Introduction to Organic and
Biochemistry
(CHE 124)
Reading Assignment
General, Organic, and Biological Chemistry: An Integrated Approach
4rd. Ed. Ramond
Chapter 1
Science and Measurements
Answers to odd numbered problems in textbook are found in
the book’s index.
What is Chemistry?
• Chemistry - the study of matter and the
changes that it undergoes (e.g. reactions).
Chemistry is the central science. It unifies the
sciences in biology, physics, engineering, medicine,
pharmacy, etc.
What is the scientific Method?
The scientific method is a way of gathering and
interpreting information about chemistry (see
next slide).
– Hypothesis
• tentative explanation (educated guess) for observations
and known facts.
– Theory
• an experimentally tested explanation of an observed
behavior. (a well tested hypothesis)
– Law
• statement that describe things that are consistently and
reproducibly observed. (a well tested theory)
Scientific Method
Creativity
Observation
Form (New) Hypothesis
Experiments
Revise Hypothesis
Discard Hypothesis
Accept Hypothesis
Theory
• Matter
What is Matter?
– anything that has mass and occupies space.
• Weight – measure of gravitational pull against matter.
• Mass – measure of amount of material.
• Phases of Matter
– Solids
• Fixed volume and shape
– Liquids
• Fixed volume, indefinite shape.
• A liquids takes on the shape of the container.
– Gases
• Indefinite shape and volume.
• A gas takes on both the shape and volume of the
container.
Properties of Substances
• Every pure substance has its own unique set of properties that serve to
distinguish it from all other substances.
– Look them up in the chemical literature
– These properties are classified as:
• Intensive properties – independent of amount of substance. ( E.g. m.p., b.p.,
density)
• Extensive properties – dependent on amount of substance. (e.g. mass, volume)
• Chemical properties observed when the substance takes part in a
chemical reaction
• becomes a new substance
– E.g. Heating Mercury (II) oxide to produce oxygen.
– Does the substance react with oxygen?
• Physical properties
• No chemical change is required
• Examples:
– Melting point (m.p.) – temp. when substance changes from solid to liquid.
– Boiling point (b.p.) – temp. when substance changes from liquid to gas.
m
– Density =
d
v
– Solubility = amount of solute that dissolves in a give amount (100g) of solvent at a
specific temp.
– Color
What is Energy?
• Energy
• The ability to do work and / or to transfer heat.
– Potential energy
• stored energy
– Kinetic energy
• energy of motion
Measurements
• Chemistry is a quantitative science. We use the SI system
of measures.
• SI Units International System of Measure
– Common name: the metric system
– Based on the decimal (powers of ten)
– Kg, L, K, °C
• English system is used in the United States.
Measuring Length
• SI unit of Length = meter (m)
– Definition of meter - the distance light travels in
1/299,792,458 of one second
• 1 m = 39.37 in.
• English units
– mile (m), yard (yd.), foot (ft.),inch (in.)
• Instruments used to measure length
– Meter stick
– Micrometer
Measuring Volume
• SI Unit of Volume = Liter (L)
– Volume is derived from SI unit of length.
– Units of volume
• cubic meter (m3) = 1000 L
• cubic centimeter (cm3 or cc)= milliliters (mL)
• English Units
– Gallon (gal.), quart (qt.), pint
(pt.), cup (c), teaspoon (tsp.) ,
table spoon (tbsp.), fluid ounce
(oz.)
• Instruments used to measure
volume
– Graduated cylinder, pipet or buret, digital
micropipet
Measuring Mass
• SI Unit = kilogram (kg)
– Definition of kg
• The kilogram is the unit of mass; it is equal to the
mass of the international prototype of the
kilogram.
• Kilogram (kg) = 1000 g
• 1 gram (g) = 1000 mg
• English Units
– Ton (ton.), pound (lb.), ounce (oz.)
• Instruments used to measure
mass
– Balance
– Scale
Mass vs Weight
• Mass
– Amount of matter in a sample.
• Weight
– The effect of gravity on the matter.
