1. Atoms, Molecules and Ions Chapter 2 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Atomic Theory Democritus – teacher, 4 th Century BC Atomist school of thought Matter is composed of tiny particles called atoms These atoms are invisible, indestructible fundamental units of matter

3. John Dalton English school teacher Studied the ratios in which elements combine in chemical reactions formulated Dalton’s Atomic Theory

4. 4 Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. 2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. 3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

5. 5 Dalton’s Atomic Theory Law of Definite Proportions (Proust-1799) Mass Ratio 16:12 Mole Ratio 1:1 3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.

6. 6 Dalton’s Atomic Theory Law of Multiple Proportions Mass Ratio 16:12 Mole Ratio 1:1 Mass Ratio 32:12 Mole Ratio 2:1

7. 7 8 X 2 Y 16 X8 Y + Law of Conservation of Mass 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

8. John Joseph Thomson (1856-1940) English Physicist 1897 – discovered electrons: negatively charged subatomic particles Experiments using flow of electric current through gases

9. 9 J.J. Thomson, measured mass/charge of e - (1906 Nobel Prize in Physics) Cathode Ray Tube

10. 10 Cathode Ray Tube

11. 11 Thomson’s Model

12. Robert Andrews Millikan (1868-1953) American Physicist 1923 – Nobel prize in physics for determining the charge of the electron. Experiments using a single drop of oil with a static charge.

13. 13 e - charge = -1.60 x 10 -19 C Thomson’s charge/mass of e - = -1.76 x 10 8 C/g e - mass = 9.10 x 10 -28 g Measured mass of e - (1923 Nobel Prize in Physics) Millikan’s Experiment

14. Wilhelm Konrad Röntgen (1845-1923) German Physicist 1901 – Nobel prize in physics for discovery of x-rays. Observed that the cathode ray caused glass and metals to emit unusual rays. Becquerel and Curie continued experiments with radioactivity

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16. 16 Thomson’s Model

17. Ernest Rutherford (1871-1937) University of Manchester, England Tested theory of atomic structure Bombarded thin gold foil with a beam of ‘alpha’ particles. If the positive charge was evenly spread out, the beam should have easily passed through.

18. 18 1.atoms positive charge is concentrated in the nucleus 2.proton (p) has opposite (+) charge of electron (-) 3.mass of p is 1840 x mass of e - (1.67 x 10 -24 g)  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) Rutherford’s Experiment “It was as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.” (1908 Nobel Prize in Chemistry)

19. 19 Chadwick’s Experiment (1932) (1935 Noble Prize in Physics) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4  + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10 -24 g

20. 20 atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

21. 21 mass p ≈ mass n ≈ 1840 x mass e -

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23. 23 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei X A Z H 1 1 H (D) 2 1 H (T) 3 1 U 235 92 U 238 92 Mass Number Atomic Number Element Symbol Atomic number, Mass number and Isotopes

24. Isotopes 24 Elements can also be represented using name and mass number Carbon – 12 How many protons? How many neutrons?

25. Problem 2.14 Calculate the number of neutrons in an atom of. 25

26. 26 The Isotopes of Hydrogen

27. 27 How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ?

28. Problem 2.16 Indicate the number of protons, neutrons, and electrons in each of the following species: ___ protons ____neutrons ____electrons 28

29. Problem 2.16 Indicate the number of protons, neutrons, and electrons in each of the following species: ___ protons ____neutrons ____electrons 29

30. Rubidium has two common isotopes, 85 Rb and 87 Rb. If the abundance of 85 Rb is 72.2% and the abundance of 87 Rb is 27.8%, what is the average atomic mass of rubidium?

31. Uranium has three common isotopes. If the abundance of 234 U is 0.01%, the abundance of 235 U is 0.71%, and the abundance of 238 U is 99.28%, what is the average atomic mass of uranium? 31

32. Titanium has five common isotopes: 46 Ti (8.0%), 47 Ti (7.8%), 48 Ti (73.4%), 49 Ti (5.5%), 50 Ti (5.3%) What is the average atomic mass of titanium? 32

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34. The Periodic Table Dmitri Mendeleev (1834-1907) –Russian chemist –Listed elements in columns in order of increasing atomic mass. –Arranged columns so that elements with similar properties were side by side.

