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Atom Models Timeline
Transcript of Atom Models Timeline
Problems/Limitations Experimented with other elements and found that elements always combine in the same ratio by weight in making a given compound (Law of Definite Proportions). Dalton performed a series of experiments on gaseous mixtures to determine what effect properties of the individual gases had on the properties of the mixture as a whole.
He theorized that the sizes of the particles making up different gases must be different.
Experiments 1 -John Dalton (1803) Dalton was the first to prepare a table of relative atomic weights and associate the ancient idea of atoms with stoichiometry. 1. All matter consists of tiny particles called atoms.
2. Atoms cannot be created, divided into smaller particles, nor destroyed in the chemical process; a chemical reaction simply changes the way atoms are grouped together.
3. Elements are characterized by the mass of their atoms and every element has a unique atomic weight.
4. When elements react, their atoms combine in simple, whole-number ratios to form chemical compounds.
The main concepts of Dalton’s Atomic Theory (chemical reactions can be explained by the union and separation of atoms, and that these atoms have characteristic properties) provided chemists with a basis for describing and explaining the behavior of matter during chemical reactions—foundation of modern physical science. Dalton’s Atomic Theory (1803): Contributions Atomic Theory was inadequate for explaining the behavior of substances:
5 - Niels Bohr (1913) Contributions Problems/Limitations Line Spectrum (atomic spectrum): Bohr's Atomic Model (1913):
So far, up to Rutherford’s findings, scientists could not explain line spectra of the hydrogen atom.
Bohr proposed a new model for the hydrogen atom. This model retained some of the features from Rutherford’s model. It explained the line spectrum for hydrogen because it incorporated several new ideas about energy. Visible light includes electromagnetic radiation of different wavelengths. Each element or compound emits energy when heated or put in a spectroscopic machine. The spectrum that is emitted for a specific element or compound is separated based on the wavelengths emitted. This is called a line spectrum. 1) An atom only has specific, allowable energy levels, called stationary states.
2) While in their stationary states, atoms do not emit energy.
3) An atom can change in stationary states by emitting or absorbing a specific quantity of energy that is equal to the difference between states.
Unlike Rutherford’s model, Bohr’s model imposed restrictions on moving electrons in a volume of space.
Bohr’s claims were made on the basis of Planck’s ideas of 1900 which stated matter can absorb or emit only discrete quantities of energy called quanta. Bohr applied the idea of quanta in his model. The change in energy when an electron moves from a higher to lower state of energy is not continuous.
The line spectra of hydrogen shows coloured bands that correspond with the change in energy of an electron as it moves to higher or lower energy level.
His realization that atom’s energy is quantized, and that electrons are restricted to specific energy levels successfully predicted lines in line spectra. His work also predicted UV and infrared portions in line spectrum.
Bohr’s theory only explains one-electron systems like H, He+, Li2+, Be3+ and cannot explain the emission spectra of atoms with 2 or more electrons. The model does not explain heavier elements well. cannot predict correctly the formulas for water, ammonia, and methane
atoms are not the most basic unit of matter because they are divisible
6 - Louis DeBroglie (1926/7) Contributions Experiments Problems/Limitations DeBroglie as a French scientist who hypothesized that matter had wave-like properties. His experiments of 1927 provided support to this theory. DeBroglie argued that if light which was normally a wave motion could take on a corpuscular (particle) form, then small particles such as electrons could also have wavelike characteristics associated with them.
Ultimately his findings helped scientists understand that the electrons did not move like a solar system around a central body in an atom because electrons do not have regular orbits. He developed an equation that enabled him to calculate the wave-length associated with any object of any size. The smaller an object was, the more significant the effect of its wavelength motion was. He established an idea that the wavelength of an electron moving at a speed of 5.9 x 106 m/s has a wavelength of 1 x 10-10 m. This happens to be greater than the wavelength of the hydrogen atom’s atomic radius.
