Chemistry: The Study of Matter

Chemistry: The Study of Matter
Chapter 1

Ch. 1 Homework Ch. 1a on Matter Classification
2, 3a-d, 5-9, (both editions) Ch. 1b on Measurements and Conversions 18, 20, 22-23, 29-31, 33, 35, 39a-c, 40, 43, (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? Career Foundation Pharmacy
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

GFP: Green fluorescent protein
Chemistry: A Science for the 21st Century Materials and Technology Plastics, ceramics, liquid crystals Room-temperature superconductors? Molecular computing? Binary data stored in DNA GFP: Green fluorescent protein Food and Agriculture Genetically modified crops “Natural” pesticides Specialized fertilizers

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)

2 Cu + H2O + CO2 + O2 → Cu(OH)2 (s) + CuCO3 (s)
The Study of Chemistry: We observe the Macroscopic Macroscopic Microscopic Chemistry explains what’s happening Macroscopically on the Microscopic scale 2 Cu + H2O + CO2 + O2 → Cu(OH)2 (s) + CuCO3 (s) 1886 Today 2-6 year Process Oxidized mixture called “Patina”

List 2 observations of each type you could make
Making Observations List 2 observations of each type you could make 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

Microscopic/Symbolic
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 *

A theory is a unifying principle that attempts to explain a body of experimental observations.
Theories offer explanations for what we observe. Theories tell us why we should expect it. Atomic Theory Cell Theory 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.

(∝ = directly proportional)
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) Charles’s Law: V ∝ T Newton's 2nd Law: Force = mass x acceleration 2nd law of thermodynamics: Entropy > 0

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. Homogenous mixture – composition of the mixture is the same throughout Solutions (soft drink), gas mixtures (air), solder (Sb/Pb alloy) 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

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. Hydrogen burns in air to form water 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.
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 possess subatomic particles:
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 *Many are derived from their Latin names
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. Quartz: SiO4 dry ice – CO2 (carbon dioxide) Lithium fluoride: LiF *Non-compounds are not necessarily always monoatomic (C, He): Can have many element atoms in a substance: P4, S8, Cl2

The Nucleus: Crash Course Chemistry #1
Review of the Nucleus: The Nucleus: Crash Course Chemistry #1

Classifications of Matter
CO2 C O2 H2 ex. Carbonated Water + H2O

The Three States of Matter: Effect of a Hot Poker on a Block of Ice
solid liquid gas

Lab Glassware Borosilicate Glass (SiO2 + B2O3) Non-Quantitative
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 Erlenmeyer Flask Beaker Used Quantitatively Crucible used instead Graduated Cylinder Volumetric Flask Buret

A Comparison: The Three States of Matter
Undefined shape, incompressible Undefined shape Compressible Defined shape, incompressible

A Comparison: The Three States of Matter
Plasma: Charged Gas, effected by magnetic field, Interacting particles Four Kinetic Molecular Theory: describes motion of particles in states of matter Random, fast movement of particles (non-interacting) Greater freedom of motion “particles shift/slide” Little particle motion “locked in place” only vibrations

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)

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

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 = x °C + 32 9 5 K = 0C 0 K = C 0 K = -460 ° F Absolute Scale (1848) “Water based” (1742) (1724) “Weather/human based” Absolute Zero: Theoretical temp where all atomic movements stops

1.3 Temperature conversion
Example 1.3 Temperature conversion 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? Helium has the lowest boiling point of all the elements at -452°F. Convert this temperature to degrees Celsius. 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 This conversion is carried out by writing
Here we have 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 (l) + energy H2O (g) 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) H2O energy + NH4NO3 (s) NH4+1 (aq) + NO3-1 (aq) energy + H2O (s) H2O (l)

Energy and Chemistry: Crash Course Chemistry #17
Review of Energy: Energy and Chemistry: Crash Course Chemistry #17

World’s Roundest Object
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 La Grande K 1 Kg Pt/Ir alloy World’s Roundest Object 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 A 1 kg bar will weigh 1 kg on earth 0.1 kg on moon

Used as Relative Standards for comparison
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 7 cm = ____________ m 300 g = ____________ kg
4.7 m = ____________mm 9,000 sec = ________Msec 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

(cm3 is more commonly used)
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 = mass volume d = m V

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 = m V

Failed to convert English to metric units
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. Failed to convert English to metric units 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.”

