2. Chapter 2 Atoms, Molecules, and Ions

3. History n Greeks –Democritus –Leucippus - atomos –Aristotle- elements n Alchemy n Robert Boyle –1661 The Skeptical Chemist – 1 st “chemist” experimental definition of element. n Lavoisier- Father of modern chemistry. –Verified Law of Conservation of Mass

4. History - Laws n Joseph Proust –Law of Definite Proportion- compounds have a constant composition. –They react in specific ratios by mass. n John Dalton –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.

5. Multiple Proportions Example n H 2 O 8 g of oxygen 1g of hydrogen. –8:1 n H 2 O 2 16 g of oxygen 1g of hydrogen n Ratio of oxygen in H 2 O and H 2 O 2. 16:8 = 2:1 n Small whole number ratios.

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

7. History - Dalton’s Atomic Theory 1) Elements are made up of atoms 2) Atoms of each element are identical. Atoms of different elements are different. 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.  Created 1 st table of atomic masses

8. n Gay-Lussac: –Under STP, compounds always react in whole number ratios by volume. –2 vol. H (g) + 1 vol. O (g) form 2 vol. H 2 O (g) n Avogadro: –Interpreted Gay-Lussac to mean at STP, equal volumes of gas contain the same number of particles. –Avagadro’s Hypothesis Volume of gas is determined by number of particles NOT size of particles. History

9. History - Experiments To Determine Parts Of An Atom. n J. J. Thomson –Used Cathode ray tubes –Determined charge to mass ratio of an electron –e/m = -1.76 X 10 8 C/g e = charge in coulombs e = charge in coulombs m = mass of electron in grams m = mass of electron in grams

10. Thomson’s Experiment Voltage source +-

11. Thomson’s Experiment Voltage source +-

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

13. Thomson’s Experiment n By adding an electric field

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

15. Thomsom’s Model n Found the electron. n Assumed atoms had positive charge n Said the atom was like plum pudding. n A bunch of positive stuff, with the electrons able to be removed.

16. Robert Millikan’s Experiment Oil Atomizer Oil droplets Microscope - +

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

18. 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

19. History - Radioactivity n Bequerel - Discovered by accident n Three types –Alpha (  ) = helium nucleus (+2 charge, large mass) (+2 charge, large mass) –Beta (  ) - high speed electron –Gamma (  ) - high energy light

20. Rutherford’s Experiment n Used uranium to produce alpha particles. n Aimed alpha particles at gold foil by drilling hole in lead block. n Hypothesis: Since the mass is evenly distributed in gold atoms, alpha particles should go straight through based on Thomson’s model. n Used gold foil because it’s very malleable

21. Lead block Uranium Gold Foil Florescent Screen

22. What he expected

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

25. What he got

26. How he explained it + n Atom is mostly empty n Small dense, positive piece at center. n Alpha particles are deflected by it if they get close enough.

27. +

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

29. Sub-atomic Particles n Protons Positive particles in nucleus, determines type of element n Neutrons No charge, same mass as protons, in nucleus, determine isotopes n Electrons negative particles, in orbitals around nucleus, determines reactivity of atom n Z - atomic number = number of protons determines type of element. n A - mass number = number of protons + neutrons. n Number of protons = number of electrons if neutral.

30. Symbols X A Z Na 23 11

31. Chemical Bonds n The forces that hold atoms together. n Covalent bonding –Sharing electrons –Makes molecules n Ionic Bonding –Transferring electrons –Make Ionic compounds n Chemical formula- the number and type of atoms in a molecule. –CO 2 indicates 1 carbon atom and 2 oxygen atoms

32. H H HH H HCC n Structural formula shows the connections, but not necessarily the shape n Ball and stick show three dimensional shape Structural Formula

33. Ions n Atoms or groups of atoms with a charge. n Cations- positive ions - get by losing electrons(s). n Anions- negative ions - get by gaining electron(s). n Ionic bonding- held together by the opposite charges. n Ionic solids are called salts.

34. Polyatomic Ions n Groups of atoms that have a charge. n Yes, you have to memorize them. n Important polyatomic ions are on handout

35. Periodic Table

36. Metals n Conductors n Lose electrons n Malleable and ductile

37. Nonmetals n Brittle n Gain electrons n Covalent bonds

38. Semi-metals or Metalloids

39. Alkali Metals

40. Alkaline Earth Metals

41. Halogens

42. Transition metals

43. Noble Gases

44. Inner Transition Metals

45. +1+2-2-3

46. Naming compounds n Two types n Ionic - metal and non metal or polyatomics. n Covalent- two non-metals.

47. Binary Ionic Compounds Type I n Positive (cation) & negative (anion) n Cation is monoatomic with single charge, written 1 st, name the metal n Anion is monoatomic, written 2 nd, name it but change the ending to –ide.