35. The Periodic Table Medeleev left blank spaces where there were no known elements with the appropriate properties or mass. Predicted the properties of the missing elements.

36. The Periodic Table Henry Mosely (1887-1915) –British Physicist –Determined the atomic number of the atoms of the elements. –Arranged elements in table by atomic number instead of mass.

37. The Modern Periodic Table Each horizontal row is a period –Seven periods –From 2 to 32 elements in a period –Properties of the elements change as you move across a period. –This pattern repeats from period to period The Periodic Law

38. The Modern Periodic Table Each column is a group or family –Elements in a group have similar physical and chemical properties. –Groups are identified by a number and the letter A or B –Group A are the representative elements –Group A can be divided into three broad classes

39. The Modern Periodic Table Metals –High electrical conductivity –High luster when clean –Ductile –malleable

40. The Modern Periodic Table Metals –Alkali metals –Alkaline earth metals –Transition metals –Inner transition metals

41. The Modern Periodic Table Non-Metals –poor electrical conductivity –Non-lustrous

42. The Modern Periodic Table Non-Metals –halogens –Noble gases

43. The Modern Periodic Table Metalloids –Elements with properties that are intermediate between those of metals and non-metals.

44. 44 The Modern Periodic Table Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal Transition Metals Inner Transition Metals

45. 45 Definitions Orbital – Area around the nucleus where electrons reside Valence electrons – electrons in the outermost orbital Octet Rule – to be “happy” atoms want 8 valence electrons – LOW ENERGY An ion is an atom, or group of atoms, that has a net positive or negative charge due to a loss or gain of electrons Where are the electrons in the atom?

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47. 47 An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na ___ protons ___ electrons P ___ protons ___ electrons P ___ protons ___ electrons 11 10 + 15 18 3-

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49. 49 Ar

50. 50 Common Ions Shown on the Periodic Table H+H+ In 3+ Ga 3+

51. 51 Common Ions Shown on the Periodic Table H+H+ In 3+ Ga 3+ 1+ 2+ 1-2-3- 3+

52. 52 Common Ions Shown on the Periodic Table H+H+ In 3+ Ga 3+ 1+ 2+ 1-2-3- 3+

53. 53 A monatomic ion contains only one atom A polyatomic ion contains more than one atom Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- OH -, CN -, NH 4 +, NO 3 -

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55. 55 ionic compounds consist of a combination of cations and anions (metal and a nonmetal) The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl

56. 56 Na Cl Will magnesium and chlorine form the octet together?

57. 57 The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds.

58. 58 Formula of Ionic Compounds Al 2 O 3 2 x +3 = +63 x -2 = -6 Al 3+ O 2- CaBr 2 1 x +2 = +22 x -1 = -2 Ca 2+ Br - Na 2 CO 3 1 x +2 = +21 x -2 = -2 Na + CO 3 2-

59. 59 Write the formula for the ionic compound chromium(III) sulfate. Write the formula for the ionic compound titanium(IV) oxide.

60. 60 A molecule is an aggregate of two or more atoms in a definite arrangement held together by covalent bonds H2H2 H2OH2ONH 3 CH 4 A diatomic molecule contains only two atoms H 2, N 2, O 2, Br 2, HCl, CO A polyatomic molecule contains more than two atoms O 3, H 2 O, NH 3, CH 4 diatomic elements

61. 61 Formulas and Models

62. 62 1 - carbon 1 – oxygen 4 - hydrogen Write the molecular formula for methanol based on the figure below. Write the molecular formula for chloroform based on the figure below. CH 4 O or CH 3 OH 1 - carbon 3 – chlorine 1 - hydrogen CHCl 3

63. 63 A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2OH2O H2OH2O molecular empirical O3O3 O N2H4N2H4 NH 2 C 6 H 12 O 6 CH 2 O Covalent Compounds Nonmetal + nonmetal

