1. Chapter 2 Atoms, Molecules, and Ions

2. Historical Discovery of Atom
Democritus – Laughing Philosopher
John Dalton – solid sphere, combine
Thomson – Plum Pudding Model
Rutherford – Gold Foil Experiment
Millikan – Oil Drop Experiment
Bohr – Orbital Model
Chadwick – Neutrons
Wave Mechanical Model
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4. Atomic Theory of Matter
The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.
Figure 2.1 John Dalton ( )
4 Postulate Theory
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5. Dalton's Postulates
1. Each element is composed of extremely small particles called atoms.
Figure 2.1 John Dalton ( )
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6. Dalton's Postulates
2. All atoms of a given element are identical in mass and properties, but the atoms of one element are different from the atoms of another element.
Figure 2.1 John Dalton ( )
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7. Dalton's Postulates
3. Atoms of an element are not changed by chemical reactions; atoms are neither created nor destroyed in chemical reactions – just REARRANGED.
Supports the Law of Conservation of Mass
Figure 2.1 John Dalton ( )
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8. Dalton’s Postulates
4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.
Supports the Law of Constant Composition
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9. Law of Constant Composition Joseph Proust (1754–1826)
also known as the law of definite proportions
states that the elemental composition of a pure substance never varies.

10. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

11. Subatomic Particles
Landmark Discoveries lead to the current model of atomic theory

12. The Busy Electron
Figure 2.4
Streams of negatively charged particles were found to emanate from cathode (- electrode) tubes.
J. J. Thompson is credited with their discovery (1897).

13. The Busy Electron Electron Charge (C) = 1.76  108 coulombs
Electron Mass (g) gram
Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.
Figure 2.4
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14. Millikan Oil Drop Experiment
He examined the relationship between voltage of plates affected rate of electron fall.
Open text to pg. 40
Fig. 2.5 Explanation
Figure 2.5
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15. Millikan Oil Drop Experiment
Robert Millikan (University of Chicago) determined the charge on the electron in 1909.
Charge of Electron = 9.10 x g
Figure 2.5
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16. Subatomic Particles Relative Mass and Charge
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17. Relative Mass Image
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18. Radioactivity
Radioactivity is the spontaneous emission of radiation by an atom.
It was first observed by Henri Becquerel
Followed by Marie and Pierre Curie.
Shared 1903 Nobel Prize in Physics

19. Radioactivity Three types were discovered by Rutherford:
 particles, Alpha, bent by electric field, +2 charge
 particles, Beta, bent by electric field, - 1 charge
 rays, Gamma, unaffected by electric field, no charge
Figure 2.8

20. The Atom, circa 1900
The prevailing theory was that of the “plum pudding” model, put forward by Thompson.
It featured a positive sphere of matter with negative electrons imbedded in it.
Figure 2.9

21. Discovery of the Nucleus
Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
Particles were deflected by foil.
Figure 2.10

22. The Nuclear Atom
Since some particles were deflected at large angles, Thompson’s model could not be correct.
Rutherford proposed the positive nucleus containing protons.
Figure 2.11

23. The Nuclear Atom
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.
Most of the volume of the atom is empty space.
Figure 2.12

24. Other Subatomic Particles
Protons discovered by Rutherford, 1919.
Neutrons discovered by James Chadwick,1932.

25. Subatomic Particles Protons and electrons are have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.
Table 2.1

26. Symbols of Elements Nuclear Symbol Hyphen Notation C -12
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27. Atomic Number
Atoms of the same element have the same number of protons: the atomic number (Z)

28. Atomic Mass
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

29. Isotopes 11 6 C 12 6 C 13 6 C 14 6 C
Isotopes are atoms have different masses.
Isotopes have different numbers of neutrons.

30. Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Turn to pg. 48 in text
Figure 2.13
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31. Average Atomic Mass
Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.
Average mass is calculated from the isotopes of an element weighted by their relative abundances.
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32. Calculating Average Atomic Mass
EX. You have a box containing two sizes of marbles.
25% of the marbles have a mass of 2.0 g.
75% of the marbles have a mass of 3.0 g.
Calculate the average weight of the marble.

