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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)
The overall perspective with which one sees and interprets the world.
Worldviews
Christian Worldview Science is the discovery of God's Handiwork
in creating matter and all the universe
Naturalistic Worldview Matter is everything and science is the only
path to “truth”.
“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
Gen. 4:22 Metallurgy – Extracted pure metals from
their minerals/ores (raw earth material). Created alloys (mixing of metals for desirable
properties)
Old Testament Chemistry
Exod. 30:25
Apothecary - Used chemicals and herbs for medicinal purposes
“The original pharmacists”
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
Greek Chemistry ~ 430 BC
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
Creation Mandate Gen. 1:26, 28
Why Study Chemistry?
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)
• Critical thinking Skills
• Problem solving
• Deductive logic
• Scientific inquiry
Career FoundationPharmacyMedical
EngineeringDietician
Agriculture Environmental
Material science
Chemistry: A Science for the 21st 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
2 Cu + H2O + CO2 + O2 → Cu(OH)2 (s) + CuCO3 (s)
Oxidized mixture called “Patina”
1886 Today
2-6 yearProcess
Chemistry explains what’s happening Macroscopically on the Microscopic scale
Qualitative observations Describes the quality of an object Color, taste, texture, appearance,
smell, etc. Think Adjectives
Making Observations
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.
A hypothesis is a tentative explanation for a set of observations.
tested modified
Macroscopic Microscopic/Symbolic Explain Observations
The scientific method is a systematic approach to research.
* http://www.wired.com/wiredscience/2013/04/whats-wrong-with-the-scientific-method/
HypotheticalMethod
ActualMethod
A theory is a unifying principle that attempts to explain a body of experimental observations.
• Atomic Theory
• Cell Theory
• Big bang theory
Theories offer explanations for what we observe.
Theories tell us why we should expect it.
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.
Newton's 2nd Law: Force = mass x acceleration
2nd law of thermodynamics: Entropy > 0
Laws describe observations Often mathematical equations
Laws tell us what we should expect
Charles’s Law: V ∝ T
( = ∝ directly proportional)
Matter is anything that occupies space and has mass.
A substance is a form of matter that has a definite composition and distinct properties.
Chemistry is the study of matter and thechanges it undergoes.
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
2. Heterogeneous mixture – composition is not uniform throughout
Solutions (soft drink), gas mixtures (air), solder (Sb/Pb alloy)
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.
Separating Sand/Iron via a magnet
Mixture
*Distillation - Separating two liquid substances by their differing boiling points
Pure
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)
An extensive property depends upon how much matter is being considered.
An intensive property of a material does not depend upon how much matter is being considered.
• mass
• length
• volume
• Density
• Temperature
• Color
•Viscosity
Extensive and Intensive Properties of matter
A physical change does not alter the composition or identity of a substance.
A chemical change alters the composition or identity of the substance(s) involved.
ice meltingsugar dissolving in water
Hydrogen burns in air to form water
Types of Changes
Metal rusting
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.
• 20 elements have been synthetically created by scientists
• 118 elements have been identified• 98 elements occur naturally (some only in trace amounts)
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
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
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
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-
When an atom has equal Protons and Electrons it is Neutralex. a neutral Helium atom contains 2 P+ and 2 e-
aurum
Plumbum
*Hydrargyros "water silver"
Argentum
KaliumFerrum
Elemental symbols
*Many are derived from
their Latin names
*
Natrium
*Wolframite: W containing ore
*
Si ≠ SI
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: SiO4dry 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
CO2
H2
The Three States of Matter: Effect of a Hot Poker on a Block of Ice
solidliquid
gas
Lab Glassware
Borosilicate Glass (SiO2 + B2O3)
• Withstands higher temperatures
• Lower thermal expansion (hot to cold)
• Less likely to shatter
• Used to contain chemicals/reactions
• Used to heat liquids
• Not used to heat solids
Buret Graduated Cylinder
Volumetric Flask
Used Quantitatively
Non-Quantitative
BeakerErlenmeyerFlask
Crucible used instead
A Comparison: The Three States of Matter
Defined shape, incompressible
Undefined shape, incompressible
Undefined shapeCompressible
A Comparison: The Three States of Matter
Kinetic Molecular Theory: describes motion of particles in states of matter
Plasma: Charged Gas,effected by magnetic field,Interacting particles
Little particle motion“locked in place”only vibrations
Greater freedom of motion“particles shift/slide”
Random, fast movement of particles (non-
interacting)
• 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.
Energy is the capacity to do work or produce heat.
• 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)
1st 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
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)
Thermochemistry is the study of heat change in chemical reactions.
Thermal Energy includes all collective kinetic movements & vibrations of particles.
K = 0C + 273.15
°F = x °C + 3295
0 K = -273.15 0C
0 K = -460 ° F
A Comparison of Temperature Scales
Absolute Zero: Theoretical temp where all atomic movements stops
(1724) “Weather/human based”“Water based”(1742)
Absolute Scale(1848)
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)
Endothermic process is any process in which heat has to be supplied to the system from the surroundings. (Feels cool to the touch)
2H2 (g) + O2 (g) 2H2O (l) + energy
H2O (g) H2O (l) + energy
energy + NH4NO3 (s) NH4+1 (aq) + NO3
-1 (aq)
energy + H2O (s) H2O (l)
H2O
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
A 1 kg bar will weigh1 kg on earth
0.1 kg on moon
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 Objecthttps://www.youtube.com/watch?v=ZMByI4s-D-Y
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 UnitsVelocity: m/s
Force: 1 Newton = 1 kg•m/s2
Volume: m3
Used most often in this class, be sure to memorize.
