1. ATOMS, MOLECULES, AND IONS
Chapter Two:
ATOMS, MOLECULES, AND IONS

2. Early History of Chemistry
Greeks were the first to attempt to explain why chemical changes occur.
Alchemy dominated for 2000 years:
Several elements discovered
Mineral acids prepared
Robert Boyle was the first “chemist”:
Performed quantitative experiments
2.1
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3. Three Important Laws Law of conservation of mass (Lavoisier):
Mass is neither created nor destroyed
Law of definite proportion (Proust):
A given compound always contains exactly the same proportion of elements by mass
2.2
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4. Three Important Laws (continued)
Law of multiple proportions (Dalton):
When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers
2.2
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5. Dalton’s Atomic Theory (1808)
Each element is made up of tiny particles called atoms.
2.3
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6. Dalton’s Atomic Theory (continued)
The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
2.3
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7. Dalton’s Atomic Theory (continued)
Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.
2.3
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8. Dalton’s Atomic Theory (continued)
Chemical reactions involve reorganization of the atoms—changes in the way they are bound together.
The atoms themselves are not changed in a chemical reaction.
2.3
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9. Concept Check
Which of the following statements regarding Dalton’s atomic theory are still believed to be true?
Elements are made of tiny particles called atoms.
All atoms of a given element are identical.
A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
Statements I and III are true. Statement II is not true (due to isotopes and ions). Statement IV is not true (due to nuclear chemistry).
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10. Early Experiments to Characterize the Atom
J. J. Thomson ( ):
Postulated the existence of electrons using cathode-ray tubes
Determined the charge-to-mass ratio of an electron
The atom must also contain positive particles that balance exactly the negative charge carried by the electrons
2.4
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11. Figure 2.8 Deflection of Cathode Rays by an Applied Electric Field

12. Figure 2.9 The Plum Pudding Model of the Atom

13. Early Experiments to Characterize the Atom
Robert Millikan (1909):
Performed experiments involving charged oil drops
Determined the magnitude of the electron charge
Calculated the mass of the electron
2.4
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14. Figure A Schematic Representation of the Apparatus Millikan Used to Determine the Charge on the Electron

15. Milikan’s Experiment
Perform Millilkan’s Experiment yourself
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16. Early Experiments to Characterize the Atom
Ernest Rutherford (1911):
Explained the nuclear atom
Atom has a dense center of positive charge called the nucleus
Electrons travel around the nucleus at a relatively large distance
2.4
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17. Figure 2.12 Rutherford's Experiment On a-Particle Bombardment of Metal Foil

18. Rutherford’s experiment
Animation with narration
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19. Rutherford’s Observations
Most α particles passed straight through the foil
Some α particles were deflected to the side
A few α particles were bounced backward
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20. Figure 2.13 a & b (a) Expected Results of the Metal Foil Experiment if Thomson's Model Were Correct (b) Actual Results

21. Rutherford’s Conclusions
Atoms are mostly empty space
Atoms contain positively charged particles
Atoms contain a dense nucleus
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22. The Modern View of Atomic Structure- Nuclear Model
The atom contains:
Electrons
Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
Neutrons – found in the nucleus; no charge; virtually same mass as a proton
2.5
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23. The Modern View of Atomic Structure
The nucleus is:
Small compared with the overall size of the atom
Extremely dense; accounts for almost all of the atom’s mass
2.5
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24. Nuclear Atom Viewed in Cross Section
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25. The Modern View of Atomic Structure
Isotopes:
Atoms with the same number of protons but different numbers of neutrons
Show almost identical chemical properties; chemistry of atom is due to its electrons
In nature most elements contain mixtures of isotopes
2.5
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26. Two Isotopes of Sodium
2.5
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27. Isotope Symbols X = Element symbol Z = Atomic Number A = Mass Number
= number of protons
A = Mass Number
= protons + neutrons
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28. Exercise A certain isotope X contains 54 electrons and 78 neutrons.
What is the mass number of this isotope?
What is the symbol of this isotope?
The mass number is 133. The plus charge in X+ means that the ion has lost an electron, therefore the number of protons is 55 (54+1). The ion is Cs+ with a mass number of 133 (55+78).
Note: Use the red box animation to assist in explaining how to solve the problem.
2.5/2.6
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29. Ions Atoms that have gained or lost electrons Gained electrons
Negative charge
Anion
Lost electrons
Positive charge
Cation
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30. Cations Mg2+ K+ Anions F- S2- Lost 2 electrons Lost 1 electron
Gained 1 electron
S2-
Gained 2 electrons
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31. Chemical Bonds Covalent Bonds:
Bonds form between atoms by sharing electrons
Resulting collection of atoms is called a molecule
2.6
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32. Covalent Bonding
2.6
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33. Chemical Bonds Ionic Bonds:
Bonds form due to force of attraction between oppositely charged ions
Ion – atom or group of atoms that has a net positive or negative charge
Cation – positive ion; lost electron(s)
Anion – negative ion; gained electron(s)
2.6
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34. Molecular vs. Ionic Compounds
2.6
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35. The Periodic Table Metals vs. Nonmetals
Groups or Families – elements in the same vertical columns
Alkali metals, alkaline earth metals, halogens, and noble gases
Periods – horizontal rows of elements
2.7
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36. The Periodic Table
2.7
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37. Naming Compounds Binary Compounds: Binary Ionic Compounds:
Composed of two elements
Binary Ionic Compounds:
Metal-nonmetal
Binary Covalent Compounds:
Nonmetal to nonmetal
2.8
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38. Binary Ionic Compounds (Type I)
Cation is always named first and the anion second.
Monatomic cation has the same name as its parent element.
Monatomic anion is named by taking the root of the element name and adding –ide.
2.8
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39. Binary Ionic Compounds (Type I)
Examples:
KCl Potassium chloride
MgBr2 Magnesium bromide
CaO Calcium oxide
2.8
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40. Chemical Formulas Identify the elements present
Element symbols
Identify the number of each atom
Subscripts
Aluminum sulfide Al2S3
2 Al atoms for every 3 S atoms

