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Contributions Made to the Advancement of the Atomic Model

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Transcript of Contributions Made to the Advancement of the Atomic Model

This discovery disproved the raisin-bun model and led to the introduction of the planetary model of the atom. In this model, electrons orbit the nucleus like planets orbiting the Sun and are kept in their orbit by centripetal force provided by the electrostatic attraction between the positive nucleus and the negative electrons. Advancements Toward the Quantum Mechanical Model The Atomic Model Dalton's Atomic Theory Quantum Model E. Rutherford Determinating the Charge-to-Mass Ratio of an Electron Millikan was able to determine the charge of an electron with his oil drop experiment and show that it was a fundamental unit of electrical charge. He used an atomizer to spray tiny drops of oil into the top of a closed vessel containing two parallel metal plates. Due to friction from the spraying process, some oil drops were given a small electrical charge. The oil drops were usually spherical, so he was able to use the diameter and density to calculate the mass each drop. Connecting a high voltage battery to the plates, Millikan observed the motion of the oil drops in the uniform electric field and taking into account the air resistance, Millikan calculated the electric force acting on each drop which allowed him to then determine the electric charge.

Millikan found that the charged oil drops had a charge of 1.60 x 10^-19 C or a multiple of this value and reasoned that this value was the smallest possible charge of an electron. Thus, the charge of a single electron is 1.60 x 10^ -19 C. The quantum mechanical model is based on quantum theory, the idea that matter has wave-like properties. Also according to this theory, it is impossible to know the exact location and momentum of an electron at the same time (Heisenberg's Uncertainly Principle). It uses complex orbitals or electron clouds to illustrate the volumes of space where there is likely an electron. The quantum mechanical model is based on probability rather than certainty. Bohr applied quantum theory to Rutherford's planetary atomic model.
Electrons orbit the nucleus in a set size and energy (energy of orbit is related to its size; lowest energy in smallest orbit). The distances are multiples of the radius of the smallest orbit. The orbits in an atom are quantized.
The kinetic energy and the electric potential energy of an electron orbiting a nucleus depends on thedistrance it is from the nucleus. energy of an electron in an atom is quantized and each orbit corresponds to a different energy level for that electron.
Energy is absorbed or emitted when an electron move from one orbit to another. Electrons travel in stationary orbits and do not radiate energy when they stay in a particular orbit. These orbits are called stationary states because the size and shape of the orbit remains constant along with the electron's energy. 1803 Advancements Limitations Dalton's atomic theory was based on the idea that atoms of different elements could be identified by their differences in weight. His theory proposed a number of basic ideas.
All matter is composed of atoms
Atoms cannot be created or destroyed
All atoms of the same element are identical
Different elements have different atoms
When atoms are rearranged, a chemical reaction occurs
Compounds are formed from atoms of different elements
Dalton's atomic theory became the theoretical foundation for modern atomic physics.Though much of his theory has been modified, the essence of it remains accurate. Solid Sphere Dalton was the first person to propose an atomic theory. It led to the evolution of the atomic model, each building upon the foundation Dalton had built with his theory. Being the first to propose an atomic theory, there were many errors that were identified later on.
Dalton made the mistake in assuming that simplest compound of two elements is formed from each element as a 1:1 ratio. (Ex. Believed water was formed from one hydrogen atom and one oxygen atom.) As a result, the atomic weight that was determined for elements was inaccurate.
Atoms can be destroyed through nuclear reactions, but cannot be destroyed through chemical reactions
There are isotopes of the same element, so atoms of the same element is not always identical. However, each has the same chemical properties such as charge.
Dalton idea of the structure of the atom was that it was an indivisible sphere that had no subatomic particles such as the electron, neutron, and proton. 1897 1909 1911 1913 1924 Through a series of cathode ray experiments, Thomson discovered the negatively charged particle of the atom, the electron. At that time, it was found that cathode rays could be deflected by a magnetic field, but no one had been able to deflect it with an electric field. He found that he was able to deflect the cathode ray with an electric field by using an extremely low pressure in the discharge tube. By observing the way the particle deflected, he was able to determine that the particle was negatively charged. 1904 The Raisin-bun Model Discovery of the Electron John Dalton J.J. Thomson Advancements Millikan made numerous momentous discoveries.
He was able to use Thomson's charge to mass ratio to determine the charge of an electron and show that it was a quantity constant for all electrons.
Showed that charge only occurred in discrete amounts (quantized) and was not a continuous quantity (parallels Planck's discovery that energy can be quantized).
Determining the charge allowed him to calculate a reasonably accurate value for the mass of the electron. It showed that the mass of the electron was roughly 1800 times less than the mass of hydrogen (the lightest atom), verifying Thomson's prediction. Limitations Thomson's theory still provided no understanding of how electrons in an atom were arranged. Also, Thomson's model of the atom attributed mass to be evenly distributed throughout the atom. This failed to explain what Rutherford observed in his gold foil experiments where some particles deflected when it struck the foil. Later, Rutherford determined that the mass was concentrated in the center of the atom while electrons orbited it.

