Transcript Chemistry: The Study of Change

Chemistry:
The Study of
Matter
Chapter 1
Ch. 1 Homework
• Ch. 1a on Matter Classification
2, 3a-d, 5-9, 12-15 (both editions)
• Ch. 1b on Measurements and Conversions
18, 20, 22-23, 29-31, 33, 35, 39a-c, 40, 43, 47
(both editions)
Worldviews
The overall perspective with which
one sees and interprets the world.
Naturalistic Worldview

Matter is everything and science is the only
path to “truth”.
Christian Worldview

Science is the discovery of God's Handiwork
in creating matter and all the universe
“They exchanged the truth about God for a lie, and worshiped
and served created things rather than the Creator-who is forever
praised.” Romans 1:25
Old Testament Chemistry

Gen. 4:22
 Metallurgy – Extracted pure metals from

their minerals/ores (raw earth material).
Created alloys (mixing of metals for desirable
properties)

Exod. 30:25
Apothecary - Used chemicals and
herbs for medicinal purposes
“The original pharmacists”
Greek Chemistry ~ 430 BC

Democritus's theory – Philosophical atomism (no evidence)
 All matter is made up of tiny identical atoms and the difference
in materials is based on the shape, position, and arrangement of
these atoms.

"Atomos" - indivisible
Alchemists


The "original" chemists
Attempted to make gold from other substances.
 Impossible challenge
 (without nuclear fusion/fission reactions)
 Nevertheless, resulted in organized approach to science
 Laboratory techniques
 Equipment
 Terminology
Why Study Chemistry?
Creation Mandate


Gen. 1:26, 28
God blessed them, and God said to them, "Be fruitful, multiply, fill the earth,
and subdue it. Rule the fish of the sea, the birds of the sky, and every creature
that crawls on the earth." (Genesis 1:28 HCSB)
Career Foundation
Pharmacy
Medical
Engineering
Dietician
Agriculture
Environmental
Material science
• Critical thinking Skills
• Problem solving
• Deductive logic
• Scientific inquiry
Chemistry: A Science for the
st
21
Century
• Materials and Technology
• Plastics, ceramics, liquid crystals
• Room-temperature superconductors?
• Molecular computing?
Binary data stored in DNA
• Food and Agriculture
• Genetically modified crops
• “Natural” pesticides
• Specialized fertilizers
GFP: Green fluorescent protein
Fields of Chemistry
• Organic - carbon containing compounds (synthesis,
plastics, drugs)
• Inorganic - all elements minus carbon (metals and
coordinating elements)
Fields of Chemistry
• Biochemistry – organic chemical processes in living
things (Biomolecules: proteins, DNA, lipids, carbohydrates)
• Analytical – Create/improve chemical techniques used in all
branches for precise quantitative measurements. (Purification,
sample analysis, water/soil testing)
• Physical - foundational theories, detailed study of
interaction and energy changes (e- probability,
thermodynamics, quantum)
The Study of Chemistry: We observe the Macroscopic
Macroscopic
Microscopic
Chemistry explains
what’s happening
Macroscopically on
the Microscopic scale
1886
2-6 year
Process
Today
2 Cu + H2O + CO2 + O2 →
Cu(OH)2 (s) + CuCO3 (s)
Oxidized mixture called “Patina”
Making Observations
Qualitative observations



Describes the quality of an object
Color, taste, texture, appearance,
smell, etc.

Think Adjectives
Quantitative observations



Describes an object using numbers
Count, length, weight, volume

Think Units
List 2 observations of each
type you could make
The scientific method is a systematic approach to
research.
Macroscopic
Microscopic/Symbolic
Explain Observations
A hypothesis is a tentative explanation for a set of
observations.
tested
modified
The scientific method is a systematic approach to
research.
Hypothetical
Method
Actual
Method
* http://www.wired.com/wiredscience/2013/04/whats-wrong-with-the-scientific-method/
A theory is a unifying principle that attempts to
explain a body of experimental observations.

