2. Three Fundamental Laws
Law of conservation of mass (Lavoisier):
Mass is neither created nor destroyed in a chemical reaction.
Law of definite proportion (Proust):
A given compound always contains exactly the same proportion of elements by mass.
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.

3. 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. Dalton's Postulates
1. Each element is composed of extremely small indivisible particles called atoms.
Figure 2.1 John Dalton ( )

5. Dalton's Postulates
2. All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.
Figure 2.1 John Dalton ( )

6. Dalton's Postulates
3. Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
Figure 2.1 John Dalton ( )

7. 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.

8. Avogadro’s Hypothesis
Gay—Lussac
Measured (under same conditions of T and P) the volumes of gases that reacted with each other.
Avogadro’s Hypothesis
At the same T and P, equal volumes of different gases contain the same number of particles.
Volume of a gas is determined by the number, not the size, of molecules.

9. Early experiments to characterize the atom
The Electron
Voltage source
Figure 2.4
In 1897 J.J. Thomson discovered that passing an electric current through an evacuated sealed glass tube, makes a stream of particles move from the negative to the positive end.

10. Thomson’s cathode ray tube experiment
Voltage source + -
By adding an electric field, he found that the moving pieces were negative.
Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.

11. Plum Pudding Model
The atom was thought to be a positive sphere of matter with negative electrons imbedded in it.
Figure 2.9
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12. Robert Milikan’s oil drop experiment
Performed experiments involving charged oil droplets.
Determined the charge on the electron to be 1.60  C.
Determined the mass of an electron to be 9.11  kg.

13. Henri Becquerel discovers radioactivity
Radioactivity is the spontaneous emission of radiation by an atom. Also studied by Marie and Pierre Curie.
Ernest Rutherford noted three spots on the detector:
a spot in the direction of the positive plate ( particles)
a spot which is not affected by the electric field ( rays)
a spot in the direction of the negative plate ( particles)

14. Rutherford’s gold foil experiment
A source of α-particles was placed at the mouth of a circular detector. The α -particles were shot through a piece of gold foil.

15. Rutherford’s nuclear atom
Since some particles were deflected at large angles, but most went straight through, Thompson’s model could not be correct.
Evidence  Theory
Some alpha particles deflected greatly  presence of very small, dense positive nucleus (with the electrons around the outside)
2) Most alpha particles went straight through  most of the volume of the atom is empty space.
Figure 2.11
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16. Niels Bohr Electrons travel in orbit around nucleus
Idea of electron transitions between energy levels, emitting/absorbing photons

17. Bohr’s energy levels
The energy levels in an atom are sort of like steps of a ladder.
The more energy an electron has, the farther away from the nucleus it usually will be.
The energy levels are not evenly spaced. They get closer together as you travel farther away.
To move from one “rung” to another requires a “quantum” of energy.

18. Bohr Atomic Model

19. The difference between continuous and quantized energy levels.
continuous energy levels
quantized energy levels

20. James Chadwick discovers the neutron
A thin sheet of beryllium was bombarded with alpha particles, and a very high energy radiation was emitted.
The radiation was electrically neutral, had no charge but a slightly larger mass than a proton.

21. Heisenberg’s Uncertainty Principle

23. The Modern View

25. Quantum Mechanical Model

27. Quantum Mechanical Model

28. Modern View of Atomic Structure
Subatomic particles
We define: mass of 12C = exactly 12 amu.
Using atomic mass units:
1 amu = x g
1 g = x 1023 amu

29. Symbols of Elements Elements are symbolized by one or two letters.
First letter is always capital.
Second letter (if there is one) is lower case.
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30. Atomic Number
All atoms of the same element have the same number of protons:
The atomic number (Z)
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31. Mass Number
The mass of an atom in atomic mass units (amu) is the total number of nucleons (protons + neutrons)
The mass number (A)
By convention, for element X, we write ZAX.
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32. Isotopes
Isotopes are atoms of the same element with different masses.
In terms of subatomic particles, isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C
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33. 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.
Average atomic mass =
Σ (isotope mass)( relative abundance)
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34. Atomic Mass Calculation 1
Isotope
Atomic Mass
Relative Abundance
28Si
27.976 %
29Si
28.976
4.6832%
30Si
29.973
3.0872%
Using the isotope information for Silicon. Find the average atomic mass.

35. Atomic Mass Calculation 2
Naming and Formula Writing
Atomic Mass Calculation 2
Silver consists of two isotopes 107Ag and 109Ag. Its average atomic mass is Calculate the percentage of each isotope in naturally occurring silver. (Assume that the masses are and respectively.)

36. Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Figure 2.13
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37. The Periodic Table

38. Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
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39. Periodic Table Metals are on the left side of the chart.
Good conductors
Lose e- to form cation
Malleable
Ductlie
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40. Periodic Table Metalloids border the stair-step line
Intermediate properties
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41. Periodic Table
Nonmetals are on the right side of the periodic table (except H).
Poor conductors
Gain e- to form anion
Brittle
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42. Alkali Metals
Very reactive 1+

43. Alkaline Earth Metals
Reactive 2+

44. Transition metals
Form colored compounds
Variable charges

45. Inner Transition Metals
Lanthanides or rare earth
Actinides are all radioactive
Z ≥ 92 are manmade

46. Chalcogens
Oxygen family
6 valence electrons

47. Halogens Very reactive Only group containing elements in s, l, g @STP
Form acids when bonded with H

48. Noble Gases
Low reactivity
Full octet
Colorless, odorless gas

49. Molecules and Molecular Compounds
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.
Molecular compounds are composed of molecules and almost always contain only nonmetals.
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50. Allotropes Different forms of the same element.
Tin: White (metallic tin), Gray Tin
Oxygen: Oxygen gas, Ozone
Carbon:
Graphite, Diamond, “Bucky Balls”

51. Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms.
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52. Molecular vs. Empirical formula
Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. NO2
Molecular formulas give the exact number of atoms of each element in a compound. N2O4
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53. Picturing molecules
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54. Ions and Ionic Compounds
Ions – charged particle
monoatomic or polyatomic
Cations – positively
charged particle
Anions – negatively

55. Ionic Compounds
Ionic “bond” is the attraction between oppositely charged ions (generally metals and nonmetals)
Called salts
Examples: Table salt (NaCl), baking soda (NaHCO3), Epsom salts (MgSO4)
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57. 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.
IONIC ONLY: If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.
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58. Naming Inorganic Compounds
Binary Ionic Compounds (Type I)
1. The cation is always named first and the anion second. 2. A monatomic cation takes its name from the name of the parent element. 3. A monatomic anion is named by taking the root of the element name and adding –ide. KCl MgBr2 CaO Gallium bromide Sodium iodide

59. Binary Ionic Compounds (Type II)
Metals in these compounds form more than one type of positive ion.
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.
Elements that form only one cation do not need to be identified by a roman numeral.
CuBr FeS PbO2
Manganese IV oxide Copper I chloride

60. Name of polyatomic ion does not change (even if it is anion)
Polyatomic Ions
Name of polyatomic ion does not change (even if it is anion)
Examples of compounds containing polyatomic ions:
NaOH Mg(NO3)2 (NH4)2SO4

61. Ionic Compounds & Formulas
Name the compound Write the formula
KCl Sodium fluoride
Mg3N2 Calcium nitrate
Na2SO4 Barium nitrite
(NH4)2CO3 Lead (II) hydroxide
CuO Managanese (IV) sulfide
Cu2O Tin (II) sulfite
FePO4 Sodium nitride

62. Names and Formulas of Binary Molecular Compounds (covalent)
2 non metals
First element name unchanged. Second element named as if it were an anion
No roman numerals
Prefixes (no mono on cation)
mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
Sulfur dioxide CO2 XeF6
difluorine monoxide CO P2O10
nitrogen trichloride CCl4 N4O4
diphosphorus pentoxide N2O4

63. Copyright © Cengage Learning. All rights reserved
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
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64. Names and Formulas of Acids

65. Name these acids
HCN HNO3
HNO2 HClO4
HClO3 H2SO3
HCl HBr
HI HC2H3O2

66. Writing formulas for acids
If it has hydro- in the name it has no oxygen (Anion ends in –ide)
No hydro in name, anion ends in -ate or -ite
Write anion and add enough H to balance the charges.
hydrofluoric acid dichromic acid
carbonic acid hydrophosphoric acid
phosphorous acid sulfuric acid
perchloric acid hydrosulfuric acid

67. Naming and Formula Writing
Hydrates
Hydrates are compounds that contain discrete water molecules as part of the crystal lattice structure.
CuSO4•5H2O is called copper(II)sulfate pentahydrate.
In the formula you put a dot and then write the number of molecules of water, using the Greek prefixes.
Calcium chloride dihydrate
Cr(NO3)3· 6H2O