Your SlideShare is downloading. ×

Adv chem chapt 2

1,092

Published on

Published in: Technology, Education
0 Comments
0 Likes
Statistics
Notes
  • Be the first to comment

  • Be the first to like this

No Downloads
Views
Total Views
1,092
On Slideshare
0
From Embeds
0
Number of Embeds
1
Actions
Shares
0
Downloads
33
Comments
0
Likes
0
Embeds 0
No embeds

Report content
Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
No notes for slide

Transcript

  • 1. Advanced Chemistry Chapter 2 Atoms, Molecules, and Ions
  • 2. 2.1 Early Chemistry
  • 3. Early Chemistry
    • The Greeks – the first to try and explain chemical changes
      • Proposed that all matter was composed of four fundamental substances.
        • Fire
        • Earth
        • Water
        • Air
  • 4. Early Chemistry
    • Demokritos (also Greek) – believed that matter was composed of small, indivisible particles.
      • Atomos – term used to describe these particles.
      • Greeks were without the ability to test these hypothesis and no definite conclusion was reached.
    • Alchemy – predominate pseudoscience that dominated the next 2000 years.
  • 5. Modern Chemistry
    • The foundations were laid in the sixteenth century
    • The first chemist to perform quantitative experiments was Robert Boyle (1627-1691)
      • Published relationship between pressure and volume
        • Boyle’s Law
  • 6. 2.2 Fundamental Chemistry Laws
  • 7. Fundamental Chemical Laws
    • Law of Conservation of Mass – Mass is neither created nor destroyed in a chemical reaction.
      • Verified by Lavoisier (1743 – 1794) through his quantitative studies of combustion.
      • Lavoisier presented the most unified and complete knowledge of chemistry to date.
  • 8. Fundamental Chemical Laws
    • Law of Definite Proportion – a given compound always contains exactly the same proportion of elements by mass.
      • Proust’s (1754-1826) discovery was made through careful experiments regarding composition.
  • 9. 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 mass of the first element can always be reduced to small whole numbers.
      • Dalton (1766-1844) discovered that carbon and oxygen form two different compounds that contain different relative amounts of carbon and oxygen.
  • 10. Dalton’s Atomic Theory
  • 11. Dalton’s Atomic Theory
    • 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.
  • 12. Dalton’s Atomic Theory
    • 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.
    • 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.
  • 13. Dalton’s Table of Atomic Masses
    • Dalton prepared the first table of atomic masses
      • Many of the masses were later found to be wrong
      • Dalton still provided an important step with table construction.
  • 14. Gay-Lussac
    • Joseph Gay-Lussac (1778-1850) – provided the keys to determining the absolute formulas for compounds by his experimental work.
      • Performed experiments under the same conditions of temperature and pressure and determined the amount of gases that would react.
  • 15. Amadeo Avogadro
    • Avogadro’s hypothesis – at the same temperature and pressure, equal volumes of different gases contain the same number of particles.
      • The volume of gas is determined by the number of molecules present and not the size of the individual particles.
      • Avogadros’ number – 6.02 x 10 23
  • 16. 2.4 Early Experiments to Characterize the Atom
  • 17. The Electron
    • 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.
  • 18. Deflection of Cathode Rays by an Applied Electric Field
  • 19. Video
    • http://www.youtube.com/watch?v=7YHwMWcxeX8
  • 20. The Electron
    • Millikan discovered the Charge-to-mass ratio of an electron with his oil drop experiment.
      • e/m = -1.76 x 10 8 C/g
  • 21. A Schematic Representation of the Apparatus Millikan Used to Determine the Charge on the Electron
  • 22. The Electron
    • Plum Pudding Model – Thomson postulated that atoms must have a ‘cloud’ of positive charge in order to counter the random negatively charged electrons.
  • 23. The Plum Pudding Model of the Atom
  • 24. Radioactivity
    • Three types of radioactivity:
      • Gamma ( ϒ) rays – is high-energy light
      • Beta (β) particles – high speed electron
      • Alpha (α) particles – Helium nucleus (+2 charge)
  • 25. The Nuclear Atom
    • Rutherford tested Thomson’s plum pudding model by sending a stream of α particles through a thin sheet of metal foil.
      • The alpha particles were scattered and reflected concluding that a dense center of the atom existed.
  • 26. Rutherford's Experiment On a -Particle Bombardment of Metal Foil
  • 27. (a) Expected Results of the Metal Foil Experiment if Thomson's Model Were Correct (b) Actual Results
  • 28. 2.5 The Modern View of Atomic Structure: Intro
  • 29. Atomic Structure
    • Protons – 1 amu (1.673 x 10 -27 kg); +1 charge
    • Neutrons – 1 amu ; no charge
    • Electrons –(9.109 x 10 -31 kg) ; -1 charge
  • 30. Two Isotopes of Sodium
  • 31. Writing Symbols for Atoms
    • Atomic number (number of protons) – is written as a subscript.
    • Mass number ( total number of protons and neutrons) – written as a superscript.
  • 32. 2.6 Molecules and Ions
  • 33. Molecules
    • Chemical Bonds - the forces that hold atoms together in compounds
    • 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.
  • 34. Molecules
    • Structural formula – individual bonds are shown (indicated by lines) between element symbols.
    • Space- filling model – shows the relative sized of the atoms as well as their relative orientation in the molecule.
