2. Atomic Structure Atoms and their structure Mr. Bruder

3. Laws n Conservation of Mass n Law of Definite Proportion- compounds have a constant composition by mass. n They react in specific ratios by mass. n 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?! n Water has 8 g of oxygen per g of hydrogen. n Hydrogen peroxide has 16 g of oxygen per g of hydrogen. n 16/8 = 2/1 n Small whole number ratios

5. Example of Law Of Multiple Proportions 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.

6. Your Turn n Nitrogen and oxygen form two compounds. Show that they follow the law of multiple proportions Amount N Amount O Compound A 1.206 g 2.755 g Compound B 1.651g 4.714 g

7. Dalton’s Atomic Theory n John Dalton (1766-1844) had four theories 1. All elements are composed of submicroscopic indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of anyone element are different from those of any other element 3. Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds 4. Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

8. Atoms & Subatomic Particles n Atom- smallest particle of an element that retains the properties of that element

9. n Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. n Avagadro- interpreted that to mean n at the same temperature and pressure, equal volumes of gas contain the same number of particles n (called Avagadro’s hypothesis) A Helpful Observation

10. Electron n J.J Thomson (1856-1940) – discovered the electron in 1897 n Electron is the negative charged subatomic particle n An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

11. Cathode Ray n The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

12. © 2009, Prentice-Hall, Inc. The Electron n Streams of negatively charged particles were found to emanate from cathode tubes. n J. J. Thompson is credited with their discovery (1897).

13. © 2009, Prentice-Hall, Inc. The Electron Thompson measured the charge/mass ratio of the electron to be 1.76  10 8 coulombs/g.

14. Thomson’s Atomic Model n Thomson’s Atomic Model n Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

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

16. Mass of Electron n In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment n He determined the mass to be 9.109 x 10 -31 kg n The oil drop apparatus

17. Millikan’s Experiment Atomizer Microscope - + Oil

18. Millikan’s Experiment Oil Atomizer Microscope - + Oil droplets

19. Millikan’s Experiment X-rays X-rays give some drops a charge by knocking off electrons

20. Millikan’s Experiment +

21. They put an electric charge on the plates + + --

22. Millikan’s Experiment Some drops would hover + + --

23. Millikan’s Experiment + ++ +++++ -- -----

24. Measure the drop and find volume from 4/3πr 3 Find mass from M = D x V + + --

25. Millikan’s Experiment From the mass of the drop and the charge on the plates, he calculated the charge on an electron + + --

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

27. Ernest Rutherford n Rutherford (1871-1937) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom n He used the Gold-Foil Experiment to prove his theory n In 1911 he discovered the Nucleus n Nucleus- central core of an atom, composed of protons and neutrons n The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

28. Rutherford’s Experiment n Used uranium to produce alpha particles n Aimed alpha particles at gold foil by drilling hole in lead block n Since the mass is evenly distributed in gold atoms alpha particles should go straight through. n Used gold foil because it could be made atoms thin

29. Lead block Uranium Gold Foil Florescent Screen

30. What he expected

31. Because

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

33. What he got

34. 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

35. +

36. Modern View n The atom is mostly empty space n Two regions n Nucleus- protons and neutrons n Electron cloud- region where you have a chance of finding an electron

37. Neutron n James Chadwick (1891-1974) – discovered the neutron in 1932 n Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

38. Quark n Protons & Neutrons can still be broken down into a smaller particle called the Quark n The Quark is held together by Gluons

39. Density and the Atom n Since most of the particles went through, it was mostly empty. n Because the pieces turned so much, the positive pieces were heavy. n Small volume, big mass, big density n This small dense positive area is the nucleus

40. Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n0 +1 0 1/1840 1 1 9.11 x 10 -28 1.67 x 10 -24

41. Symbols n Contain the symbol of the element, the mass number and the atomic number X Mass number Atomic number

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

43. Symbols X A Z Na 23 11

44. Atomic Structure Symbols n Proton = p + n Electron = e - n Neutron = n 0 n Atomic # - Subscript n Mass # - Superscript

