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Transcript
ATOMIC AND NUCLEAR PHYSICS Topic 7.1 The Atom Courtesy: https://sites.google.com/site/ndhssciencerevisionstorage/ib-physics-summarypowerponts Physics for the IB Diploma by Tsokos ATOMIC STRUCTURE_ DALTON’S MODEL John Dalton said that atoms were tiny indivisible spheres, but in 1897 J. J. Thomson discovered that all matter contains tiny negatively-charged particles. He showed that these particles are smaller than an atom. He had found the first subatomic particle - the electron. ATOMIC STRUCTURE_ THOMSON’S MODEL Scientists then set out to find the structure of the atom. Thomson thought that the atom was a positive sphere of matter and the negative electrons were embedded in it as shown here This `model' was called the `plum-pudding' model of the atom. RUTHERFORD’S EXPERIMENT Ernst Rutherford decided to probe the atom using fast moving alpha (α) particles. He got his students Geiger and Marsden to fire the positively-charged α-particles at very thin gold foil and observe how they were scattered. Simulation RUTHERFORD’S EXPERIMENT_ FINDINGS Most of the α-particles passed straight through the foil, but to his surprise a few were scattered back towards the source at very large scattering angles. Rutherford said that this was rather like firing a gun at tissue paper and finding that some bullets bounce back towards you! INTERPRETING THE RESULTS The very large deflection was indicative of a large repulsive force between the alpha positive particles and the carrier of the positive charge in the atom. Such force could not be produced if the positive charge was distributed over the entire atomic volume (Thomson’s model). Rutherford concluded that the positive charge in the atom must be very tiny so that the alpha particles could come very close, and massive; otherwise, the atom would have recoiled. Thus, the alpha particles can approach the positive charge at very small distance and the electric repulsive force, being proportional to the inverse of the square of the distance, would be very large This explains the large deflection. Most α-particles are hardly deflected because they are far away from the nucleus and the field is too weak to repel them much. The electrons do not deflect the α-particles because the effect of their negative charge is spread thinly throughout the atom. RUTHERFORD’S MODEL OF THE ATOM_NUCLEAR MODEL Rutherford nuclear model of the atom: All of an atom's positive charge and most of its mass is concentrated in a tiny core. The electrons orbit the nucleus, like planets orbit the sun Rutherford called this the nucleus. Planetary model The atom is mainly empty space! Using this model Rutherford calculated that the diameter of the gold nucleus could not be larger than 10-14 m. This diagram is not to scale. With a 1 mm diameter nucleus the diameter of the atom would have to be 10000 mm or 10 m! The nucleus is like a pea at the centre of a football pitch. DIFFICULTIES WITH THE RUTHERFORD’S MODEL Atoms emit certain discrete characteristic frequencies of electromagnetic radiation The Rutherford model is unable to explain this phenomena According to Maxwell’s theory of electromagnetism, an accelerated charge should radiate electromagnetic waves and thus lose energy. The electrons in Rutherford’s model move in circular orbits and are subject to centripetal acceleration. Accordingly, they must radiate and lose energy, and then spiral into the nucleus. This is not the case! BOHR’S MODEL OF THE ATOM Bohr was the first to address the problems with the Rutherford Model. He benefitted from earlier studies done on the emission spectrum of gases. CONTINUOS SPECTRUM OF WHITE LIGHT The light coming from the sun, from a candle, or from an incandescent lamp is called white light because it contains all the wavelengths of the visible light. If you pass this white light through a prism or a through a spectroscope, then the white light is broken down into all the visible colors. We obtain a continuous spectrum LIGHT SPECTRA Not all spectra of light are continuous. If you pass the light coming from a Fluorescent lamp through a prism, the obtained spectrum is not continuous. There are three types of spectra: Continuous spectrum Emission spectrum Absorption spectrum EMISSION SPECTRA OF GASES Thomas Melville was the first to study the light emitted by various gases. He used a flame as a heat source, and passed the light emitted by the heated gas through a prism. Melville discovered that the pattern produced by light from heated gases is very different from the continuous rainbow pattern produced when sunlight passes through a prism. EMISSION SPECTRA OR LINE SPECTRA The new type of spectrum consisted of a series of bright lines separated by dark gaps. This spectrum became known as a line spectrum. Melville also noted the line spectrum produced by a particular gas was always the same. The line spectra are produced from atoms of gases that have been excited in some way, either