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Chapter 2

Chapter 2. Atoms, Molecules, and Ions. Chapter #2 – Atoms, Molecules and Ions. 2.1 The Early History of Chemistry 2.2 Fundamental Chemical laws 2.3 Dalton’s Atomic Theory 2.4 Cannizzaro’s Interpretation 2.5 Early experiments to Characterize the Atom

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Chapter 2

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. Chapter #2 – Atoms, Molecules and Ions 2.1 The Early History of Chemistry 2.2 Fundamental Chemical laws 2.3 Dalton’s Atomic Theory 2.4 Cannizzaro’s Interpretation 2.5 Early experiments to Characterize the Atom 2.6 The Modern View of Atomic Structure: An Introduction 2.7 Molecules and Ions 2.8 An Introduction to the Periodic Table 2.9 Naming Simple Compounds

  3. Computer simulation of the interior view of a twisted nanotube.

  4. Priestley Medal Source: Roald Hoffman, Cornell University

  5. Laws of Mass Conservation & Definite Proportions (Composition) Law of Mass Conservation: The total mass of substances does not change during a chemical reaction. Law of Definite ( or constant ) Composition: No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass.

  6. Figure 2.2: John Dalton Source: Manchester Literary & Philosophical Society

  7. Mass of Oxygen that Combines with 1.00g of Carbon Compound #1 1.33g Compound #2 2.66g

  8. Law of Multiple Proportions If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers: Example: Nitrogen Oxides I & II Nitrogen Oxide I : 46.68% Nitrogen and 53.32% Oxygen Nitrogen Oxide II : 30.45% Nitrogen and 69.55% Oxygen in 100 g of each Cpd: g O = 53.32 g & 69.55 g g N = 46.68 g & 30.45 g g O /g N = 1.142 & 2.284 2.284 2 = 1.142 1

  9. Mass of Nitrogen that Combines with 1.00g of Oxygen Compound #1 1.750 g Compound #2 0.8750 g Compound #3 0.4375 g I 1.750 2 Cpd #1 N2O NO N4O2 II 0.8750 1 II 0.8750 2 Cpd #2 NO or NO2 or N2O2 III 0.4375 1 I 1.750 4 Cpd #3 NO2 NO4 N2O4 III 0.4375 1 = = = = = =

  10. Dalton’s Atomic Theory Postulates: 1. Each element is made up of tiny particles called atoms. 2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. 3. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. 4. 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.

  11. Figure 2.3 (P19): Combining gases on a Molecular Level

  12. Avogadro’s Hypothesis At the same temperature and Pressure, equal volumes of different gases contain the same number of particles (Molecules).

  13. Stanislao Cannizzaro Source: Corbis

  14. Cannizzaro’s Relative Atomic(Molecular) Masses of Carbon and Hydrogen Compound Relative Percent Carbon Relative mass of Molecular Mass by Mass Carbon Present Methane 16 75 12 Ethane 30 80 24 Propane 44 82 36 Butane 58 83 48 Carbon Dioxide 44 27 12 Compound Relative Percent Hydrogen Relative mass of Molecular Mass by Mass Hydrogen Present Methane 16 25 4 Ethane 30 20 6 Propane 44 18 8 Butane 58 17 10

  15. Comparison of Several of Berzelius’s Atomic Masses with Current Values Element Atomic Mass Berzelius’s Value Current Value Chlorine 35.41 35.45 Copper 63.00 63.55 Hydrogen 1.00 1.01 Lead 207.12 207.2 Nitrogen 14.05 14.01 Oxygen 16.00 16.00 Potassium 39.19 39.10 Silver 108.12 107.87 Sulfur 32.18 32.07

  16. Figure 2.4: An STM image of nickel atoms placed on a copper surface. Source: IBM Research

  17. Figure 2.5: Image of a ring of cobalt atoms placed on a copper surface. Source: IBM Research

  18. Figure 2.6: A cathode-ray tube. The fast-moving electrons

  19. Figure 2.7: Deflection of cathode rays by an applied electric field.

  20. Figure 2.8: (P24)Thomson’sPlum Pudding model

  21. Figure 2.9: Schematic representation of the apparatus Millikan

  22. Marie Sklodowska Curie Source: Corbis

  23. Rutherford Experiment • Alpha particles bombarding the atom. • Rationale - to study the internal structure of the atom, and to know more about the mass distribution in the atom! • Bombarded a thin Gold foil with Alpha particles from Radium.

  24. Figure 2.11 (P25): Rutherford’s experiment

  25. Figure 2.12: The expected results of the metal foil experiment

  26. Ernest Rutherford (1871-1937) • Won the Nobel Prize in Chemistry in 1908 • “It was quite the most incredible event..... It was almost as if a gunner were to fire a shell at a piece of tissue and the shell bounced right back!!!!! ”

  27. Figure 2.13 (P26): Nuclear atom cross section

  28. Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. Although atoms are composed of smaller particles (electrons, protons, and neutrons), the atom is the smallest body thatretains the unique identity of the element. 2. Atoms of one element cannot be converted into atoms of another element ina chemical reaction. Elements can only be converted into other elements in Nuclear reactions in which protons are changed. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number, but a sample of the element is treated as though its atoms have an average mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios, as originally stated by Dalton.

  29. Atomic Definitions I: Symbols, Isotopes,Numbers A X The Nuclear Symbol of the Atom, or Isotope Z X = Atomic symbol of the element, or element symbol A = The Mass number; A = Z + N Z = The Atomic Number, the Number of Protons in the Nucleus N = The Number of Neutrons in the Nucleus Isotopes = atoms of an element with the same number of protons, but different numbers of Neutrons in the Nucleus

  30. Table 2.2 (P 27) The Masses and Charges of the Electron Proton and Neutron • Particle Mass Charge* • Electron 9.11 x 10 –31 kg -1 • Proton 1.67 x 10 – 27 kg +1 • Neutron 1.67 x 10 – 27 kg none • The magnitude of the charge on the electron and proton is • 1.60 x 10-19 coulombs .

  31. Figure 2.14(P28) Isotopes of sodium

  32. Neutral ATOMS • 51 Cr = P+ (24), e- (24), • N (27) • 239 Pu = P+(94), e- (94), • N (145) • 15 N = P+(7), e-(7), N(8) • 56 Fe = P+(26), e-(26), • N (30) • 235 U =P+(92), e-(92), • N (143)

  33. Definitions for Components of Matter Pure Substances - Their compositions are fixed! Elements and compounds are examples of Pure Substances. Element - Is the simplest type of substance with unique physical and chemical properties. An element consists of only one type of atom. It cannot be broken down into any simpler substances by physical or chemical means. Molecule - Is a structure that is consisting of two or more atoms that are chemically bound together and thus behaves as an independent unit. Compound - Is a substance composed of two or more elements that are chemically combined. Mixture - Is a group of two or more elements and/or compounds that are physically intermingled.

  34. Figure 2.15: Space-filling model of the methane molecule

  35. Figure 2.17 : Ball-and-stick model

  36. Chemical Formulas Empirical Formula- Shows the relative number of atoms of each element in the compound. It is the simplest formula, and is derived from masses of the elements. Molecular Formula - Shows the actual number of atoms of each element in the molecule of the compound. Structural Formula - Shows the actual number of atoms, and the bonds between them ; that is, the arrangement of atoms in the molecule.

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