Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Chapter 5 – Early Atomic Theory & Structure 5.1 Early Thoughts 5.6 Subatomic Parts of the Atom 5.2 Dalton's Model of the Atom 5.7 The Nuclear Atom 5.3 Composition of Compounds 5.8 General Arrangement of Subatomic Particles 5.4 The Nature of Electric Charge 5.9 Atomic Numbers of the Elements 5.5 Discovery of Ions 5.10 Isotopes of the Elements 5.11 Atomic Mass 1 Early Thoughts • The earliest models of the atom were developed by the ancient Greek philosophers. Empedocles stated that matter was made of 4 elements: earth, air, fire, and water. Democritus (about 470-370 B.C.) thought that all forms of matter were divisible into tiny indivisible particles. He called them “atoms” from the Greek “atomos” indivisible. 2 Early Thoughts Aristotle (384-322 B.C.) rejected the theory of Democritus and Empedoclean theory. advanced the – Aristotle’s influence dominated the thinking of scientists and philosophers until the beginning of the 17th century. 3 Dalton’s Model of the Atom 2000 years after Aristotle, John Dalton an English schoolmaster, proposed his model of the atom–which was based on experimentation. 4 Dalton’s Atomic Theory 1. Elements are composed of minute indivisible particles called atoms. 2. Atoms of the same element are alike in mass and size. 3. Atoms of different elements have different masses and sizes. Modern research has demonstrated that Atoms under special circumstances can 4. Chemical compounds are formed by atoms are composed of subatomic be decomposed. the union of two or atoms of different particles. elements. 5 Dalton’s Atomic Theory 5. Atoms combine to form compounds in simple numerical ratios, such as one to one , two to two, two to three, and so on. 6. Atoms of two elements may combine in different ratios to form more than one compound. 6 Dalton’s atoms were individual particles. Atoms of each element are alike in mass and size. 7 5.1 Dalton’s atoms were individual particles. Atoms of different elements are not alike in mass and size. 8 5.1 H 2 = O 1 H 1 = O 1 Daltons atoms combine in specific ratios to form compounds. 9 Composition of Compounds The Law of Definite Composition A compound always contains two or more elements combined in a definite proportion by mass. 10 Composition of Water • Water always contains the same two elements: hydrogen and oxygen. The percent by mass of hydrogen in water is 11.2%. The percent by mass of oxygen in water is 88.8%. Water always has these percentages. If the percentages were different the compound would not be water. 11 Composition of Compounds The Law of Multiple Proportions Atoms of two or more elements may combine in different ratios to produce more than one compound. 12 Composition of Hydrogen Peroxide • Hydrogen peroxide always contains the same two elements: hydrogen and oxygen. The percent by mass of hydrogen in hydrogen peroxide is 5.9%. The percent by mass of oxygen in hydrogen peroxide is 94.1%. Hydrogen peroxide always has these percentages. If the percentages were different the compound would not be hydrogen peroxide. 13 Combining Ratios of Hydrogen and Oxygen • Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water. The formula for water is H2O. The formula for hydrogen peroxide is H2O2. 15 Composition of Compounds 16 The Nature of Electric Charge Properties of Electric Charge Charge may be of two types: positive and negative. Unlike charges attract (positive attracts negative), and like charges repel (negative repels negative and positive repels positive). Charge may be transferred from one object to another, by contact or induction. The less the distance between two charges, the greater the force of attraction between unlike charges (or repulsion between identical charges). kq1q 2 F= 2 r q1 and q2 are charges, r is the distance between charges and k is a constant. 17 Discovery of Ions • Michael Faraday discovered that certain substances when dissolved in water conducted an electric current. He found that atoms of some elements moved to the cathode (negative electrode) and some moved to the anode (positive electrode). He concluded they were electrically charged and called them ions (Greek wanderer). 18 Discovery of Ions • Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge. Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-. Δ NaCl → + Na + Cl 19 Discovery of Ions NaCl → + Na + Cl In the melt the positive Na+ ions moved to the cathode (negative electrode). Thus positive ions are called cations. In the melt the negative Cl- ions moved to the anode (positive electrode). Thus negative ions are called anions. 20 Subatomic Parts of the Atom The diameter an billionths atom is 0.1ofto a This is 1 to 5often 0.5 nm. meter. If the diameter of this dot is 1 Even smaller particles than atoms mm, then 10 million hydrogen exist. These are called subatomic atoms would form a line across particles. the dot. 21 Subatomic Particle - Electron • In 1875 Sir William Crookes invented the Crookes tube. Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom. Crookes tube emissions are called cathode rays. 22 Subatomic Particle - Electron In 1897 Sir Joseph Thompson demonstrated that cathode rays: travel in straight lines. are negative in charge. are deflected by electric and magnetic fields. produce sharp shadows are capable of moving a small paddle wheel. 23 Subatomic Particle - Electron This was the discovery of the fundamental unit of charge – the electron. 24 Subatomic Particle - Proton • Eugen Goldstein, a German physicist, first observed protons in 1886: Thompson determined characteristics. the proton’s Thompson showed that atoms contained both positive and negative charges. This disproved the Dalton model of the atom which indivisible. held that atoms were 25 Subatomic Particle - Neutron • James Chadwick discovered the neutron in 1932. Its actual mass is slightly greater than the mass of a proton. 