33. EX. Copper has two naturally occurring isotopes:
Cu-63 (69.17% abundant) with amu =
Cu-65 (30.83% abundant) with amu =
Calculate the average atomic mass.
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34. Periodic Table It is a systematic catalog of the elements.
Elements are arranged in order of atomic number.
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35. Periodicity or Periodic Law
Notices a repeating pattern of reactivities!
Function of the atomic number called periodicity.

36. Periodic Table The rows on the periodic chart are periods.
Columns are groups or families.
Elements in the same group have similar chemical properties.
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37. Groups These five groups are known by their names.
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39. Periodic Table
Nonmetals are on the right side of the periodic table (with the exception of H).

40. Periodic Table
Metalloids border the stair-step line (with the exception of Al and Po).

41. Periodic Table Metals are on the left side of the chart.
80% of elements are metals.

42. Naming and Writing Inorganics
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.
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43. Chemical Formulas
Molecular compounds are composed of molecules and almost always contain only nonmetals.
- covalent compounds
- NM + NM
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44. Diatomic Molecules Br2 I2 N2 Cl2 H2 O2 F2
These seven elements occur naturally as molecules containing two atoms.
Br2 I2 N2 Cl2 H2 O2 F2
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45. Types of Formulas
Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.
Molecular formulas give the exact number of atoms of each element in a compound.
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46. Types of Formulas
Structural formulas show the order in which atoms are bonded.
Perspective drawings show the three-dimensional array of atoms in a compound.
Ball+Stick models show the atoms as spheres and bonds as sticks.
Space filling model shows relative sizes of atoms.
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47. Ions When atoms lose or gain electrons, they become ions.
Ca+ions are positive, formed as atoms lose e-, metals
Anions are negative, formed as atoms gain e-, nonmetals

48. Ionic Bonds Ionic compounds = Metal + Nonmetal
Ionic compounds = cation + anion

49. Writing Formulas
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.

50. RIP Kris Kross
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51. Writing Chemical Formulas
Assign Charges then criss x cross the numbers
DO NOT criss x cross the charges, just the #s
Ex. Magnesium Chloride
Assign Charges Mg Cl-1
Criss-Cross Mg Cl2
Chemical Formula MgCl2
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52. Common Cations
Pg. 62 or on back fold of text

53. Common Anions Pg. 64 or on back fold of text
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54. Inorganic Nomenclature
Write the name of the metal cation.
If the metal cation can have more than one charge, write the charge as a Roman Numeral in parentheses. (CLIMT)
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.
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55. Patterns in Oxyanion Nomenclature
When there are two oxyanions involving the same element:
The one with fewer oxygens ends in -ite.
NO2− : nitrite; SO32− : sulfite
The one with more oxygens ends in -ate.
NO3− : nitrate; SO42− : sulfate
-ate is GREAT and –ite is LIGHT

56. Patterns in Oxyanion Nomenclature
The one with the second fewest oxygens ends in -ite.
ClO2− : chlorite
The one with the second most oxygens ends in -ate.
ClO3− : chlorate
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57. Patterns in Oxyanion Nomenclature
The one with the fewest oxygens has the prefix hypo- and ends in -ite.
ClO− : hypochlorite
The one with the most oxygens has the prefix per- and ends in -ate.
ClO4− : perchlorate

58. Acid Nomenclature
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
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59. Acid Nomenclature
If the anion in the acid ends in -ate, change the ending to -ic acid.
HClO3: chloric acid
HClO4: perchloric acid
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60. Acid Nomenclature
If the anion in the acid ends in -ite, change the ending to -ous acid.
HClO: hypochlorous acid
HClO2: chlorous acid
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61. Molecular Nomenclature
The less electronegative atom is usually listed first.
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) .
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62. Molecular Nomenclature
The ending on the more electronegative element is changed to -ide.
CO2: carbon dioxide
CCl4: carbon tetrachloride
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63. Molecular Nomenclature
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.
N2O5: dinitrogen pentoxide
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64. Nomenclature of Organic Compounds
Organic chemistry is the study of carbon.
Organic chemistry has its own system of nomenclature.
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65. Nomenclature of Organic Compounds
The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.
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66. Nomenclature of Organic Compounds
The first part of the names above correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
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67. Nomenclature of Organic Compounds
When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.
The ending denotes the type of compound.
An alcohol ends in -ol.
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