64. 64 Write the empirical formula for each of the following molecules: CH C2H2C2H2 molecularempirical N2ON2ON2ON2O C2H6C2H6 CH 3 C 8 H 10 N 4 O 2 C4H5N2OC4H5N2O Acetylene Nicotine Nitrous oxide Ethane

65. Problem 2.46 What are the empirical formulas of the following compounds? a)Al 2 Br 6 b)Na 2 S 2 O 4 c)N 2 O 5 d)K 2 Cr 2 O 7 65

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67. Ask Yourself… 67 Binary or tertiary? –Binary – composed of 2 different elements MgBr 2, H 2 O, CrO –Tertiary – composed of 3 or more different elements BaSO 4, AgNO 3 Ionic, Covalent, or Acid? –Ionic = metal + nonmetal - MgBr 2 –Covalent = two nonmetals - SiO 2 –Acid (first element is hydrogen) Binary Acids = acid that contains 2 elements (hydrogen and a nonmetal) Oxyacids = hydrogen, oxygen, and a third element (nonmetal)

68. 68 Chemical Nomenclature Ionic Compounds –Often a metal + nonmetal –Anion (nonmetal), add “ide” to element name BaCl 2 barium chloride K2OK2O potassium oxide Mg(OH) 2 magnesium hydroxide KNO 3 potassium nitrate

69. 69 Some transition metals can have more than 1 charge –indicate charge on metal with Roman numerals FeCl 2 2 Cl - -2 so Fe is +2iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3iron(III) chloride Cr 2 S 3 3 S -2 -6 so Cr is +3 (6/2)chromium(III) sulfide

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72. 72 Cu(NO 3 ) 2 Name the following compounds: copper (II) nitrate or cupric Nitrate KH 2 PO 4 potassium dihydrogen phosphate NH 4 ClO 3 ammonium chlorate

73. 73 Molecular compounds −Nonmetals or nonmetals + metalloids −Common names −H 2 O, NH 3, CH 4, −Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula −If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom −Last element name ends in ide

74. 74 HIhydrogen iodide NF 3 nitrogen trifluoride SO 2 sulfur dioxide N 2 Cl 4 dinitrogen tetrachloride NO 2 nitrogen dioxide N2ON2Odinitrogen monoxide Molecular (covalent) Compounds

75. 75 B2H6B2H6 diborane CH 4 methane SiH 4 silane NH 3 ammonia H2OH2Owater H2SH2Shydrogen sulfide Exceptions – compounds containing Hydrogen

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77. 77 An acid can be defined as a substance that yields hydrogen ions (H + ) when dissolved in water. For example: HCl gas and HCl in water Pure substance, hydrogen chloride Dissolved in water (H 3 O + and Cl − ), hydrochloric acid

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79. 79 An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO 3 nitric acid H 2 CO 3 carbonic acid H 3 PO 4 phosphoric acid

80. 80 Naming Oxoacids and Oxoanions

81. 81 The rules for naming oxoanions, anions of oxoacids, are as follows: 1. When all the H ions are removed from the “-ic” acid, the anion’s name ends with “-ate.” 2. When all the H ions are removed from the “-ous” acid, the anion’s name ends with “-ite.” 3. The names of anions in which one or more but not all the hydrogen ions have been removed must indicate the number of H ions present. For example: –H 2 PO 4 - dihydrogen phosphate –HPO 4 2- hydrogen phosphate –PO 4 3- phosphate

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83. 83 A base can be defined as a substance that yields hydroxide ions (OH - ) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide

84. 84 Hydrates are compounds that have a specific number of water molecules attached to them. BaCl 2 2H 2 O LiClH 2 O MgSO 4 7H 2 O Sr(NO 3 ) 2 4H 2 O barium chloride dihydrate lithium chloride monohydrate magnesium sulfate heptahydrate strontium nitrate tetrahydrate CuSO 4 5H 2 O CuSO 4

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