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Chapter 3 – Atoms: The Building Blocks of Matter I. The Atom: From Philosophical Idea to Scientific Theory A. The Philosophers 1. Democritus (~400 BC) – 2. Another Greek philosopher of the time, Aristotle, believed * B. Foundations of Atomic Theory Law of Conservation of Mass – Law of Definite Proportions – Law of Multiple Proportions- C. John Dalton’s Atomic Theory (1808) 1. Basic Points a. b. c. d. e. D. 2. The Law of Conservation of Mass is explained by point e 3. The Law of Multiple Proportions is explained by point d Modern Atomic Theory 1. Some of Dalton’s points were not exactly accurate and have been modified throughout the years 2. Two important concepts remain unchanged: a. b. II. The Structure of the Atom A. Discovery of the Electron 1. Cathode Rays and Electrons (Late 1800s) a. b. When current was passed through the cathode ray tube, the surface of the tube opposite the cathode glowed c. 1. Hypothesized that the glow was caused by a stream of particles called “cathode rays” 2. Rays seemed to travel from the cathode to the anode Further experiments supported the hypothesis and added information: 1. Included a paddle wheel which, when placed on rails between the electrodes, rolled along the rails from the cathode to the anode 2. Cathode rays, when exposed to electricallycharged plates, were deflected by the negative plate and attracted to the positive plate * This information led to the hypothesis that 2. 3. Two Important Scientists a. JJ Thomson (1897) – b. Robert Milliken (1909) – Two Other Inferences a. Because atoms are electrically neutral, they must contain b. Because electrons have such a small mass, atoms must contain B. Discovery of the Atomic Nucleus 1. 2. 3. The Gold Foil Experiment (1911) a. Carried out by Ernest Rutherford, Hans Geiger, and Ernest Marsden (among others) b. The Experimental Design 1. A thin, gold foil was bombarded with alpha particles which were 2. The foil was surrounded by a screen that would detect where the alpha particles hit after passing through the foil The Results a. Most of the alpha particles went straight through the foil or were only slightly deflected (this was expected) b. HOWEVER, Rutherford’s Explanation (1913) a. b. c. C. Composition of the Atomic Nucleus 1. Except for the simplest type of hydrogen atom, all atomic nuclei are made up of two kinds of particles, Particle 2. Electrons travel 3. Summary of Subatomic Particles Symbols Relative Electric Charge Mass Number (A) Relative Mass (amu) Actual Mass (g) Electron Proton Neutron D. The Sizes of Atoms 1. The region of the atom occupied by the electron can be thought of as 2. The radius of an atom is defined as 3. Atomic radii are expressed in units of picometers (pm)…or, in other words, III. Counting Atoms A. Atomic Number (Z) – 1. Every element has a different atomic number…this identifies the element! 2. Found on the periodic table…in each element block near the symbol of the element 3. B. In a neutral atom, there will be Mass Number (A) – 1. MUST BE GIVEN TO YOU!!! NOT on the periodic table! 2. Atoms of the same element may have different numbers of neutrons and, therefore, different mass numbers a. These atoms are referred to as isotopes… b. Example - Hydrogen Protium (Most Deuterium common form of “Heavy H… Hydrogen” 99.985%) (0.015%) Tritium “Radioactive Hydrogen” (trace) # protons # neutrons Mass Number c. C. Although isotopes have different masses, they do not significantly differ in chemical behavior Designating Isotopes 1. Nuclear Symbol 2. Hyphen Notation 3. Using either designation, one can determine the number of protons, neutrons, and electrons… 4. Nuclear Symbol a. # protons = atomic number (Z) = # electrons b. # neutrons = Examples Hyphen Notation # Protons # Neutrons # Electrons 20 21 20 35Cl 131I Uranium-235 5. D. Nuclide – Relative Atomic Masses of the Elements 1. Because atoms are so small, using mass in units of grams is inconvenient and hard to understand 1 atom hydrogen-1 = 1.673 x 10-24 grams 2. Scientists set up a relative scale of atomic masses…known as the atomic mass scale a. Carbon-12 was arbitrarily assigned a mass of 12 atomic mass units (amus) b. One atomic mass unit (amu), or 1 amu, is c. E. The mass of all other nuclides are determined by comparing each one with the mass of a carbon-12 atom Average Atomic Masses of the Elements 1. The mass found in each element block on the periodic table must account for all isotopes of that element that are in existence and how abundant they are in nature 2. Average atomic mass - a. To do this, use the following formula for each isotope: (% occurrence 100) (mass) b. 3. Add all values together to get the average atomic mass Example a. In a box containing 2 sizes of marbles, 25% of the marbles have a mass of 2.00 g each and 75% of the marbles have a mass of 3.00 g each. What is the weighted average of the marbles? b. Find the average atomic mass for oxygen using the following information: Percentage Natural Abundance 99.762 0.038 0.200 Isotope Oxygen-16 Oxygen-17 Oxygen-18 4. Atomic Mass (amu) 15.994915 16.999131 17.999160 In our textbook, an element’s atomic mass is usually rounded to two decimal places before it’s used in a calculation IV. Basic Bonding A. Atoms are held together (forming compounds) by chemical bonds; there are two types 1. Ionic Bonding - 2. Covalent Bonding - B. Determining the Type of Bonding Based on the Chemical Formula 1. If the compound is ionically bonded… 2. If the compound is covalently bonded…