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8/31/08 2.1 The Early History of Chemistry Atoms, Molecules, and Ions It all started with the Greeks. (400 B.C.) Chapter 2 Demokritos theorized atomos = indivisible particles. Alchemy dominated for 2000 years. 2.2 Fundamental Chemical Laws 2.1 The Early History of Chemistry Systematic metallurgy is the foundation of modern chemistry. (1500’s) Extraction of metals from ore. Robert Boyle (1627-1691) Antoine Lavoisier (1743-1794) “Father of Modern Chemistry” Explained the true nature of combustion Verified the law of conservation of mass First to perform truly quantitative experiments. “Boyles Law” pressure v. volume 2.2 Fundamental Chemical Laws Joseph Proust (1754-1826) Law of definite proportion given compound always contains the same proportion of elements by mass 2.3 Dalton’s Atomic Theory John Dalton (1766-1844) Four Postulates (1808) Prepared the first table of atomic masses A John Dalton (1766-1844) “Psuedoscience” that tried to turn common metals into gold. We will get back to these later. Many values later proven wrong. Law of multiple proportions (see p. 43) Dalton could not determine absolute chemical formulas. 1 8/31/08 2.3 Dalton’s Atomic Theory Dalton’s Four Postulates (see p. 45) 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(s). 2.3 Dalton’s Atomic Theory 2.3 Dalton’s Atomic Theory Dalton’s Four Postulates Chemical compounds are formed when atoms of different elements combine with each other. A given chemical compound always has the same relative numbers and types of atoms. Chemical reactions involve reorganization of atoms—changes in the way they are bound together. The atoms themselves remain unchanged. 2.3 Dalton’s Atomic Theory Joseph Gay-Lussac (1778-1850) Performed experiments combining measured volumes of gases at constant temperature and pressure. 2.3 Dalton’s Atomic Theory Avogadro’s Hypothesis Avagadro’s Hypothesis (1811) At the same temperature and pressure, equal volumes of different gases contain the same number of particles. The volume of a gas is determined by the number of molecules present not the size of the individual particles. Figure 2.5 2 8/31/08 2.4 Early Atomic Experiments 2.4 Early Atomic Experiments J.J. Thomson (1856-1940) Used cathode-ray tubes When electricity is applied to the tube a “ray” is emitted from the cathode Figure 2.7 2.4 Early Atomic Experiments Thomson noted the ray was repelled by a negative electric field. 2.4 Early Atomic Experiments Thomson’s “Plum Pudding Model” He postulated the ray was a steam of negatively charged particles He named these particles “electrons.” Atoms are a diffuse cloud of positive charge with negative electrons embedded in it. Atoms are electrically neutral Must have positive particles = “protons” 2.4 Early Atomic Experiments Robert Milikan (1868-1953) “Oil Drop Experiment” Determined the mass of an electron 9.11 2.4 Early Atomic Experiments Ernest Rutherford (1871-1937) x 10-31 kg Tested the plum pudding model “Gold Foil Experiment” Shot radioactivity (α particles) at a thin piece of gold foil. He expected the heavy α particles to pass right through the foil with little, if any, deflection. What happened? 3 8/31/08 2.4 Early Atomic Experiments 2.4 Early Atomic Experiments Rutherford dismissed the plum pudding model. 2.5 Modern Atomic Structure Atoms must be mostly space with a center of positive charge. 2.5 Modern Atomic Structure Nucleus Contains “protons” Positively charged, equal in magnitude to the negative charge of an electron. Contains “neutrons” Very small Equal in mass to protons, but no charge. 2.5 Modern Atomic Structure 2.5 Modern Atomic Structure Why are atoms different from one another? Numbers of protons and electrons. Protons = identity = reactivity Electrons Numbers of neutrons Isotopes – atoms with the same number of protons, but different numbers of neutrons Figure 2.15 4 8/31/08 2.6 Molecules and Ions Chemical Bonds Covalent Bonds – electrons are shared between atoms Molecules 2.6 Molecules and Ions Structural Formula – indicates atoms and individual bonds have covalent bonds. into the page Chemical Formula – indicates the kind and numbers of atoms in a compound e.g. CO2 C6H12O6 H 2O out of page 2.6 Molecules and Ions Ion – an atom or group of atoms that has a net positive or negative charge Results from a gain (-) or loss (+) of electrons Cation – positive ion Anion – negative ion Na+ Cl- 2.6 Molecules and Ions Ionic Bonding – attraction between oppositely charged ions Called “ionic solids” or “salts” Cu2+ Fe3+ O2 - N3 - 2.6 Molecules and Ions Sodium Chloride 2.7 The Periodic Table Polyatomic Ions – “more than one atom ions” See Table 2.5, p. 66 You will need to memorize these! Name, formula, and Charge 5 8/31/08 2.7 The Periodic Table Metals Left and middle of table Shiny, good conductors, malleable Tend to lose electrons to form positive ions Non-metals Right side of table (and hydrogen) Gain electrons to form negative ions 2.7 The Periodic Table Families or groups – vertical columns Have similar chemical properties Alkali metals – Group 1A Alkali earth metals – Group 2A Halogens – Group 7A Noble gases – Group 8A Periods – horizontal rows 6