John Dalton Robert Andrews Millikan Niels Bohr (1803) (1909) (1912) Joseph John Thomas Ernest Rutherford Quantum Model (1897) (1912) (1924) TIMELINE
JOHN DALTON- Around 1803, Dalton developed an atomic theory to explain the ratios inwhich elements combine to form compounds. It was the cornerstone formodern atomic theory.- Five main points of Dalton’s atomic theory:1. Elements are made of extremely small particles called atoms.2. Atoms of a given element are identical in size, mass, and other properties:atoms of different elements differ in size, mass, and other properties.3. Atoms cannot be subdivided, created, or destroyed.4. Atoms of different elements combine in simple whole-number ratios toform chemical compounds.5. In chemical reactions, atoms are combined, separated, or rearranged.- He also developed an assumption based on the faith of nature’s simplicitythat when atoms combine in only one ratio, it must be presumed to be abinary one, unless some cause appear to the contrary.- Limitations included that nowadays, Dalton’s second and third points to theatomic theory and proven wrong.
J. J. THOMSON- In 1896 , he took the cathode ray experiments a step further by ﬁrstlyimproving Perrin’s version to more clearly prove cathode rays do carry negativecharges. With this , Thomson then went on to discover the electron through hisdemonstration of cathode rays responding to electrode ﬁelds just as negativelycharged particles would.- Thomson had ﬁgured out a way to determine the charge of the mass by usingboth an electric and magnetic ﬁeld.- Used mutually perpendicular electric and magnetic ﬁelds to determine thespeed of cathode rays. Then with only one ﬁeld turned on, he measured thedeﬂection of rays. These deﬂections depended on magnitude of ﬁeld, length ofpath in the ﬁeld, and the speed, mass, and charge of cathode-ray particles.- With calculation, he found reasonably consistent values for the charge-massratio, which allowed him to conclude that all cathode rays consist of identicalparticles with exactly the same negative charge.- q/m for an electron is roughly 1011 C/kg- Reasoned electrons are much smaller than atoms.- Put forward the daring theory that atoms were divisible and tiny particles incathode rays were “the substance from which all the chemical elements are builtup.” Although he was wrong about electrons being the only constituents ofatoms, recognizing electrons are subatomic particles was a major advance inatomic physics.- Atoms are neutral and since electrons are constituents of atoms, atoms mustcontain some form of positive charge.- Suggested that atoms might consist of electrons embedded in a blob ofmassless positive charge; raisin-bun model.- Limitations included how Thomson had made no tests to indicate thatelectrons were embedded in a positive blob, but merely suggested the idea.
ROBERT ANDREW MILLIKAN- In 1909, he determined the unit charge ofthe electron with his oil drop experiment tobe 1.60 x 10-19.- Millikan calculated the mass of each dropfrom its diameter, then observed the motionof the oil drops in a uniform ﬁeld. Byanalyzing this motion, Millikan calculated theelectric force acting on each drop. He foundthat the charge of each oil drop was amultiply of 1.60 x 10-19.- Since others had already determined thecharge-mass ratio, Millikan could nowcalculate a reasonably accurate valuable forthe mass of the electron.- Limitations included how out of 175measurements, he only reported 58.
ERNEST RUTHERFORD- By 1909, he had shown that some radioactive elements, such as radium andthorium, emitted positively charged helium ions, which are also known as alphaparticles and when passed through a thing sheet of mica, a beam of alpha particleswill spread out.- He had a technique that allowed him and his assistants, Hans Geiger and ErnestMarsden, to measure the proportion of alpha particles scattered at different anglesfrom various materials. They would produce a pencil-shaped beam of alpha particlesand position a thin sheet of gold foil at a right angle to the beam. Then they woulduse a screen coated with zinc sulﬁde, which would detect the scattered particles byletting off faint ﬂashes of light visible with a microscope. By moving the screen andmicroscope around the foil, they were able to measure the rates at which alphaparticles appear at various angles. Eventually, they concluded the positive charge in agold atom must be concentrated in an incredibly tiny volume, so most of gold wasactually empty space.- When using aluminium foil instead of gold, they proved that the positive chargeand most of the mass of an atom are contained in a radius less than 10 -14.- Discovered nucleus and disproved the raisin-bun model.- Lead to planetary model of atom consisting of electrons orbiting the nucleus of anatom and there being an electrostatic attraction between positive nucleus andnegative electrons, which provides the centripetal force to keep the electrons inorbit.- Limitations included the fact that Rutherford’s model was later adjusted by NielsBohr because in Rutherford’s model, the electrons should spiral into the nucleus in afew microseconds due to a constant acceleration, which would emit electromagneticwaves that would take energy from the orbiting electrons.
