1. Chemistry
In this science we
study matter and the
changes it
undergoes.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

2. Matter
We define matter as anything that has mass
and takes up space.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

3. Matter
• Atoms are the building blocks of matter.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

4. Matter
• Atoms are the building blocks of matter.
• Each element is made of the same kind of atom.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

5. Matter
• Atoms are the building blocks of matter.
• Each element is made of the same kind of atom.
• A compound is made of two or more different kinds of
elements. Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

6. States of Matter
Matter
And
Measurement
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7. Classification of Matter
Matter
And
Measurement
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8. Classification of Matter
Matter
And
Measurement
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9. Classification of Matter
Matter
And
Measurement
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10. Classification of Matter
Matter
And
Measurement
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11. Classification of Matter
Matter
And
Measurement
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12. Classification of Matter
Matter
And
Measurement
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13. Classification of Matter
Matter
And
Measurement
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14. Classification of Matter
Matter
And
Measurement
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15. Classification of Matter
Matter
And
Measurement
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16. Classification of Matter
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

17. Properties and
Changes of
Matter
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

18. Types of Properties
• Physical Properties…
– Can be observed without changing a
substance into another substance.
• Boiling point, density, mass, volume, etc.
• Chemical Properties…
– Can only be observed when a substance is
changed into another substance.
• Flammability, corrosiveness, reactivity with
acid, etc.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

19. Types of Properties
• Intensive Properties…
– Are independent of the amount of the
substance that is present.
• Density, boiling point, color, etc.
• Extensive Properties…
– Depend upon the amount of the substance
present.
• Mass, volume, energy, etc.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

20. Types of Changes
• Physical Changes
– These are changes in matter that do not
change the composition of a substance.
• Changes of state, temperature, volume, etc.
• Chemical Changes
– Chemical changes result in new substances.
• Combustion, oxidation, decomposition, etc.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

21. Chemical Reactions
In the course of a chemical reaction, the
reacting substances are converted to new
Matter
substances. And
Measurement
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22. Compounds
Compounds can be
broken down into
more elemental
particles.
Matter
And
Measurement
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23. Separation of
Mixtures
Matter
And
Measurement
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24. Filtration
In filtration solid
substances are
separated from liquids
and solutions.
Matter
And
Measurement
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25. Distillation
Distillation uses
differences in the
boiling points of
substances to
separate a
homogeneous
mixture into its
components.
Matter
And
Measurement
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26. Chromatography
This technique separates substances on the
basis of differences in solubility in a solvent.
Matter
And
Measurement
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27. What do these countries have
in common?
US, Liberia and Burma
Matter
And
Measurement
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28. What do these countries have
in common?
US, Liberia and Burma
• They use the imperial system
Matter
And
Measurement
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29. View of Countries using Metric
USA
Berma Liberia
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And
Measurement
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30. Units of
Measurement
Matter
And
Measurement
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31. SI Units
• Système International d’Unités
• A different base unit is used for each quantity.
Matter
And
Measurement
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32. Metric System
Prefixes convert the base units into units that
are appropriate for the item being measured.
Matter
And
Measurement
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33. Volume
• The most commonly
used metric units for
volume are the liter (L)
and the milliliter (mL).
– A liter is a cube 1 dm
long on each side.
– A milliliter is a cube 1 cm
long on each side.
Matter
And
Measurement
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34. Uncertainty in
Measurement
Matter
And
Measurement
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35. Uncertainty in Measurements
Different measuring devices have different
uses and different degrees of accuracy.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

36. Uncertainty in Measurements
Different measuring devices have different
uses and different degrees of accuracy.
1 ml
Matter
And
Measurement
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37. Uncertainty in Measurements
Different measuring devices have different
uses and different degrees of accuracy.
0.1 ml
Matter
And
Measurement
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38. Accuracy versus Precision
• Accuracy refers to the proximity of
a measurement to the true value
of a quantity.
• Precision refers to the proximity of
several measurements to each
other.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

39. Significant Figures
• The term significant figures refers to
digits that were measured.
• When rounding calculated numbers, we
pay attention to significant figures so we
do not overstate the accuracy of our
answers.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

40. Significant Figures
1. All nonzero digits are significant.
2. Zeroes between two significant figures
are themselves significant.
3. Zeroes at the beginning of a number
are never significant.
4. Zeroes at the end of a number are
significant if a decimal point is written
in the number.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