    • Ball-and-stick model – also used to represent molecules.
  • 35. Space-Filling Model of Methane
  • 36. Ball-and-Stick Model of Methane
  • 37. Ball-and-Stick Models of the Ammonium Ion and the Nitrate Ion
  • 38. The Structural Formula for Methane
  • 39. Molecules
    • Covalent bonds – sharing electrons to form a molecule .
    • Ionic bond – a force of attraction between oppositely charged ions that form a compound.
  • 40. Ions
    • Ion – is an atom or group of atoms that has a net positive or negative charge.
      • Anion – ion with a negative charge.
      • Cation – ion with a positive charge.
    • Polyatomic Ion – a compound consisting of many ions.
  • 41. 2.7 An Introduction to the Periodic Table
  • 42. Periodic Table
    • Most elements are metals.
      • Efficient conduction so heat and electricity
      • Malleability
      • Ductility
      • Often lustrous
      • Tend to lose electrons
  • 43. Periodic Table
    • Nonmetals – the relative few appear on the upper right corner of the table. (right of the heavy line)
      • Tend to gain electrons
      • Bond together forming covalent bonds.
  • 44. Periodic Table
    • Groups/Families – (vertical columns in table) Have similar chemical properties due to their similar atomic structure.
      • Alkali metals – group 1A – readily form +1 ions
      • Alkaline earth metals – group 2A – readily form +2 ions
      • Halogens – group 7A – all form diatomic molecules and react with metals to form salts – readily form -1 ions.
      • Noble gases – all exist under normal conditions as monatomic and have little chemical reactivity.
  • 45. Periodic Table
    • Periods – horizontal rows that represent the energy levels of the elements.
  • 46. The Periodic Table
  • 47. 2.8 Naming Simple Compounds
  • 48. Binary Ionic Compounds Type I
    • The cation is always named first and the anion second
    • A monatomic (“one atom”) cation takes its name from the name of the element.
      • Na + is sodium
    • A monatomic anion is named by taking the root of the element name and adding –ide.
      • Cl - is chloride
  • 49. Binary Ionic Examples
    • NaCl
    • KI
    • CaS
    • Li 3 N
    • CsBr
    • MgO
    • Sodium Chloride
    • Potassium Iodide
    • Calcium Sulfide
    • Lithium Nitride
    • Cesium Bromide
    • Magnesium Oxide
  • 50. Binary Examples
    • H – Hyrdrogen
      • H - = Hydride
      • LiH
        • Lithium Hydride
  • 51. Common Monatomic Cations and Anions
  • 52. Binary Ionic Compounds Type 2
    • The charge on the metal ion must be specified.
    • Roman numerals indicate the charge of the cation.
  • 53. Binary (Type 2) Examples
    • CuCl
    • HgO
    • Fe 2 O 3
    • MnO 2
    • PbCl 2
    • Copper (I) chloride
    • Mercury (II) oxide
    • Iron (III) oxide
    • Manganese (IV) oxide
    • Lead (II) chloride
  • 54. Common Cations and Anions
  • 55. Common Type 2 Cations
  • 56. Ionic Compounds: Polyatomic Ions
    • Oxyanions – anions that contain an atom of a given element and different numbers of oxygen atoms.
  • 57. Oxyanions
    • -ite – the name of the one with the smaller number of oxygen atoms.
      • SO 3 - Sulfite
    • -ate –the name of the one with the larger number of oxygen atoms
      • SO 4 - Sulfate
  • 58. Oxyanions
    • When more than two oxyanions make up a series the following prefixes apply:
        • Hypo- (less than) – the least amount of oxygen atoms
          • Hypochorite (note -ite)
        • Per- (more than) – the most amount of oxygen atoms
          • Perchlorate (note – ate)
  • 59. Polyatomic Ions
    • Na 2 SO 4
    • KH 2 PO 4
    • CsClO 4
    • NaOCl
    • Fe(NO 3 ) 3
    • Sodium sulfate
    • Potassium dihydrogen phosphate
    • Cesium perchlorate
    • Sodium hypochlorite
    • Iron (III) nitrate
  • 60. Common Polyatomic Ions
  • 61. Binary Covalent Compounds
    • The first element in the formula is named first, using the full element name.
    • The second element is named as if it were an anion
    • Prefixes are used to denote the numbers of atoms present
    • The prefix mono- is never used for naming the first element.
  • 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. Binary Covalent Compounds
    • N 2 O
    • NO 2
    • N 2 O 4
    • N 2 O 5
    • NO
    • Dinitrogen monoxide
    • Nitrogen dioxide
    • Dinitrogen tetroxide
    • Dinitrogen pentoxide
    • Nitrogen monoxide
  • 64. A Flowchart for Naming Binary Compounds
  • 65. Naming Chemical Compounds
  • 66. Acids
    • Binary acids – the acid is named with the prefix hydro- and the anion ends in ic. Add the name acid on the end.
  • 67. Names of Acids* that Do Not Contain Oxygen
  • 68. Acids with Oxygen
    • If the anion ends in –ate, the suffix –ic is added to the root.
    • If the anion has an –ite, the suffix –ous is added to the root.
    • Add the name acid to the end.
  • 69. Acids
    • HClO 4
    • HClO 3
    • HClO 2
    • HClO
    • HCl
    • Perchloric acid
    • Chloric acid
    • Chlorous acid
    • Hypochlorous acid
    • Hydrochloric
  • 70. Names of Some Oxygen-Containing Acids
  • 71. Figure 2.25 Naming Acids
  • 72. Figure 2.7 A Cathode-Ray Tube
  • 73. Plant is Newly Discovered Source of Gold
  • 74. Figure 2.1 The Priestley Medal is the Highest Honor Given by the American Chemical Society
  • 75. A Silicon Chip
  • 76. Atomic Nucleus
  • 77. Crystals of Copper(II) Sulfate
  • 78. Various Chromium Compounds Dissolved in Water
  • 79. Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron
  • 80. Table 2.2 The Symbols for the Elements That Are Based on the Original Names
  • 81. Table 2.6 Prefixes Used to Indicate Number in Chemical Names

×