Chapter 2 Atoms, Molecules and Ions

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.

The Early History of Chemistry
Before 16th Century Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold 17th Century Robert Boyle: First “chemist” to perform quantitative experiments 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.

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

Other Fundamental Chemical Laws
Law of Definite Proportion A given compound always contains exactly the same proportion of elements by mass. Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine: CCl4 Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Other Fundamental Chemical Laws
Law of Multiple Proportions 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. The ratio of the masses of oxygen in H2O and H2O2 will be a small whole number (“2”). Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Law of Multiple Proportions
2 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2.1

Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Mass of Nitrogen That Combines With 1 g Oxygen Compound A g Compound B g Compound C g A/B = 1.750/ = 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.

Dalton’s Atomic Theory (1808)
Each element is made up of tiny particles called atoms. The 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.

Dalton’s Atomic Theory (continued)
Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Chemical 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.

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.

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

Early Experiments to Characterize the Atom
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 108 C/g m Copyright©2000 by Houghton Mifflin Company. All rights reserved.

2.2

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

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.

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

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

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.

The Mass and Change of the Electron, Proton, and Neutron

Summary J.J. Thompson (1897) “cathode rays are electrons” (e–) and finds e/m ratio Robert Millikan (1909) measures e and hence melectron known at 9.1110-31 kg E. Rutherford (1906) bounces  (He2+) 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.

The Chemists’ Shorthand: Atomic Symbols
39 Mass number  K  Element Symbol 19 Atomic number  Copyright©2000 by Houghton Mifflin Company. All rights reserved.

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.

The Chemists’ Shorthand: Formulas
Chemical Formula: Symbols = types of atoms Subscripts = relative numbers of atoms CO2 Structural Formula: Individual bonds are shown by lines. O=C=O Copyright©2000 by Houghton Mifflin Company. All rights reserved.

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.

Figure 2.18: Ball-and-stick model of methane.

Ions Cation: A positive ion Mg2+, NH4+ Anion: A negative ion Cl, SO42 Ionic Bonding: Force of attraction between oppositely charged ions. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

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.

Earth’s Most Abundant Elements Oxygen 0 46% Silicon Si 27%
Aluminum Al 8% Iron Fe 5% Calcium Ca 3% Sodium Na 2.8% Potassium K 2.5% Magnesium Mg 2.0% Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Naming Compounds Binary Ionic Compounds: 1. Cation first, then anion
2. Monatomic cation = name of the element Ca2+ = calcium ion 3. Monatomic anion = root + -ide Cl = chloride CaCl2 = calcium chloride HI = hydrogen iodide Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Naming Compounds (continued)
Binary Ionic Compounds (Type II):  metal forms more than one cation  use Roman numeral in name PbCl2 Pb2+ is cation PbCl2 = lead (II) chloride Copyright©2000 by Houghton Mifflin Company. All rights reserved.

FeCl2 iron(II) chloride 2 Cl- -2 so Fe is +2 FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 2.22: The common cations and anions

Naming Compounds (continued)
Binary compounds (Type III):  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- P2O5 = diphosphorus pentoxide Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Molecular Compounds NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride TOXIC! NO2 nitrogen dioxide N2O dinitrogen monoxide Laughing Gas Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2.7

Figure 2.23: A flowchart for naming binary compounds.

Figure 2.24: Overall strategy for naming chemical compounds.

Naming Exercise Al2(S2O3)3 P4O10 Cu(NO2)2 NaMnO4 CS2 Fe2(CrO4)3 HCl (gas) PH4BrO2 Aluminum thiosulfate Tetraphosphorous decaoxide Copper(II) nitrite Sodium permanganate Carbon disulfide Iron(III) chromate Hydrogen chloride Phosphonium bromite Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Chapter 2 Atoms, Molecules and Ions

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