Chapter Two Lecture for AP Chemistry

1. 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.

2. Dalton’s Postulates
Each element is composed of
extremely small particles called atoms.
Still true!

3. Dalton’s Postulates
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.
True except for
the discovery of
isotopes!

4. Dalton’s Postulates
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.
Still true!

5. Dalton’s Postulates
Compounds are formed when atoms of
more than one element combine; a
given compound always has the same
relative number and kind of atoms.
Still true!

6. Law of Constant Composition
Joseph Proust (1754–1826)
• Also known as the law of definite
proportions.
• The elemental composition of a pure
substance never varies.

7. 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.

8. The Electron
• Streams of negatively charged particles were
found to emanate from cathode tubes.
• J. J. Thompson is credited with their
discovery (1897).

9. J.J. Thompson’s Cathode Tube
the electron

10. The Electron
Thompson measured the charge/mass ratio
of the electron to be 1.76 × 108 coulombs/g.

11. Millikan Oil Drop Experiment
Once the charge/mass
ratio of the electron
was known,
determination of either
the charge or the mass
of an electron would
yield the other.

12. Millikan Oil Drop Experiment
Robert Millikan
(University of Chicago)
determined the charge
on the electron in
1909.

13. Radioactivity:
• The spontaneous emission of radiation
by an atom.
• First observed by Henri Becquerel.
• Also studied by Marie and Pierre Curie.

14. Radioactivity
• Three types of radiation were discovered by
Ernest Rutherford:
 α particles
 β particles
 γ rays

15. The Atom, circa 1900:
• “Plum pudding” model,
put forward by
Thompson.
• Positive sphere of matter
with negative electrons
imbedded in it.

16. Discovery of the Nucleus
Ernest Rutherford
shot α particles at a
thin sheet of gold foil
and observed the
pattern of scatter of
the particles.

17. The Nuclear Atom
Since some particles
were deflected at
large angles,
Thompson’s model
could not be correct.

18. 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.

19. Rutherford’s Gold Foil
the nucleus

20. Other Subatomic Particles
• Protons were discovered by Rutherford
in 1919.
• Neutrons were discovered by James
Chadwick in 1932.

21. Subatomic Particles
• Protons and electrons are the only particles that
have a charge.
• Protons and neutrons have essentially the same
mass.
• The mass of an electron is so small we ignore it.

22. Symbols of Elements
Elements are symbolized by one or two
letters.

23. Atomic Number
All atoms of the same element have the same
number of protons:
The atomic number (Z)

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

25. Isotopes:
• Atoms of the same element with different masses.
• Isotopes have different numbers of neutrons.
11 12 13 14
6C 6C 6C 6C

26. SAMPLE EXERCISE 2.2 Determining the Number of Subatomic Particles in Atoms
How many protons, neutrons, and electrons are
in (a) an atom of 197Au (b) an atom of
strontium-90?

27. SAMPLE EXERCISE 2.2 Determining the Number of Subatomic Particles in Atoms
How many protons, neutrons, and electrons are
in (a) an atom of 197Au (b) an atom of
strontium-90?
Solution (a) The superscript 197 is the mass number, the sum of the number of protons plus the number of
neutrons. According to the list of elements given on the front inside cover, gold has an atomic number of 79.
Consequently, an atom of 197Au has 79 protons, 79 electrons, and 197 – 79 = 118 neutrons. (b) The atomic
number of strontium (listed on the front inside cover) is 38. Thus, all atoms of this element have 38 protons
and 38 electrons. The strontium-90 isotope has 90 – 38 = 52 neutrons.

28. Atomic Mass
Atomic and
molecular masses
can be measured
with great accuracy
with a mass
spectrometer.

29. Average 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.

30. SAMPLE EXERCISE 2.4
Naturally occurring chlorine is 75.78% 35Cl, which has an
atomic mass of 34.969 amu, and 24.22% 37Cl, which has an
atomic mass of 36.966 amu. Calculate the average atomic
mass (that is, the atomic weight) of chlorine.

31. SAMPLE EXERCISE 2.4
Naturally occurring chlorine is 75.78% 35Cl, which has an
atomic mass of 34.969 amu, and 24.22% 37Cl, which has an
atomic mass of 36.966 amu. Calculate the average atomic
mass (that is, the atomic weight) of chlorine.
The average atomic mass is found by multiplying the
Solution
abundance of each isotope by its atomic mass and summing
these products.

32. SAMPLE EXERCISE 2.4
Naturally occurring chlorine is 75.78% 35Cl, which has an
atomic mass of 34.969 amu, and 24.22% 37Cl, which has an
atomic mass of 36.966 amu. Calculate the average atomic
mass (that is, the atomic weight) of chlorine.
The average atomic mass is found by multiplying the
Solution
abundance of each isotope by its atomic mass and summing
these products.

