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BERLIN TOWNSHIP SCHOOL DISTRICT VISION UPDATE: DECEMBER 2011 1 UPDATE EVALUATION TOOL • The administration is looking at the McCrell assessment tool as an initial step in this process • While other options may be available, McCrell offers the easiest usage at an appropriate cost • The state may be pushing this initiative back one year due to the 9 pilot districts getting up and running late • The Council on Instruction will look at options • An assessment tool will be chosen 2/2012 to be used 9/2012 if the state mandate is still required by 9/2012 USE DATA TO ASSESS INSTRUCTION
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Atomic Nature of Matter
Course: NUCLEAR PHYSICS Lesson: ATOMIC NATURE OF MATTER TABLE OF CONTENTS INTRODUCTION OBJECTIVES 1.0 PROPERTIES OF SUBSTANCES 1.1 CHEMICAL AND PHYSICAL PROPERTIES 2.0 COMPOSITION OF ATOMS 2.1 ATOMIC STRUCTURE 2.2 NUCLEAR NOTATION 2.3 ATOMIC MASS AND WEIGHT 2.4 ELECTRON STRUCTURE 3.0 THE NUCLEUS 4.0 ISOTOPES 4.1 ISOTOPIC ABUNDANCE 5.0 SUMMARY INTRODUCTION This is the first of five lessons in nuclear physics. This lesson, will present the basic atomic nature of matter. Included in the discussion is information on Atomic Structure, Atomic Mass and Atomic Weight. At the completion of this lesson, the Contractor Health Physics Technician should have a good understanding of the construction of an atom, and the standard nomenclature for describing atoms. The information presented in this lesson provides the foundation for more complex ideas. OBJECTIVES TERMINAL OBJECTIVE The Contractor Health Physics Technician will describe the structure and properties of atoms and their constituents. The Contractor Health Physics Technicians will describe the idea of ionization. ENABLING OBJECTIVES Upon completion of this lesson, the Contractor Health Physics Technician will be able to: ·Describe the components of atoms including each component's symbol, relative mass, and relative electrical charge. ·Recognize and manipulate atomic symbols to determine the number of protons, neutrons, or chemical symbol. ·Define isotope. 1.0 PROPERTIES OF SUBSTANCES Atoms are the basic building blocks of all matter. There are presently 105 elements (92 occur naturally, 13 made artificially), each element having atoms with the same number of protons, but possibly with different number of neutrons. The atom is the smallest
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particle of an element that can enter into a chemical reaction. Figure NP11 lists 22 elements common to nuclear power along with their symbols. Figure NP11 ELEMENTS COMMON TO NUCLEAR PLANTS When atoms combine and join, they form molecules. A molecule is the smallest particle of a substance that still retains the characteristics of that substance. In many cases, a single atom displays the properties of the substance, so that a molecule consists of a single atom. This is true for metals such as iron, lead, aluminum, etc. Molecules of compounds consist of one or more atoms of at least two elements in combination. The properties of a substance depend on the structure of the molecule and frequently do not resemble the properties of the elements from which they are produced. Sodium (an explosive metal) and chlorine (a poisonous gas) combine to form sodium chloride, common table salt. 1.1 CHEMICAL AND PHYSICAL PROPERTIES Chemical properties are those which deal with a substance entering into a reaction with other substances. These properties include the ability to burn or corrode and form compounds. Chemical reactions concern the transfer or sharing of outer shell electrons, as shown in Figure NP12 and NP13. Chemical reactions do not involve the atom's nucleus and the energy involved in chemical reactions is about a million times less than nuclear reactions. Observing the physical properties of a substance does not create a chemical change in the material. Some physical properties are: color, density, hardness, elasticity, melting point, boiling point and crystalline structure. Figure NP12 Formation of Sodium Chloride Figure NP13 Formation of a Water Molecule 2.0 COMPOSITION OF ATOMS Atoms, the building blocks of all substances, contain even smaller (subatomic) particles. All atoms consist of a positively charged nucleus surrounded by a number of negatively charged electrons. An adequate model for our purposes includes electrons as very small particles revolving around a central nucleus in orbits (similar to planets revolving around the sun). Figure NP14 illustrates the model. Atoms are normally electrically neutral because the number of electrons equals the number of protons. The coulombic force (opposite charges attract) attracts electrons toward the nucleus. They continue to orbit the nucleus because of the kinetic energy they possess. They can escape from the atom only if they obtain sufficient energy to overcome the forces that hold them in orbit. 14 Atoms are mostly empty space; the nucleus has a diameter of about 10 meters and the 10 entire atomic radius is about 10 meters. The nucleus contains about 99.97% of the mass 10 of the atom in about 10 % of the volume. If you could increase the size of the nucleus to that of a baseball, the distance between it and the nearest orbital electron would be about onehalf mile! STRUCTURE2.1 ATOMIC