by heating or by an electrical discharge. ABSORPTION SPECTRA Suppose that you pass white light (containing all visible light wavelengths) through a cool gas element. If you analyze the light that passes through the cool gas, you will notice that the spectrum of the light is not continuous. It is missing some wavelengths. This spectrum is called absorption spectrum because the cool gas absorbed these wavelengths. EMISSION VERSUS ABSORPTION SPECTRA The dark lines on an absorption spectrum will fall in exactly the same position as the bright lines on an emission spectrum for a given element. For example, the emission spectrum of sodium shows a pair of characteristic bright lines in the yellow region of the visible spectrum. An absorption spectrum will show 2 dark lines in the same position. Emission Spectra Absorption Spectra BOHR’S MODEL OF HYDROGEN In 1913 Bohr provided an explanation of atomic spectra that includes some features of the currently accepted theory His model includes both classical and non-classical ideas His model included an attempt to explain why the atom was stable THE BOHR THEORY OF HYDROGEN In 1913 Bohr provided an explanation of atomic spectra that includes some features of the currently accepted theory His model includes both classical and nonclassical ideas His model included an attempt to explain why the atom was stable BOHR’S ASSUMPTIONS FOR HYDROGEN The electron moves in circular orbits around the proton under the influence of the Coulomb force of attraction The Coulomb force produces the centripetal acceleration Only certain electron orbits are stable These are the orbits in which the atom does not emit energy in the form of electromagnetic radiation Therefore, the energy of the atom remains constant and classical mechanics can be used to describe the electron’s motion BOHR’S MODEL OF HYDROGEN The hydrogen single electron can only exist in specific orbits of specific energy (Potential energy). These orbits are called energy levels. The electron energy is discrete (quantum) and not continuous The electron has its lowest energy state when it is in the ground level (least PE). When the electron is in an energy level other than the ground level, it is said to be in an excited state of energy BOHR’S MODEL OF HYDROGEN The electron can lose energy only when it jumps from one energy state level to a lower one. The emitted energy is equal to the difference in energy between the initial and final states. Similarly, an electron can jump from one energy state level to a higher one only if it receives an amount of energy equal to the difference in energy between the final and initial states. In this sense, Bohr’s model provided an explanation for the emission and absorption spectra NUCLEAR STRUCTURE MASS NUMBER The total number of protons and neutrons in the nucleus is called the mass number (or nucleon number). NUCLEON Protons and neutrons are called nucleons. Each is about 1800 times more massive than an electron, so virtually all of an atom's mass is in its nucleus. ATOMIC NUMBER All materials are made from about 100 basic substances called elements. An atom is the smallest `piece' of an element you can have. Each element has a different number of protons in its atoms: it has a different atomic number (sometimes called the proton number). The atomic number also tells you the number of electrons in the atom. ISOTOPES Every atom of oxygen has a proton number of 8. That is, it has 8 protons (and so 8 electrons to make it a neutral atom). Most oxygen atoms have a nucleon number of 16. This means that these atoms also have 8 neutrons. This is 168O. Some oxygen atoms have a nucleon number of 17. These atoms have 9 neutrons (but still 8 protons). This is 178O. 168O and 178O are both oxygen atoms. They are called isotopes of oxygen. There is a third isotope of oxygen 188O. How many neutrons are there in the nucleus of an 188O atom? Isotopes are atoms with the same proton number, but different nucleon numbers. Since the isotopes of an element have the same number, of electrons, they must have the same chemical properties. The atoms have different masses, however, and so their physical properties are different. EVIDENCE FOR NEUTRONS The existence of isotopes is evidence for the existence of neutrons because there is no other way to explain the mass difference of two isotopes of the same element. By definition, two isotopes of the same element must have the same number of protons, which means the mass attributed to those protons must be the same. Therefore, there must be some other particle that accounts for the difference in mass, and that particle is the neutron. INTERACTIONS IN THE NUCLEUS Electrons are held in orbit by the force of attraction between opposite charges. Protons and neutrons (nucleons) are bound tightly together in the nucleus by a different kind of force, called the strong, short-range nuclear force. There are also Coulomb interaction between protons. Due to the fact that they are charged particles.