41. Significant Figures
• When addition or subtraction is
performed, answers are rounded to the
least significant decimal place.
• When multiplication or division is
performed, answers are rounded to the
number of digits that corresponds to the
least number of significant figures in
any of the numbers used in the
calculation.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

42. Temperature
By definition
temperature is a
measure of the
average kinetic
energy of the
particles in a
sample.
Matter
And
Measurement
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43. Temperature
• In scientific
measurements, the
Celsius and Kelvin
scales are most often
used.
• The Celsius scale is
based on the
properties of water.
– 0°C is the freezing point
of water.
– 100°C is the boiling
point of water. Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

44. Temperature
• The Kelvin is the SI
unit of temperature.
• It is based on the
properties of gases.
• There are no
negative Kelvin
temperatures.
• K = °C + 273.15
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

45. Temperature
• The Fahrenheit
scale is not used in
scientific
measurements.
∀ °F = 9/5(°C) + 32
∀ °C = 5/9(°F − 32)
Matter
And
Measurement
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46. Density
Density is a physical property of a
substance.
m
d=
V
Matter
And
Measurement
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47. Dimensional Analysis
• We use dimensional analysis
to convert one quantity to
another.
• Most commonly dimensional
analysis utilizes conversion
factors (e.g., 1 in. = 2.54 cm)
1 in. 2.54 cm
or
2.54 cm 1 in.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

48. Dimensional Analysis
Use the form of the conversion factor
that puts the sought-for unit in the
numerator.
desired unit
Given unit × = desired unit
given unit
Conversion factor Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

49. Dimensional Analysis
• For example, to convert 8.00 m to inches,
– convert m to cm
– convert cm to in.
100 cm 1 in.
8.00 m × × = 315 in.
1m 2.54 cm
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

50. 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.
Matter
And
Measurement
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51. Dalton's Postulates
Each element is composed of extremely small
particles called atoms.
Matter
And
Measurement
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52. 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.
Matter
And
Measurement
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53. 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.
Matter
And
Measurement
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54. 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.
Matter
And
Measurement
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55. Law of Constant Composition
Joseph Proust (1754–1826)
• This is also known as the law of definite
proportions.
• It states that the elemental composition
of a pure substance never varies.
Matter
And
Measurement
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56. 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.
Matter
And
Measurement
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57. The Electron
• Streams of negatively charged particles were
found to emanate from cathode tubes.
• J. J. Thompson is credited with their
discovery (1897). Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

58. The Electron
Thompson measured the charge/mass ratio
of the electron to be 1.76 × 108 coulombs/g.
Matter
And
Measurement
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59. 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.
Matter
And
Measurement
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60. Millikan Oil Drop Experiment
Robert Millikan
(University of Chicago)
determined the charge
on the electron in
1909.
Matter
And
Measurement
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61. Radioactivity
• Radioactivity is the spontaneous
emission of radiation by an atom.
• It was first observed by Henri
Becquerel.
• Marie and Pierre Curie also studied it.
Matter
And
Measurement
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62. Radioactivity
• Three types of radiation were discovered by
Ernest Rutherford:
α particles
β particles
γ rays
Matter
And
Measurement
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63. 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.
Matter
And
Measurement
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64. 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.
Matter
And
Measurement
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65. The Nuclear Atom
Since some particles
were deflected at
large angles,
Thompson’s model
could not be correct.
Matter
And
Measurement
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66. 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.
Matter
And
Measurement
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67. Other Subatomic Particles
• Protons were discovered by Rutherford
in 1919.
• Neutrons were discovered by James
Chadwick in 1932.
Matter
And
Measurement
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68. 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.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

69. Symbols of Elements
Elements are symbolized by one or two
letters.
Matter
And
Measurement
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70. Atomic Number
All atoms of the same element have the same
number of protons:
The atomic number (Z)
Matter
And
Measurement
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71. Atomic Mass
The mass of an atom in atomic mass units
(amu) is the total number of protons and
neutrons in the atom.
Matter
And
Measurement
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72. Isotopes
• Isotopes are atoms of the same element with
different masses.
• Isotopes have different numbers of neutrons.
11 12 13 14
6 C 6 C 6 C 6 C
Matter
And
Measurement
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73. Atomic Mass
Atomic and
molecular masses
can be measured
with great accuracy
with a mass
spectrometer.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

74. 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.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

