GENERAL CHEMISTRY
General, Organic, and Biochemistry
9th Edition
Katherine J. Denniston
Joseph J. Topping
Danaè R. Quirk Dorr
Robert L. Caret
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1.1 The Discovery Process
Chemistry:
• the study of matter
• its chemical and physical properties
• the chemical and physical changes it undergoes
– As matter undergoes changes, it gains or
loses energy
• Matter - anything that has mass and occupies
space
• Energy - the ability to do work to accomplish
some change
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1.1 The Discovery Process
Role of Chemistry
public health
pharmaceutical industry
CHEMISTRY
food science
medical practitioners
forensic sciences
1.1 The Discovery Process
THE SCIENTIFIC METHOD
• a systematic approach to the discovery of new
information
2
Characteristics:
 Observation
 Formulation of a question
 Pattern recognition
 Developing theories
 Experimentation
• Data: the individual result of a single measurement
• Results: the outcome of an experiment
 Summarizing information
3
1.1 The Discovery Process
Representation of the Scientific Method
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1.1 The Discovery Process
Models in Chemistry
• Aid in the understanding of
a chemical unit or system
– often based on everyday
experience
• Ball and stick model of
methane
– color coded balls (atoms)
– sticks (attractive forces
holding atoms together)
1.2 The Classification of Matter
• Properties - characteristics of matter
scientists can use to categorize different
types of matter
• Ways to Categorize matter:
1. By State
2. By Composition
1.2 The Classification of Matter
Three States of Matter
1.Gas - particles widely separated, no
definite shape or volume solid
2. Liquid - particles closer together,
definite volume but no definite shape
3. Solid - particles are very close together,
define shape and definite volume
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1.2 The Classification of Matter
Composition of Matter
5
• Pure substance - a substance that has only one
component
• Mixture - a combination of two or more pure
substances in which each substance retains its
own identity, not undergoing a chemical reaction
1.2 The Classification of Matter
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Pure
Substances
• Element - a pure substance that cannot be
changed into a simpler form of matter by any
chemical reaction
• Compound - a pure substance resulting from
the combination of two or more elements in a
definite, reproducible way, in a fixed ratio
5
1.2 The Classification of Matter
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5
Mixture
• Mixture - a combination of two or more pure
substances in which each substance retains its own
identity
• Homogeneous - uniform composition, particles well
mixed, thoroughly intermingled
• Heterogeneous – nonuniform composition, random
placement
1.2 The Classification of Matter
Classes of Matter
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(a) Pure Substance
(b) Homogenous Mixture
(c) Heterogeneous Mixture
1.2 The Classification of Matter
Physical Property vs. Physical Change 6
• Physical property - is observed without
changing the composition or identity of a
substance
• Physical change - produces a
recognizable difference in the appearance
of a substance without causing any
change in its composition or identity
- conversion from one physical state to
another
- melting an ice cube
Physical Properties and Physical
Change
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(a) Solid
(b) Liquid
(c) Gas
Separation by Physical Properties
Magnetic iron is separated from other nonmagnetic
substances, such as sand. This property is used as
a large-scale process in the recycling industry.