48. Binary Ionic Compounds Type II n Metal forms more than one type of cation n Name must include Roman numeral to indicate the charge on cation. n Be aware of older system for naming –Name with –ic ending is has higher charge –Name with –ous ending has lower charge –Example FeCl 3 is ferric chloride and FeCl 2 is ferrous chloride

49. Ionic Compounds w/ Polyatomic Ions n Polyatomic ions contain more than one atom & have a charge n Must memorize n Oxyanions: series of anions with different number of oxygens atoms –Anion with fewer number of oxygens ends in –ite SO 3 is sulfite –Anion with greater number of oxygens ends in –ate SO 4 is sulfate

50. Ionic Compounds Practice n CaS nK2SnK2SnK2SnK2S n Al(PO 4 ) n K 2 ( SO 4 ) n FeS n CoI 3

51. Ionic Compounds Practice n Fe 2 (C 2 O 4 ) n MgO n MnO n KMnO 4 n NH 4 NO 3 n (Hg 2 ) Cl 2 n Cr 2 O 3

52. Ionic Compounds Practice n K(ClO 4 ) n Na(ClO 3 ) n Y(BrO 2 ) n Cr(ClO) 6

53. Binary Covalent Compounds Type III n Two words, with prefixes n Prefixes tell you how many. n Mono (1), di (2), tri (3), tetra (4), penta (5), hexa (6), septa (7), octa (8), nona (9), deca (10) n First element whole name with the appropriate prefix, except mono n Second element, -ide ending with appropriate prefix

54. n CO 2 n CO n CCl 4 nN2O4nN2O4nN2O4nN2O4 n XeF 6 nN4O4nN4O4nN4O4nN4O4 n P 2 O 10 Naming Covalent Compounds

55. Writing Formulas n Two sets of rules, ionic and covalent n To decide which to use, decide what the first element is. n If is a metal or polyatomic use ionic. n If it is a non-metal use covalent.

56. Ionic Formulas n Charges must add up to zero. n Get charges from table, name of metal ion, or memorized from the list. n Use parenthesis on all polyatomics.

57. Ionic Formulas n Sodium nitride –sodium Na is always +1 –Nitride -ide tells you it comes from the table –nitride is N -3 –One of each element doesn’t add up to zero charge –Need 3 Na to balance 1 nitride Na +1 N -3 Na 3 N

58. Ionic Compounds Practice n Sodium sulfite n Calcium iodide n Lead (II) oxide n Lead (IV) oxide n Mercury (I) sulfide n Barium chromate n Aluminum hydrogen sulfate n Cerium (IV) nitrite

59. Covalent compounds n The name tells you how to write the formula n Sulfur dioxide n diflourine monoxide n nitrogen trichloride n diphosphorus pentoxide

60. Acids n Substances that produce H + ions when dissolved in water. n Acids usually with H n Two types of acids: –Oxyacids Table 2.8 page 67 –Non-oxyacids Table 2.7 page 67

61. Naming acids n If the formula has oxygen in it –write the name of the anion, but change »ate polyatomic ion to -ic acid »ite polyatomic ion to -ous acid n Watch out for sulfuric and sulfurous n H 2 ( CrO 4 ) n H(MnO 4 ) n H(NO 2 )

62. Naming acids n If the acid doesn’t have oxygen n add the prefix hydro- n change the suffix -ide to -ic acid n HCl nH2SnH2SnH2SnH2S n HCN

63. Formulas for acids n Backwards from names. n If it has hydro- in the name it has no oxygen n Anion ends in -ide n No hydro, anion ends in -ate or -ite n Write anion and add enough H to balance the charges.

64. Formulas for acids n hydrofluoric acid n dichromic acid n carbonic acid n hydrophosphoric acid n hypofluorous acid n perchloric acid n phosphorous acid

65. Hydrates n Some salts trap water crystals when they form crystals. n These are hydrates. n Both the name and the formula needs to indicate how many water molecules are trapped. n In the name we add the word hydrate with a prefix that tells us how many water molecules.

66. Hydrates n In the formula you put a dot and then write the number of molecules. Calcium chloride dihydrate = CaCl 2  2  Calcium chloride dihydrate = CaCl 2  2  Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O