75. Periodic Table
• It is a systematic
catalog of the
elements.
• Elements are
arranged in order
of atomic number.
Matter
And
Measurement
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76. Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
reactivities.
Matter
And
Measurement
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77. Periodic Table
• The rows on the
periodic chart are
periods.
• Columns are groups.
• Elements in the same
group have similar
chemical properties.
Matter
And
Measurement
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78. Groups
These five groups are known by their names.
Matter
And
Measurement
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79. Periodic Table
Nonmetals are
on the right
side of the
periodic table
(with the
exception of
H).
Matter
And
Measurement
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80. Periodic Table
Metalloids
border the
stair-step line
(with the
exception of
Al, Po, and
At).
Matter
And
Measurement
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81. Periodic Table
Metals are
on the left
side of the
chart.
Matter
And
Measurement
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82. 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.
Matter
And
Measurement
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83. Chemical Formulas
Molecular compounds
are composed of
molecules and almost
always contain only
nonmetals.
Matter
And
Measurement
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84. Diatomic Molecules
These seven elements occur naturally as
molecules containing two atoms.
Matter
And
Measurement
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85. 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.
Matter
And
Measurement
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86. 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.
Matter
And
Measurement
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87. 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 Matter
on the right side of the periodic chart. And
Measurement
© 2009, Prentice-Hall, Inc.

88. Ionic Bonds
Ionic compounds (such as NaCl) are
generally formed between metals and
nonmetals.
Matter
And
Measurement
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89. 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
Matter
factor. And
Measurement
© 2009, Prentice-Hall, Inc.

90. Common Cations
Matter
And
Measurement
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91. Common Anions
Matter
And
Measurement
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92. 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.
Matter
And
Measurement
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93. 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
Matter
And
Measurement
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94. 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
Matter
And
Measurement
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95. 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
Matter
And
Measurement
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96. 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
Matter
And
Measurement
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97. Acid Nomenclature
• If the anion in the acid
ends in -ite, change
the ending to -ous
acid.
– HClO: hypochlorous
acid
– HClO2: chlorous acid
Matter
And
Measurement
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98. Acid Nomenclature
• If the anion in the acid
ends in -ate, change
the ending to -ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid
Matter
And
Measurement
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99. 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) .
Matter
And
Measurement
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100. Nomenclature of Binary
Compounds
• The ending on the more
electronegative element
is changed to -ide.
– CO2: carbon dioxide
– CCl4: carbon tetrachloride
Matter
And
Measurement
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101. 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
Matter
And
Measurement
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102. Nomenclature of Organic
Compounds
• Organic chemistry is the study of carbon.
• Organic chemistry has its own system of
nomenclature.
Matter
And
Measurement
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103. Nomenclature of Organic
Compounds
The simplest hydrocarbons (compounds
containing only carbon and hydrogen) are
alkanes.
Matter
And
Measurement
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104. Nomenclature of Organic
Compounds
The first part of the names above correspond
to the number of carbons (meth- = 1, eth- = 2,
prop- = 3, etc.).
Matter
And
Measurement
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105. 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.
Matter
– An alcohol ends in -ol. And
Measurement
© 2009, Prentice-Hall, Inc.

106. Law of Conservation of Mass
“We may lay it down as an
incontestable axiom that, in all
the operations of art and nature,
nothing is created; an equal
amount of matter exists both
before and after the experiment.
Upon this principle, the whole art
of performing chemical
experiments depends.”
--Antoine Lavoisier, 1789
Matter
And
Measurement
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107. Chemical Equations
Chemical equations are concise
representations of chemical reactions.
Matter
And
Measurement
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108. Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Matter
And
Measurement
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109. Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Reactants appear on the left
side of the equation. Matter
And
Measurement
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110. Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Products appear on the
right side of the equation. Matter
And
Measurement
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111. Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
The states of the reactants and products
are written in parentheses to the right of Matter
And
each compound. Measurement
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112. Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Coefficients are inserted
to balance the equation. Matter
And
Measurement
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113. Subscripts and Coefficients
Give Different Information
• Subscripts tell the number of atoms of
each element in a molecule.
Matter
And
Measurement
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114. Subscripts and Coefficients
Give Different Information
• Subscripts tell the number of atoms of
each element in a molecule
• Coefficients tell the number of Matter
And
molecules. Measurement
© 2009, Prentice-Hall, Inc.