1.2 The Classification of Matter
Chemical Property vs. Chemical Reaction 6
• Chemical property - results in a change in
composition and can be observed only
through a chemical reaction
• Chemical reaction (chemical change) - a
chemical substance is converted in to one or
more different substances by rearranging,
removing, replacing, or adding atoms
hydrogen + oxygen  water
reactants
products
1.2 The Classification of Matter
Classification of Properties
Classify the following as either a
chemical or physical property:
a. Color
b. Flammability
c. Hardness
d. Odor
e. Taste
1.2 The Classification of Matter
Classification of Changes
Classify the following as either a
chemical or physical change:
a. Boiling water becomes steam
b. Butter turns rancid
c. Burning of wood
d. Mountain snow melting in spring
e. Decay of leaves in winter
1.2 The Classification of Matter
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Intensive and Extensive Properties
• Intensive properties - a property of
matter that is independent of the
quantity of the substance
- Color
- Melting Point
• Extensive properties - a property of
matter that depends on the quantity of
the substance
- Mass
- Volume
1.3 The Units of Measurement
Units - the basic quantity of mass, volume or
whatever quantity is being measured
– A measurement is useless without its units
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• English system - a collection of functionally unrelated
units
– Difficult to convert from one unit to another
1 foot = 12 inches = 0.33 yard = 1/5280 miles
• Metric System - composed of a set of units that are
related to each other decimally, systematic
– Units relate by powers of tens
1.3 The Units of
Measurement
Metric System Units
8
• Mass - the quantity of matter in an object
– not synonymous with weight
• Weight = mass x acceleration due to gravity
– standard unit is the gram (g)
– The pound (lb) is the common English unit
1 lb = 454 g
• Mass must be measured on a balance (not a scale)
1.3 The Units of
Measurement
Metric System Units - continued
• Length - the distance between two points
– standard unit is the meter (m)
– The yard is the common English unit
1 yd = 0.914 m
• Volume - the space occupied by an object
– standard unit is the liter (L)
– The quart is the common English unit
1 qt = 0.946 L
• Time
• The metric unit is the second (s)
Metric System Prefixes
• Basic units are the units of a quantity without any
metric prefix
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1.3 The Units of
Measurement
Relationship among
various volume units
Volume =
Length x width x height
Volume =
1mx1mx1m=
1 m3
1 m3 = 1 L
1.4 The Numbers of Measurement
• Information-bearing digits or figures in a
number are significant figures
9
• The measuring device used determines the
number of significant figures in a measurement
• The degree of uncertainty associated with a
measurement is indicated by the number of
figures used to represent the information
1.4 The Numbers of
Measurement
Significant Figures Example
Significant figures - all digits in a number
representing data or results that are known
with certainty plus one uncertain digit
1.4 The Numbers of
Measurement
Recognition of Significant
Figures
• All nonzero digits are significant
• 7.314 has four significant digits
• The number of significant digits is independent
of the position of the decimal point
• 73.14 also has four significant digits
• Zeros located between nonzero digits are
significant
• 60.052 has five significant digits
1.4 The Numbers of
Measurement
Use of Zeros in Significant
Figures
• Zeros at the end of a number (trailing zeros) are:
– Significant if the number contains a decimal point
• 4.70 has three significant digits
– Insignificant if the number does not contain a decimal
point
• 100 has one significant digit; 100. has three
• Zeros to the left of the first nonzero integer are not
significant
• 0.0032 has two significant digits
1.4 The Numbers of
Measurement
How many significant figures are in
the following?
1. 3.400
2. 3004
3. 300.
4. 0.003040
1.4 The Numbers of
Measurement
Scientific Notation
• Used to express very large or very small
numbers easily and with the correct number
of significant figures
• Represents a number as a power of ten
• Example:
4,300 = 4.3  1,000 = 4.3  103
1.4 The Numbers of
Measurement
Scientific Notation Rules
• To convert a number greater than 1 to
scientific notation, the original decimal
point is moved x places to the left, and the
resulting number is multiplied by 10x
• The exponent x is a positive number equal to
the number of places the decimal point moved
6200 = 6.2  103
• What if you want to express the above
number with three significant figures?
= 6.20  103
1.4 The Numbers of
Measurement
Scientific Notation Rules continued
• To convert a number less than 1 to scientific
notation, the original decimal point is
moved x places to the right, and the
resulting number is multiplied by 10–x
• The exponent x is a negative number equal to
the number of places the decimal point moved
0.0062 = 6.2  10–3
1.4 The Numbers of
Measurement
Scientific Notation Example
• When a number is exceedingly large or small,
scientific notation must be used to input the
number into a calculator:
0.000000000000000000000006692 g
must be entered into calculator as:
6.692 x 10−24
1.4 The Numbers of
Measurement
Represent the following numbers in
scientific notation:
1. 0.00018
2. 3004
3. 300.