Prefixes can be used to simplify for extremely large or small quantities of base units
“mu”
*not true
• 7 cm = ____________ m
• 300 g = ____________ kg
• 4.7 m = ____________mm
• 9,000 sec = ________Msec
Prefix examples
0.07 m
0.3 kg
4,700 mm
0.009 sec
100 cm = 1 m
1000 g = 1 kg
1 m = 1000 mm
1,000,000 sec = 1 Msec
Volume – SI derived unit for volume is cubic meter (m3)
1 cm3 = (0.01 m)3 = 1 x 10-6 m3
1 L = 1000 mL = 1000 cm3 = 1 dm3
1 mL = 1 cm3
(cm3 is more commonly used)
Density – SI derived unit for density is kg/m3
1 g/cm3 = 1 g/mL = 1000 kg/m3
density = mass
volume
d = mV
*more commonly used
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.
d = mV
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
“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.”
Failed to convert English to
metric units
Scientific NotationThe 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
N x 10nN is the base number between 1 and 10
Exponent (n) is a positive or negative integer
We can factor out powers of 10 to simplify very large or small numbers
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 Decimal moved left so (+) exponent
• 3.2 x 10-8 = 0.000000032 Decimal moved right so (–) exponent
• The base number must be between 1 and 9
.
.
Scientific Notation Practice
Write these in long notation
• 2.0 x 103
• 3.58 x 10-4
• 4.651 x 107
• 9.87 x 10-2
Write these in scientific notation
• 0.00578
• 579
• 0.63
• 96,000
• 0.0140
Mathematics in Scientific Notation568.762
+ n
568.762 = 5.68762 x 102
move decimal left
0.00000772
- n
0.00000772 = 7.72 x 10-6
move decimal right
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
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
Mathematics in Scientific Notation
Bell Ringer
b) Rewrite above numbers using the nearest SI prefix
c) Perform the below mathematics in Sci. Notation
a) Write in scientific 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 significant1.234 kg 4 significant figures
34,000 m 2 significant figures
2) Zeros between nonzero digits are significant606 cm 3 significant figures
50,050 s 4 significant figures
3) Zeros to the left of the first nonzero digit are not significant0.08 mL 1 significant figure
0.00054 ML 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 2 significant figures20.000 g 5 significant figures
5) If a number is less than 1, then only the zeros at the end are significant0.00420 g 3 significant figures0.1000 g 4 significant figures
Significant Figures
Every significant figure is shown when using Scientific notation.
0.001400 m____ 4 significant figures
1.400 x 10-3 Not 1.4 x 10-3
500 mL__ 2 significant figures
5.0 x 102 Not 5 x 102
Example
1.4 Unit Conversions
Determine the number of significant figures in the following measurements:
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 L1.1+
90.492 round off to 90.5
one significant figure after decimal point
3.70-2.91330.7867
two significant figures after decimal point
round off to 0.79
± 50 mL
± 1.0 mL
XX
XX
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
3 sig figsround to 3 sig figs
6.8 ÷ 112.04 = 0.0606926
2 sig figs round to 2 sig figs
= 0.061
Example 1.5Carry 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 NumbersNumbers 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 + 6.703
= 6.67333 = 7
Because 3 is an exact number, not a measured number; It is not used for sigfigs.
= 6.673
• How many feet are in 6.82 yards?
6.82 yards x 3 ft/yard
1 yard = exactly 3 ft by definition
= 20.5 ft = 20 ft
Accuracy – how close a measurement is to the true value
Precision – how close a set of measurements are to each other
accurate&
precise
precisebut
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
3.0x 100% = 3.3% error
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 g) ÷ (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 desired unit
given unit
Train Track ExampleHow many inches are in 3.0 miles?
Identify beginning information
Write measurement as a fraction
3 miles 1
Draw a train track
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
3 miles 1
5280 feet 1 mile
• Write your conversion factor on the track so that miles cancels out and you are left with the unit feet.
Always need same units on opposite sides to cancel out
3.0 miles 1
5280 feet 1 mile
12 inches 1 foot
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.
Train Track Unit Conversions
How many inches are in 3.0 miles?
3.0 miles 1
5,280 feet 1 mile
Multiply all numbers on the topDivide all numbers on the bottom
3.0 x 5,280 x 12 1 x 1 x 1
12 inches 1 foot
= 190,080 inches
Inches are the only remaining unit ✔
= 1.9 x 105 inches
Example
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.)
1.6
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?
70.0 miles 1 hour
1609 meter 1 mile
70.0 x 1,609 x 1 x 1 1 x 1 x 60 x 60
1 hour 60 min
= 31.286 m/s
Meter/sec are the only remaining units ✔
= 31.3 m/s
1 min 60 sec
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
Example 1.7
An average adult has 5.2 L of blood. What is the volume of blood in m3?
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 low-temperature 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.
Liquid nitrogen0.808 g/cm3 = ? kg/m3
Example 1.8 Solution
Solution In Example 1.7 we saw that 1 cm3 = 1 ×10−6 m3. The conversion factors are:
10-6 Mm = 1 m10-3 Km = 1 m
meter (base)102 cm = 1 m
103 mm = 1 m106 m = 1 m
Conversion factors can be written/used 2 ways
1 Mm = 106 m1Km = 103 mmeter (base)
1 cm = 10-2 m1 mm = 10-3 m1 m = 10-6 m
Or
I favor the forms using (+) exponents
Practice Conversion problems• Convert 3 mL to ounces
(33.8 oz = 1 L)
• 1.67 Mm to m
• 2.35 x 1012 inches to cm (1 ft = 0.305 m)
• 3.50x104 L to cLReview
Unit Conversion & Significant Figures: Crash Course Chemistry #2www.youtube.com/watch?v=hQpQ0hxVNTg
• 42.0 km/h to ft/ms
• 0.55 Acres to m2 (247 acre = 1 km2)
• 106 g/mm3 to kg/m3