115. Reaction
Types
Matter
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Measurement
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116. Combination Reactions
• In this type of
reaction two
or more
substances
react to form
one product.
• Examples:
– 2 Mg (s) + O2 (g) → 2 MgO (s)
– N2 (g) + 3 H2 (g) → 2 NH3 (g)
Matter
– C3H6 (g) + Br2 (l) → C3H6Br2 (l) And
Measurement
© 2009, Prentice-Hall, Inc.

117. Decomposition Reactions
• In a decomposition
one substance breaks
down into two or more
substances.
• Examples:
– CaCO3 (s) → CaO (s) + CO2 (g)
– 2 KClO3 (s) → 2 KCl (s) + O2 (g)
Matter
– 2 NaN3 (s) → 2 Na (s) + 3 N2 (g) And
Measurement
© 2009, Prentice-Hall, Inc.

118. Combustion Reactions
• These are generally
rapid reactions that
produce a flame.
• Most often involve
hydrocarbons
reacting with oxygen
in the air.
• Examples:
– CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)
Matter
– C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (g) And
Measurement
© 2009, Prentice-Hall, Inc.

119. Formula
Weights
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And
Measurement
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120. Formula Weight (FW)
• A formula weight is the sum of the
atomic weights for the atoms in a
chemical formula.
• So, the formula weight of calcium
chloride, CaCl2, would be
Ca: 1(40.1 amu)
+ Cl: 2(35.5 amu)
111.1 amu
• Formula weights are generally reported
for ionic compounds. Matter
And
Measurement
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121. Molecular Weight (MW)
• A molecular weight is the sum of the
atomic weights of the atoms in a
molecule.
• For the molecule ethane, C2H6, the
molecular weight would be
C: 2(12.0 amu)
+ H: 6(1.0 amu)
30.0 amu
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122. Percent Composition
One can find the percentage of the
mass of a compound that comes from
each of the elements in the compound
by using this equation:
(number of atoms)(atomic weight)
% element = x 100
(FW of the compound)
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123. Percent Composition
So the percentage of carbon in ethane
is…
(2)(12.0 amu)
%C =
(30.0 amu)
24.0 amu
= x 100
30.0 amu
= 80.0%
Matter
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124. Moles
Matter
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125. Avogadro’s Number
• 6.02 x 1023
• 1 mole of 12C has a
mass of 12 g.
Matter
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126. Molar Mass
• By definition, a molar mass is the mass
of 1 mol of a substance (i.e., g/mol).
– The molar mass of an element is the mass
number for the element that we find on the
periodic table.
– The formula weight (in amu’s) will be the
same number as the molar mass (in
g/mol).
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127. Using Moles
Moles provide a bridge from the molecular
scale to the real-world scale.
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128. Mole Relationships
• One mole of atoms, ions, or molecules contains
Avogadro’s number of those particles.
• One mole of molecules or formula units contains
Avogadro’s number times the number of atoms or
ions of each element in the compound. Matter
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129. Finding
Empirical
Formulas
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130. Calculating Empirical
Formulas
One can calculate the empirical formula from
the percent composition.
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131. Calculating Empirical
Formulas
The compound para-aminobenzoic acid (you may have
seen it listed as PABA on your bottle of sunscreen) is
composed of carbon (61.31%), hydrogen (5.14%),
nitrogen (10.21%), and oxygen (23.33%). Find the
empirical formula of PABA.
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132. Calculating Empirical
Formulas
Assuming 100.00 g of para-aminobenzoic acid,
C: 61.31 g x 1 mol = 5.105 mol C
12.01 g
1 mol
H: 5.14 g x = 5.09 mol H
1.01 g
1 mol
N: 10.21 g x = 0.7288 mol N
14.01 g
1 mol
O: 23.33 g x = 1.456 mol O
16.00 g
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133. Calculating Empirical
Formulas
Calculate the mole ratio by dividing by the smallest number
of moles:
5.105 mol
C: = 7.005 ≈ 7
0.7288 mol
5.09 mol
H: = 6.984 ≈ 7
0.7288 mol
0.7288 mol
N: = 1.000
0.7288 mol
1.458 mol
O: = 2.001 ≈ 2 Matter
0.7288 mol And
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134. Calculating Empirical
Formulas
These are the subscripts for the empirical formula:
C7H7NO2
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135. Combustion Analysis
• Compounds containing C, H and O are routinely
analyzed through combustion in a chamber like this.
– C is determined from the mass of CO2 produced.
– H is determined from the mass of H2O produced.
– O is determined by difference after the C and H have been
determined. Matter
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136. Elemental Analyses
Compounds
containing other
elements are
analyzed using
methods analogous
to those used for C,
H and O.
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137. Stoichiometric Calculations
The coefficients in the balanced equation give
the ratio of moles of reactants and products.
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138. Stoichiometric Calculations
Starting with the
mass of Substance
A you can use the
ratio of the
coefficients of A and
B to calculate the
mass of Substance
B formed (if it’s a
product) or used (if
it’s a reactant).