The raisin-bun model could not explain radioactivity or the atomic spectra. Limitations Millikan stated that he had used all of his data to come to the conclusion that the charge of an electron was quantized. However, he had only reported 58 of the 175 measurements found in his notebook. When all of the data was used, his evidence was not as conclusive. It is still debated as to whether or not Millikan was guilty of scientific fraud or if he had a deep scientific insight that led him to select data that allowed him to decisively prove the quantization of charge. Advancements Advancements Bohr's atomic model describes most of today's accepted features for atomic theory. Stationary orbits overcame the problem of an atom's stability as electrons accelerated around the nucleus. His model explained the emission and absorption of the spectra and led to the possibility of calculating orbiting energy levels. It also explains the Rydberg Formula (emission of hydrogen radiation) by determining that the spectral lines corresponded to differences between quantized energy levels in the hydrogen atom. Limitations Bohr's theory has several serious flaws.
It violates the Heisenberg Uncertainty Principle because it states that electrons have a known radius and orbit.
The model is not accurate at predictin the spectra for larger atoms.
Provides an inaccurate value for the ground state orbital angular momentum.
Does not explain why energy is quantized or why orbiting electrons do not emit electromagnetic radiation.
Does not explain why a magnetic field splits the main spectral line into multiple closely spaced lines (Zeeman effect). Advancements Rutherford discovered the nucleus and was able to determine that most of the atom's mass is concentrated at its center. He disproved Thomson's raisin bun model and proposed a new model as a result of his discovery, the planetary model. Rutherford was also able to calculate the size of the nucleus by applying the law of conservation of energy and an equation for electric potential energy. The equation can be derived from Coulomb's law using basic calculus. Limitations Unlike planets, electrons are charged particles. According to Maxwell's theory of electromagnetic radiation, the accelerating motion of the electrons should emit electromagnetic waves which would take energy from the orbiting electrons. The loss of energy would cause the electrons to spiral toward the nucleus. However, evidence indicated that electrons did not spiral into their nucleus and the stability of the atom could not be explained with this model.

This model was also not able to explain how atoms only emitted light at certain wavelengths or frequencies. Advancements Today, this model is the most comprehensive and accurate model of atoms. The wave nature of particles provides a natural explanation for why electrons do not radiate electromagnetic energy continuously (quantized energy levels) and can be used to explain the observations of larger atoms. It follows the Heisenberg Uncertainty Principle. Limitations Though it is the most accurate model to date, this model does not predict the precise location of the electron, but rather the probability of its location. Many scientists such as Einstein and Schrodinger, had troubles accepting this theory due to its uncertain nature. This discovery led Thomson to develop a new model called the Raisin-bun Model. Since the nucleus had not yet been discovered, Thomson proposed that atoms were uniform spheres of positive matter in which electrons were embedded into by electrostatic force. Thomson built upon Dalton's theory and atomic model. He discovered a small piece of Dalton's atoms, the electron, a particle that was much less massive than the atom it was a part of and determined that all matter contained these negatively charged particles. His model was also able to explain the overall neutrality of the atom.

