards Quantum Mechanical Model of The Atom.

• Dualities/ Uncertainties = Help with formulas by

Quantum Mechanical of the Atom

• Newton + Macroscopic is where this is seen as successful
• Does not work where concepts don't matter
• Quantum Mechanics= Take double nature in regards with matter

Problems With Hydrogen atoms and Schrodinger Equations

• Schrodinger equation cannot happen in multi areas atoms

Mechanics/Orbitals

• Schrodinger theory can describe how to solve atoms for energy - Three quantum numbers;N,l,ml

Features

• Energies are measured when related to an atom example is Electrons locked into nucleus – Existence for quantified EM levels leads as electrons are Wavelike to meet conditions for wave equation
• . Velocities/positions not found easily. Leads to statements regarding position of electrons Atomic orbital’s are wave lengths electrons have, since wave functions possible for electrons atoms have various orbital’s Energy must exist.