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139. Stoichiometric Calculations
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O
Starting with 1.00 g of C6H12O6…
we calculate the moles of C6H12O6…
use the coefficients to find the moles of H2O…
Matter
and then turn the moles of water to grams. And
Measurement
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140. Limiting
Reactants
Matter
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141. How Many Cookies Can I
Make?
• You can make cookies
until you run out of one
of the ingredients.
• Once this family runs
out of sugar, they will
stop making cookies
(at least any cookies
you would want to eat).
Matter
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142. How Many Cookies Can I
Make?
• In this example the
sugar would be the
limiting reactant,
because it will limit the
amount of cookies you
can make.
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143. Limiting Reactants
• The limiting reactant is the reactant present in
the smallest stoichiometric amount.
– In other words, it’s the reactant you’ll run out of first (in
this case, the H2).
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144. Limiting Reactants
In the example below, the O2 would be the
excess reagent.
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145. Theoretical Yield
• The theoretical yield is the maximum
amount of product that can be made.
– In other words it’s the amount of product
possible as calculated through the
stoichiometry problem.
• This is different from the actual yield,
which is the amount one actually
produces and measures.
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146. Percent Yield
One finds the percent yield by
comparing the amount actually obtained
(actual yield) to the amount it was
possible to make (theoretical yield).
Actual Yield
Percent Yield = x 100
Theoretical Yield
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147. Solutions
• Solutions are defined as
homogeneous mixtures
of two or more pure
substances.
• The solvent is present in
greatest abundance.
• All other substances are
solutes.
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148. Dissociation
• When an ionic
substance dissolves
in water, the solvent
pulls the individual
ions from the crystal
and solvates them.
• This process is called
dissociation.
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149. Dissociation
• An electrolyte is a
substances that
dissociates into ions
when dissolved in
water.
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150. Electrolytes
• An electrolyte is a
substances that
dissociates into ions
when dissolved in
water.
• A nonelectrolyte may
dissolve in water, but
it does not dissociate
into ions when it does
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151. Electrolytes and
Nonelectrolytes
Soluble ionic
compounds tend
to be electrolytes.
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152. Electrolytes and
Nonelectrolytes
Molecular
compounds tend to
be nonelectrolytes,
except for acids and
bases.
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153. Electrolytes
• A strong electrolyte
dissociates completely
when dissolved in
water.
• A weak electrolyte
only dissociates
partially when
dissolved in water.
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154. Strong Electrolytes Are…
• Strong acids
• Strong bases
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155. Strong Electrolytes Are…
• Strong acids
• Strong bases
• Soluble ionic salts
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156. Precipitation Reactions
When one mixes ions
that form compounds
that are insoluble (as
could be predicted by
the solubility
guidelines), a
precipitate is formed.
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157. Metathesis (Exchange) Reactions
• Metathesis comes from a Greek word that
means “to transpose.”
AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)
Matter
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158. Metathesis (Exchange) Reactions
• Metathesis comes from a Greek word that
means “to transpose.”
• It appears the ions in the reactant
compounds exchange, or transpose, ions.
AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)
Matter
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159. Solubility of different compounds
(NS = non soluble in water, S = soluble in water)
Cl- Br- I- NO3- SO42- CO32- PO43-
Li+ S S S S S S S S
Na+ S S S S S S S S
K+ S S S S S S S S
Mg2+ NS S S S S S NS NS
Ca2+ S S S S S S NS NS
Sr2+ S S S S S NS NS NS
Ba2+ S S S S S NS NS NS
Fe2+ NS S S S S S NS NS
Fe3+ NS S S S S S NS NS
Ni2+ NS S S S S S NS NS
Cu+ NS S S S S S NS NS
Cu2+ NS S S S S S NS NS
Al3+ NS S S S S S NS NS
Zn2+ NS S S S S S NS NS
Ag+ NS NS NS NS S S NS NS
Pb2+ NS NS NS NS S NS NS NS
Matter
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160. Solution Chemistry
• It is helpful to pay attention to exactly
what species are present in a reaction
mixture (i.e., solid, liquid, gas, aqueous
solution).
• If we are to understand reactivity, we
must be aware of just what is changing
during the course of a reaction.
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161. Molecular Equation
The molecular equation lists the reactants
and products in their molecular form.
AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)
Matter
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162. Ionic Equation
• In the ionic equation all strong electrolytes (strong