45. Rules for Atomic Structure 1. Atomic # = # of Protons 2. # of Protons = # of Electrons 3. Mass # = # of Protons + # of Neutrons n # of Neutrons = Mass # - # of Protons n If you know the Mass # & Atomic # you know the composition of the element

46. Symbols n Find n Find the –number –number of protons of neutrons of electrons –Atomic –Atomic number –Mass –Mass Number Br 80 35

47. Isotopes n Isotope- atoms that have the same number of protons but different number of neutrons n Since isotopes have a different number of neutrons the isotope has a different mass number. n Isotopes are still chemically alike because they have the same number of protons and electrons

48. Examples of Isotopes

49. © 2009, Prentice-Hall, Inc. Isotopes n Isotopes are atoms of the same element with different masses. n Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

50. Naming Isotopes n Put the mass number after the name of the element n carbon- 12 n carbon -14 n uranium-235

51. Electrical Charges n Electrical charges are carried by particles of matter n Atoms have no net electrical charges n Given the number of negative charges combines with the number of positive charges = Electrically Neutral n All elements are Electrically Neutral

52. Atomic Mass vs. Atomic Weight n Atomic Mass is for a single element n Most elements are Isotopes n How do we find their mass? n We use Atomic Weight

53. © 2009, Prentice-Hall, Inc. Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

54. Measuring Atomic Mass n Unit is the Atomic Mass Unit (amu) n One twelfth the mass of a carbon-12 atom n Each isotope has its own atomic mass. We need the average from the percent abundance n Each isotope of an element has fixed mass and a natural % abundance n You need both of these values to find the Atomic Weight

55. Calculating Atomic Weight n Cl-35 34.969amu and 75.77% abundance n Cl-37 36.966amu and 24.23% abundance n To solve for Cl-35 1. AMU x Abundance 2. 34.969 x.7577 3. = 26.496 n You solve for Cl-37

56. Atomic Weight Cont. n Cl-37 1. AMU x Abundance 2. 36.966 x.2423 3. = 8.957 n Now you combine your two answers n 26.496 + 8.957= n 35.453 n Look at Cl on the table. What is the Atomic Weight?

57. Example n Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of 62.93 amu. The other has a mass of 64.93 amu. What is the atomic weight???

58. Atomic Weight & Decimals n Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element n Atomic Weights use decimal points because it is an average of an element

59. Periodic Table

60. © 2009, Prentice-Hall, Inc. Periodic Table n It is a systematic catalog of the elements. n Elements are arranged in order of atomic number.

61. © 2009, Prentice-Hall, Inc. Periodic Table n The rows on the periodic chart are periods. n Columns are groups. n Elements in the same group have similar chemical properties.

62. © 2009, Prentice-Hall, Inc. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

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

64. Nonmetals n Brittle n Gain electrons n Covalent bonds

65. Semi-metals or Metalloids

66. Chemical Nomenclature

67. © 2009, Prentice-Hall, Inc. Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

68. © 2009, Prentice-Hall, Inc. Chemical Formulas Molecular compounds are composed of molecules and almost always contain only nonmetals.

69. © 2009, Prentice-Hall, Inc. Types of Formulas n Empirical formulas give the lowest whole- number ratio of atoms of each element in a compound. n Molecular formulas give the exact number of atoms of each element in a compound.

70. © 2009, Prentice-Hall, Inc. Types of Formulas n Structural formulas show the order in which atoms are bonded. n Perspective drawings also show the three-dimensional array of atoms in a compound.

71. Octet Rule n Atoms tend to achieve electron configuration of Noble Gases n Octet = Eight n Noble Gases have eight electrons in their highest energy level n General Equation for Noble Gases is S 2 P 6

72. n Atoms of Metallic Elements tend to lose valence electron/s, leaving an octet in the next lowest energy level n Atoms of a Non-Metallic Element tend to gain a valence electron/s to achieve an Octet n There are EXCEPTIONS to the Octet Rule

73. © 2009, Prentice-Hall, Inc. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms.

74. 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

75. © 2009, Prentice-Hall, Inc. Ions n When atoms lose or gain electrons, they become ions. –Cations are positive and are formed by elements on the left side of the periodic chart. –Anions are negative and are formed by elements on the right side of the periodic chart.