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The three basic particles that make up an atom are the proton, neutron and electron. A description of each particle's physical properties is as follows: Proton The proton is an elementary particle located in the atom's nucleus. It has a positive 14 electrostatic charge and a mass of 1.6724 x 10 grams, symbol p. Neutron The neutron is an elementary particle also located in the atom's nucleus. It has no 24 charge. Its' mass is slightly greater than the proton at 1.6747 x 10 grams. Presently the character0is in use to symbolize it.. Electron The electron is an elementary particle that orbits the nucleus. It has a negative electrostatic charge, equal in charge intensity to the proton. Its mass is 1/1838 that of a proton, symbol e. An atom consists of the protons and neutrons tightly bound in a central nucleus, surrounded by the electrons orbiting in "shells" or energy levels. The nucleus carries a positive electric charge due to the presence of protons, so that the total charge is the sum of the total number of protons. This positive charge balances the total negative charge of the orbiting electrons, so that the whole atom is electrically neutral. Thus, the number of orbital electrons equals the number of protons in the nucleus for a neutral atom. The number of electrons determines the atom's chemical properties, which are largely uninfluenced by the nucleus itself. 2.2 NUCLEAR NOTATION Figure NP15 illustrates a simple method for designating any nuclide. This accepted notation is in common usage. Recall again the definitions of atomic mass number and atomic number. Atomic number (Z) is the number of protons in the nucleus. It defines the element. Atomic mass number (A) is the sum of the number of protons and neutrons in the nucleus. Therefore, the number of neutrons in a given nucleus is the difference between the atomic mass number and atomic number. Number of neutrons = A Z Hydrogen, helium and uranium atoms are examples of this designation and are as follows: 1 0 neutron 1=H = 1 proton 1 protonH 4 2 neutrons 2neutrons+ 2 2 protons He = =2 protonsHe 235 143 neutrons 9292 protons U = neutrons =+ 143 92 protonsU To determine the number of particles in any atom, the following rules apply:
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The number of protons in the nucleus is always equal to the atomic number of the element (Z). The number of neutrons is equal to the difference between the atoms atomic mass number and atomic number (AZ). The number of orbital electrons is equal to the number of protons in electrically neutral atoms. 2.3 ATOMIC MASS AND WEIGHT 14 Since atomic masses are about 10 grams, a need for a more convenient unit exists. The chosen standard is the "atomic mass unit" (AMU). One AMU is equal to 1/12 the mass of a carbon 12 nucleus (carbon 12 contains six protons and six neutrons). Present evidence indicates that 1 AMU is very nearly the mass of a proton or neutron. The accurately determine value is: 14 gm1 AMU = 1.6605 x 10 The atomic weight of an atom is its mass in AMUs. The nuclear mass differs from the atomic weight by the total masses of the electrons. Further, the atomic weight of an element must be considered in detail. A naturally occurring element consists of all those atoms whose nuclei contain a given number of protons, which is the atomic number. For example, the element carbon contains all naturally occurring atoms with atomic number 6. However, the atomic number does not completely describe the nucleus, since there may be differing numbers of neutrons included in all nuclei with an atomic number of 6. Each Nuclei of a given element, has the same atomic number, but possess different numbers of neutrons, and are known as the isotopes of the element. The electron structure of an atom determines the chemical properties of an atom. The isotopes of an element make up a set of different atoms having the same chemical nature, but differing in atomic weight. Thus, the atomic weight of an element is the average atomic weight of all the naturally occurring isotopes of that element. Taking into account the relative isotopic abundance, most elements consist of 2 or more isotopic