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U N I T L E C T U R E 1 N O T E S LECTURE 11 Atomic Models (3.3, 3.5 & 11.1) GENERAL CHEMISTRY Fall 2009 Updated: 5/3/2017 3.3 Dalton’s Atomic Theory John Dalton (1766 – 1844) explained observations such as the law of constant composition (a compound always has the same composition) using his theory. The predictive value of the theory led to its eventual acceptance. A. Defining the Atom 1. Atomic Theory a. All matter is made up of very tiny particles called atoms b. Atoms of the same element are chemically alike c. Individual atoms of an element may not all have the same mass. However, the atoms of an element have a definite average mass that is characteristic of the element d. Atoms of different elements have different average masses e. Atoms are not subdivided, created, or destroyed in chemical reactions 3.5 The Structure of the Atom 1e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass. The volume of the hydrogen nucleus is about one trillion times less than the volume of the hydrogen atom, yet the nucleus contains almost all the mass in the form of one proton. The diameter of an atom of any one of the elements is about 10,000 to 100,000 times greater than the diameter of the nucleus. The mass of the atom is densely packed in the nucleus. The electrons occupy a large region of space centered around a tiny nucleus, and so it is this region that defines the volume of the atom. If the nucleus (proton) of a hydrogen atom were as large as the width of a human thumb, the electron would be on the average about one kilometer away in a great expanse of empty space. The electron is almost 2,000 times lighter than the proton; therefore, the large region of space occupied by the electron contains less than 0.1 percent of the mass of the atom. 2. Sizes of Atoms a. Atomic radius i. 40 to 270 picometers (pm) 1. 1 pm = 10-12m ii. Most of the atomic radius is due to the electron cloud b. Nuclear radius i. 0.001 pm ii. density is 2x108 metric tons/cm3 1. 1 metric ton = 1000kg Updated: 5/3/2017 B. Models of the Atom Scientist Year Democritus Dalton ~ 400 B.C.E. 1808 Model Thomson 1897 Plank 1900 Rutherford 1911 Bohr 1913 Einstein Schrödinger 1905 1926 Energy emitted in discrete quantities Nuclear Atom; also called the planetary model Bohr Model, electrons travel in discrete orbits Wave mechanical model Wave mechanical model Heisenberg 1929 Wave mechanical model Solid sphere, tiny, indivisible, indestructible particles Plum pudding Experiment Focus None Suggested Atom Weather data Cathode Ray Tube; also invented mass spectrometer Radiation from solids Electrons Gold foil Nucleus Spectrum of Hydrogen Excited and Ground state Photoelectric Effect Schrödinger cat; thought experiment Photons Schrödinger equation Quanta Heisenberg uncertainty principle Quarks and leptons (matter) Murray Gell1970s Standard Model Mann; George Zweig * There were many other models developed during this time period but we’ll only focus on these particular ones. C. Contributed to the Models of the Atom Maxwell 1873 Planck 1900 Chadwick 1932 Visible light consists of electromagnetic waves Energy emitted in discrete quantities Radiation from solids Identified subatomic particle Provides description of light Quanta; Plank’s constant Neutron Updated: 5/3/2017 11.1 Rutherford’s Atom 1a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass. An atom consists of a nucleus made of protons and neutrons that is orbited by electrons. The number of protons, not electrons or neutrons, determines the unique properties of an element. This number of protons is called the atomic number. Elements are arranged on the periodic table in order of increasing atomic number. Historically, elements were ordered by atomic mass, but now scientists know that this order would lead to misplaced elements (e.g., tellurium and iodine) because differences in the number of neutrons for isotopes of the same element affect the atomic mass but do not change the identity of the element. D. Structure of the Nuclear Atom 1. The Electron a. Discovery i. Joseph John Thomson (1897) 1. Cathode ray tube produces a ray with a constant charge to mass ratio 2. All cathode rays are composed of identical negatively charged particles (electrons) ii. Plum-pudding model b. Inferences from the properties of electrons i. Atoms are neutral, so there must be positive charges to balance the negatives ii. Electrons have little mass, so atoms must contain other particles that account for most of the mass 2. The Nucleus a. The Rutherford Experiment (1911) b. Alpha particles (helium nuclei) fired at a thin sheet of gold i. Assumed that the positively charged particles were bounced back if they approached a positively charged atomic nucleus head-on (Like charges repel one another) Updated: 5/3/2017 Results from gold foil experiment 1. Very few particles were greatly deflected back from the gold sheet a. nucleus is very small, dense and positively charged b. most of the atom is empty space 2. Structure of the Nucleus a. Protons i. Positive charge, mass of 1.673x10-27kg ii. The number of protons in the nucleus determines the atom's identity and is called the atomic number (Z) b. Neutrons i. James Chadwick (1932) ii. No charge, mass of 1.675x10-27kg c. Nuclear Forces i. Short range attractive forces: a. neutron-to-neutron, proton-to-proton, neutron proton-to- Unanswered Questions What are the electrons doing? How are the electrons arranged? How do electrons move? Why aren’t electrons (negatively charged) attracted to the positive nucleus? Updated: 5/3/2017