4. 0.00304
1.4 The Numbers of
Measurement
Accuracy and Precision
• Accuracy - the degree of
agreement between the true
value and the measured
value
– Error - the difference
between the true value and
our estimation
• Random
• Systematic
• Precision - a measure of
the agreement of replicate
measurements
– Deviation – amount of
variation present in a set of
replicate measurements
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1.4 The Numbers of
Measurement
Exact (Counted) and Inexact Numbers
• Inexact numbers have uncertainty (degree
of doubt in final significant digit)
• Exact numbers are a consequence of
counting
– A set of counted items (beakers on a shelf)
has no uncertainty
– Exact numbers by definition have an infinite
number of significant figures
1.4 The Numbers of
Measurement
Rules for Rounding Numbers
• When the number to be dropped is less
than 5, the preceding number is not
changed
• When the number to be dropped is 5 or
larger, the preceding number is increased
by one unit
• Round the following number to 3
significant figures: 3.34966  104
=3.35  104
1.4 The Numbers of
Measurement
Round off each number to three
significant figures:
1. 61.40
2. 6.171
3. 0.066494
Significant Figures in Calculation of Results
Rules for Addition and Subtraction
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• The result in a calculation cannot have greater
significance than any of the quantities that
produced the result
• Consider:
37.68
6.71862
108.428
152.82662
liters
liters
liters
liters
correct answer 152.83 liters
1.4 The Numbers of
Measurement
Report the result of each to the
proper number of significant figures:
1. 4.26 + 3.831
2. 8.321 − 2.4
Adding and Subtracting in Scientific Notation
• There are two ways to solve the following:
9.47 x 10−6 + 9.3 x 10−5
SOLUTION 1: convert both numbers to
standard form and add
0.000009 47
+
0.000093
0.000102 47
correct answer 1.02 x 10−4
Addition Example
• There are two ways to solve the following:
9.47 x 10−6 + 9.3 x 10−5
SOLUTION 2: change one of the exponents so
that both have the same power of 10, then add
9.47 x 10−6 changes to 0.947 x 10−5
0.9 47 x 10−5
+
9.3 x 10−5
10.2 47 x 10−5
correct answer 1.02 x 10−4
Rules for Multiplication and Division
• The answer can be no more precise than the least
precise number from which the answer is derived
• The least precise number is the one with the
fewest significant figures
(4.2 103 )(15.94)
 2.96886918 (on calculator)
4
2.255 10
Which number has the fewest
significant figures? 4.2  103 has only 2
The answer is therefore, 3.0
1.5 Unit Conversion
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• Factor-Label Method (Dimensional
Analysis)
–Uses Conversion Factors to:
- Convert from one unit to another within
the same system
- Convert units from one system to another
1.5 Unit Conversion
English Unit Conversion - Example
• To convert from one unit to another
you must know the conversion factor,
which is the relationship between the
two units
– The Relationship:
1 gal = 4 qt
– The Conversion Factor:
1 gal
or
4 qt
4 qt
1 gal
1.5 Unit Conversion
Using Conversion Factors
Convert 12 gallons to quarts
– The Relationship (English system):
1 gal = 4 qt
– The Conversion Factor:
1 gal
or
4 qt
4 qt
1 gal
Data Given: 12 gal
Use Conversion Factor with gal in denominator
1.5 Unit Conversion
Using Conversion Factors - Solution
Convert 12 gallons to quarts
Solution:
• Write the Data Given
• Multiply by the Conversion Factor
with the unit of the Data Given (gal) in
the denominator
12 gal x
4 qt
1 gal
=
48 qt
Desired Result
1.5 Unit Conversion
Unit Conversion - Example
Convert 360 feet to miles
– The Relationship (English system):
5280 ft = 1 mi
– The Conversion Factor:
5280 ft
or
1 mi
1 mi
5280 ft
Data Given: 360 ft
Use Conversion Factor with ft in denominator
1.5 Unit Conversion
Unit Conversion - Solution
Convert 360 feet to miles
Solution:
• Write the Data Given
• Multiply by the Conversion Factor
with the unit of the Data Given (ft)
in the denominator
360 ft x
1 mi
5280 ft
=
0.068 miles
Desired Result
1.5 Unit Conversion
Multistep Conversion - Example
Convert 0.0047 kilograms to milligrams
– The Relationships (metric system):
1 kg = 103 g and 103 mg = 1 g
– The Conversion Factors:
1 kg or 103 g and 1 g or 103 mg
103 g
1 kg
103 mg
1g
Data Given: 0.0047 kg
1. Use Conversion Factor with kg in denominator