Theories offer explanations for
what we observe.
• Atomic Theory
• Cell Theory

Theories tell us why we should
expect it.
• Big bang theory
Do not confuse scientific theories as improbable explanations filled
with inconsistency. They are often incapable of absolute proof, but
all available data are still in support of them.
A law is a concise statement of a relationship
between phenomena that is always the same
under the same conditions.

Laws describe observations
 Often mathematical equations

Laws tell us what we should expect
(∝ = directly proportional)
Newton's 2nd Law:
Force = mass x acceleration
2nd law of thermodynamics: Entropy > 0
Charles’s Law: V ∝ T
Chemistry is the study of matter and the
changes it undergoes.
Matter is anything that occupies space and has
mass.
A substance is a form of matter that has a definite
composition and distinct properties.
liquid nitrogen
gold ingots
silicon crystals
A mixture is a combination of two or more substances in
which the substances retain their distinct identities.
1.
Homogenous mixture – composition of the
mixture is the same throughout
Solutions (soft drink), gas mixtures (air),
solder (Sb/Pb alloy)
2.
Heterogeneous mixture – composition is
not uniform throughout
cement, oil and water,
iron filings in sand,
insoluble compounds
Heterogeneous or Homogenous?
• Chicken Broth
• Vegetable beef soup
• Air
• Vinaigrette dressing (oil/water base)
• Salt water
Mixtures can be separated into their pure components by
some physical means.
*Distillation - Separating two liquid substances by
their differing boiling points
Separating Sand/Iron via a magnet
Pure
Mixture
Physical Properties: can be measured or observed without
changing the composition or identity of a substance.
•
Density: amount of mass per volume of space
•
Malleability: Hammered into a thin sheet
•
Ductility: Drawn into long thin strings
•
Conductivity: Ability to transfer either heat and/or electricity
•
Phase transition temperatures: temp. where melting/boiling occurs
•
Appearance: color, luster
•
Solubility: amount dissolvable in solvent (water)
•
Hardness: measured by Mohs scale (1: Talc - 10: diamond)
Extensive and Intensive Properties of matter
An extensive property depends upon how much matter is
being considered.
• mass
• length
• volume
An intensive property of a material does not depend
upon how much matter is being considered.
• Density
• Temperature
• Color
•Viscosity
Types of Changes
A physical change does not alter the composition or
identity of a substance.
ice melting
sugar dissolving
in water
A chemical change alters the composition or
identity of the substance(s) involved.
Metal rusting
Hydrogen burns in
air to form water
Chemical Properties: A chemical change must
occur to observe:
 Temperature change
 Color change
 Gas production (effervescence) (at constant P&T)
 Solid production from solution (precipitation)
 Flammability
 Toxicity
 Acidity/Basicity
Physical or Chemical Change?
• Grinding coffee beans
• Food rotting
• Lighting a match
• Cutting paper in half
• water evaporating to vapor
• Jewelry tarnishing
• Dissolving salt in water
An element is a substance that cannot be separated
into simpler substances by chemical means.
The first person to discover an element using scientific inquiry was Hennig
Brand, a German scientist who discovered phosphorus (P) in 1649.