His experiments of 1927 supported his hypothesis. Other researchers observed that moving electrons produced diffraction patterns to those of waves of electromagnetic radiation. Diffraction was the spreading of waves through a material and thus electrons (matter) were said to have wave-like properties. DeBroglie’s Hypothesis On Large Objects: valid only when the mass is small, and the velocity of the object is very large
large objects such as moving baseballs do not exhibit very clear wave-like movement because the wavelength is so small that it is undetectable. 7 - Erwin Schrödinger (1926/7) Contributions Problems/Limitations An Austrian physicist who viewed electrons as continuous clouds and introduced "wave mechanics" as a mathematical model of the atom. He used DeBroglie’s idea of matter moving in waves and Einstein’s idea of quantized energy particles as photons. His findings led to the start of a new field of science known as quantum mechanics which involved mathematic equations to describe the wave properties of sub-microscopic particles. He attempted to prove a model of the atom that involved particle waves which travelled in circles.
Schrodinger's view of the atom can be seen as "layers within layers" in terms of the electron shells. This means that each electron shell is made up of a number of subshells which are the same number in a shell is the same as the shell number (the first shell has only one subshell etc). These subshells can be further subdivided into orbitals, and each orbital is a distinct region of space that can contain a maximum of two electrons. Schrodinger Model: He could ultimately describe where electrons were in terms of probability. His model stated the importance of energy levels in orbitals rather than specific location of electrons in these orbitals. This wasn’t necessary a problem but a limitation of the theory in the complex field he was studying. 8 - Werner Heisenburg (1927) Contributions Experiments Problems/Limitations A German physicist by the name of Werner Heisenberg had a theory of his own called matrix mechanics which also explained the behavior of atoms, published in 1925. Compared Schrodinger’s theory, this theory worked on a basis of different assumptions. Both theories were said to “work”. Heisenberg based his theory on mathematical quantities called “matrices” that fit with the conception of electrons as particles whereas Schrödinger based his theory on waves. Both theories resulted in same mathematic results.
Visual concept of the atom now appeared as an electron "cloud" which surrounds a nucleus. The cloud consists of a probability distribution map which determines the most probable location of an electron. For example, if one could take a snap-shot of the location of the electron at different times and then superimpose all of the shots into one photo, then the electrons would look like a cloud. Heisenberg’s Uncertainty Principle: Around the year 1927, Heisenberg formulated an idea, which agreed with his Uncertainty Principle tests: that no experiment can measure the position and momentum of a quantum particle simultaneously. Scientists call this the "Heisenberg uncertainty principle." Quantum Mechanical Model 9 - James Chadwick (1932) Matrix Mechanics: Heisenberg argued for the uncertainty principle by using an imaginary microscope as a measuring device. He stated that an experimenter would attempt to measure the position and momentum of an electron in an atom by shooting a photon at it. However, this imaginary microscope experiment poses problems for describing position of the said electron:
Problem 1 - If the photon has a short wavelength, and therefore a large momentum, the position can be measured accurately. But the photon scatters in a random direction, transferring a large and uncertain amount of momentum to the electron according to the law of conservation of momentum.
Problem 2 - If the photon has a long wavelength and low momentum, the collision does not disturb the electron's momentum very much, but the scattering will reveal its position only vaguely.
Ultimately, independent of the wavelength of the photon, the result of any experiment to find the position of an electron would not give a certain result. Each orbital has its own associated energy, and each represents information about where, inside the atom, the electrons would spend most of their time. Scientists cannot determine the actual paths of moving electrons. The limitations of Heisenberg’s findings are that it allowed orbitals to simply indicate where there is a high probability of finding electrons. Contributions Experiments Many of the methods used by the scientists who contributed to the discovery of the understanding of an atom’s model used experiments to detect charged particles. It was a unique process used by scientist James Chadwick which ultimately led to his discovery of a neutron as an subatomic particle. Using alpha particles, he discovered a neutral atomic particle with a mass close to a proton. He found it to measure slightly heavier than the proton with a mass of 1840 electrons and with no charge (neutral). The proton-neutron together, received the name, "nucleon."
