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1 No. Atomic Theory / Model Development of Atomic Theory Phenomena that Diagram illustrating the Can’t Be Explained by Old Atomic Theory / Model Theory / Model 1 (i) 400 B.C. Democritus (A Greek Philosopher (ii) 1807 – John Dalton Atomic Theory Solid-Particle Model (1805) Democritus coined the word “Atom”. Dalton’s Atomic Theory: 1. Atoms are the basic particles of matter. 2. They are indivisible and cannot be created or destroyed. 3. Atoms of a given element are identical, having the same mass and the same chemical properties. 4. Atoms of different elements combine with one another in simple whole number ratios to form molecules of compounds. 5. In chemical reactions, atoms are combined, separated, or re-arranged. Phenomena or Laws Explained by Current Theory / Model Using deductive logic, the Greeks hypothesized that matter cut into smaller and smaller pieces would eventually reach what they called the “atom”, the word coined by Democritus literally meaning indivisible. (i). Law of Conservation of Matter (Antoine Lavoisier) (ii). Law of Constant Composition Static electricity was observed. Electron was discovered, providing the crucial clue that atom is not indivisible, but have a sub-structure. (Read page 7 of this notes for the details of J.J. Thomson’s and Robert Milliken’s Experiments and their discovery of the e/m ratio and the mass of an electron (i). Charge-to-Mass (1799-Joseph Proust) (iii). Law of Multiple Proportions (John Dalton) Billiard Ball Model 2 1800’s (i).1897-J.J Thomson: Experiments with Cathode Ray tubes (ii). 1909Robert Milliken: Oil Drop Experiment Thomson Atomic Theory (1897): Plum Pudding Model – Negatively charged electrons are imbedded in a large positively charged atom, just like plums embedded in the famous English dessert, pudding. Atoms may gain or lose electrons to form ions in solution. Particular atoms and ions gain or lose a specific number of electrons. Electricity is composed of negatively charged particles. Electrons are a component of all matter. Electrons have a specific fixed electric charge. ratio was found by experiment: e / m = −176 . ×108 C / g (ii). Charge & Mass of Electron were found by experiment: (a) e = −160 . × 10 −19 C (b) m = 911 . × 10 −28 g Raison Cookie Model 3 1910 Ernest Rutherford: Gold Foil Experiment (See page 172 of this notes for details of this experiment) 4 1913 Niels Bohr: Theorized Atomic Spectra of Hydrogen (Bohr Atomic Theory) 5 1927 Erwin Schrödinger : Rutherford Atomic Theory (1911): Nuclear Model – An atom is made of an equal number of negatively charged electrons and positively charged protons. Electrons move in orbits around a very dense and positively charged nucleus. The proton in the nucleus carries the positive charge. A very strong nuclear force holds the positive charge within the nucleus. Most of the atom is empty space Planetary / Nuclear Model Radioactivity was first observed and proton was discovered: While most of the +vely charged alpha particles went through the foil with no visible effect, a few were deflected from their path and some actually bounced back in the direction from which they came, suggesting a dense nucleus with a +ve charge. Atom is mostly empty space with an extremely dense, +vely charged nucleus. The +ve charge is due to the proton located in the nucleus. (The neutral particle, neutron, also in the nucleus was found by James Chadwick later in 1937) Solar System Model – The nucleus According to the Rutherford Model, an orbiting electron should continuously emit electromagnetic radiation, lose energy, and collapse the atom. But it didn’t. Atomic Spectra of discrete wavelengths and frequencies were observed, that couldn’t be explained by Thomson’s Atomic Theory. In 1885, J. Balmer found an empirical mathematical relationship between the wavelengths of the lines in the visible region of the spectrum. Johannes Rydberg extended Electrons do not radiate energy as they orbit the nucleus. Each orbit corresponds to a state of constant energy based on Planck’s Theory of quantized energy. The Model was able to derive Rydberg’s constant in calculating the wave numbers of the spectral lines and to successfully predict the infrared and ultraviolet spectra for hydrogen. It offered a reasonable explanation for Mendeleev’s Periodic Law. lies in the center of the atom and the electrons move in certain “allowed” orbits around it. These circular orbits represent the quantized energy levels of the electrons. The orbit closest to the nucleus represents the ground state of the energy level designated by a quantum number n =1. As the electron absorbs a quantum of energy, it jumps to a higher level, excited state. When the electron jumps back down to the n =1 level, it is accompanied by an emission of radiation. There are a maximum number of electrons allowed in each orbit. Wave-Particle Model–An electron is regarded as a wave that has quantized its energy. Electrons have no precise orbits. Instead, their motion can only be described by the probability of finding them in certain regions surrounding the nucleus. These regions are called orbitals. These orbitals form an electron cloud around the nucleus. The cloud is most dense where the probability of finding the electron is highest: electron density. According to the Uncertainty Principle, the position and the momentum of an electron cannot both be known exactly at the same time. Solar System/Orbit Model his Equation so that all of the wavelengths could be predicted. It couldn't explain why the protons in the nucleus don't fly apart and it couldn't totally account for the total mass of the nucleus. In Bohr’s model, there is no way to observe or to measure the orbit of an electron in an atom. The Bohr’s model was unable to explain completely (i) the Zeeman Effect (ii) the spectral details of the atoms that have several electrons. Electron Cloud / Quantum Mechanical Model (i) Heisenberg’s Uncertainty Principle (1927): ( ∆x )( ∆mv) ≥ h 4π (ii) In 1924 Louis De Broglie suggested electron could behave as a wave as well as a particle: hf = mc 2 2 3 Figure 2 The solid lines show the intensity of the colours of light emitted Figure 3 Scientists used to think that as the intensity or by a red-hot wire and a white-hot wire. Notice how the curve becomes brightness of light changes, the total energy increases higher and shifts toward the higher-energy UV as the temperature increases. continuously, like going up the slope of a smooth hill. The dotted line represents the predicted curve for a white-hot object, As a consequence of Planck’s work, Einstein suggested according to the existing classical theory before Planck. that the slope is actually a staircase with tiny steps, where each step is a quantum of energy. 4 Figure 6 In the photoelectric effect, light shining on a metal liberates electrons from the metal surface. The ammeter (A) records the electric current (the number of electrons per second) in the circuit. 5 Figure 2 The original Bohr orbit for n = 2 is circular (red line). Sommerfeld’s revision is three, slightly elliptical orbits for n = 2 (blue lines). [The secondary quantum number relates primarily to the shape of the electron orbit. The number of values for l equals the volume of the principal quantum number.]