In 1789, a French scientist, Antoine Lavoisier defined what was meant by a
chemical element and drew a table that contained 33 known elements
• 118 elements have been identified
• 98 elements occur naturally (some only in trace amounts)
• 20 elements have been synthetically created by scientists
Atoms: the basic particles that make up the different elements
• Ex. Li, Be, B, C, N, O, F, Ne, Au
• Either 1 or 2 letter symbol; first letter capitalized
Atoms possess subatomic particles:
Neutrons (N0) - no charge, but have mass
Protons (P+) - positively charged and have mass
Electrons (e-) - negatively charged, but little mass
When an atom has equal Protons and Electrons it is Neutral
ex. a neutral Helium atom contains 2 P+ and 2 e-
Ion: When P+ and e- are unbalanced in an atom, it is Charged.
ex. an ionized Sodium ion (Na+1) has 11 P+ and 10 e-
Elemental symbols
Si ≠ SI
aurum
Kalium
Ferrum
Plumbum
Argentum
Natrium
*
*
*Many are derived from
their Latin names
*Hydrargyros
"water silver"
*Wolframite:
W containing ore
A compound is a substance composed of atoms of
two or more different elements chemically united
(bonded) in fixed proportions.
Compounds can only be separated (broken down) into their
pure components (elements) by chemical means.
Lithium fluoride: LiF
Quartz: SiO4
dry ice – CO2 (carbon dioxide)
*Non-compounds are not necessarily always monoatomic (C, He):
Can have many element atoms in a substance: P4, S8, Cl2
Review of the Nucleus:
The Nucleus: Crash Course Chemistry #1
https://www.youtube.com/watch?v=FSyAehMdpyI
Classifications of Matter
ex. Carbonated
Water
CO2
+
H2O
C
O2
H2
The Three States of Matter: Effect of a Hot Poker on
a Block of Ice
gas
liquid
solid
Lab Glassware
Non-Quantitative
Borosilicate Glass (SiO2 + B2O3)
• Withstands higher temperatures
Erlenmeyer
Flask
Beaker
• Lower thermal expansion (hot to cold)
• Less likely to shatter
Used Quantitatively
• Used to contain chemicals/reactions
• Used to heat liquids
• Not used to heat solids
Crucible
used instead
Buret
Graduated
Cylinder
Volumetric
Flask
A Comparison: The Three States of Matter
Defined shape,
incompressible
Undefined shape,
incompressible
Undefined shape
Compressible
A Comparison: The Three States of Matter
Kinetic Molecular Theory: describes
motion of particles in states of matter
Random, fast
movement of
Greater freedom of motion particles (nonLittle particle motion
“particles shift/slide”
interacting)
“locked in place”
only vibrations
Plasma:
Charged Gas,
effected by
magnetic field,
Interacting
particles
Energy is the capacity to do work
or produce heat.
• Mechanical energy is sum of kinetic
(energy of motion) and potential
(energy of position).
• Thermal energy is the energy
associated with the random motion of
atoms and molecules (heat).
• Encompasses all kinetic energy of
particles.
• Electrical energy – derived from electron potential energy
(Ni → Ni+2 + 2e-)
• Chemical energy is the energy stored
within the bonds of chemical substances
•
Nuclear energy is the energy stored within the collection of
neutrons and protons in the atom (very exothermic)
• Electromagnetic energy is the energy associated
with electricity and magnetic fields. (visible
light, Infrared, UV, Gamma, radiowaves).
• Acoustic energy is the movement of
particles (kinetic) moving in periodic
waves (sound waves).
Mass – Energy Equivalence

