Chapter 2

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. • They 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. What?! • 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.

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. If you have 100 g of the substances, then: Compound 1 contains 96.2 g Hg and 3.8 g O Compound 2 contains 92.6 g Hg and 7.4 g O Ratio of O in compound 1 to compound 2 is 3.8 : 7.4 or 1:2

6. Proof • Mercury has two oxides. One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Speculate on the formula of the two oxides. HgO and HgO2

7. Dalton’s Atomic Theory • Elements are made up of atoms • Atoms of each element are identical. Atoms of different elements are different. • Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. • Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

8. 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)

9. Experiments to determine what an atom was • J. J. Thomson- used Cathode ray tubes

10. Voltage source Thomson’s Experiment - +

11. Voltage source Thomson’s Experiment - +

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

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

14. Thomson’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. • 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. Because, he thought the mass was evenly distributed in the atom.

23. What he got

24. + How he explained it • Atom is mostly empty • Small dense, positive pieceat center. • Alpha particlesare deflected by it if they get close enough.

25. +

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

27. Sub-atomic Particles • Z - atomic number = number of protons determines type of atom. • A - mass number = number of protons + neutrons. • Number of protons = number of electrons if neutral.

28. A X Z 23 Na 11 Symbols

29. 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 - 2 carbon atoms, 6 hydrogen atoms, • Structural formula shows the connections, but not necessarily the shape.

30. H H • Structural Formula • There are also other models that attempt to show three dimensional shape. H C C H H H

31. Ball and stick. • Space-filling model

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.

33. Polyatomic Ions • Groups of atoms that have a charge. • Yes, you have to memorize them. • List given on first day – some on page 54

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

39. Alkaline Earth Metals

40. Halogens

41. Transition metals

42. Noble Gases

43. Inner Transition Metals

44. +1 +2 -3 -2 -1

45. Naming compounds • Two types • Ionic - metal and non metal or polyatomics. • Covalent- we will just learn the rules for 2 non-metals.

46. Naming Practice

47. Ionic compounds • If the cation is an element- Name the metal (cation) just write the name. • If the cation is polyatomic- name it • If the anion is an element- name it but change the ending to -ide • If the anion is poly atomic- just name it • practice

48. Ionic Compounds • Have to know what ions they form • off table, polyatomic, or figure it out • CaS • K2S • AlPO4 • K2SO4 • FeS • CoI3

49. Ionic Compounds • FeC2O4 • MgO • MnO2 • KMnO4 • NH4NO3 • HgCl2 • Cr2O3

50. Ionic Compounds • KClO4 • NaClO3 • YBrO2 • Cr(ClO)3