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Index Honors Chemistry Unit 02 Atoms and Elements Based on the PowerPoints by Kevin Boudreux. Honors Unit 02 Lesson 01 1 Index Unit 02 Module 01: THE ROAD TO THE ATOMIC THEORY Honors Unit 02 Lesson 01 2 Index Atomos • The ancient Greek philosopher Democritus (c. 460 – 370 BC) reasoned that if you cut a lump of matter into smaller and smaller pieces, you would eventually cut it down to a particle which could not be subdivided any further. He called these particles atoms (from the Greek atomos, “uncuttable”) • Aristotle (384-322 BC) believed that matter was continuous, and elaborated the idea that everything was composed for four elementary substances, assembled in varying proportions – earth, air, fire, and water, which possessed four properties – hot, dry, wet, and cold. • The idea of atoms did not surface again until the 17th and 18th centuries. Honors Unit 02 Lesson 01 3 Index Law of Conservation of Mass • In 1661, Robert Boyle redefined an element as a substance that cannot be chemically broken down further. • Law of Conservation of Mass – Mass is neither created nor destroyed in chemical reactions (i.e. the total mass of a system does not change during a reaction). Massreactants = Massproducts #Atomsreactant= #Atomsproduct Honors Unit 02 Lesson 01 4 Index Conservation of mass demo (Alka-Seltzer) Honors Unit 02 Lesson 01 5 Index Balancing Reactions by Mass CaCO3( s ) CaO( s ) CO2( g ) If 25g of calcium carbonate (CaCO3) is heated to produce 14 g of calcium oxide (CaO), what is the mass of the carbon dioxide (CO2) produced? Index Balancing Reactions by Mass Guided Practice Na + H2O NaO + H2 If 23 g of sodium reacts with 18 g of water to produce 39 g of sodium oxide how much hydrogen gas is produced? Index Balancing Reactions by Mass Guided Practice 6CO2 + 6H2O C6H12O6 + 6O2 Photosynthesis If 30. g of glucose and 32 g of oxygen are produced from 44 g of carbon dioxide, how much water is needed? Index Law of Definite Proportions • Law of Definite Proportions – All samples of a pure chemical substance, regardless of their source or how they were prepared, have the same proportions by mass of their constituent elements. (Joseph Proust, 1754-1826) – Calcium carbonate, which is found in coral, marble, the white cliffs of Dover, seashells, chalk, limestone, and Coleman tap water, is always 40.04% by mass calcium, 12.00% carbon, and 47.96% oxygen. (We now know that this results from the fact that calcium carbonate is CaCO3.) Honors Unit 02 Lesson 01 9 Index Limestone Honors Unit 02 Lesson 01 10 Index Oolitic limestone Honors Unit 02 Lesson 01 11 Index Marble Honors Unit 02 Lesson 01 12 Index Honors Unit 02 Lesson 01 13 Index Honors Unit 02 Lesson 01 14 Index The Law of Multiple Proportions • Law of Multiple Proportions – Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other. (John Dalton, 1804) Honors Unit 02 Lesson 01 15 Index The Law of Multiple Proportions Compound Sample Mass of Size Sulfur Mass of Oxygen Sulfur oxide “I” 2.00 g 1.00 g 1.00 g Sulfur oxide “II” 2.50 g 1.00 g 1.50 g mass of oxygen in sulfur oxide "II" per gram of sulfur 1.50g 3 = = mass of oxygen in sulfur oxide "I" per gram of sulfur 1.00g 2 Honors Unit 02 Lesson 01 16 Index Honors Unit 02 Lesson 01 17 Index Dalton’s Atomic Theory • John Dalton (1766-1844) explained these observations in 1808 by proposing the atomic theory: – Each element consists of tiny indivisible (not quite) particles called atoms. – All atoms of the same element have the same mass (not quite), but atoms of different elements have different masses. – Atoms combine in simple, whole-number ratios to form compounds. – Atoms of one element cannot change into atoms of another element (not quite). In a chemical reaction, atoms change the way they are bound together with other atoms to