Matter can be converted to energy
 First proposed relationship by
Isaac Newton (1717)
 Related by a constant

Einstein was the first to derive the
equivalence (1905)
st
1


Law of Thermodynamics: Conservation of Energy
Energy (or matter) is never gained or lost in a closed system
Energy is only converted from one form to another


Gasoline – Stored chemical energy (in the bonds)
Gasoline + Oxygen combusts → CO2 + H2O + ...
 Produces steam to drive pistons (kinetic energy)
 Produces heat (thermal energy)
 Produces light
 Produces sound
Thermochemistry is the study of heat change in
chemical reactions.
Thermal Energy includes all collective kinetic
movements & vibrations of particles.
Heat is the transfer of thermal energy between
two bodies that are at different temperatures.
• Always flows from high to low energy
Temperature is a relative measure of the
thermal energy.
Temperature  Thermal Energy
As Temp ↑, Thermal Energy ↑
(particles move faster/collide more)
A Comparison of Temperature Scales
°F
9
=
x °C + 32
5
K = 0C + 273.15
0 K = -273.15 0C
0 K = -460 ° F
Absolute Scale
(1848)
“Water based” (1724) “Weather/human based”
(1742)
Absolute Zero: Theoretical temp where all atomic movements stops
Example 1.3 Temperature conversion
a) Solder is an alloy made of tin and lead that is used in
electronic circuits. A certain solder has a melting point
of 224°C. What is its melting point in degrees
Fahrenheit?
(b) Helium has the lowest boiling point of all the elements
at -452°F. Convert this temperature to degrees Celsius.
(c) Mercury, the only metal that exists as a liquid at room
temperature, melts at -38.9°C. Convert its melting point
to Kelvins.
Example
1.3 Solution
(a)This conversion is carried out by writing
(b)Here we have
(c)The melting point of mercury in Kelvins is given by
Exothermic process is any process that gives off heat –
transfers thermal energy from the system to the surroundings.
(Feels hot to the touch)
2H2 (g) + O2 (g)
H2O (g)
2H2O (l) + energy
H2O (l) + energy
Endothermic process is any process in which heat has to
be supplied to the system from the surroundings. (Feels
cool to the touch)
energy + NH4NO3 (s)
energy + H2O (s)
H2O
NH4+1 (aq) + NO3-1 (aq)
H2O (l)
Review of Energy:
Energy and Chemistry: Crash Course Chemistry #17
www.youtube.com/watch?v=GqtUWyDR1fg
Matter - anything that occupies space and has mass
mass – measure of the quantity of matter
SI unit of mass is the kilogram (kg)
1 kg = 1000 g = 1 x 103 g
weight – force that gravity exerts
on an object
weight = mass x g (F = m•a)
on earth, g = 1.0
on moon, g ~ 0.1
La Grande K
1 Kg Pt/Ir alloy
World’s Roundest Object
https://www.youtube.com/watch?v=ZMByI4s-D-Y
A 1 kg bar will weigh
1 kg on earth
0.1 kg on moon
International System of Units (SI) Base Units
Utilized in this class
Used as Relative Standards for comparison
All other units are derived from these units and are known as Derived Units
Velocity: m/s
Force: 1 Newton = 1 kg•m/s2
Volume: m3
Prefixes can be used to simplify for extremely large
or small quantities of base units
“mu”
Used most often in this class, be sure to memorize.
*not true
Prefix examples
0.07 m
• 7 cm = ____________
m
100 cm = 1 m
0.3 kg
• 300 g = ____________
kg
1000 g = 1 kg
4,700 mm
• 4.7 m = ____________mm
1 m = 1000 mm
sec
• 9,000 sec = 0.009
________Msec
1,000,000 sec = 1 Msec
Volume – SI derived unit for volume is cubic meter (m3)
(cm3 is more commonly used)
1 cm3 = (0.01 m)3 = 1 x 10-6 m3
1 L = 1000 mL = 1000 cm3 = 1 dm3
1 mL = 1 cm3
Density – SI derived unit for density is kg/m3
1 g/cm3 = 1 g/mL = 1000 kg/m3
*more commonly used
density =
m
d= V
mass
volume
Example
1.1
Gold is a precious metal that is chemically unreactive.
It is used mainly in jewelry, dentistry, and electronic devices.
A piece of gold ingot with a mass of 301 g has a volume of
15.6 cm3. Calculate the density of gold.
gold ingots