types. Isotopes will be discussed in greater detail in Section 4.0 STRUCTURE2.4 ELECTRON The configuration of the orbiting electrons determines the chemical properties of an atom. Specifically, these properties depend upon the number of electrons in the outermost shell (valence electrons). Each shell represents a different energy level or energy state, the innermost shell representing the lowest energy state the outermost shell the highest. If a shell has less than the allowed number of electrons, an electron from a higher energy shell can "drop" into the vacant spot (with Xrays or visible light range energy given off in the process). An atom is in the "ground" state when its electrons are in the innermost or lowest energy level shells. An atom in an "excited" state has an electron in a higher energy shell. Each shell has a maximum number of electrons that it will hold. Table 1 gives the maximum number of electrons for any given shell. Table 1 demonstrates that if there are 2 electrons they will occupy the K shell. In an atom with more than 2 and less than 10 electrons, the first 2 electrons are in the K shell and the remaining electrons are in the L shell. In atoms with more than 10 but less than 28
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electrons, the first 2 electrons are in the K shell, the next 8 are in the L shell, and the remaining electrons are in the M shell, etc... . However, the outermost shell (valence shell) cannot contain more than 8 electrons. Atoms with only 1 or 2 valence shell electrons are highly reactive in chemical reactions. Atoms with a full valence shell normally do not react at all. See Figures NP16 and NP17. NUCLEUS3.0 THE As mentioned earlier, the atomic nucleus contains protons and neutrons closely packed together. Clearly some force must be present that holds the nucleus together and overcomes the strong coulombic repulsive forces of the tightly packed protons. The "nuclear force" is the name for this very strong force. The nuclear force acts between fundamental particles that are neither electrical nor gravitational in nature. The nuclear force is a "shortrange" force, and when protons or neutrons are within about 1013cm the nuclear force binds them together strongly, overcoming any electrostatic repulsion between protons. Nuclear forces have several important characteristics: They are charge independent; that is, the force is just as strong whether neutrons or protons are being considered. They are extremely strong, much stronger than gravitational or electrical forces. 13 cm.They have very short range, about 10 They are saturable, one particle can only exert a nuclear force on a limited number of other particles. Although nuclear forces are much stronger than electrostatic forces, the electrostatic forces play a role in tending to force the nucleus apart. Because neutrons can add nuclear force without electrostatic repulsion, their presence in the nucleus is very important. This gives rise to the following two important facts: ·there is no stable nuclei consisting of two or more protons with no neutrons ·as the number of protons in a nucleus becomes greater, a relatively greater number of neutrons needed to stabilize the nucleus. The "belt of stability" shown in Figure NP18 (a plot of stable nuclei) demonstrates the latter point. (The terminology used here is that an "unstable" nucleus is a combination of protons and neutrons that will not remain a unit indefinitely. An unstable nucleus emits energy or particles in a process called "radioactive decay". Each box in Figure NP18 represents a known nucleus containing Z number of protons and n number of neutrons. The straight line represents the line along which nuclei would lie if they contained equal numbers of protons and neutrons. For nuclei with mass numbers less than 40, the stable isotopes contain approximately equal numbers of protons and neutrons. However, the heavier stable isotopes contain considerably more neutrons than protons. The "belt of stability" would encompass a line drawn through the average location of stable nuclei on this graph. Observations indicate that nuclei on either side of the line have too low or too high a N/P ratio for stability and will undergo radioactive decay to bring the nucleus closer to stability. 4.0 ISOTOPES
