Lecture D Atomic Structure PDF

Fundamentals of Radiation & Radiation Safety Lecture D Topic: Fundamentals of Radiation Subject: Atomic Physics Lecturer: Ciara Mc Nally Read Chapter 2 ‘Atoms and matter’ Graham's Principles and Applications of Radiologi...

Fundamentals of Radiation & Radiation Safety Lecture D Topic: Fundamentals of Radiation Subject: Atomic Physics Lecturer: Ciara Mc Nally Read Chapter 2 ‘Atoms and matter’ Graham's Principles and Applications of Radiological Physics, 7th Edition  Any attempt to understand the universe around us must start with the fundamental question – What is matter made of ? What is Matter Made of? The atom as the fundamental building block of matter has been the subject of a great deal of both theoretical debate and experimental study by physicists Many of the modern theories concerning atomic and subatomic structures are extremely complex. However most of the phenomena encountered in radiography can be explained using a relatively simple planetary model of the atom The History of Atomic Theory 400 BC – Concept of Particles Atoms The word atom comes from a Greek word meaning something which cannot be split John Dalton 1766 - 1844 Following his experimental work on gasses, Dalton developed the first Atomic Theory of Matter - Elements were composed of tiny indestructible particles called atoms http://www.slcc.edu/schools/ Each elements atoms hum_sci/physics/whatis/ were all identical and biography/dalton.html had the same atomic weight John Joseph Thomson 1856 - 1940 Announced his discovery of the electron at his evening lecture at the Royal Institute on the 30th April 1897 Following research on cathode rays  Eugen Goldstein 1850 - 1930  Experimented with a cathode ray tube in which the cathode was perforated.  He produced so called canal rays which travelled in the opposite direction to the cathode rays. He concluded they must have the opposite charge. Ernest Rutherford 1871 - 1937  In 1911 Rutherford concluded that the atom was composed of a tiny positively charged nucleus, where nearly all of its mass was concentrated, with the negatively charged electrons some distance away.  An atom is mostly empty space!  Rutherford did not explain why the electrons remained separated from the nucleus. Niels Bohr 1884 - 1962 Bohr worked with Rutherford between 1911 and 1913. Before publishing his atomic theory. In the Bohr model: electrons orbit the nucleus at discrete energy levels, often called shells or orbits Neutrons In 1932 James Chadwick discovered a third type of sub-atomic particle which he named the neutron. Neutrons, which have no charge, help to reduce the repulsion between protons within the nucleus. Atomic Structure - Models Planetary model describes a Quantum model describes central nucleus with orbiting electrons behaving as particles electrons and waves The planetary model of the atom Subatomic particle properties The atomic nucleus is a very small and dense structure in the centre of the atom and contains protons and neutrons. Nuclear sizes vary between 10 -15 and 10-14 metres. The diameter of the outermost shell of atoms varies between 1 x 10-10 and 3 x 10-10 metres. See TABLE 2.1 in Graham's Principles and Applications of Radiological Physics, 7th Edition Subatomic particle properties Rest mass Electric Particle Symbol Rest mass energy charge (MeV) 1.672 x 10-27 kg Proton p 938 +1 1.007 amu 1.675 x 10-27 kg Neutron n 939 0 1.009 amu 9.109 x 10-31 kg Electron e 0.511 -1 0.00055 amu The atomic nucleus  The number of protons and neutrons in the atomic nucleus determines the mass and the charge of the nucleus and the configuration of electron orbitals of the atom. E = element name A Z E A = atomic mass number Z = atomic number A - Atomic mass See TABLE 2.2 in Z - Atomic number is the Graham's number is the number of Principles and number of protons and Applications of protons in the neutron in the Radiological nucleus. nucleus Physics, 7th Edition The atomic nucleus Z - Atomic number is the number of protons in the nucleus A - Atomic mass number is the number of protons and neutron in the nucleus How do you know the number of electrons in an atom? Nuclide’s/Isotopes A nuclide is an atom, or a number of atoms of an element, characterised by a specific number of neutrons in the nucleus Example Each of the following nuclides are isotopes of carbon A 12 14 Z E 6 C 13 6 C 6 C All of the nuclides have the same atomic number (atomic number = number of protons) Nuclide’s/Isotopes 12 6 C 13 6 C 14 6 C All of the nuclides have the same atomic number number of protons Nuclide’s/Isotopes 6 protons 6 protons 6 protons + + + 6 neutrons 7 neutrons 8 neutrons 12 6 C 13 6 C 14 6 C All of the nuclides have the same atomic number (atomic number = number of protons) The stability of the nucleus  The atomic nucleus contains protons which have a positive charge and neutrons which have no charge.  In atoms other than hydrogen the nucleus should blow apart because of the electrostatic forces between the protons. The stability of the nucleus  In reality many nuclei are very stable whilst others readily decay.  This is the result of the opposing electrostatic and short range nuclear forces.  The energy expended in keeping the nucleus together is called the nuclear binding energy. Electron orbitals Electrons are found in the orbital In a stable atom, the number of Therefore all atoms are shells. protons in the nucleus is equal to electrically neutral. the number of electrons orbiting the nucleus. Electron orbitals  Electrons orbit the nucleus of the atom in stable paths  The electron orbits are grouped into ‘shells’ where there is a particular number of electrons of approximately the same energy in each orbit.  