Example
1.2
The density of mercury, the only metal that is a liquid at room
temperature, is 13.6 g/mL. Calculate the mass of 5.50 mL of the
liquid.
m
d= V
Chemistry In Action
On 9/23/99, $125M Mars Climate Orbiter entered
Mars’ atmosphere 100 km (62 miles) lower than
planned and was destroyed by heat.
1 lb = 1 N
1 lb = 4.45 N
Failed to convert
English to
metric units
“This is going to be the
cautionary tale that will be
embedded into introduction to
the metric system in
elementary school, high
school, and college science
courses till the end of time.”
Scientific Notation
The number of atoms in 12 g of carbon:
602,200,000,000,000,000,000,000
6.022 x 1023
The mass of a single carbon atom in grams:
0.0000000000000000000000199
1.99 x 10-23
We can factor out powers of 10 to simplify very large or small numbers
N is the base number
between 1 and 10
N x 10n
Exponent (n) is a positive
or
negative integer
Scientific Notation
• A base number that is multiplied by a factor of 10
• Base x 10exponent
• To write the number out in long notation, move the decimal to
left or right according to the exponent.
• 3.2 x 108 = 320,000,000
.
• 3.2 x 10-8 = 0.000000032.
Decimal moved left so (+) exponent
Decimal moved right so (–) exponent
• The base number must be between 1 and 9
Scientific Notation Practice
Write these in scientific
notation
• 0.00578
• 579
• 0.63
• 96,000
• 0.0140
Write these in long
notation
• 2.0 x 103
• 3.58 x 10-4
• 4.651 x 107
• 9.87 x 10-2
Mathematics in Scientific Notation
568.762
0.00000772
move decimal left
+n
568.762 = 5.68762 x 102
move decimal right
-n
0.00000772 = 7.72 x 10-6
Addition or Subtraction: Must have same exponent
1. Write each quantity with the
same exponent n
2. Combine N1 and N2
3. The exponent, n, remains the
same
4.31 x 104 + 3.9 x 103 =
4.31 x 104 + 0.39 x 104 =
4.70 x 104
Mathematics in Scientific Notation
Multiplication: Add exponents
1. Multiply N1 and N2
2. Add exponents n1 and n2
(4.0 x 10-5) x (7.0 x 103) =
(4.0 x 7.0) x (10-5+3) =
28 x 10-2 =
2.8 x 10-1
Division: Subtract exponents
1. Divide N1 and N2
2. Subtract exponents n1 and n2
8.5 x 104 ÷ 5.0 x 109 =
(8.5 ÷ 5.0) x 104-9 =
1.7 x 10-5
Bell Ringer
a) Write in scientific notation
• 8,705,000 m
• 0.0000045 L
• 0.00237 sec
• 9,300 g
b) Rewrite above numbers using the nearest SI prefix
c) Perform the below mathematics in Sci. Notation
• 9.01 x 103 g + 3.8 x 102 g
• 3.98 x 10-2 m – 8.2 x 10-3 m
• (2.61 x 107 m) x (9.87 x 10-2 m)
• (8.4 x 109 g) ÷ (2.0 x 104 L)
Significant Figures: Used to prevent uncertainty from rounding of
various measured quantities with various levels of precision.
1) Any digit that is not zero is significant
1.234 kg
34,000 mm
4 significant figures
2 significant figures
2) Zeros between nonzero digits are significant
606 cm
50,050 s
3 significant figures
4 significant figures
3) Zeros to the left of the first nonzero digit are not significant
0.08 mL
0.00054 ML
1 significant figure
2 significant figures
4) If a number is greater than 1, then all zeros to the right of the decimal point are
significant
2.0 mg
20.000 g
2 significant figures
5 significant figures
5) If a number is less than 1, then only the zeros at the end are significant
0.00420 g 3 significant figures
0.1000 g 4 significant figures
Significant Figures
Every significant figure is shown when using
Scientific notation.
____ m
0.001400
4 significant figures
1.400 x
-3
10
Not 1.4 x
-3
10
__ mL
500
2 significant figures
5.0 x
2
10
Not 5 x
2
10
Exampl 1.4 Unit Conversions
e
Determine the number of significant figures in the following
measurements:
(a) 478 cm
(d) 0.0430 kg
22
10
(b) 600,001 g
(e) 1.310 ×
(c) 0.85 m
(f) 7000 mL
atoms
Example
1.4 Solution
(a) 478 cm -- Three, because each digit is a nonzero digit.
(b) 600,001- Six, because zeros between nonzero digits are significant.
(c) 0.825 m -- Three, because zeros to the left of the first nonzero