Although scientists knew that atoms of a particular element have the same number of protons, they discovered that some of these atoms have slightly different masses. Chadwick and others concluded that the variations in mass result, more or less, from the number of neutrons in the nucleus of the atom. Atoms of an element having the same atomic number but different atomic masses get called "isotopes" of that element. A sample of the element beryllium, which was bombarded with alpha particles (another type of naturally occurring radiation) which are technically just ionized helium nuclei. Beryllium began to emit mysterious radiation. Other scientists had previously discovered that this radiation, after striking protons, would discharge some of the protons, which could then be detected using a Geiger counter (a device that measures radiation).
Chadwick’s experiments and hypothesis was based on this observation. He determined that the “mysterious” radiation observed was neutral. This was because it was not affected by proximity to any magnetic fields. Unlike standard gamma radiation, did not invoke the photoelectric effect. The photoelectric effect was when photons, such as gamma rays, strike certain surfaces; they discharge electrons, which can be simply measured. In this case, protons were discharged instead. Chadwick ultimately concluded and discovered the neutron in 1932. Contributions Worked with cathode rays and stated that the cathode rays are negatively charged and that they are composed of identical negatively charged particles, thus discovering the existence of electrons using a cathode ray tube.
Discovered an electron’s mass is less than 1/1837 of a hydrogen atom and electron has a very high charge despite its tiny mass. Plum Pudding Model (1904):
Experiments Problems/Limitations Plum Pudding Model:
cannot explain radioactivity (which was discovered in the coming years)
could not predict why atoms absorbed and emitted spectral lines
atoms have nuclei 3 - Joseph John Thomson (1904) suggested that an atom is composed of electrons surrounded by a cloud of positive charge to balance the electron's negative charge.
the electrons are stationed all around the atom (he saw the atom as a positively charged sphere embedded with sufficient numbers of electrons to balance the total charge)
made other scientists realize that atoms must also contain a positive charge to balance the negative charge (which later led to the discovery of protons)
contradicted Dalton’s idea of an indivisible atom Experiment 1:
built a cathode ray tube to measure electrical charge
Thomson knew from Jean Perrin’s experiment in 1895 that cathode rays deposited an electric charge, and he wanted to see if he can separate the charge from the rays be bending them with a magnet
found that the negative charge and the cathode rays must somehow be stuck together: you cannot separate the charge from the rays (cannot bend the cathode rays)
Thomson suspected that the traces of gas remaining in the tube were acting like electrical conductors inhibiting the cathode rays from bending.
To test this idea, he extracted nearly all of the gas from a tube, and found that now the cathode rays did bend in an electric field
Thomson tried to determine the basic properties of the particles
using the measurements Thomson obtained of how much the rays were bent by a magnetic field and how much energy they carried, he calculated the ratio of the mass of a particle to its electric charge (m/e)
collected data using a variety of tubes and using different gases 2 - Max Planck (1900) Contributions proposed that energy is radiated in tiny and discrete quantized amounts (or packets), rather than in a continuous wave
“packets” of energy called quanta (quantum – singular)
revolutionized scientists’ understanding of atomic and subatomic processes
a fundamental theory of 20th-century physics that led to industrial and military applications which affect every aspect of modern life
Planck was able to determine that the energy of each quantum is equal to the frequency of the radiation multiplied by a universal constant
this number (expressed in erg-seconds—the amount of energy needed to raise a milligram of mass by a distance of 1 centimeter) measures the energy of an individual quantum
Planck's constant (h) is approximately 6.63 x 10-27 erg-second has become one of the basic constants of physics
used to describe the behavior of particles and waves at the atomic scale
forms the basis of quantum theory
