Adv chem chapt 2


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Adv chem chapt 2

  1. 1. Advanced Chemistry Chapter 2 Atoms, Molecules, and Ions
  2. 2. 2.1 Early Chemistry
  3. 3. Early Chemistry <ul><li>The Greeks – the first to try and explain chemical changes </li></ul><ul><ul><li>Proposed that all matter was composed of four fundamental substances. </li></ul></ul><ul><ul><ul><li>Fire </li></ul></ul></ul><ul><ul><ul><li>Earth </li></ul></ul></ul><ul><ul><ul><li>Water </li></ul></ul></ul><ul><ul><ul><li>Air </li></ul></ul></ul>
  4. 4. Early Chemistry <ul><li>Demokritos (also Greek) – believed that matter was composed of small, indivisible particles. </li></ul><ul><ul><li>Atomos – term used to describe these particles. </li></ul></ul><ul><ul><li>Greeks were without the ability to test these hypothesis and no definite conclusion was reached. </li></ul></ul><ul><li>Alchemy – predominate pseudoscience that dominated the next 2000 years. </li></ul>
  5. 5. Modern Chemistry <ul><li>The foundations were laid in the sixteenth century </li></ul><ul><li>The first chemist to perform quantitative experiments was Robert Boyle (1627-1691) </li></ul><ul><ul><li>Published relationship between pressure and volume </li></ul></ul><ul><ul><ul><li>Boyle’s Law </li></ul></ul></ul>
  6. 6. 2.2 Fundamental Chemistry Laws
  7. 7. Fundamental Chemical Laws <ul><li>Law of Conservation of Mass – Mass is neither created nor destroyed in a chemical reaction. </li></ul><ul><ul><li>Verified by Lavoisier (1743 – 1794) through his quantitative studies of combustion. </li></ul></ul><ul><ul><li>Lavoisier presented the most unified and complete knowledge of chemistry to date. </li></ul></ul>
  8. 8. Fundamental Chemical Laws <ul><li>Law of Definite Proportion – a given compound always contains exactly the same proportion of elements by mass. </li></ul><ul><ul><li>Proust’s (1754-1826) discovery was made through careful experiments regarding composition. </li></ul></ul>
  9. 9. Fundamental Chemical Laws <ul><li>Law of Multiple Proportions – when two elements form a series of compounds, the ratios of the masses of the second element that combine with mass of the first element can always be reduced to small whole numbers. </li></ul><ul><ul><li>Dalton (1766-1844) discovered that carbon and oxygen form two different compounds that contain different relative amounts of carbon and oxygen. </li></ul></ul>
  10. 10. Dalton’s Atomic Theory
  11. 11. Dalton’s Atomic Theory <ul><li>Each element is made up of tiny particles called atoms. </li></ul><ul><li>The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. </li></ul>
  12. 12. Dalton’s Atomic Theory <ul><li>Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms. </li></ul><ul><li>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. </li></ul>
  13. 13. Dalton’s Table of Atomic Masses <ul><li>Dalton prepared the first table of atomic masses </li></ul><ul><ul><li>Many of the masses were later found to be wrong </li></ul></ul><ul><ul><li>Dalton still provided an important step with table construction. </li></ul></ul>
  14. 14. Gay-Lussac <ul><li>Joseph Gay-Lussac (1778-1850) – provided the keys to determining the absolute formulas for compounds by his experimental work. </li></ul><ul><ul><li>Performed experiments under the same conditions of temperature and pressure and determined the amount of gases that would react. </li></ul></ul>