digit do not count as significant figures.
(d) 0.0430 kg -- Three. The zero after the nonzero is significant
because the number is less than 1.
(e) 1.310 × 1022 atoms -- Four, because the number is greater than one
so all the zeros written to the right of the decimal point count as
significant figures.
Example
1.4 solution
(f)7000 mL -- This is an ambiguous case. The number of significant
figures may be four (7.000 × 103), three (7.00 × 103), two (7.0 × 103),
or one (7 × 103).
This example illustrates why scientific notation must be used to
show the proper number of significant figures.
If no decimal is present it is usually assumed only non-zeros are
significant. If a decimal is present, than all zero’s are significant.
7,000 mL ≠ 7,000. mL
They display differing degrees of precision.
Significant Figures
Addition or Subtraction
The answer cannot have more digits to the right of
the decimal point than any of the original numbers.
Use the least precise number.
89.392 L
+ 1.1 XX
90.492
± 50 mL
one significant figure after decimal point
round off to 90.5
± 1.0 mL
3.70 XX
-2.9133
0.7867
two significant figures after decimal point
round off to 0.79
Significant Figures
Multiplication or Division
The number of significant figures in the result is set by the original
number that has the smallest number of significant figures.
4.51 x 3.0006 = 13.532706 = 13.5
round to 3 sig figs
3 sig figs
6.8 ÷ 112.04 = 0.0606926 = 0.061
2 sig figs
round to 2 sig figs
Example
1.5
Carry out the following arithmetic operations to the correct number of
significant figures:
(a) 11,254.1 g + 0.1983 g
(b) 66.59 L − 3.113 L
(c) 8.16 m × 5.1355 kg
(d) 0.0154 kg ÷ 88.3 mL
(e) 2.64 × 103 cm + 3.27 × 102 cm
Example
1.5 Solution
Solution In addition and subtraction, the number of decimal places in
the answer is determined by the number having the lowest number of
decimal places. In multiplication and division, the significant number
of the answer is determined by the number having the smallest number
of significant figures.
(a)
(b)
Example
1.5 Solution
(c)
(d)
(e) First we change 3.27 × 102 cm to 0.327 × 103 cm and then carry
out the addition (2.64 cm + 0.327 cm) × 103. Following the procedure
in (a), we find the answer is 2.97 × 103 cm.
Significant Figures
Exact Numbers
Numbers from definitions or numbers of objects are
considered to have an infinite number of significant figures.
•The average of three measured lengths: 6.64, 6.68 and 6.70?
6.64 + 6.68 + 6.70
3
= 6.67333 = 7
= 6.673
Because 3 is an exact number, not a
measured number; It is not used for sigfigs.
• How many feet are in 6.82 yards?
6.82 yards x 3 ft/yard = 20.5 ft = 20 ft
1 yard = exactly 3 ft by definition
Accuracy – how close a measurement is to the
true value
Precision – how close a set of measurements are
to each other
accurate
&
precise
precise
but
not accurate
not accurate
&
not precise
Percent Error
A way to determine how accurate your measurements are
to a known value.
|Obtained value – Actual value| x 100%
Actual Value
Ranges between 0 and 100%
Ex. I weigh a 3 kg block on three different scales:
3.2 kg, 3.0 kg, 3.1 kg = 3.1 kg average
3.1 – 3.0 x 100% = 3.3% error
3.0
Precision also indicates to what degree we
know our measurement. (Arithmetic precision)
A measurement of 8.0 grams could be made on an average
countertop food scale (balance). (~$20)
A high-precision milligram scale could weigh the same sample
with a much higher precision (8.0235 grams) (~$1,500)
Bell Ringer
a) Perform the below mathematics in Sci. Notation. using
Significant Figures in your answer.
1. (9.8 x 105 g) + (6.75 x 104 g)
2. (5.98 x 10-6 m) – (7 x 10-8 m)
3. (2.612 x 1010 m) x (9.87 x 10-3 m)
4. (7 x 102 mg) ÷ (1.875 x 104 mL)
b) Rewrite the first 2 solutions
using the nearest SI prefix
Dimensional Analysis Method of Solving
Problems (Train-Track)
1. Determine which unit conversion factors are needed