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Not all atoms of a particular element are exactly alike. Hydrogen, for example, exists in three forms. The most abundant form of hydrogen exists as one proton and one electron. The deuterium atom, an isotope of hydrogen, differs from ordinary hydrogen in that deuterium contains a neutron together with the proton in the nucleus. Tritium, another isotope of hydrogen, has two neutrons and one proton in the nucleus. It follows that the three isotopes of hydrogen differ in their atomic mass. Figure NP19 illustrates the three isotopes of hydrogen. We can the define isotopes as: atoms that have the same atomic number, but different atomic weights. To identify individual isotopes of an element, use the appropriate mass number; e.g., hydrogen1, hydrogen2 and hydrogen3 identify the isotopes of hydrogen.. Most elements exist in nature as a mixture of two or more isotopes. For example, there are two stable isotopes of nitrogen with mass numbers 14 and 15, as shown in Figure NP 110. These isotopes of nitrogen differ only in the fact that N15 has one more neutron in the nucleus. The electron structure is the same, otherwise the two isotopes would have different chemical properties. Since all isotopes of an element have the same number of protons in the nucleus and the same number and distribution of electrons, we can properly define an element as a substance consisting entirely of atoms having the same atomic number; that is, the same number of protons. 4.1 ISOTOPIC ABUNDANCE Isotopic abundance refers to the percentage of each isotope in an element's natural state. For instance, ordinary copper has an abundance of 69.2% of copper63 and 30.8% of copper65. The composition of natural silver is of 51.83% of silver107 and 48.17% silver109. Carbon has an isotopic abundance of 98.89% carbon12 and 1.11% carbon13. abundance is the amount of the isotope (percentage) present in aDefinition: Isotopic normal natural mixture of the element. As an example of an element existing in a natural state, consider uranium. Natural uranium consists of a mixture of isotopes having mass numbers 234, 235 and 238. Table 2 lists the isotopic abundance of natural uranium isotopes. Table 2 Composition of Natural Uranium Abundance Atomic Isotope Percent Mass U234 0.006% 234.0409 U235 0.714% 235.0439 238.0508U238 99.28% 238.03Natural U The atomic mass given in Table 2 for natural uranium is a weighted average based on the percent abundance of the atomic mass of each isotope. It is the atomic mass given for uranium on any modern periodic chart. This is the case for all atomic masses given in periodic tables in that the table gives the atomic mass of the natural element.
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5.0 SUMMARY Atoms consist of 3 basic particles: protons, neutrons and electrons. The protons and neutrons located in the nucleus about which the electron(s) orbit. The nucleus contains almost all the mass of the atom, but occupies only a very small part of the atom's entire volume. Atomic mass, atomic weight, and the atomic mass unit (AMU), were the subjects of this lesson. The AMU is equal to 1/12 the mass of a carbon12 nucleus. All the atoms of a particular element have the same number of protons (Z). However, not all atoms of an element have the same atomic mass (A). Atoms with the same Z, but different A numbers are isotopes. The following notation can identify an isotope: ZXA, 235 element name, or symbol and atomic mass number. For example,92U or uranium235 or U235. Electrons orbit the nucleus in "shells". Each shell represents a discrete energy level. Each shell has a limit on the number of electrons it may contain. The shell closest to the nucleus (K shell) represents the lowest energy state, and the outermost shell, the highest energy state. The nucleus is a tightly bound cluster of neutrons and protons called nucleons. The "nuclear force" holds the nucleons within the nucleus. This force is charge independent, extremely strong, has a very short range, and is saturable. Although nuclear forces are much stronger than electrostatic forces, the electrostatic forces tend to force the nucleus apart. Neutrons add nuclear force without adding electrostatic repulsion. Current theory is that the stability of the nucleus is dependent on the neutron to proton ratio of the nucleus. A nucleus with too low or too high of an n/p ratio for stability will undergo radioactive decay.
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REFERENCES Freeman, Ira M. Physics Made Simple, Garden City, N.Y.: Doubleday and Co., 1965. Shortly, George; William Dudley. Principles of College Physics, Second Edition. Englewood Cliffs, N.J.: PrenticeHall, Inc., 1967. Tulley, Donald E.; Thumm, Walter. Physics for College Students with Applications to the Life Sciences, Menlo Park, CA: Cummings Publishing Co., 1974. LIST OF FIGURES NP11Chemical Symbols NP12Formation of Sodium Chloride NP13Formation of a Water Molecule NP14Composition of Atoms NP15Nuclear Notation NP16Electron Structure NP17Electron Structure NP18Nuclei for Z and n NP19Isotopes of Hydrogen of NitrogenNP110 Isotopes LIST OF TABLES Table 1 Maximum Number of Electrons per Shell Table 2 Composition of Natural Uranium