In his cathode ray experiments he was able to measure the deflection of the rays and determine the charge to mass ratio of the electron. Later, this charge to mass ratio allowed Robert Millikan to determine the elementary electric charge. The Quantization of the Electric Charge Bohr's Atomic Model The Planetary Model The Quantum Mechanical Model R. Millikan Niels Bohr Three physicists made significant contributions to the quantum mechanical model.
Lousis de Broglie - In 1924, he suggested that particles have wave properties.
Erwin Schrodinger - In 1926, he derived an equation for determining how electron waves behave in the electric field surrounding the nucleus and calculate electron energy levels. Defined electrons as circular standing waves that are multiples of whole numbers.
Werner Heisenberg - Heisenberg's Uncertainty Principle Bibliography + + + + + + _ _ _ _ _ _ x x x x x x x x x x x x x x Potential Difference e- Deflected: Once Thomson determined the velocity of the electron, he measured the deflection of the rays after removing the electric field. The magnetic force acting perpendicular to the electron caused uniform circular motion.
Fnet = Fm
Fc = Fm
mv^2/r = qvB
q/m = v/(Br) Undeflected: The electron passes through a potential difference undeflected when the electric force is equal to the magnetic force.
Fe = Fm
IEIq = qvB
Rearranging the formula to solve for the velocity of the electron, we get:
v = IEI/B Example: Find the charge to mass ratio for an ion that travels in an arc of radius 0.50 cm when moving at 1.0 x 10^6 m/s perpendicular to a 1.0 T magnetic field.
Fc = Fm
q/m = v/(Br)
q/m = (1.0 x 10^6 m/s) / (1.0 T x 0.005 m)
q/m = 2.0 x 10^8 C/kg Determination of the Elementary Charge + + + + + + _ _ _ _ _ _ Potential Difference O Accelerated Down: An oil droplet with a mass of 3.0 x 10^-16 kg accelerates downward at 6.8 m/s^2 in an electric field of 2.8 x 10^3 N/C. What is the charge?

q = m(g-a)/IEI
q = (3.0 x 10^-16 kg) (9.81 m/s^2 - 6.8m/s^2) / (2.8 x 10^3 N/C)
q = 3.2 x 10^-19 C Suspended Oil Drop: An oil droplet with a mass of 6.9 x 10^-15 kg is suspended motionless in a uniform electric field of 4.23 x 10^4 N/C [down]. What is the charge?

Fe = Fg
q = mg/IEI
q = (6.9 x 10^-15kg) (9.81 m/s^2) / (-4.23 x 10^4 N/C)
q = -1.6 x 10^-18 C Accelerated Upward: An oil droplet with a mass of 2.7 x 10^-16 kg accelerates upward at 3.2 m/s^2 between two horizontal plates 0.15 m apart. The potential difference is 4.5 x 10^3 V. What is the charge?

IEI = V/d
IEI= (4.5 x 10^3 V) / (0.15 m)
IEI = 3.0 x 10^4 V/m

q = m(g+a)/IEI
q = (2.7 x 10^-16kg) (9.81 m/s^2 - 3.2 m/s^2) / 3.0 x 10^4 V/m
q = 5.9 x 10^-20 C Oil Drops O O Fe = Fg
Suspended
IEIq = mg
q = mg/IEI Fe > Fg
Accelerated Up
q = m(g+a)/IEI Fg > Fe
Accelerated Down
q = m(g-a)/IEI Determining the Mass of the Electron Once Millikan determined the fundamental unit of electric charge, he was able to calculate the mass of the electron using the charge-to-mass ratio Thomson had determined earlier. q/m = v/Br
m = qBr/v In his gold-foil experiment, Rutherford fired alpha particles through a sheet of thin gold foil and used screens coated with zinc sulfide to measure the angles of deflection. Most of the of particles passed straight through the foil with little deflection. However, some were deflected at large angles or would just bounce straight back. The particles were not supposed to deflect at such wide angles by the repulsion of the positive charge in the cold atoms if the charge was distributed evenly according to Thomson's atomic model. This led Rutherford to believe that most of the gold foil was actually empty space and that the mass of the atom was contained at its center. The Atomic Spectra Continuous Spectrum An emission spectrum that exhibits all the wavelengths or frequencies of visible light. Usually emitted by a hot, dense material. Emission Spectrum A hot gas at low pressure will produce a pattern of bright lines. Heating a gas excites electrons, causing some to jump up to higher energy levels. These lines represent the frequencies of visible light that are emitted by the element. As the gas cools, some electrons move down to a lower energy level, emitting a photon in the process at one of the elements special frequencies. Absorption Spectrum When light passes through a gas at low pressure, a pattern of dark lines is produced. When photons from the light source interact with the atoms, the photons with the right frequency are absorbed by the electrons in order for them to jump to higher energy levels. The lines produced match the bright lines in the emission spectrum for the gas. Elements absorb the same frequencies they emit. The bright lines of the emission spectrum will correspond exactly to the dark lines of the absorption spectrum for a given substance.