  15. 15. Amadeo Avogadro <ul><li>Avogadro’s hypothesis – at the same temperature and pressure, equal volumes of different gases contain the same number of particles. </li></ul><ul><ul><li>The volume of gas is determined by the number of molecules present and not the size of the individual particles. </li></ul></ul><ul><ul><li>Avogadros’ number – 6.02 x 10 23 </li></ul></ul>
  16. 16. 2.4 Early Experiments to Characterize the Atom
  17. 17. The Electron <ul><li>Cathode-Ray Tube – a high voltage applied to a tube produced a ray (cathode ray). The ray was produced at the negative electrode and was repelled by the negative pole of an applied electric current. Thomson (1898-1903) concluded that the ray was a stream of negatively charged particles. </li></ul>
  18. 18. Deflection of Cathode Rays by an Applied Electric Field
  19. 19. Video <ul><li> </li></ul>
  20. 20. The Electron <ul><li>Millikan discovered the Charge-to-mass ratio of an electron with his oil drop experiment. </li></ul><ul><ul><li>e/m = -1.76 x 10 8 C/g </li></ul></ul>
  21. 21. A Schematic Representation of the Apparatus Millikan Used to Determine the Charge on the Electron
  22. 22. The Electron <ul><li>Plum Pudding Model – Thomson postulated that atoms must have a ‘cloud’ of positive charge in order to counter the random negatively charged electrons. </li></ul>
  23. 23. The Plum Pudding Model of the Atom
  24. 24. Radioactivity <ul><li>Three types of radioactivity: </li></ul><ul><ul><li>Gamma ( ϒ) rays – is high-energy light </li></ul></ul><ul><ul><li>Beta (β) particles – high speed electron </li></ul></ul><ul><ul><li>Alpha (α) particles – Helium nucleus (+2 charge) </li></ul></ul>
  25. 25. The Nuclear Atom <ul><li>Rutherford tested Thomson’s plum pudding model by sending a stream of α particles through a thin sheet of metal foil. </li></ul><ul><ul><li>The alpha particles were scattered and reflected concluding that a dense center of the atom existed. </li></ul></ul>
  26. 26. Rutherford's Experiment On a -Particle Bombardment of Metal Foil
  27. 27. (a) Expected Results of the Metal Foil Experiment if Thomson's Model Were Correct (b) Actual Results
  28. 28. 2.5 The Modern View of Atomic Structure: Intro
  29. 29. Atomic Structure <ul><li>Protons – 1 amu (1.673 x 10 -27 kg); +1 charge </li></ul><ul><li>Neutrons – 1 amu ; no charge </li></ul><ul><li>Electrons –(9.109 x 10 -31 kg) ; -1 charge </li></ul>
  30. 30. Two Isotopes of Sodium
  31. 31. Writing Symbols for Atoms <ul><li>Atomic number (number of protons) – is written as a subscript. </li></ul><ul><li>Mass number ( total number of protons and neutrons) – written as a superscript. </li></ul>
  32. 32. 2.6 Molecules and Ions
  33. 33. Molecules <ul><li>Chemical Bonds - the forces that hold atoms together in compounds </li></ul><ul><li>Chemical formula – symbols for the elements are used to indicate the types of atoms present and subscripts are used to indicate the relative numbers of atoms. </li></ul>
  34. 34. Molecules <ul><li>Structural formula – individual bonds are shown (indicated by lines) between element symbols. </li></ul><ul><li>Space- filling model – shows the relative sized of the atoms as well as their relative orientation in the molecule. </li></ul><ul><li>Ball-and-stick model – also used to represent molecules. </li></ul>
  35. 35. Space-Filling Model of Methane
  36. 36. Ball-and-Stick Model of Methane
  37. 37. Ball-and-Stick Models of the Ammonium Ion and the Nitrate Ion
  38. 38. The Structural Formula for Methane