76. © 2009, Prentice-Hall, Inc. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

77. © 2009, Prentice-Hall, Inc. Writing Formulas n Because compounds are electrically neutral, one can determine the formula of a compound this way: –The charge on the cation becomes the subscript on the anion. –The charge on the anion becomes the subscript on the cation. –If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

78. © 2009, Prentice-Hall, Inc. Common Cations

79. © 2009, Prentice-Hall, Inc. Common Anions

80. Polyatomic Ions n Polyatomic Ion- Tightly bound groups of atoms that behave as a unit and carry a charge n Unlike monatomic ions; Sulfate anion is composed of 1 Sulfur atom and 4 oxygen atoms n These five atoms form a Sulfate Anion n It has a –2 charge an is written SO 4 2-

81. n Polyatomic anions either end in ITE or ATE n Out of the two similar polyatomic ions, the polyatomic with less Oxygens ends in ite n Example: n Sulfite and Sulfate n Sulfite; SO 3 2- n Sulfate; SO 4 2-

82. n There are three exceptions to the Polyatomic Rule n 1) Ammonium NH 4 + ---- The only positive polyatomic ion n 2) Cyanide CN - ---- Ends in IDE n 3) Hydroxide OH - --- Ends in IDE

83. Stock System n The stock system uses roman numerals in ( ). The ( ) indicate the numerical charge of the cation. Example: Fe 2+ Name: Iron(II) There is no space between the name and the parenthesis Example: Cu 1+ Name: Copper(1)

84. Classical System n The classical system uses the root word with different suffixes as the end of the word n OUS- is used to name the cation with the lower of the two ionic charges n IC- is used to name the cation with the higher of the two ionic charges

85. n Example: n Fe 2+ and Fe 3+ n Name: Ferrous n Name: Ferric n What is the problem with the classical system?

86. © 2009, Prentice-Hall, Inc. Inorganic Nomenclature n Write the name of the cation. n If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. n If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

87. © 2009, Prentice-Hall, Inc. Patterns in Oxyanion Nomenclature n When there are two oxyanions involving the same element: –The one with fewer oxygens ends in -ite. »NO 2 − : nitrite ; SO 3 2− : sulfite –The one with more oxygens ends in -ate. »NO 3 − : nitrate; SO 4 2− : sulfate

88. © 2009, Prentice-Hall, Inc. Patterns in Oxyanion Nomenclature n The one with the second fewest oxygens ends in -ite. –ClO 2 − : chlorite n The one with the second most oxygens ends in -ate. –ClO 3 − : chlorate

89. © 2009, Prentice-Hall, Inc. Patterns in Oxyanion Nomenclature n The one with the fewest oxygens has the prefix hypo- and ends in -ite. –ClO − : hypochlorite n The one with the most oxygens has the prefix per- and ends in -ate. –ClO 4 − : perchlorate

90. © 2009, Prentice-Hall, Inc. Nomenclature of Binary Compounds n The less electronegative atom is usually listed first. n A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however).

91. © 2009, Prentice-Hall, Inc. Nomenclature of Binary Compounds n The ending on the more electronegative element is changed to -ide. –CO 2 : carbon dioxide –CCl 4 : carbon tetrachloride

92. © 2009, Prentice-Hall, Inc. Nomenclature of Binary Compounds n If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N 2 O 5 : dinitrogen pentoxide

93. © 2009, Prentice-Hall, Inc. Acid Nomenclature n If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro-. –HCl: hydrochloric acid –HBr: hydrobromic acid –HI: hydroiodic acid

94. © 2009, Prentice-Hall, Inc. Acid Nomenclature n If the anion in the acid ends in -ite, change the ending to -ous acid. –HClO: hypochlorous acid –HClO 2 : chlorous acid

95. © 2009, Prentice-Hall, Inc. Acid Nomenclature n If the anion in the acid ends in -ate, change the ending to -ic acid. –HClO 3 : chloric acid –HClO 4 : perchloric acid

96. 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.

98. 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

99. 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