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History of Atoms and Molecules

Explore the fascinating history of atoms and molecules, from ancient Greek theories to modern scientific discoveries. Learn about the different laws and experiments that laid the foundation for our understanding of chemistry.

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History of Atoms and Molecules

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. History • Greeks • Democritus and Leucippus - atomos • Aristotle - elements. • Alchemy • 1660 - Robert Boyle - experimental definition of element. • Lavoisier- Father of modern chemistry. • He wrote the book.

  3. Laws • Conservation of Mass • Law of Definite Proportion - compounds have a constant composition. • The react in specific ratios by mass. • Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

  4. Simple Ratios in Compounds • Water has 8 g of oxygen per g of hydrogen. • Hydrogen peroxide has 16 g of oxygen per g of hydrogen. • 16/8 = 2/1 • Small whole number ratios. • So the formula of Hydrogen Peroxide is H2O2

  5. Proof • Mercury has two oxides. One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Show that these compounds follow the law of multiple proportion. • Speculate on the formula of the two oxides.

  6. Dalton’s Atomic Theory 1. Elements are made up of atoms 2. Atoms of each element are identical. Atoms of different elements are different. (We need to address this one a bit further) 3. Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4. Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

  7. A Helpful Observation • Gay-Lussac - Under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. • Avogadro - interpreted that to mean at the same temperature and pressure, equal volumes of gas contain the same number of particles. • (called Avogadro’s Hypothesis)

  8. Experiments to determine what an atom was • J. J. Thomson - used Cathode ray tubes to study the electrical charges found in atoms.

  9. Voltage source Thomson’s Experiment - +

  10. Voltage source Thomson’s Experiment - +

  11. Voltage source Thomson’s Experiment - + • Passing an electric current makes a beam appear to move from the negative to the positive end.

  12. Voltage source Thomson’s Experiment • By adding an electric field

  13. Voltage source Thomson’s Experiment + - • By adding an electric field, he found that the moving pieces were negative, because they moved away from the negative charge.

  14. Thomsom’s Model • Found the electron. • Couldn’t find positive (for a while). • Said the atom was like plum pudding. • A bunch of positive stuff, with the electrons able to be removed.

  15. Atomizer Oil droplets + - Oil Telescope Millikan’s Experiment

  16. Millikan’s Experiment X-rays X-rays give some electrons a charge.

  17. Millikan’s Experiment Some drops would hover From the mass of the drop and the charge on the plates, he calculated the mass of an electron

  18. Radioactivity • Discovered by accident • Bequerel • Three types • alpha- helium nucleus (+2 charge, large mass) • beta- high speed electron • gamma- high energy light

  19. Rutherford’s Experiment • Used uranium to produce alpha particles. • Aimed alpha particles at gold foil by drilling hole in lead block and inserting a small bit of alpha emitting particles. • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin.

  20. Florescent Screen Lead block Uranium Gold Foil

  21. What he expected

  22. Notice that there is no deflection in the beam.

  23. Because, he thought the mass was evenly distributed in the atom.

  24. What he got was much different

  25. + How he explained it • Atom is mostly empty • Small dense, positive pieceat center. • Alpha particles are deflected by it if they get close enough. • Refer to the animations on the website for more detail.

  26. + The beam was often deflected

  27. Modern View • The atom is mostly empty space. • Two regions • Nucleus - protons & neutrons. • Electron cloud - region where you might find an electron.

  28. Sub-atomic Particles • Z - atomic number = number of protons determines type of atom. • A - mass number = number of protons + neutrons. (Sometimes you will see N) • Number of protons = number of electrons if neutral.

  29. Symbols A Charge goes here X Z +1 23 Na 11

  30. Chemical Bonds • The forces that hold atoms together. • Covalent bonding - sharing electrons. • Makes molecules. • Chemical formula - the number and type of atoms in a molecule. • C2H6 has 2 carbon atoms, 6 hydrogen atoms, • Structural formula shows the connections, but not necessarily the shape.

  31. H H • There are also other model that attempt to show three dimensional shape. • This is called a Ball and stick model. H C C H H H

  32. Ions • Atoms or groups of atoms with a charge. • Cations - positive ions - get by losing electrons(s). • Anions - negative ions - get by gaining electron(s). • Ionic bonding - held together by the opposite charges. • Ionic solids are called salts, the most notable one being Sodium Chloride or Table Salt (NaCl).

  33. Polyatomic Ions • Groups of atoms that have a charge. • Yes, you have to memorize them. • Lists are on the website and on the given sheet.

  34. Periodic Table

  35. Metals • Conductors • Lose electrons • Malleable and ductile

  36. Nonmetals • Brittle • Gain electrons • Covalent bonds

  37. Semi-metals or Metalloids

  38. Alkali Metals or Group I Metals

  39. Alkaline Earth Metals or Group II Metals

  40. Halogens or Group VII Metals

  41. Transition metals

  42. Noble Gases or Inert Gases or Group VIII

  43. Inner Transition Metals or Rare Earth Metals

  44. You can use the chart to help you with charges +1 +2 -3 -2 -1

  45. Naming compounds • There are basically 2 types of compounds but we can combine the rules. Ionic and Covalent • We start by looking at the location of the first element in the compound on the periodic chart. If it is to the left of the metalloid line, they we need to use prefixes. They are on another sheet. If it is between the metalloid line up to the first two columns (IIA) then we need Roman Numerals. These are summarized on your Ion Sheet. If it is in the first two columns on the right (IA and IIA), then we just write the ions as shown.

  46. More nomenclature

  47. Ionic Compounds - Practice • Have to know what ions they contain: • CaS K2S AlPO4 • K2SO4 FeS CoBr3 • Fe2(SO4)3 MgO MnO • KMnO4 NH4NO3 Hg2Cl2 • Cr2O3 KClO4 NaClO3 • AgNO3 Cr(ClO)6

  48. Prefixes • 1 – Mono - • 2 – Di - • 3 – Tri - • 4 – Tetra - • 5 – Penta - • 6 – Hexa - • 7 – Hepta - • 8 – Octa - • 9 – Nona - • 10 – Deca -

  49. Naming Covalent Compounds • Use the same rules as before ! • CO2 • CO • CCl4 • N2O4 • XeF6 • N4O4 • P2O10

  50. Guidelines • Charges must add up to zero. • This is true for every single compound. • Get charges from table, name of metal ion, or memorized from the list. • Use parenthesis to indicate multiple polyatomic ions. Example: Fe(NO3)2

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