  39. 39. Molecules <ul><li>Covalent bonds – sharing electrons to form a molecule . </li></ul><ul><li>Ionic bond – a force of attraction between oppositely charged ions that form a compound. </li></ul>
  40. 40. Ions <ul><li>Ion – is an atom or group of atoms that has a net positive or negative charge. </li></ul><ul><ul><li>Anion – ion with a negative charge. </li></ul></ul><ul><ul><li>Cation – ion with a positive charge. </li></ul></ul><ul><li>Polyatomic Ion – a compound consisting of many ions. </li></ul>
  41. 41. 2.7 An Introduction to the Periodic Table
  42. 42. Periodic Table <ul><li>Most elements are metals. </li></ul><ul><ul><li>Efficient conduction so heat and electricity </li></ul></ul><ul><ul><li>Malleability </li></ul></ul><ul><ul><li>Ductility </li></ul></ul><ul><ul><li>Often lustrous </li></ul></ul><ul><ul><li>Tend to lose electrons </li></ul></ul>
  43. 43. Periodic Table <ul><li>Nonmetals – the relative few appear on the upper right corner of the table. (right of the heavy line) </li></ul><ul><ul><li>Tend to gain electrons </li></ul></ul><ul><ul><li>Bond together forming covalent bonds. </li></ul></ul>
  44. 44. Periodic Table <ul><li>Groups/Families – (vertical columns in table) Have similar chemical properties due to their similar atomic structure. </li></ul><ul><ul><li>Alkali metals – group 1A – readily form +1 ions </li></ul></ul><ul><ul><li>Alkaline earth metals – group 2A – readily form +2 ions </li></ul></ul><ul><ul><li>Halogens – group 7A – all form diatomic molecules and react with metals to form salts – readily form -1 ions. </li></ul></ul><ul><ul><li>Noble gases – all exist under normal conditions as monatomic and have little chemical reactivity. </li></ul></ul>
  45. 45. Periodic Table <ul><li>Periods – horizontal rows that represent the energy levels of the elements. </li></ul>
  46. 46. The Periodic Table
  47. 47. 2.8 Naming Simple Compounds
  48. 48. Binary Ionic Compounds Type I <ul><li>The cation is always named first and the anion second </li></ul><ul><li>A monatomic (“one atom”) cation takes its name from the name of the element. </li></ul><ul><ul><li>Na + is sodium </li></ul></ul><ul><li>A monatomic anion is named by taking the root of the element name and adding –ide. </li></ul><ul><ul><li>Cl - is chloride </li></ul></ul>
  49. 49. Binary Ionic Examples <ul><li>NaCl </li></ul><ul><li>KI </li></ul><ul><li>CaS </li></ul><ul><li>Li 3 N </li></ul><ul><li>CsBr </li></ul><ul><li>MgO </li></ul><ul><li>Sodium Chloride </li></ul><ul><li>Potassium Iodide </li></ul><ul><li>Calcium Sulfide </li></ul><ul><li>Lithium Nitride </li></ul><ul><li>Cesium Bromide </li></ul><ul><li>Magnesium Oxide </li></ul>
  50. 50. Binary Examples <ul><li>H – Hyrdrogen </li></ul><ul><ul><li>H - = Hydride </li></ul></ul><ul><ul><li>LiH </li></ul></ul><ul><ul><ul><li>Lithium Hydride </li></ul></ul></ul>
  51. 51. Common Monatomic Cations and Anions
  52. 52. Binary Ionic Compounds Type 2 <ul><li>The charge on the metal ion must be specified. </li></ul><ul><li>Roman numerals indicate the charge of the cation. </li></ul>
  53. 53. Binary (Type 2) Examples <ul><li>CuCl </li></ul><ul><li>HgO </li></ul><ul><li>Fe 2 O 3 </li></ul><ul><li>MnO 2 </li></ul><ul><li>PbCl 2 </li></ul><ul><li>Copper (I) chloride </li></ul><ul><li>Mercury (II) oxide </li></ul><ul><li>Iron (III) oxide </li></ul><ul><li>Manganese (IV) oxide </li></ul><ul><li>Lead (II) chloride </li></ul>
  54. 54. Common Cations and Anions
  55. 55. Common Type 2 Cations