to convert to Initial Data Result in g
2. Use Conversion Factor with g in denominator
1.5 Unit Conversion
Multistep Conversion - Solution
Convert 0.0047 kilograms to milligrams
0.0047 kg x 103 g = 4.7 g
1 kg
Data Given x Conversion Factor = Initial Data Result
4.7 g x 1 mg = 4.7 x 103 mg
10−3 g
Initial Data Result x Conversion Factor = Desired Result
1.5 Unit Conversion
Multistep Conversions Alternate Solution
Convert 0.0047 kilograms to milligrams
Alternatively, Solve in a single step:
0.0047 kg x 103 g x 1 mg = 4.7 x 103 mg
1 kg
10−3 g
Data Given x Conversion Factor x Conversion Factor
= Desired Result
1.5 Unit Conversion
Practice Unit Conversions
1. Convert 5.5 inches to millimeters
2. Convert 50.0 milliliters to pints
3. Convert 1.8 in2 to cm2
1.6 Additional Experimental
Quantities
• Temperature - the degree of “hotness” of an
object
1.6 Additional Experimental
Quantities
Conversions Between Fahrenheit
and Celsius
ToC =
ToF − 32
1.8
ToF = 1.8 x ToC + 32
1. Convert 75oC to oF
2. Convert -10oF to oC
1. Ans. 167 oF
2. Ans. -23oC
12
1.6 Additional Experimental
Quantities
Kelvin Temperature Scale
• The Kelvin (K) scale is another temperature
scale
• It is of particular importance because it is
directly related to molecular motion
• As molecular speed increases, the Kelvin
temperature proportionately increases
TK = ToC + 273.15
1.6 Additional Experimental
Quantities
Energy
• Energy - the ability to do work
• kinetic energy - the energy of motion
(energy of action)
• potential energy - the energy of position
(stored energy)
• Energy is also categorized by form:
•
•
•
•
•
light
heat
electrical
mechanical
chemical
1.6 Additional Experimental
Quantities
Characteristics of Energy
• Energy cannot be created or destroyed
• Energy may be converted from one form to
another
• Energy conversion always occurs with less
than 100% efficiency
• All chemical reactions involve either a
“gain” or “loss” of energy
1.6 Additional Experimental
Quantities
Units of Energy
• Basic Units:
• calorie or joule
• 1 calorie (cal) = 4.18 joules (J)
• kilocalorie (kcal) = food Calorie
1 kcal = 1 Calorie = 1000 calories
• 1 calorie = amount of heat energy
required to increase the temperature of 1
gram of water 1oC
1.6 Additional Experimental
Quantities
Concentration
Concentration:
– the number or mass of particles of a
substance contained in a specified volume
Often used to represent the mixtures of
different substances
– Concentration of oxygen in the air
– Pollen counts
– Proper dose of an antibiotic
1.6 Additional Experimental
Quantities
Density and Specific Gravity
• Density
13
– the ratio of mass to volume
– an extensive property
– use to characterize a substance as
each substance has a unique
density
– Units for density include:
• g/mL
• g/cm3
• g/cc
d=
mass
m
=
volume V
1.6 Additional Experimental
Quantities
Density
Examples
cork
water
brass nut
liquid mercury
Densities of Some Common
Materials
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Calculating Density
A 2.00 cm3 sample of aluminum is found to weigh
5.40 g. Calculate the density in g/cm3 and g/mL.
Use the expression:
• Density (d) = m/V
Substitute information given into the expression
d = 5.40 g = 2.70 g/cm3
2.00 cm3
• Since 1 cm3 = 1 mL,
= 2.70 g/mL
Use Density in Calculation
Calculate the volume, in mL, of a liquid that has
a density of 1.20 g/mL and a mass of 5.00 g.
• Density can be written as a Conversion Factor
1.20 g
1 mL
1 mL
1.20 g
or
• Multiply the Data Given (g) by the Conversion
Factor with the unit g in the denominator
5.00 g x
1 mL
1.20 g
=
4.17 mL
1.6 Additional Experimental
Quantities
Density Calculations
• Air has a density of 0.0013 g/mL. What is
the mass of 6.0-L sample of air?
• Calculate the mass in grams of 10.0 mL if
mercury (Hg) if the density of Hg is 13.6
g/mL.
• Calculate the volume in milliliters, of a liquid
that has a density of 1.20 g/mL and a mass of
5.00 grams.
1.6 Additional Experimental
Quantities
Specific Gravity
• Values of density are often related to a standard
• Specific gravity - the ratio of the density of the
object in question to the density of pure water at
4oC
• Specific gravity is a unitless term because the 2
units cancel
• Often the health industry uses specific gravity
to test urine and blood samples
density of object (g/mL)
specific gravity =
density of water (g/mL)