Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Chapter 2 Atoms, Molecules and Ions Preview: Fundamental Chemical Laws and Atom. Modern View of Atomic Structure, Molecules, and Ions. Periodic Table. Naming Simple compounds, Ionic compounds, Formula from names.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 The Early History of Chemistry 4 Before 16th Century – Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold 4 17th Century –Robert Boyle: First “chemist” to perform quantitative experiments 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 Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine: CCl 4 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

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Mass of Nitrogen That Combines With 1 g Oxygen Compound A1.750 g Compound B g Compound C g A/B = 1.750/2 = 2/1 B/C = 0.875/ = 2/1 A/C = 1.750/ = 4/1 i.e. amount of nitrogen in A is twice that in B, etc.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 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. 9 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. 10 Dalton's theory lead to: 1gm hydrogen + 8gm of oxygen water he assumed that water formula is "OH" and the mass of hydrogen is "1" and of oxygen is "8". Using the same concepts, Dalton's proposed the first table of atomic masses. It has been proved later that Dalton's table contain incorrect. Gay-lussac ( ) found experimentally that:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Figure 2.4: A representation of some of Gay-Lussac's experimental results on combining gas volumes. Interpreted in 1811 by Avogadro

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 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. 13 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. 14 Early Experiments to Characterize the Atom H J. J. Thomson - postulated the existence of electrons using cathode ray tubes. Measured mass/charge of e- Received 1906 Nobel Prize in Physics e = x 10 8 C/g m

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Figure 2.7: A cathode-ray tube. The fast-moving electrons excite the gas in the tube, causing a glow between the electrodes.

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 (Uranium compound)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Figure 2.9: Thomson plum pudding model of the atom.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 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. 20 Figure 2.12: Rutherford's experiment on  -particle bombardment of metal foil.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Figure 2.13: (a) The expected results of the metal foil experiment if Thomson's model were correct. (b) Actual results.Actual results

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Millikan oil drop experiment Millikan did another experiment to determine the mass of the –ve particles (electrons). The experiment used mainly to determine the magnitude of the electron charge and using e/m to get m- value.value

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 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. 24 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. 25 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. 26 The Mass and Change of the Electron, Proton, and Neutron

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Summary J.J. Thompson (1897) “cathode rays are electrons” (e – ) and finds e/m ratio Robert Millikan (1909) measures e and hence m electron known at 9.11  kg E. Rutherford (1906) bounces  (He 2+ ) off Au tissue proving protons (p + ) in nucleus F.A. Aston (1919) “weighs” atomic ions J. Chadwick (1939) observes neutrons (no charge) by decomposition (to p +, e –, and ).

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 The Chemists’ Shorthand: Atomic Symbols K  Element Symbol Mass number  Atomic number 

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Chemical Bonds The forces that hold atoms together in compounds. Covalent bonds result from atoms sharing electrons. Molecule: a collection of covalently-bonded atoms.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 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

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Figure 2.16: The structural formula for methane.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Figure 2.17: Space-filling model of methane. This type of model shows both the relative sizes of the atoms in the molecule and their spatial relationships.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Figure 2.18: Ball-and-stick model of methane.

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 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.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 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. 37 Figure 2.21: The Periodic Table.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 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 HI = hydrogen iodide Binary Ionic Compounds:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Figure 2.19: Sodium metal reacts with chlorine gas to form solid sodium chloride.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 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. 41 FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S so Cr is +3 (6/2)chromium(III) sulfide

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Figure 2.22: The common cations and anions

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 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):

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 NF 3 nitrogen trifluoride SO 2 sulfur dioxide N 2 Cl 4 dinitrogen tetrachloride NO 2 nitrogen dioxide N2ON2Odinitrogen monoxide Molecular Compounds 2.7 TOXIC ! Laughing Gas

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Figure 2.23: A flowchart for naming binary compounds.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Figure 2.24: Overall strategy for naming chemical compounds.

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Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Figure 2.25: A flowchart for naming acids. An acid is best considered as one or more H+ ions attached to an anion.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 Naming Exercise Al 2 (S 2 O 3 ) 3 P 4 O 10 Cu(NO 2 ) 2 NaMnO 4 CS 2 Fe 2 (CrO 4 ) 3 HCl (gas) PH 4 BrO 2 Aluminum thiosulfate Tetraphosphorous decaoxide Copper(II) nitrite Sodium permanganate Carbon disulfide Iron(III) chromate Hydrogen chloride Phosphonium bromite