The bright or dark lines produced is useful in identifying the substance using a spectrometer. Each element emits their own unique spectra. Bohr's Atomic Model and the Atomic Spectra The Bohr model explains why absorption and emission spectra occur for hydrogen and other elements and why electrons can only jump energy levels by emitting or absorbing energy in fixed quanta. An electron must gain energy if it jumps to a higher energy level by absorbing a photon. The photon's energy must match the difference between the electron's initial energy and the higher one. Since energy and frequency of the photon is related by the equation E = hf, the atom can only absorb frequencies that correspond to differences between the atom's energy levels. Conversely, if an electron jumps to a lower energy level, it must emit energy equal to the difference of the two energy levels. American Institute Physics. (2003). Robert Andrews Millikan. Retrieved May 21, 2013, from http://www.aip.org/history/gap/Millikan/Millikan.html
Bohr Atomic Model. (n.d.). Retrieved May 21, 2013, from 2013: http://abyss.uoregon.edu/~js/glossary/bohr_atom.html
Chemical Hertiage Foundation. (2010). Joseph John Thomson. Retrieved May 21, 2013, from Chemical Heritage Foundation: http://www.chemheritage.org/discover/online-resources/chemistry-in-history/themes/atomic-and-nuclear-structure/thomson.aspx
Dufferin-Peel Catholic District School Board. (2013). The Quantum Mechanical Model of the Atom. Retrieved May 26, 2013, from Dufferin-Peel Catholic District School Board: http://www.dpcdsb.org/NR/rdonlyres/9B6E4EF6-77F0-407F-A858-F0DF78B6896C/104314/33TheQuantumMechanicalModeloftheAtom.pdf
Encyclopædia Britannica, Inc. (2013). Sir J.J. Thomson. Retrieved May 21, 2013, from Encyclopædia Britannica: http://www.britannica.com/EBchecked/topic/593074/Sir-JJ-Thomson
John Wiley & Sons, Inc. (2013). Atomic Structure: The Quantum Mechanical Model. Retrieved May 26, 2013, from For Dummies: http://www.dummies.com/how-to/content/atomic-structure-the-quantum-mechanical-model.html
National Academy of Engineering. (2006, July 20). Case Study 2: The Millikan Case - Discrimination Versus Manipulation of Data . Retrieved May 21, 2013, from Online Ethics Center for Engineering and Research: http://www.onlineethics.org/cms/9726.aspx
Royal Society of Chemistry. (2013). Chemsoc Timeline. Retrieved May 21, 2013, from Dalton - Atomic Theory: http://www.rsc.org/chemsoc/timeline/pages/1803.html
Royal Society of Chemistry. (2013). Niels Bohr - Atomic Model. Retrieved May 21, 2013, from Chemsoc Timeline: http://www.rsc.org/chemsoc/timeline/pages/1913.html
Royal Society of Chemistry. (2013). Rutherford - Atomic Theory. Retrieved May 21, 2013, from Chemsoc Timeline: http://www.rsc.org/chemsoc/timeline/pages/1911.html
Thomson Atom. (n.d.). Retrieved May 21, 2013, from Thomson Atom: http://abyss.uoregon.edu/~js/21st_century_science/lectures/lec11.html
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