GENERAL CHEMISTRY
Spring 2022-2023
College of Arts and Sciences
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
General, Organic and Biochemistry
1.1 The Discovery Process
Self reading
• Chemistry - The study of matter…
– Matter - Anything that has mass and
occupies space
• A table
• A piece paper
– What about air?
• Yes, it is matter
1.1 The Discovery Process
Chemistry:
• the study of matter
• its chemical and physical properties
• the chemical and physical changes it
undergoes
• the energy changes that accompany
those processes
• Energy - the ability to do work to
accomplish some change
1.1 The Discovery Process
MAJOR AREAS OF CHEMISTRY
• Biochemistry - the study of life at the 1
molecular level
• Organic chemistry - the study of matter
containing carbon and hydrogen
• Inorganic chemistry - the study of matter
containing elements, not organic
• Analytic chemistry - analyze matter to
determine identity and composition
1.1 The Discovery Process
• Physical chemistry - attempts to
explain the way matter behaves
public health
pharmaceutical industry
CHEMISTRY
food science
medical practitioners
forensic sciences
1.1 The Discovery Process
THE SCIENTIFIC METHOD
• The scientific method - a systematic 2
approach to the discovery of new
information
Characteristics of the scientific process
 Observation
 Formulation of a question
 Pattern recognition
 Developing theories
 Experimentation
 Summarizing information
1.1 The Discovery Process
1.1 The Discovery Process
Models in Chemistry
• To aid in understanding of
a chemical unit or system
– a model is often used
– good models are based on
everyday experience
• Ball and stick methane
model
– color code balls
– sticks show attractive forces
holding atoms together
1.2 Matter and Properties
• Properties - characteristics of matter
– chemical vs. physical
• Data – the individual result of a single
measurement.
• Results – the outcome of an experiment
3
1.2 Matter and Properties
Three States of Matter
1.gas - particles widely separated, no
4
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
Three States of Water
(a) Solid
(b) Liquid
(c) Gas
1.2 Matter and Properties
Comparison of the Three
Physical States
1.2 Matter and Properties
• 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
5
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 Matter and Properties
• Chemical property - result in a
change in composition and can be
observed only through a chemical
reaction
• Chemical reaction (chemical
change) - a process of rearranging,
removing, replacing, or adding atoms
to produce new substances
hydrogen + oxygen  water
reactants
products
5
1.2 Matter and Properties
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 pack melting in spring
e. Decay of leaves in winter
1.2 Matter and Properties
• Intensive properties - a property of
6
matter that is independent of the
quantity of the substance
- Density
- Specific gravity
• Extensive properties - a property of
matter that depends on the quantity of
the substance
- Mass
- Volume
1.2 Matter and Properties
Classification of Matter
• Pure substance - a substance that has only one
component
7
• Mixture - a combination of two or more pure
substances in which each substance retains its
own identity, not undergoing a chemical reaction
1.2 Matter and Properties
Classification of Matter
• Element - a pure substance that cannot be
changed into a simpler form of matter by any
chemical reaction
7
• Compound - a substance resulting from the
combination of two or more elements in a
definite, reproducible way, in a fixed ratio
1.2 Matter and Properties
Classification of Matter
• Mixture - a combination of two or more pure
substances in which each substance retains its own
identity
7
• Homogeneous - uniform composition, particles well
mixed, thoroughly intermingled
• Heterogeneous – nonuniform composition, random
placement
1.2 Matter and Properties
Classes of Matter
1.3 Significant Figures and
Scientific Notation
• Information-bearing digits or figures in a 8
number are significant figures
• The measuring devise used determines the
number of significant figures a
measurement has
• The amount of uncertainty associated with a
measurement is indicated by the number of
digits or figures used to represent the
information
1.3 Significant Figures and
Scientific Notation
Significant figures - all digits in a number
representing data or results that are known
with certainty plus one uncertain digit
1.3 Significant Figures and
Scientific Notation
Recognition of Significant
Figures
1. 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
2. Zeros located between nonzero digits are
significant
• 60.052 has five significant digits
1.3 Significant Figures and
Scientific Notation
Use of Zeros in Significant
Figures
3. Zeros at the end of a number (trailing zeros) are
significant if the number contains a decimal
point. 4.70 has three significant digits
4. Trailing zeros are insignificant if the number
does not contain a decimal point.
100 has one significant digit; 100. has three
5. Zeros to the left of the first nonzero integer are
not significant.
0.0032 has two significant digits
1.3 Significant Figures and
Scientific Notation
How many significant figures are in
the following?
1. 3.400
2. 3004
3. 300.
4. 0.003040
1.3 Significant Figures and
Scientific Notation
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
1.3 Significant Figures and
Scientific Notation
• 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
5340 = 5.34  103
• What if you want to show the above number
has four significant figures?
= 5.340  103
1.3 Significant Figures and
Scientific Notation
• 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
0.0534 = 5.34  10–2
1.3 Significant Figures and
Scientific Notation
Types of Uncertainty
• Error - the difference
between the true value
and our estimation
– Random
– Systematic
• Accuracy - the degree
of agreement between
the true value and the
measured value
• Precision - a measure
of the agreement of
replicate measurements
1.3 Significant Figures and
Scientific Notation
Significant Figures in Calculation of
Results
Rules for Addition and Subtraction
• 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
1.3 Significant Figures and
Scientific Notation
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)
8