SI Units
Base Quantity
Name
Symbol
* Length
* Mass
* Time
Electric current
* Temperature
* Amount of substance
Luminous intensity
Meter
Kilogram
Second
Ampere
Kelvin
Mole
Candela
m
kg
s
A
K
mol
cd
Source: http://physics.nist.gov/cuu/Units/units.html
Derived SI Units
Derived
Quantity
* Area
* Volume
* Speed, velocity
Acceleration
Name
Symbol
square meter
cubic meter
meter per second
meter per second
squared
m2
m3
m/s
m/s2
Derived SI Units with
Special Names
Derived quantity
force
* pressure, stress
* energy, work, quantity of heat
power, radiant flux
* Celsius temperature
Name
Expression
Expression
Symbol in terms of
in terms of
other SI units SI base units
newton
N
-
m·kg·s-2
pascal
Pa
N/m2
m-1·kg·s-2
joule
J
N·m
m2·kg·s-2
watt
W
J/s
m2·kg·s-3
degree
Celsius
°C
-
K
Metric Prefixes
Prefix
mega
kilo
hecto
deca
Unit
deci
centi
milli
micro
nano
Symbol
Multiple
1,000,000
M
1,000
k
100
h
10
da
1
------0.1
d
0.01
c
0.001
m
0.000001
μ
0.000000001
n
1 x 106
1 x 103
1 x 102
1 x 10
1
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-6
1 x 10-9
English Conversions
Length
Volume
1 mile (m) = 5280 feet (ft.) 1 gallon (gal) = 4 quarts (qt.)
1 ft. = 12 inches (in.)
1 qt. = 2 pints (pt.)
1 yard (yd.) = 3 ft.
Mass
1 pt. = 2 cups (c.)
1 pt. = 16 fluid ounces (fl. oz.)
1 c. = 8 fl. oz.
1 ton = 2000 pounds (lbs.) 1 in3 = 16.387 cm3
1 lb. = 16 ounces (oz.)
2 tablespoons (T or tbsp) = 1 fl. oz
English to Metric Conversions
Length
1 in. = 2.54 cm (exact)
1 m = 39.37 in
1 mi = 1.609 km
Volume
1 ft3 = 28.32 L
1 L = 1.057 qt
1 gal. = 3.785 L
1 tsp. = 5 mL
1 tbsp = 15 mL
1 mL = 15 drops (gtt)
Mass
1 lb. = 453.6 g = 0.4536 kg
1 g = 0.03527 oz.
1 kg = 2.2 lbs.
Typical Conversions Problems
• The dosage on a bottle of medicine reads:
“Take 2 tablespoons every twelve hours.”
Convert this volume to mL.
• The box at the pet store states: “Aquarium
volume is 55 gal.”. Convert this volume to
liters.
• The distance (length) from Clinton to
Vicksburg is appr. 32 miles. What is this
distance in cm?
• A marathon is defined as 42.195 km. What
is this distance in miles?
Measuring Temperature
• Factor that determines the direction of heat flow
• SI units = °C or K
– Fahrenheit (°F)
• Named after German instrument maker Daniel
Fahrenheit (1686 -1736)
– Celsius (°C) (old name centigrade)
• Named after Swedish astronomer Anders Celsius
(1701-1744)
– Kelvin (K)
• Do not use (°) or degree in relation to K.
• Defined as 1/273.16 of the difference between the
lowest attainable temp. (0K) and the triple point of
water (0.01 °C)
•
Instruments to measure
– Mercury thermometer
• Mercury expands and contracts as temperature changes.
Tube contains only 2% of the Hg in thermometer.
– Digital thermometer
•
•
Water boils = 212 °F = 100 °C = 373.15 K
Water freezes = 32 °F = 0 °C = 273.15 K
– Triple point is the temp. /pressure combination at which
water is capable of coexisting as a solid, liquid and gas.
Converting Temperature Scales
• Comparing °F to °C
– 0 °C is 32 °F
• Freezing point of water
F  1.8C  32
– 100 °C is 212 °F
• Boiling point of water
– There are 180
Fahrenheit degrees for
every 100 Celsius
degrees, so each °C is
1.8 times larger than
each °F
• Comparing °C to K
– Celsius degree and
Kelvin degree are the
same size.
K  C  273.15
Scientific Notation
• A way of dealing with very large or very
small numbers.
• Show examples on the board:
– Avogadro number 6.022 X 1023
– 83,000
– 0.000056
– See table 1.3 p.12
Accuracy vs Precision
• Accuracy
– How close a reported value is to the real
value. (see next slide about error).
• Precision
– A measure of how close repeated
measurements are to one another.
Uncertainties in Measurements
• How much solution is in the large graduated cylinder?
• How much solution is in the small graduated cylinder?
• Do these two measurements have the same uncertainty?