acids, strong bases, and soluble ionic salts) are
dissociated into their ions.
• This more accurately reflects the species that are
found in the reaction mixture.
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq) →
AgCl (s) + K+ (aq) + NO3- (aq)
Matter
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163. Net Ionic Equation
• To form the net ionic equation, cross out anything
that does not change from the left side of the
equation to the right.
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) →
Matter
AgCl (s) + K+(aq) + NO3-(aq) And
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164. Net Ionic Equation
• To form the net ionic equation, cross out anything
that does not change from the left side of the
equation to the right.
• The only things left in the equation are those things
that change (i.e., react) during the course of the
reaction.
Ag+(aq) + Cl-(aq) → AgCl (s)
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165. Net Ionic Equation
• To form the net ionic equation, cross out anything
that does not change from the left side of the
equation to the right.
• The only things left in the equation are those things
that change (i.e., react) during the course of the
reaction.
• Those things that didn’t change (and were deleted
from the net ionic equation) are called spectator ions.
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) →
Matter
AgCl (s) + K+(aq) + NO3-(aq) And
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166. Writing Net Ionic Equations
1. Write a balanced molecular equation.
2. Dissociate all strong electrolytes.
3. Cross out anything that remains
unchanged from the left side to the
right side of the equation.
4. Write the net ionic equation with the
species that remain.
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167. Acids
• Arrhenius defined acids
as substances that
increase the
concentration of H+ when
dissolved in water.
• Brønsted and Lowry
defined them as proton
donors.
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168. Acids
There are only seven
strong acids:
• Hydrochloric (HCl)
• Hydrobromic (HBr)
• Hydroiodic (HI)
• Nitric (HNO3)
• Sulfuric (H2SO4)
• Chloric (HClO3)
• Perchloric (HClO4) Matter
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169. Bases
• Arrhenius defined bases
as substances that
increase the
concentration of OH−
when dissolved in water.
• Brønsted and Lowry
defined them as proton
acceptors.
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170. Bases
The strong bases
are the soluble
metal salts of
hydroxide ion:
• Alkali metals
• Calcium
• Strontium
• Barium
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171. Acid-Base Reactions
In an acid-base
reaction, the acid
donates a proton
(H+) to the base.
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172. Neutralization Reactions
Generally, when solutions of an acid and a base are
combined, the products are a salt and water.
CH3COOH (aq) + NaOH (aq) →CH3COONa (aq) + H2O (l)
Matter
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173. Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
Matter
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174. Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) →
Na+ (aq) + Cl- (aq) + H2O (l)
Matter
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175. Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) →
Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq) → H2O (l)
Matter
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176. Gas-Forming Reactions
• Some metathesis reactions do not give the
product expected.
• In this reaction, the expected product (H2CO3)
decomposes to give a gaseous product
(CO2).
CaCO3 (s) + HCl (aq) →CaCl2 (aq) + CO2 (g) + H2O (l)
Matter
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177. Gas-Forming Reactions
When a carbonate or bicarbonate reacts with
an acid, the products are a salt, carbon
dioxide, and water.
CaCO3 (s) + HCl (aq) →CaCl2 (aq) + CO2 (g) + H2O (l)
NaHCO3 (aq) + HBr (aq) →NaBr (aq) + CO2 (g) + H2O (l)
Matter
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Measurement
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178. Gas-Forming Reactions
Similarly, when a sulfite reacts with an acid,
the products are a salt, sulfur dioxide, and
water.
SrSO3 (s) + 2 HI (aq) →SrI2 (aq) + SO2 (g) + H2O (l)
Matter
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Measurement
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179. Gas-Forming Reactions
• This reaction gives the predicted product, but
you had better carry it out in the hood, or you
will be very unpopular!
• But just as in the previous examples, a gas is
formed as a product of this reaction.
Na2S (aq) + H2SO4 (aq) → Na2SO4 (aq) + H2S (g)
Matter
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180. Oxidation-Reduction Reactions
• An oxidation occurs
when an atom or ion
loses electrons.
• A reduction occurs
when an atom or ion
gains electrons.
• One cannot occur
without the other.
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181. Oxidation Numbers
To determine if an oxidation-reduction
reaction has occurred, we assign an
oxidation number to each element in a
neutral compound or charged entity.
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182. Oxidation Numbers
• Elements in their elemental form have
an oxidation number of 0.
• The oxidation number of a monatomic
ion is the same as its charge.
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183. Oxidation Numbers