Electrons fill up the inner shells first as this is the lowest energy state and therefore most stable situation. The planetary model of the atom Xwww.britannica.com/nobel/micro/514_59.html The Quantum Mechanics Perspective  Mathematical equation called the Wave Function describes ‘electron probability density’ for different atoms  Describes the liklihood of locating an electron in space at different energy levels Electron orbitals CANNOT PREDICT PRECISE THE ‘PROBABILITY’ OF FINDING AN EACH LOCATION HAS A SPECIFIC SPATIAL LOCATION OF AN ELECTRON AT A GIVEN SPATIAL ENERGY. THINK OF ELECTRONS AS ELECTRON LOCATION IS DETERMINED BY THE BEING IN QUANTISED ENERGY MATHEMATICAL PROBABILITY LEVELS RATHER THAN SHELLS. DENSITY. Electronic Energy Levels What determines electron configuration? n=3 n=2  An electron energy level cannot contain n=1 more than 2n2 electrons  The outer electronic energy level cannot contain more than eight electrons A sodium atom  Numbers of electrons in atomic shells Shell Shell letter Maximum number (n) number of electrons 1 K 2 2 L 8 3 M 18 4 N 32 Number of electrons equals 2n2, where n is the shell number See TABLE 2.3 in Graham's Principles and Applications of Radiological Physics, 7th Edition Electron binding energy n=3  Electrons can either be bound to an atom or free.  The electron binding energy is n=2 the amount of work which must be done to remove an electron n=1 from its energy level.  Electrons in the innermost level have the highest electron binding energy  The actual binding energy of electrons is is expressed in electron volts (eV) or keV (1keV = 1000 eV)  1eV = 1.6022 x 10-19 joules Electron binding energy Increase in the atomic number = increase in the binding energy of the electrons (there are more protons and, therefore, more energy is needed to release the electrons from the greater positive pull). Increase in the distance between the nucleus and the electron = decrease in the binding energy of the electron (decrease in the positive pull of the protons in the nucleus). Pauli’s exclusion principle states that no two electrons can occupy the same quantum state at the same time. Dmitre Mendeleev 1834 - 1907 www.aip.org/history/curie/periodic.htm The Periodic Table When elements arranged in order of increasing atomic number, their valency and physical properties occur in a periodic manner, as determined by electronic configuration. Electron Excitation  Outer orbitals are still present for smaller elements but are normally unoccupied.  An electron can absorb an amount of energy that is less than its binding energy but is sufficient to promote the electron to a higher energy level. This process is called excitation.  Atoms in an excited state are unstable and quickly return to the ground state (lowest energy state) with the emission of the excess energy in the form of electromagnetic radiation Ionisation If an electron absorbs sufficient energy to overcome the binding energy, then the electron can escape from the atom leaving behind a charge imbalance. This process is called ionisation.  When atoms are ionised through the removal of an electron they are left with a net positive charge.  Atoms with fewer (or more) electrons are called ions. Ionisation Ionic bonding The chemical properties of an element are controlled by the electron configuration of its atoms. Atoms with a filled outer electron level are chemically inert. Atoms with an almost full outer electron level are chemically active electron acceptors Atoms with an almost empty outer electron level are chemically active electron donors Donation and acceptance of electrons will produce positive and negative ions Sodium Chlorine has an has an atomic atomic number of number of 11 17 Sodium donates an electron to the chlorine atom + - The two ions are attracted to each other and an ionic bond is formed to produce Sodium Chloride  Molecular Compounds Atoms of more than one element are chemically bound in fixed ratios to make compounds O2 CH4 Oxygen Methane  Covalent bonding Atoms with half full outer electron levels satisfy their need for a full outer level by sharing electrons rather than donating or accepting them. By sharing electrons a covalent bond is formed. O2 CH4 Oxygen Methane Other Types of Matter Crystalline compounds consist of atoms held together in an extended array or ‘lattice’ of atoms In medical imaging caesium iodide (CsI) is a crystalline material used in x-ray detectors Single element materials such as carbon and silicon form extensive covalently bound structures The structure dictates the material property, for example carbon can exist as diamond, graphite and graphene  Other Types of Matter Inert gases or ‘noble’ gases are monoatomic and chemically unreactive due to a filled outer electron shell Helium is a noble gas which in its liquid form is about -269 C. It is used in MRI scanners to maintain superconductivity Metals are lustrous, malleable and electrically conductive. The structures are held together by metallic bonds attributed to electrostatic attractions between positively charged metal ions and delocalised electrons  States of Matter The three common states of matter are gases, liquids and solids. Matter transforms in state according to temperature Gases have no intermolecular bonds, liquids have weak intermolecular bonds and solids have strong intermolecular bonds Low density materials are Image Pixel Brightness represented as black and high density materials are represented as white on radiographs Bodily tissues are varying shades of grey depending on density and thickness Summary