form a new substance. Honors Unit 02 Lesson 01 18 Index Dalton’s Atomic Theory • Dalton’s atomic hypothesis had an uphill struggle – many scientists didn’t like the idea of using small, invisible entities to explain phenomena. • Most (but not all) chemists had accepted the existence of atoms by the early 20th century; however, many influential physicists did not accept the atomic theory until Einstein’s landmark paper on Brownian motion (1905). Honors Unit 02 Lesson 01 19 Index Dalton’s Atomic Theory • Dalton’s original formulation of atoms as miniature billiard balls, however, could not explain how atoms combined to form compounds, or anything about their internal structure. The theory was modified greatly once charge particles coming from inside the atom were discovered in the late 19th and early 20th century. Honors Unit 02 Lesson 01 20 Index Brownian Motion Source: http://protonsforbreakfast.wordpress.com/2012/04/11/brownian-motion-observedin-milk/ Honors Unit 02 Lesson 01 21 Index The Electron (e-) • In 1897, J.J. Thomson (1856-1940) discovered that cathode rays, which are produced by passing an electric current through two electrodes within a vacuum tube, were composed of negatively charge particles, which have a very low mass. These particles were called electrons. • His experiments showed that electrons were emitted by many types of metals, so electrons must be present in most if not all samples of matter. • Although Thomson was unable to measure the mass of the electron directly, he was able to determine that the chargeto-mass ration, e/m, was -1.758820x108 C/g. This meant that the electron was about 2000 times lighter than hydrogen, and that atoms were not the smallest unit of matter. Honors Unit 02 Lesson 01 22 Index Cathode Ray Tubes Honors Unit 02 Lesson 01 23 Index High voltage (+) High voltage (-) Cathode Anode Honors Unit 02 Lesson 01 24 Index This is how J.J. Thomson discovered that the electron was negatively charged. Honors Unit 02 Lesson 01 25 Index High Voltage (+) on N End View High Voltage (-) on Cathode Anode S This is how J.J. Thomson discovered that the electron had mass. Honors Unit 02 Lesson 01 26 Index Honors Unit 02 Lesson 01 27 Index Source: http://www.youtube.com/watch?v=0pWBXcnP5Po Honors Unit 02 Lesson 01 28 Index _ N S + Honors Unit 02 Lesson 01 29 Index Inversely Prop to the mass 2.37 North South _ 1.96 Prop to the charge + Honors Unit 02 Lesson 01 30 Index Plate Charge (Conversion Factors) Magnetic Field Charge of Electron e = = Mass of Electron m e 8 coulombs 8 c = -1.76 x10 gram = -1.76 x10 g m Cathode Tube Demo Honors Unit 02 Lesson 01 31 Index The Mass of the Electron • In 1909, Robert Millikan (1868-1953) measured the charge on the electron by observing the movement of tiny ionized droplets of oil passing between two electrically charge plates. Knowing the e/m ratio from Thomson’s work, the mass of the electron could then be determined: Charge of an electron: e = - 1.60218 x 10-19 C Charge to Mass Ratio: e/m = -1.758820 x 108 C/g Mass of an electron: me = 9.1093897 x 10-28 g Honors Unit 02 Lesson 01 32 Index Millikan’s Oil Drop Experiment Honors Unit 02 Lesson 01 33 Index Millikan’s Oil Drop Experiment Honors Unit 02 Lesson 01 34 Index Source:http://www.youtube.com/watch?v=XMfYHag7Liw Honors Unit 02 Lesson 01 35 Index So Where’s the Positive Charge? • If there is a negatively charged thing somewhere inside an electrically neutral atom, then there must also be positively charged portion. • The model for the atom that Thomson proposed was of a diffuse, positively charged lump of matter with electrons embedded in it like “raisins in a plum pudding” (seeds in a watermelon, blueberries in a muffin, or chocolate chips in a cookie might be more familiar analogies) J.J. Thomson’s Plum Pudding Model of the Atom Honors Unit 02 Lesson 01 36