• Nonmetals tend to have negative
oxidation numbers, although some are
positive in certain compounds or ions.
Oxygen has an oxidation number of −2,
except in the peroxide ion in which it has
an oxidation number of −1.
Hydrogen is −1 when bonded to a metal,
+1 when bonded to a nonmetal.
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184. Oxidation Numbers
• Nonmetals tend to have negative
oxidation numbers, although some are
positive in certain compounds or ions.
Fluorine always has an oxidation number
of −1.
The other halogens have an oxidation
number of −1 when they are negative; they
can have positive oxidation numbers,
however, most notably in oxyanions. Matter
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185. Oxidation Numbers
• The sum of the oxidation numbers in a
neutral compound is 0.
• The sum of the oxidation numbers in a
polyatomic ion is the charge on the ion.
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186. Displacement Reactions
• In displacement reactions,
ions oxidize an element.
• The ions, then, are
reduced. Matter
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187. Displacement Reactions
In this reaction,
silver ions oxidize
copper metal.
Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)
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188. Displacement Reactions
The reverse reaction,
however, does not
occur.
x
Cu2+ (aq) + 2 Ag (s) → Cu (s) + 2 Ag+ (aq)
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189. Activity Series
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190. Molarity
• Two solutions can contain the same
compounds but be quite different because the
proportions of those compounds are different.
• Molarity is one way to measure the
concentration of a solution.
moles of solute
Molarity (M) =
volume of solution in liters
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191. Mixing a Solution
• To create a solution of a
known molarity, one
weighs out a known mass
(and, therefore, number of
moles) of the solute.
• The solute is added to a
volumetric flask, and
solvent is added to the line
on the neck of the flask.
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192. Dilution
• One can also dilute a more concentrated
solution by
– Using a pipet to deliver a volume of the solution to
a new volumetric flask, and
– Adding solvent to the line on the neck of the new
flask.
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193. Dilution
The molarity of the new solution can be determined
from the equation
Mc × Vc = Md × Vd,
where Mc and Md are the molarity of the concentrated and dilute
solutions, respectively, and Vc and Vd are the volumes of the
two solutions.
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194. Using Molarities in
Stoichiometric Calculations
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195. Titration
Titration is an
analytical
technique in
which one can
calculate the
concentration
of a solute in
a solution.
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196. Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 5
Thermochemistry
John D. Bookstaver
Matter
St. Charles Community College And
Cottleville, MO Measurement
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197. Energy
• Energy is the ability to do work or
transfer heat.
– Energy used to cause an object that has
mass to move is called work.
– Energy used to cause the temperature of
an object to rise is called heat.
Matter
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198. Potential Energy
Potential energy is energy an object
possesses by virtue of its position or chemical
composition.
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199. Kinetic Energy
Kinetic energy is energy an object possesses
by virtue of its motion.
1
KE =  mv2
2
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200. Units of Energy
• The SI unit of energy is the joule (J).
kg m2
1 J = 1 
s2
• An older, non-SI unit is still in
widespread use: the calorie (cal).
1 cal = 4.184 J
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201. Definitions:
System and Surroundings
• The system includes the
molecules we want to
study (here, the hydrogen
and oxygen molecules).
• The surroundings are
everything else (here, the
cylinder and piston).
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202. Definitions: Work
• Energy used to
move an object over
some distance is
work.
• w=F×d
where w is work, F
is the force, and d is
the distance over
which the force is
exerted. Matter
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203. Heat
• Energy can also be
transferred as heat.
• Heat flows from
warmer objects to
cooler objects.
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204. Conversion of Energy
• Energy can be converted from one type to
another.
• For example, the cyclist above has potential
energy as she sits on top of the hill.
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205. Conversion of Energy
• As she coasts down the hill, her potential
energy is converted to kinetic energy.
• At the bottom, all the potential energy she had
at the top of the hill is now kinetic energy.
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206. First Law of Thermodynamics
• Energy is neither created nor destroyed.
• In other words, the total energy of the universe is
a constant; if the system loses energy, it must be
gained by the surroundings, and vice versa.
Matter
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207. Internal Energy