  56. 56. Ionic Compounds: Polyatomic Ions <ul><li>Oxyanions – anions that contain an atom of a given element and different numbers of oxygen atoms. </li></ul>
  57. 57. Oxyanions <ul><li>-ite – the name of the one with the smaller number of oxygen atoms. </li></ul><ul><ul><li>SO 3 - Sulfite </li></ul></ul><ul><li>-ate –the name of the one with the larger number of oxygen atoms </li></ul><ul><ul><li>SO 4 - Sulfate </li></ul></ul>
  58. 58. Oxyanions <ul><li>When more than two oxyanions make up a series the following prefixes apply: </li></ul><ul><ul><ul><li>Hypo- (less than) – the least amount of oxygen atoms </li></ul></ul></ul><ul><ul><ul><ul><li>Hypochorite (note -ite) </li></ul></ul></ul></ul><ul><ul><ul><li>Per- (more than) – the most amount of oxygen atoms </li></ul></ul></ul><ul><ul><ul><ul><li>Perchlorate (note – ate) </li></ul></ul></ul></ul>
  59. 59. Polyatomic Ions <ul><li>Na 2 SO 4 </li></ul><ul><li>KH 2 PO 4 </li></ul><ul><li>CsClO 4 </li></ul><ul><li>NaOCl </li></ul><ul><li>Fe(NO 3 ) 3 </li></ul><ul><li>Sodium sulfate </li></ul><ul><li>Potassium dihydrogen phosphate </li></ul><ul><li>Cesium perchlorate </li></ul><ul><li>Sodium hypochlorite </li></ul><ul><li>Iron (III) nitrate </li></ul>
  60. 60. Common Polyatomic Ions
  61. 61. Binary Covalent Compounds <ul><li>The first element in the formula is named first, using the full element name. </li></ul><ul><li>The second element is named as if it were an anion </li></ul><ul><li>Prefixes are used to denote the numbers of atoms present </li></ul><ul><li>The prefix mono- is never used for naming the first element. </li></ul>
  62. 62. Prefixes Used in Covalent Compounds Prefix Number Indicated mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 nona 9 deca 10
  63. 63. Binary Covalent Compounds <ul><li>N 2 O </li></ul><ul><li>NO 2 </li></ul><ul><li>N 2 O 4 </li></ul><ul><li>N 2 O 5 </li></ul><ul><li>NO </li></ul><ul><li>Dinitrogen monoxide </li></ul><ul><li>Nitrogen dioxide </li></ul><ul><li>Dinitrogen tetroxide </li></ul><ul><li>Dinitrogen pentoxide </li></ul><ul><li>Nitrogen monoxide </li></ul>
  64. 64. A Flowchart for Naming Binary Compounds
  65. 65. Naming Chemical Compounds
  66. 66. Acids <ul><li>Binary acids – the acid is named with the prefix hydro- and the anion ends in ic. Add the name acid on the end. </li></ul>
  67. 67. Names of Acids* that Do Not Contain Oxygen
  68. 68. Acids with Oxygen <ul><li>If the anion ends in –ate, the suffix –ic is added to the root. </li></ul><ul><li>If the anion has an –ite, the suffix –ous is added to the root. </li></ul><ul><li>Add the name acid to the end. </li></ul>
  69. 69. Acids <ul><li>HClO 4 </li></ul><ul><li>HClO 3 </li></ul><ul><li>HClO 2 </li></ul><ul><li>HClO </li></ul><ul><li>HCl </li></ul><ul><li>Perchloric acid </li></ul><ul><li>Chloric acid </li></ul><ul><li>Chlorous acid </li></ul><ul><li>Hypochlorous acid </li></ul><ul><li>Hydrochloric </li></ul>
  70. 70. Names of Some Oxygen-Containing Acids
  71. 71. Figure 2.25 Naming Acids
  72. 72. Figure 2.7 A Cathode-Ray Tube
  73. 73. Plant is Newly Discovered Source of Gold
  74. 74. Figure 2.1 The Priestley Medal is the Highest Honor Given by the American Chemical Society
  75. 75. A Silicon Chip
  76. 76. Atomic Nucleus
  77. 77. Crystals of Copper(II) Sulfate
  78. 78. Various Chromium Compounds Dissolved in Water
  79. 79. Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron
  80. 80. Table 2.2 The Symbols for the Elements That Are Based on the Original Names
  81. 81. Table 2.6 Prefixes Used to Indicate Number in Chemical Names