33. SAMPLE EXERCISE 2.4
Naturally occurring chlorine is 75.78% 35Cl, which has an
atomic mass of 34.969 amu, and 24.22% 37Cl, which has an
atomic mass of 36.966 amu. Calculate the average atomic
mass (that is, the atomic weight) of chlorine.
The average atomic mass is found by multiplying the
Solution
abundance of each isotope by its atomic mass and summing
these products.
This answer makes sense: The average atomic mass of Cl is
between the masses of the two isotopes and is closer to the
value of 35Cl, which is the more abundant isotope.

34. BELLWORK
Which of the following diagrams is most
likely to represent an ionic compound, and
which a molecular one? Explain your choice.

35. Periodic Table:
• A systematic catalog
of elements.
• Elements are
arranged in order of
atomic number.

36. Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
reactivities.

37. Periodic Table
• The rows on the periodic
chart are periods.
• Columns are groups.
• Elements in the same
group have similar
chemical properties.

38. Groups
These five groups are known by their names.

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.

42. SAMPLE EXERCISE 2.5 Using the Periodic Table
Which two of the following elements would you
expect to show the greatest similarity in chemical
and physical properties: B, Ca, F, He, Mg, P?
PRACTICE EXERCISE
Locate Na (sodium) and Br (bromine) on the
periodic table. Give the atomic number of each,
and label each a metal, metalloid, or nonmetal.

43. 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.

44. Molecular Compounds
Molecular compounds are
composed of molecules and
almost always contain only
nonmetals.

45. Diatomic Molecules
These seven elements occur naturally as molecules
containing two atoms.

46. 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.

47. SAMPLE EXERCISE 2.6 Relating Empirical and Molecular Formulas
Write the empirical formulas for the following
molecules: (a) glucose, a substance also known as
either blood sugar or dextrose, whose molecular
formula is C6H12O6; (b) nitrous oxide, a substance
used as an anesthetic and commonly called laughing
gas, whose molecular formula is N2O.
PRACTICE EXERCISE
Give the empirical formula for the substance called
diborane, whose molecular formula is B2H6.

48. Types of Formulas
• Structural formulas show the
order in which atoms are
bonded.
• Perspective drawings also show
the three-dimensional array of
atoms in a compound.

49. Ions
• 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.

50. SAMPLE EXERCISE 2.7 Writing Chemical Symbols for Ions
Give the chemical symbol, including mass
number, for each of the following ions:
(a) The ion with 22 protons, 26 neutrons, and 19
electrons; (b) the ion of sulfur that has 16
neutrons and 18 electrons.
PRACTICE EXERCISE
How many protons and electrons does the Se2–
ion possess?

51. SAMPLE EXERCISE 2.8 Predicting the Charges of Ions
Predict the charge expected for the most stable
ion of barium and for the most stable ion of
oxygen.
PRACTICE EXERCISE
Predict the charge expected for the most stable
ion of aluminum and for the most stable ion of
fluorine.

52. Ionic Bonds
Ionic compounds (such as NaCl) are generally
formed between metals and nonmetals.

53. SAMPLE EXERCISE 2.9 Identifying Ionic and Molecular Compounds
Which of the following compounds would you
expect to be ionic: N2O, Na2O, CaCl2, SF4?
PRACTICE EXERCISE
Which of the following compounds are
molecular: CBr4, FeS, P4O6, PbF2 ?

54. 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.

55. Common Cations

56. Common Anions

57. If you have equal masses of the
following metals, which will have
the largest number of atoms?
• Pb, 207.2 g/mol
• Au, 197 g/mol
• Ag, 107.9 g/mol
• Cu, 63.54 g/mol
• Fe, 55.85 g/mol

58. If you have equal masses of the
following metals, which will have
the largest number of atoms?
• Pb, 207.2 g/mol
• Au, 197 g/mol
• Ag, 107.9 g/mol
• Cu, 63.54 g/mol
• Fe, 55.85 g/mol

59. The mass spectrum of chlorine is
shown below. How many different
molar masses are possible for Cl2?
• 1
• 2
• 3
• 4
• 5

60. The mass spectrum of chlorine is
shown below. How many different
molar masses are possible for Cl2?
• 1
• 2
• 3
• 4
• 5

61. Which molar mass will be most
common among Cl2 molecules?
• 70 g/mol
• 71 g/mol
• 72 g/mol
• 73 g/mol
• 74 g/mol

62. Which molar mass will be most
common among Cl2 molecules?
• 70 g/mol
• 71 g/mol
• 72 g/mol
• 73 g/mol
• 74 g/mol

63. Which combination is likely to
produce an ionic compound?
• C and H
• S and Cl
• Ca and F
• Br and I
• Xe and F

64. Which combination is likely to
produce an ionic compound?
• C and H
• S and Cl
• Ca and F
• Br and I
• Xe and F

65. What are the ions present in the
compound
(NH4)2SO4 ?
• NH3, H2, and SO2
• N3-, H+, S2-, O2-
• NH42+ and SO4-
• NH4+ and SO42-
• (NH4+)2 and SO42-