26 Subatomic Particles 27 Ions • Positive ions were explained by assuming that a neutral atom loses electrons. Negative ions were explained by assuming that extra electrons can be added to atoms. 28 When one or more electrons are lost from an atom, a cation is formed. 5.4 29 When one or more electrons are added to a neutral atom, an anion is formed. 5.4 30 The Nuclear Atom • Radioactivity was discovered by Becquerel in 1896. Radioactive elements spontaneously emit alpha particles, beta particles and gamma rays from their nuclei. By 1907 Rutherford found that alpha particles emitted by certain radioactive elements were helium nuclei. 31 The Rutherford Experiment • Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil. Most of the alpha particles passed through the foil with little or no deflection. He found that a few were deflected at large angles and some alpha particles even bounced back. 32 The Rutherford Experiment 5.5 Rutherford’s alpha particle scattering experiment. 33 The Rutherford Experiment • An electron with a mass of 1/1837 amu could not have deflected an alpha particle with a mass of 4 amu. Rutherford knew that like charges repel. Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus. 34 The Rutherford Experiment • If a positive alpha particle approached close enough to the positive mass it was deflected. Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space. 35 The Rutherford Experiment Because alpha particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense. 36 The Rutherford Experiment Deflection Scattering 37 Deflection and scattering of alpha particles by positive gold nuclei. 5.5 General Arrangement of Subatomic Particles • Rutherford’s experiment showed that an atom had a dense, positively charged nucleus. Chadwick’s work in 1932 demonstrated the atom contains neutrons. Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge. 38 General Arrangement of Subatomic Particles • Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center. The negative electrons surround the nucleus. The nucleus contains protons and neutrons 39 General Arrangement of Subatomic Particles 40 5.6 Atomic Numbers of the Elements • The atomic number of an element is equal to the number of protons in the nucleus of that element. The atomic number of an atom determines which element the atom is. 41 Atomic Numbers of the Elements Every atom with an atomic number of 1 is a hydrogen atom. Every hydrogen atom contains 1 proton in its nucleus. 42 atomic number Every atom with an atomic number of 1 is a hydrogen atom. 1 proton in the nucleus 43 Atomic Numbers of the Elements Every atom with an atomic number of 6 is a carbon atom. Every carbon atom contains 6 protons in its nucleus. 44 atomic number Every atom with an atomic number of 6 is a carbon atom. 6 protons in the nucleus 45 atomic number Every atom with an atomic number of 92 is a uranium atom. 92 protons in the nucleus 46 Atomic Numbers of the Elements • Atoms of the same element can have different masses. They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. The difference in the number of neutrons accounts for the difference in mass. These are isotopes of the same element. 47 Atomic Numbers of the Elements Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons 48 Isotopic Notation 49 Isotopic Notation 6 protons + 6 neutrons 12 C 6 6 protons 50 Isotopic Notation 6 protons + 8 neutrons 14 C 6 6 protons 51 Isotopic Notation 8 protons + 8 neutrons 16 O 8 8 protons 52 Isotopic Notation 8 protons + 9 neutrons 17 O 8 8 protons 53 Isotopic Notation 8 protons + 10 neutrons 18 O 8 8 protons 54 Hydrogen has three isotopes 1 proton 1 proton 1 proton 0 neutrons 1 neutron 2 neutrons 55 Examples of Isotopes Element Protons Electrons Neutrons Symbol Hydrogen Hydrogen Hydrogen Uranium Uranium Chlorine Chlorine 1 1 1 92 92 17 17 1 1 1 92 92 17 17 0 1 2 143 146 18 20 1 1 H 2 1 H 3 1 H 235 192 U 236 192 U 35 17 37 17 Cl Cl 56 Atomic Mass • The mass of a single atom is too small to measure on a balance. Using a mass spectrometer, the mass of the hydrogen atom was determined. 57 A Modern Mass Spectrometer Positive ions formed from sample. Electrical field From the intensity and positions at slits A mass of the lines on the Deflection of mass accelerates spectrogram spectrogram, theions different positive ions. positive is recorded. isotopes occurs and their relative at amounts can be determined. magnetic field. 5.8 58 A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given. 5.9 59 Atomic Mass Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 g. To overcome this problem of such a small mass relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers. 60 Atomic Mass The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon – carbon 12. A mass of exactly 12 atomic mass units (amu) was assigned to carbon 12. 1 1 amu is defined as exactly equal to the mass of 12 a carbon-12 atom 1 amu = 1.6606 x 10-24 g 12 6 C 61 Atomic Mass H K U Average atomic mass 1.00797 amu. Average atomic mass 39.098 amu. Average atomic mass 248.029 amu. 62 Average Relative Atomic Mass • Most elements occur as mixtures of isotopes. Isotopes of the same element have different masses. The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly 12.0000…amu). 63 To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the results. Isotope Isotopic mass (amu) Abundance (%) 63 29 Cu 62.9298 69.09 65 29 Cu 64.9278 30.91 Average atomic mass (amu) (62.9298 amu) 0.6909 = 43.48 amu (64.9278 amu) 0.3091 = 20.07 amu 63.55 amu 64 Concepts Dalton’s Atomic Mode; Law of Definite Composition Law of Multiple Proportions Three principle subatomic particles Thomson Model of the Atom Rutherford alpha-scattering experiment Atomic Number, Mass number, number of neutrons, number of Protons, Number of Electrons Three Isotopes of Hydrogen Average Atomic Mass of an Element 66