2
.
9688692

10
(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  10-8
1.3 Significant Figures and
Scientific Notation
Exact and Inexact Numbers
• Inexact numbers have uncertainty by
definition
• 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
1.3 Significant Figures and
Scientific Notation
Rules for Rounding Off 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
1.3 Significant Figures and
Scientific Notation
Round off each number to three
significant figures:
1. 61.40
2. 6.171
3. 0.066494
1.4 Units and Unit Conversion
Units - the basic quantity of mass, volume or
whatever quantity is being measured
– A measurement is useless without its units
9
• 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
1.4 Units and Unit
Conversion
Basic Units of the Metric System
Mass
Length
Volume
gram
meter
liter
g
m
L
• Basic units are the units of a quantity
without any metric prefix.
1.4 Units and Unit
Conversion
1.4 Units and Unit
Conversion
UNIT CONVERSION
• You must be able to convert between
units
- within the metric system
9
- between the English system and metric system
• The method used for conversion is called
the Factor-Label Method or Dimensional
Analysis
!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!
1.4 Units and Unit
Conversion
• Let your units do the work for you by
simply memorizing connections
between units.
– For example: How many donuts are in
one dozen?
– We say: “Twelve donuts are in a dozen.”
– Or: 12 donuts = 1 dozen donuts
• What does any number divided by
itself equal?
• ONE!
12 donuts
1
1 dozen
1.4 Units and Unit
Conversion
12 donuts
1
1 dozen
• This fraction is called a unit factor
• What does any number times one
equal?
• That number
• Multiplication by a unit factor does
not change the amount – only the unit
1.4 Units and Unit
Conversion
• We use these two mathematical facts to
use the factor label method
– a number divided by itself = 1
– any number times one gives that number
back
• Example: How many donuts are in 3.5
dozen?
• You can probably do this in your head
but try it using the Factor-Label
Method.
1.4 Units and Unit
Conversion
Start with the given information...
12 donuts
3.5 dozen 
1 dozen
= 42 donuts
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...
1.4 Units and Unit
Conversion
Common English System Units
• Convert 12 gallons to units of quarts
1.4 Units and Unit
Conversion
Intersystem Conversion Units
1.4 Units and 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
1.5 Experimental Quantities
• Mass - the quantity of matter in an object
– not synonymous with weight
– standard unit is the gram
• Weight = mass  acceleration due to
gravity
• Mass must be measured on a balance (not a
scale)
1.5 Experimental Quantities
• Units should be chosen to suit the
quantity described
–
–
–
–
A dump truck is measured in tons
A person is measured in kg or pounds
A paperclip is measured in g or ounces
An atom?
• For atoms, we use the atomic mass unit
(amu)
– 1 amu = 1.661  10-24 g
1.5 Experimental Quantities
• Length - the distance between two points
– standard unit is the meter
– long distances are measured in km
– distances between atoms are measured in nm,
1 nm = 10-9 m
• Volume - the space occupied by an object
– standard unit is the liter
– the liter is the volume occupied by 1000
grams of water at 4 oC
– 1 mL = 1/1000 L = 1 cm3
1.5 Experimental Quantities
The milliliter
(mL) and the
cubic centimeter
(cm3) are
equivalent
1.5 Experimental Quantities
• Time
- metric unit is the second
• Temperature - the degree of “hotness”
of an object
10
1.5 Experimental Quantities
Conversions Between Fahrenheit
and Celsius
o
o
o
F - 32
C
1.8
F  1.8 ( C)  32
o
1. Convert 75oC to oF
2. Convert -10oF to oC
1. Ans. 167 oF
2. Ans. -23oC
1.5 Experimental Quantities
Kelvin Temperature Scale
• The Kelvin 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.
K = oC + 273
1.5 Experimental Quantities
Energy
• Energy - the ability to do work
• kinetic energy - the energy of motion
• potential energy - the energy of position
(stored energy)
• Energy is also categorized by form:
•
•
•
•
•
light
heat
electrical
mechanical
chemical
1.5 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
1.5 Experimental Quantities
Units of Energy
• Basic Units:
• calorie or joule
• 1 calorie (cal) = 4.184 joules (J)
• A kilocalorie (kcal) also known as the
large Calorie. This is the same Calorie as
food Calories.
• 1 kcal = 1 Calorie = 1000 calories
• 1 calorie = the amount of heat energy
required to increase the temperature of 1
gram of water 1oC.
1.5 Experimental Quantities
Concentration
Concentration:
– the number of particles of a substance
– the mass of those particles
– that are 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
1.5 Experimental Quantities
Density and Specific Gravity
• Density
11
– 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
mass
m
d

volume
V
1.5 Experimental Quantities
cork
water
brass nut
liquid mercury
1.5 Experimental Quantities
Calculating the Density of a
Solid
• 2.00 cm3 of aluminum are found to weigh
5.40g. Calculate the density of aluminum
in units of g/cm3.
– Use the formula
– Substitute our values
5.40 g
2.00 cm3
= 2.70 g / cm3
mass
m
d

volume
V
1.5 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.
1.5 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)
End of Chapter
one