41. Writing the Formulas from the Names
Determine the ions present
Determine the charges on the cation and anion
Balance the charges to get the subscripts

42. Writing the Formulas from the Names (cont.)
Magnesium chloride
Mg2+ Cl-
Total positive charge must equal total negative charge
1 Mg2+ will balance 2Cl-
MgCl2

43. Polyatomic Ions Ions consisting of two or more covalently bonded atoms
Examples of compounds containing polyatomic ions:
NaOH Sodium hydroxide
Mg(NO3)2 Magnesium nitrate
(NH4)2SO4 Ammonium sulfate
2.8
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44. Compounds Containing Polyatomic Ions
Polyatomic ions are charged entities that contain more than one atom
Polyatomic compounds contain one or more polyatomic ions
Name polyatomic compounds by naming cation and anion
Al(C2H3O2)3 = aluminum acetate

45. Binary Ionic Compounds (Type II)
Metals in these compounds form more than one type of positive charge.
Charge on the metal ion must be specified.
Roman numeral indicates the charge of the metal cation.
Transition metal cations usually require a Roman numeral.
2.8
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46. Determining the Charge on a Cation – Au2S3
Determine the charge on the anion
Sulfide = -2
Determine the total negative charge
(-2) x 3 = -6
Total positive charge = total negative charge
Divide total positive charge by the number of cations
+6/2 = +3

47. Binary Ionic Compounds (Type II)
Examples:
CuBr Copper(I) bromide
FeS Iron(II) sulfide
PbO2 Lead(IV) oxide
2.8
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48. Type II Binary Ionic Compounds (cont.)
Old naming system
Uses root name or latin name for metal
Uses suffixes instead of roman numerals

49. Type II Binary Ionic Compounds (cont.)
Suffixes
_____ic for higher charge
_____ous for lower charge

50. Type II Binary Ionic Compounds (cont.)
Examples
Iron (III) chloride = ferric chloride
Iron (II) chloride = ferrous chloride

51. Type II Binary Ionic Compounds (cont.)
Common latin names
Cu – cupric, cuprous
Pb – plumbic, plumbous
Sn – stannic, stannous
Fe – ferric, ferrous

52. Writing the Formulas from the Names (cont.)
For Type II
Determine the charges on the cation and anion
Charge on cation = roman numeral
Balance the charges to get the subscripts

53. Writing the Formulas from the Names (cont.)
Copper (I) oxide
Cu+1 O-2
2 Cu+1 will balance 1 O-2
Cu2O

54. Formation of Ionic Compounds
2.8
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55. Binary Covalent Compounds (Type III)
Formed between two nonmetals.
First element is named first, using the full element name.
Second element is named as if it were an anion.
Prefixes are used to denote the numbers of atoms present.
Prefix mono- is never used for naming the first element.
2.8
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56. Table 2.6 Prefixes Used to Indicate Number in Chemical Names

57. Binary Covalent Compounds (Type III)
Examples:
CO2 Carbon dioxide
SF6 Sulfur hexafluoride
N2O4 Dinitrogen tetroxide
2.8
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58. Writing the Formulas from the Names
For Type III compounds, use the prefixes to determine the subscripts
Carbon tetrachloride = CCl4

59. Overall Strategy for Naming Chemical Compounds
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60. Flowchart for Naming Binary Compounds
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61. Acids
Acids can be recognized by the hydrogen that appears first in the formula—HCl.
Molecule with one or more H+ ions attached to an anion.
2.8
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62. Acids
If the anion does not contain oxygen, the acid is named with the prefix hydro- and the suffix -ic.
Examples:
HCl Hydrochloric acid
HCN Hydrocyanic acid
H2S Hydrosulfuric acid
2.8
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63. Acids If the anion does contain oxygen: Examples: HNO3 Nitric acid
The suffix -ic is added to the root name if the anion name ends in -ate.
Examples:
HNO3 Nitric acid
H2SO4 Sulfuric acid
HC2H3O2 Acetic acid
2.8
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64. Acids If the anion does contain oxygen: Examples: HNO2 Nitrous acid
The suffix -ous is added to the root name if the anion name ends in -ite.
Examples:
HNO2 Nitrous acid
H2SO3 Sulfurous acid
HClO2 Chlorous acid
2.8
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65. Flowchart for Naming Acids
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66. Exercise Which of the following compounds is named incorrectly?
KNO3 potassium nitrate
TiO2 titanium(II) oxide
Sn(OH)4 tin(IV) hydroxide
PBr5 phosphorus pentabromide
CaCrO4 calcium chromate
The correct answer is “b”. The charge on oxygen is -2. Since there are two oxygen atoms, the overall charge is -4. Therefore, the charge on titanium must be +4 (not +2 as the Roman numeral indicates).
2.8
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