Atomic Structure History TimelinePPT

1. Atomic Structure History TimelineDemocritusv    5th century Greek philosopherv    matter is composed of small, indivisible particles – atomos (uncuttable, indivisible)v    contemporaries, Plato and Aristotle, rejected ideas

2. John Daltonv    1808v    formulated the precise definition of atoms:Ø     Elements are formed from extremely small particles – atoms. All atoms of a given element are identical, with same size, mass and chemical properties. Atoms of one element are different from all other elements.Ø     Compounds are composed of atoms of more then one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.Ø     A chemical reaction involves only the separation, combination or rearrangement of atoms; it does not result in their creation or destruction

3. J.J. Thomsonv    Turn of the 20th centuryv    Considered to have discovered the electron, via experiments with cathode ray tubes. Ø     Obtained the ratio of electric charge to mass of an electron.v    Described an atom as a uniform positive sphere of matter in which electrons are embedded.v    Plum pudding theory:Ø     +Electrons are like plums embedded in a pudding of positive matter.

4. Robert Millikenv    Measured the charge and the mass of an electron via oil drop experiments to be -1.75 x 108 C/g (C is coulomb, measuring electric charge)Ø     Aka elementary or fundamental unit of chargeØ     Oil was sprayed out of an atomizer, the small oil drops would pick up some unknown number of electrons due to friction between the oil and the nozzle.  These oil drops were sprayed between two oppositely charged parallel plates.  Light shined upon the oil drops from the side caused them to show up like little stars when looking through a microscope to view the region between the plates.  The weight of an oil drop could be calculated by using the terminal velocity of the freely falling oil drop and the formula for terminal velocity of a sphere.  Once the weight was determined, the voltage across the plates could be adjusted until the upward electric force exactly cancelled out the weight of the oil drop.   When the exact balance was reached, the oil drop would remain suspended between the plates.

5. Wilhelm Roentgenv    Germanyv    1895 discovered x-rays (aka Roentgen rays). Ø     Unlike cathode rays, x-rays can pass through glass. Ø     Bone absorbs more x-rays then muscle.

6. Henri Becquerelv    French, late 1890’sv    Studied fluorescent properties of substances and was the first to witness radioactivity (a term later coined by Marie Curie) in uraniumØ     A compound containing uranium darkened photographic plates wrapped in paper or thin metal sheets – with stimulation of cathode rays.Ø     These rays are also able to discharge an electroscope that indicated rays that have an electric charge.v    Invented the phosphoroscope

7. Marie Curiev    Building upon the work of Roentgen and Becquerel, Marie and her husband, Pierre:Ø     discovered polonium and radiumØ     established that beta rays (now known to consist of electrons) are negatively charged particlesv    Died of leukemia caused by exposure to radiation

8. Ernest Rutherfordv    1911v    Rutherford identified the three main components of radiation and named them alpha, beta, and gamma raysv    From experiments using gold foil and other metals as targets for alpha particles emitted by a radioactive source. Some particles penetrated the foils with no or minimal deflections while some were deflected at a large angle. v    Proposed a new model of atomic structure, (other then Thomson’s, who he had studied under) Ø     Describing the atom as a dense nucleus about which electrons circulate in orbits.Ø     Most of an atom is empty spaceØ     Positive charges are concentrated in a central core within the atom – the nucleusØ     When the experimental alpha particle traveled close to the nucleus it experienced a large repulsive force – resulting in deflection. If the particle traveled into the nucleus the deflection would reverse the direction of the particle.

9. James Chadwickv    1932v    Discovered neutrons – electronically neutral, higher mass then protonsØ     Bombarded a thin sheet of beryllium with alpha particles and the beryllium emitted a very high energy radiation the resembled  rays.Ø     This discovery completed the mass ratio question.

10. Werner Heisenburgv    ~1932v    Heisenberg uncertainty principle: It is impossible to know simultaneously both the momentum (p, mass x velocity) and the position of a particle (x) with certainty.Ø     ΔxΔp≥h/4pv    A consequence of the dual nature of matter – All matter behaves as waves and particles (thanks to de Broglie’s equation. = h/mv)\

11. Max Planckv    1900v    Quantum Theory - atoms and molecules emit energy only in whole-number multiples of certain well-defined quantities. – energy can be released only in certain definite amounts – quanta.Ø     Quantum – the smallest quantity of energy that can be emitted or absorbed in the form of electro magnetic radiation. §        The energy of a single quantum is proportional to the frequency of the radiation. E v§        Planck’s constant: E = hv ·       h = 6.63 x 10-34 J§        Energy is emitted in multiples of hv: hv, 2hv, 3hv…Ø     Prior to Planck, classical physics assumed that atoms and molecules could emit or absorb any arbitrary amount of radiant energy.

12. Niels Bohr (1885)v    1913v    Accounting for spectral line positions – while electrons moved in circular orbits, the single electron in H could only be in certain orbits. Each orbit has a specific energy, so the energies associated with electron motion in the permitted orbits have a fixed value – quantized.v    The radiation emitted by an energized H atom is caused by the electron dropping from a higher-energy orbit to a lower on and giving up a quantum of E (photon)Ø     Prior the atom was envisioned having electrons whirling around a nucleus in circular orbits at high velocities. In hydrogen it was assumed that electrostatic attraction between the positive “solar” proton and the negative “planetary” electron pulls the electron inward and this force is balanced exactly by the acceleration due to the circular motion of the electron.v    Energies possessed by electrons in the H atom: En = -RH(1/n2)Ø     RH = 2.18 x 10-18 Ø     n = principal quantum number. Values n = 1, 2, 3…Ø     Negative sign implies that all allowable energies of the electron are negative. The energy of the electron in the atom is lower then the energy of a free electron (one infinitely far from the nucleus = 0)Ø     n = 1: most stable ground stateØ     excited state is any value higher then n = 1