NIELS BOHR- Corrected the critical ﬂaw in Rutherford’s model.- Focuses on the quantization of energy of electrons.- Basic principles of Bohr’s model:1. Electrons can orbit the nucleus only at certain speciﬁc distancesfrom the nucleus. These distances are particular multiples of theradius of the smallest permitted orbit meaning the orbits in an atomare quantized.2. The electron’s distance from the nucleus determines both thekinetic and electric potential energy of an electron in orbit. So forththe energy in an electron is also quantized and each orbitcorresponds to a speciﬁc energy level for the electron.3. Only by emitting or absorbing photons of equal energy to thedifference between energy levels can an electron move from oneenergy level to another. When an electron continues to orbit at aparticular energy level, no energy is radiated. Also, since the size andshape of the orbit remains the same and at a ﬁxed energy level,these orbits are often referred to as stationary states.- Limitations included him not explaining as to why energy isquantized, why orbiting electrons do not radiate electromagneticenergy, why a magnetic ﬁeld splits the main spectral lines intomultiple closely spaced lines, and the fact that it is not accurate forelectrons to have two or more electrons.
QUANTUM MODEL- In 1924, Louis de Broglie developed his theory that particles havewave properties. He concluded this through diffraction experiments.So forth, the principles of interference and standing waves apply forelectrons orbiting a nucleus.- For most sizes of orbit, successive cycles of the electron wave will beout of phase, and destructive interference will reduce the amplitude ofthe wave. For constructive interference to occur, the circumference ofthe orbit must be equal to a whole number of wavelengths.- The wave nature of matter provides a natural explanation forquantized energy levels.- In 1926, Erwin Schrödinger derived an equation for determining howelectron waves behave in the electric ﬁeld surrounding a nucleus. Thesolutions to his equation are functions that deﬁne the amplitude of theelectron wave in the space around a nucleus.- Max Born showed that the square of the amplitude of these wavefunctions at any point is proportional to the probability distribution ororbital.- In quantum model, electrons behave as waves, which do not have aprecise location.-Orbitals in the quantum model show the likelihood of an electronbeing at a given point, not a path they follow.- Electrons behaving as waves rather than orbiting particles in an atomexplains why they do not radiate electromagnetic energy continuously.
ATOMIC SPECTRA- The atomic spectra is a range of characteristic frequencies ofelectromagnetic radiation that are readily absorbed and emittedby an atom.- An electron can jump from one ﬁxed orbital to another. If theorbital it jumps to has a higher energy, the electron must absorba photon of a certain frequency. If it’s a lower energy, theelectron must give off a photon of a certain frequency.- The frequency depends on the difference in energy betweenthe orbitals.- This relates to Bohr’s model because Bohr’s model describeshow in order for electrons to move from one orbital toanother, the electron must release or absorb a photon ofappropriate energy.
MATHEMATICSMass: Charge:Thomson - Thomson - m = Bqr/v q = Fe/EMillikan - q = Fm/vB m = qE/g q = vm/BrBohr - Millikan - m = h2/4π2rke2 q = mg/E m = Eh2/2π2k2e4 q = neQuantum Model - Rutherford - m = nh/2πvr q = E d/kq 1 p 2 q = E d/kq 2 p 1