Uncertainties in Measurements
• Three volume measurements with their
uncertainties
– Large graduated cylinder, 8 ± 1 mL
– Small graduate cylinder, 8.0 ± 0.1 mL
– Pipet or buret, 8.00 ± 0.01 mL
• To denote how much uncertainty is in a
measurement, Significant figures are used.
• Significant Figures
– Every measurement carries uncertainty
– All measurements must include estimates of uncertainty
with them
– There is an uncertainty of at least one unit in the last digit
• Text convention
– Uncertainty of ± in the last digit is assumed but not stated
Significant Figures
• Significant figures are meaningful digits in
measurements
– In 8.00 mL, there are three significant figures
– In 8.0 mL, there are two significant figures
– In 8 mL, there is one significant figure
Ambiguity in Significant Figures
• Consider the measurement, 500 g
– If the measurement was made to the nearest 1 g,
all three digits are significant
– If the measurement was made to the nearest 10
g, only two digits are significant
– Resolve by using scientific notation
• 5.00 X 102 g
• 5.0 X 102 g
• See Table 1.5 p. 17
Rounding
• Rounding off numbers
– If the first digit to be discarded is 5 or greater,
round up
– If the first digit to be discarded is 4 or smaller,
round down
Significant Figures in Calculations
• Addition and Subtraction
– Count the number of decimal places in each number
– Round off so that the answer has the same number of decimal
places as the measurement with the greatest uncertainty (i.e.,
the fewer number of decimal places).
• Multiplication and Division
– When multiplying or dividing two numbers, the answer is rounded
to the number of significant figures in the less (or least in the case
of three or more) measurements
– 2.40 X 2 = 5
• Exact Numbers
– Exact numbers carry an infinite number of significant figures
• Exact numbers do not change the number of significant figures in a
calculation
– Example: The numbers 1.8 and 32 in the conversion between
Fahrenheit and Celsius
Dimensional Analysis / Factor Label /
Converting Units
• In many cases throughout your study of chemistry, the
units (dimensions) will guide you to the solution of a
problem
• A correct answer must have the NUMBER and UNITS!
• Conversion factors are used to convert one set of units to
another
– Only the units change
– Conversion factors are numerically equal to 1
• 1L = 1000 cm3 Choose a conversion factor that puts the initial units in
the denominator
– The initial units will cancel
– The final units will appear in the numerator
1L
1000 cm3

1
3
3
1000 cm
1000 cm
Some Examples
• Convert 25 mL to L.
• Convert 200 pounds to grams.
• Convert 20 miles to kilometers.
• Convert 25 microliters to liters.
Properties of Substances
• Every pure substance has its own unique set of properties that serve to
distinguish it from all other substances.
– Look them up in the chemical literature
– These properties must be intensive.
• Intensive properties – independent of amount of substance. ( E.g. m.p., b.p.,
density)
• Extensive properties – dependent on amount of substance. (e.g. mass, volume)
• Chemical properties observed when the substance takes part in a
chemical reaction
• becomes a new substance
– E.g. Heating Mercury (II) oxide to produce oxygen.
– Does the substance react with oxygen?
• Physical properties
• No chemical change is required
• Examples:
– Melting point (m.p.) – temp. when substance changes from solid to liquid.
– Boiling point (b.p.) – temp. when substance changes from liquid to gas.
– Density =
m
d
v
– Solubility = amount of solute that dissolves in a give amount (100g) of solvent at a
specific temp.
– Color
Density
• The density of a substance is its
mass divided by its volume. (the
amount of mass contained in a
given volume.)
• Units = g / mL
• Density of:
– Water is 1 g / mL
• Temperature must be stated.
– density changes with changes in temperature.
• Note table 1.8 p. 22
m
d
v
Specific Gravity
• Relates density of a substance to that of water.
– Measured using refractometeres, hydrometers or test
strips.
– Used to determine
•
•
•
•
acid level in car batteries
antifreeze level in car radiators
alcohol content in beer and wine
Urine to diagnose kidney problems
– Temperature must be specified since the density of the
substance and water vary, but not necessarily at the
same rate.
Specific Gravity = Density of substance = 0.785 g/mL = 0.785
Density of water
1.00 g/mL
Specific Heat
• Relates energy (in calories), mass (in grams),
and temperature (in degrees Celsius).
• Units = cal / g ∙ °C
• Relates the mass, temperature, and energy.