2. Carry units through calculation
3. If all units cancel except for the desired unit(s), then the
problem was solved correctly.
given quantity x conversion factor = desired quantity
given unit x
desired unit
given unit
= desired unit
Train Track Example
How many inches are in 3.0 miles?
Identify beginning information
Draw a train track
3 miles
1
Write measurement as a fraction
Train Track Example
How many inches are in 3.0 miles?
• We are going from a larger measurement to a smaller one.
• Find a conversion factor you know that changes miles into
something smaller.
Conversion Factor: 1 mile = 5,280 feet
• Write your conversion factor on the track so that miles
cancels out and you are left with the unit feet.
3 miles
1
5280 feet
1 mile
Always need same units on
opposite sides to cancel out
Train Track Example
How many inches are in 3.0 miles?
We now need another conversion factor between
Feet and Inches: 1 foot = 12 inches
Again, place conversion factor so that the previous unit
cancels out.
3.0 miles
1
5280 feet
1 mile
12 inches
1 foot
Train Track Unit Conversions
How many inches are in 3.0 miles?
3.0 miles
1
5,280 feet
1 mile
12 inches
1 foot
Inches are the only
remaining unit ✔
Multiply all numbers on the top
Divide all numbers on the bottom
3.0 x 5,280 x 12
1x1x1
= 190,080 inches
= 1.9 x 105 inches
Example
1.6
A person’s average daily intake of glucose (a form of sugar) is 0.0833
pound (lb). What is this mass in milligrams (mg)?
(1 lb = 453.6 g.)
A metric conversion is then needed to convert grams to milligrams (1
mg = 1 × 10−3 g)
(Or one could write: 1,000 mg = 1 g)
Either Conversion factor will work
Example
1.6 Solution
Solution The sequence of conversions is
Using the following conversion factors
we can now write:
2-D Conversion Problems
Convert
70.0 miles/hour to m/s?
We convert one unit followed by the other
Conversion factors: 1 mile = 1,609 meters; 1 hour = 60 min; 1 min = 60 sec
*Note: to cancel out hours (on bottom) it must appear again on the top
70.0 miles
1 hour
1609 meter
1 mile
1 hour
60 min
1 min
60 sec
Meter/sec are the only remaining units ✔
70.0 x 1,609 x 1 x 1
1 x 1 x 60 x 60
= 31.286 m/s
= 31.3 m/s
Example
1.7
An average adult has 5.2 L of blood. What is the
3
volume of blood in m ?
Strategy
How many conversion factors are needed for this problem?
L → cm3 → m3
Recall that 1 L = 1,000 cm3 and (1 cm)3 = (0.01 m)3.
Example
1.7 Solution
We need two conversion factors here: one to convert liters to cm3 and
one to convert centimeters to meters:
Because the second conversion factor deals with length
(cm and
m) and we want volume here, it must therefore be cubed to give
This means that 1 cm3 = 1 × 10−6 m3.
Example
1.7 Solution
Now we can write
Check From the preceding conversion factors you can show that 1 L
= 1 × 10−3 m3. Therefore, 5 L of blood would be equal to 5 × 10−3
m3, which is close to the answer.
Example
1.8
Liquid nitrogen is obtained from
liquefied air and is used to prepare
frozen goods and in lowtemperature research.
The density of the liquid at its
boiling point (−196°C or 77 K) is
0.808 g/cm3. Convert the density
to units of kg/m3.
0.808
g/cm3
=?
kg/m3
Liquid nitrogen
Example
1.8 Solution
Solution In Example 1.7 we saw that 1 cm3 = 1 ×10−6 m3. The
conversion factors are:
Conversion factors can be written/used 2 ways
6
10
1 Mm =
m
3
1Km = 10 m
meter (base)
-2
1 cm = 10 m
-3
1 mm = 10 m
-6
1 mm = 10 m
-6
10
Mm = 1 m
-3
10 Km = 1 m
meter (base)
Or
2
10 cm = 1 m
3
10 mm = 1 m
6
10 mm = 1 m
I favor the forms using (+) exponents
Practice Conversion problems
• Convert 3 mL to ounces
(33.8 oz = 1 L)
• 42.0 km/h to ft/ms
• 1.67 Mm to mm
• 0.55 Acres to m2
(247 acre = 1 km2)
• 2.35 x 1012 inches to cm
(1 ft = 0.305 m)
• 106 g/mm3 to kg/m3
• 3.50x104 mL to cL
Review
Unit Conversion & Significant
Figures: Crash Course Chemistry #2
www.youtube.com/watch?v=hQpQ0hxVNTg