66. What are the ions present in the
compound
(NH4)2SO4 ?
• NH3, H2, and SO2
• N3-, H+, S2-, O2-
• NH42+ and SO4-
• NH4+ and SO42-
• (NH4+)2 and SO42-

67. Bellwork Write the molecular and structural
formulas for the compounds represented

68. Common Cations

69. Common Anions

70. Inorganic Nomenclature
• Write the name of the cation.
• 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.
• If the cation can have more than one
possible charge, write the charge as a
Roman numeral in parentheses.

71. SAMPLE EXERCISE 2.12 Determining the Names of Ionic Compounds from
Their Formulas
Name the following compounds: (a) K2SO4 , (b)
Ba(OH)2 , (c) FeCl3.
(a) potassium sulfate. (b) barium hydroxide. (c) iron(III)
chloride or ferric chloride.
PRACTICE EXERCISE
Name the following compounds: (a) NH4Br, (b)
Cr2O3, (c) Ca(NO3)2.
(a) ammonium bromide, (b) chromium(III) oxide, (c)
calcium nitrate

72. SAMPLE EXERCISE 2.13 Determining the Formulas of Ionic Compounds
from
Their Names
Write the chemical formulas for the following
compounds: (a) potassium sulfide, (b) calcium
hydrogen carbonate, (c) nickel(II) perchlorate.
(a) K2S. (b) Ca(HCO3)2. (c) Ni(ClO4)2.
PRACTICE EXERCISE
Give the chemical formula for (a) magnesium
sulfate, (b) silver sulfide, (c) lead(II) nitrate.
(a) MgSO4 , (b) Ag2S, (c) Pb(NO3)2

73. 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

74. Patterns in Oxyanion
• The one with the second fewest oxygens ends in -ite
ClO2− : chlorite
• The one with the second most oxygens ends in -ate
ClO3− : chlorate

75. 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

76. SAMPLE EXERCISE 2.11 Determining the Formula of an Oxyanion from Its Name
Based on the formula for the sulfate ion, predict the
formula for (a) the selenate ion and (b) the selenite
ion. (Sulfur and selenium are both members of
group 6A and form analogous oxyanions.)
PRACTICE EXERCISE
The formula for the bromate ion is analogous to
that for the chlorate ion. Write the formula for
the hypobromite and perbromate ions.

77. 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

78. Acid Nomenclature
• If the anion in the acid
ends in -ate, change the
ending to -ic acid:
– HClO3: chloric acid
– HClO4: perchloric acid

79. Acid Nomenclature
• If the anion in the acid
ends in -ite, change the
ending to -ous acid:
– HClO: hypochlorous
acid
– HClO2: chlorous acid

80. SAMPLE EXERCISE 2.14 Relating the Names and Formulas of Acids
Name the following acids: (a) HCN, (b) HNO3,
(c) H2SO4, (d) H2SO3.
(a)hydrocyanic acid (b)nitric acid (c)sulfuric acid.
(d)sulfurous acid
PRACTICE EXERCISE
Give the chemical formulas for (a) hydrobromic
acid, (b) carbonic acid.
(a) HBr, (b) H2CO3

81. Nomenclature of Binary
Compounds
• 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.)

82. Nomenclature of Binary
Compounds
• The ending on the more
electronegative element is
changed to -ide.
– CO2: carbon dioxide
– CCl4: carbon tetrachloride

83. Nomenclature of Binary Compounds
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

84. SAMPLE EXERCISE 2.15 Relating the Names and Formulas of
Binary Molecular Compounds
Name the following compounds: (a) SO2, (b) PCl5,
(c) N2O3.
(a) sulfur dioxide (b) phosphorus pentachloride (c)
dinitrogen trioxide.
PRACTICE EXERCISE
Give the chemical formula for
(a) silicon tetrabromide, (b) disulfur dichloride.
(a) SiBr4, (b) S2Cl2

86. Consider the alkane called pentane. (a) Assuming that the
carbon atoms are in a straight line, write a structural formula
for pentane. (b) What is the molecular formula for pentane?
(a) Alkanes contain only carbon and hydrogen, and each carbon
atom is attached to four other atoms.
(b) C5H12
PRACTICE EXERCISE
Butane is the alkane with four carbon atoms. (a) What is
the molecular formula of butane? (b) What are the name
and molecular formula of an alcohol derived from
butane?
(a) C4H10 (b) butanol, C4H10O or C4H9OH