Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 The Early History of Chemistry 4 Greeks 400 B.C. - Four fundamental substances – fire, earth, water, a air. - Democritus – uses term “atomos” (atoms) to describe small, indivisible matter. No experiments to support the idea, so it is dropped. 4 Before 16th Century – Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 The Early History of Chemistry 4 17th Century –Robert Boyle: First “chemist” to perform quantitative experiments (pressure/volume) –-Incorrectly believed that the alchemist’s view that metals were not true elements and that a way would eventually be found to change one metal into another. 4 18th Century –George Stahl: Phlogiston flows out of a burning material. –Joseph Priestley: Discovers oxygen gas, “dephlogisticated air.”

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Law of Conservation of Mass 4 Discovered by Antoine Lavoisier 4 Mass is neither created nor destroyed 4 Combustion involves oxygen, not phlogiston

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Other Fundamental Chemical Laws 4 A given compound always contains exactly the same proportion of elements by mass. 4 NaCl – always 39.34% Cl and 60.66% Na (mass) 4 Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine. Law of Definite Proportion

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Other Fundamental Chemical Laws 4 When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. 4 The ratio of the masses of oxygen in H 2 O and H 2 O 2 will be a small whole number (“2”). Law of Multiple Proportions

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Law of Multiple Proportions Mass of oxygen that combines with 1 g of Carbon Compound 1 (CO)1.33 g Compound 2 (CO 2 )2.66 g

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Dalton’s Atomic Theory (1808) ÊEach element is made up of tiny particles called atoms. ËThe atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Dalton’s Atomic Theory (continued) ÌChemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. ÍChemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Avogadro’s Hypothesis (1811) 5 liters of oxygen 5 liters of nitrogen Same number of particles! At the same temperature and pressure, equal volumes of different gases contain the same number of particles.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Figure 2.5: A representation of combining gases at the molecular level. The spheres represent atoms in the molecules.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Early Experiments to Characterize the Atom H J. J. Thomson - postulated the existence of electrons using cathode ray tubes. H Ernest Rutherford - explained the nuclear atom, containing a dense nucleus with electrons traveling around the nucleus at a large distance.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Figure 2.7: A cathode-ray tube. The fast- moving electrons excite the gas in the tube, causing a glow between the electrodes.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Figure 2.8: Deflection of cathode rays by an applied electric field.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 The Modern View of Atomic Structure l Electrons l protons: found in the nucleus, they have a positive charge equal in magnitude to the electron’s negative charge. l neutrons: found in the nucleus, virtually same mass as a proton but no charge. The atom contains:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Figure 2.9: The plum pudding model of the atom.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Figure 2.10: A schematic representation of the apparatus Millikan used to determine the charge on the electron.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Figure 2.12: Rutherford's experiment on  -particle bombardment of metal foil.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Figure 2.13: (a) The expected results of the metal foil experiment if Thomson's model were correct. (b)Actual results.

Figure 2.14: A nuclear atom viewed in cross section. Note that this drawing is not to scale.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 The Mass and Change of the Electron, Proton, and Neutron

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 The Chemists’ Shorthand: Atomic Symbols K  Element Symbol Mass number  Atomic number 

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Figure 2.15: Two isotopes of sodium. Both have eleven protons and eleven electrons, but they differ in the number of neutrons in their nuclei.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 The Chemists’ Shorthand: Formulas Chemical Formula: Symbols = types of atoms Subscripts = relative numbers of atoms CO 2 Structural Formula: Individual bonds are shown by lines. O=C=OO=C=O

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Ions Cation: A positive ion Mg 2+, NH 4 + Anion: A negative ion Cl , SO 4 2  Ionic Bonding: Force of attraction between oppositely charged ions.

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Periodic Table Elements classified by:  properties  atomic number Groups (vertical) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases Periods (horizontal)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Figure 2.21: The Periodic Table.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Naming Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element Ca 2+ = calcium ion 3. Monatomic anion = root + -ide Cl  = chloride CaCl 2 = calcium chloride Binary Ionic Compounds:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Naming Compounds (continued)  metal forms more than one cation  use Roman numeral in name PbCl 2 Pb 2+ is cation PbCl 2 = lead (II) chloride Binary Ionic Compounds (Type II):

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Crystals of copper(II) sulfate.

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Various chromium compounds dissolved in water. From left to right; CrCl 2, K 2 Cr 2 O 7, Cr(NO 3 ) 3, CrCl 3, K 2 CrO 4.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Naming Compounds (continued)  Compounds between two nonmetals  First element in the formula is named first.  Second element is named as if it were an anion.  Use prefixes  Never use mono- P 2 O 5 = diphosphorus pentoxide Binary compounds (Type III):

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Figure 2.23: A flowchart for naming binary compounds.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 2.25: A flowchart for naming acids. An acid is best considered as one or more H+ ions attached to an anion.