The internal energy of a system is the sum of all
kinetic and potential energies of all components
of the system; we call it E.
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208. Internal Energy
By definition, the change in internal energy, ∆E,
is the final energy of the system minus the initial
energy of the system:
∆E = Efinal − Einitial
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209. Changes in Internal Energy
• If ∆E > 0, Efinal > Einitial
– Therefore, the system
absorbed energy from
the surroundings.
– This energy change is
called endergonic.
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210. Changes in Internal Energy
• If ∆E < 0, Efinal < Einitial
– Therefore, the system
released energy to the
surroundings.
– This energy change is
called exergonic.
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211. Changes in Internal Energy
• When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
• That is, ∆E = q + w.
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212. ∆E, q, w, and Their Signs
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213. Exchange of Heat between
System and Surroundings
• When heat is absorbed by the system from
the surroundings, the process is endothermic.
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Measurement
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214. Exchange of Heat between
System and Surroundings
• When heat is absorbed by the system from
the surroundings, the process is endothermic.
• When heat is released by the system into the
surroundings, the process is exothermic.
Matter
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215. State Functions
Usually we have no way of knowing the
internal energy of a system; finding that value
is simply too complex a problem.
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216. State Functions
• However, we do know that the internal energy
of a system is independent of the path by
which the system achieved that state.
– In the system below, the water could have reached
room temperature from either direction.
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And
Measurement
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217. State Functions
• Therefore, internal energy is a state function.
• It depends only on the present state of the
system, not on the path by which the system
arrived at that state.
• And so, ∆E depends only on Einitial and Efinal.
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Measurement
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218. State Functions
• However, q and w are
not state functions.
• Whether the battery is
shorted out or is
discharged by running
the fan, its ∆E is the
same.
– But q and w are different
in the two cases.
Matter
And
Measurement
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219. Work
Usually in an open
container the only work
done is by a gas
pushing on the
surroundings (or by the
surroundings pushing
on the gas).
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And
Measurement
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220. Work
We can measure the work done by the gas if
the reaction is done in a vessel that has been
fitted with a piston.
w = -P∆V
Matter
And
Measurement
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221. Enthalpy
• If a process takes place at constant
pressure (as the majority of processes we
study do) and the only work done is this
pressure-volume work, we can account
for heat flow during the process by
measuring the enthalpy of the system.
• Enthalpy is the internal energy plus the
product of pressure and volume:
H = E + PV Matter
And
Measurement
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222. Enthalpy
• When the system changes at constant
pressure, the change in enthalpy, ∆H, is
∆H = ∆(E + PV)
• This can be written
∆H = ∆E + P∆V
Matter
And
Measurement
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223. Enthalpy
• Since ∆E = q + w and w = -P∆V, we can
substitute these into the enthalpy
expression:
∆H = ∆E + P∆V
∆H = (q+w) − w
∆H = q
• So, at constant pressure, the change in
enthalpy is the heat gained or lost.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

224. Endothermicity and
Exothermicity
• A process is
endothermic when
∆H is positive.
Matter
And
Measurement
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225. Endothermicity and
Exothermicity
• A process is
endothermic when
∆H is positive.
• A process is
exothermic when
∆H is negative.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.

226. Enthalpy of Reaction
The change in
enthalpy, ∆H, is the
enthalpy of the
products minus the
enthalpy of the
reactants:
∆H = Hproducts − Hreactants
Matter
And
Measurement
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227. Enthalpy of Reaction
This quantity, ∆H, is called the enthalpy of
reaction, or the heat of reaction.
Matter
And
Measurement
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228. The Truth about Enthalpy
• Enthalpy is an extensive property.
• ∆H for a reaction in the forward
direction is equal in size, but opposite
in sign, to ∆H for the reverse reaction.
• ∆H for a reaction depends on the state
of the products and the state of the
reactants.
Matter
And
Measurement
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229. Calorimetry
Since we cannot
know the exact
enthalpy of the
reactants and
products, we
measure ∆H through
calorimetry, the
measurement of
heat flow.
Matter
And
Measurement
© 2009, Prentice-Hall, Inc.