describes how energy dissipates (in terms of different wavelengths of radiation) from an ideal non-reflective black object (black body)
states that the energy of an electromagnetic waves is contained in indivisible quanta that have to be radiated or absorbed as a whole
the magnitude of energy is proportional to the frequency where the constant of proportionality is given by Planck's constant Quantum Theory (1900): Planck’s Constant - 1900: Planck’s Law – 1900: Experiments Black Body Experiments A black body does not reflect any light (it absorbs most of the light that is directed its way)
He worked on the problem of how the radiation an object emits is related to its temperature
The aim of the experiment was to test thermal radiation:
apparatus was set up to detect the radiation from an object maintained at a temperature (warm body gives off radiation in all directions)
performed the experiment for several different temperatures, and obtained a range of radiation vs. wavelength curves, which yield significant results:
1. The total intensity radiated over all wavelengths increases as the temperature increases
2. The value of the wavelength at which the radiation reaches its maximum decreases as the temperature increases Quantum Theory Problems/Limitations Planck’s Quantum Theory is a paradox of the “Particle/Wave” duality of matter
when objects are very close together, the predictions of Planck’s Law would break down because it’s hard to prevent the two objects from touching (therefore only valid for large systems)
In experiments, there is extreme difficulty in measuring temperature differences over very small distances 4 - Ernest Rutherford (1911)
the model states that the atom consists of a dense positive small and heavy center called the nucleus and is orbited by negatively charged electrons (like a planetary model)
most of the atom is made up of empty space
the nucleus contains neutrons and protons. These particles are all collectively called nucleons
nucleus is 10-5 times smaller than the total size of atom
similar to solar system Nuclear Model of Atom (Planetary Model) - 1911 Nuclear Atom/Planetary model - the electrons are in motion around the atomic nucleus. Experiments Nuclear Scattering Experiment- 1909 Rutherford carried out this experiment by shooting high speed positively charged αalpha-particles emitted from Ra at a thin (10-4 mm) Au foil, and came up with the following observations:
most of the αalpha-particles passed without being deflected
some of them were deflected away from their path
only a few (one in about 10,000) were returned back to their original direction of emission
From the observations he concluded that an atom consists of:
a nucleus which is small in size but carries the entire mass (contains all the neutrons and protons)
an “extra nuclear part” which contains the electrons Gold Foil Experiment – 1911 led Rutherford to discover that the atom is made up of mainly empty space with a small massive region of concentrated charge at the centre Indivisible, Solid Sphere Model
Plum Pudding Model Planetary/Nuclear Model if the atomic nucleus was composed entirely of positive charges, it should fly apart because like charges repel
cannot explain the total mass of an atom
contradicts the laws of 19th century physics (an electron in motion around a central body must continuously give off radiation) meaning it gives off a continuous spectrum (rainbow) of light energy (Maxwell’s theory of electrodynamics) Problems/Limitations The End...for now. Works Cited "Dalton's Atomic Theory." MASTERING CHEMISTRY......... Web. 06 Apr. 2011. <http://chemistryreaders.blogspot.com/2010/04/daltons-atomic-theory.html>.
"Ernest Rutherford - Biography." Nobelprize.org. Web. 06 Apr. 2011. <http://nobelprize.org/nobel_prizes/chemistry/laureates/1908/rutherford-bio.html>.
"General Chemistry Online: Companion Notes: Atoms & Ions: Dalton's Atomic Theory." Dalton's Atomic Theory. Web. 06 Apr. 2011. <http://antoine.frostburg.edu/chem/senese/101/atoms/dalton.shtml>.
"J. J. Thomson's Cathode Ray Experiment." The Scientific Method, Science, Research and Experiments. Web. 06 Apr. 2011. <http://www.experiment-resources.com/cathode-ray.html>.
"James Chadwick." The History of Computing Project. Web. 06 Apr. 2011. <http://www.thocp.net/biographies/chadwick_james.htm>.
"Max Planck - Biography." Nobelprize.org. Web. 06 Apr. 2011. <http://nobelprize.org/nobel_prizes/physics/laureates/1918/planck-bio.html>.
"Niels Bohr - Biography." Nobelprize.org. Web. 06 Apr. 2011. <http://nobelprize.org/nobel_prizes/physics/laureates/1922/bohr-bio.html>.