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: 1.99 x 10-23 We can factor out powers of 10 to simplify very large or small numbers Exponent (n) is a positive or negative integer N is the base number between 1 and 10 N x 10n

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 = 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 579 0.63 96,000 0.0140

Mathematics in Scientific Notation
move decimal left move decimal right + n - n = x 102 = 7.72 x 10-6 Addition or Subtraction: Must have same exponent Write each quantity with the same exponent n Combine N1 and N2 The exponent, n, remains the same 4.31 x x 103 = 4.31 x x 104 = 4.70 x 104

Mathematics in Scientific Notation
Multiplication: Add exponents (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 Multiply N1 and N2 Add exponents n1 and n2 Division: Subtract exponents 8.5 x 104 ÷ 5.0 x 109 = (8.5 ÷ 5.0) x = 1.7 x 10-5 Divide N1 and N2 Subtract exponents n1 and n2

b) Rewrite above numbers using the nearest SI prefix
Bell Ringer a) Write in scientific notation 8,705,000 m L 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 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 significant figures 34,000 mm 2 significant figures 2) Zeros between nonzero digits are significant 606 cm significant figures 50,050 s 4 significant figures 3) Zeros to the left of the first nonzero digit are not significant 0.08 mL significant figure 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 significant figures g 5 significant figures 5) If a number is less than 1, then only the zeros at the end are significant g 3 significant figures g significant figures

Significant Figures 0.001400 m ____ 4 significant figures 500 mL __
Every significant figure is shown when using Scientific notation. 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 478 cm 0.0430 kg 600,001 g
Determine the number of significant figures in the following measurements: 478 cm kg 600,001 g 1.310 × 1022 atoms 0.85 m 7000 mL

Example 1.4 Solution 478 cm -- Three, because each digit is a nonzero digit. (b) 600,001- Six, because zeros between nonzero digits are significant. (c) m -- Three, because zeros to the left of the first nonzero digit do not count as significant figures. (d) kg -- Three. The zero after the nonzero is significant because the number is less than 1. (e) × 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 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 ± 50 mL ± 1.0 mL
The answer cannot have more digits to the right of the decimal point than any of the original numbers. Use the least precise number. L 1.1 + 90.492 ± 50 mL XX one significant figure after decimal point round off to 90.5 ± 1.0 mL 3.70 0.7867 XX two significant figures after decimal point round off to 0.79

Significant Figures 4.51 x 3.0006 = 13.532706 = 13.5
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 = = 13.5 round to 3 sig figs 3 sig figs 6.8 ÷ = = 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: 11,254.1 g g 66.59 L − L 8.16 m × kg (d) kg ÷ 88.3 mL (e) 2.64 × 103 cm × 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 × 103 cm and then carry out the addition (2.64 cm cm) × 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? 3 = = 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

Ex. I weigh a 3 kg block on three different scales:
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 3.0 x 100% = 3.3% error

Precision also indicates to what degree we know our measurement
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 ( grams) (~$1,500)

Bell Ringer a) Perform the below mathematics in Sci. Notation. using Significant Figures in your answer. b) Rewrite the first 2 solutions using the nearest SI prefix (9.8 x 105 g) + (6.75 x 104 g) (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)

Dimensional Analysis Method of Solving Problems (Train-Track)
Determine which unit conversion factors are needed Carry units through calculation If all units cancel except for the desired unit(s), then the problem was solved correctly. given quantity x conversion factor = desired quantity desired unit given unit given unit x = desired unit

Train Track Example How many inches are in 3.0 miles?
Identify beginning information Draw a train track Write measurement as a fraction 3 miles 1

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

Again, place conversion factor so that the previous unit cancels 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

How many inches are in 3.0 miles?
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 1 x 1 x 1 = 190,080 inches = 1.9 x 105 inches

Example 1.6 A person’s average daily intake of glucose (a form of sugar) is pound (lb). What is this mass in milligrams (mg)? (1 lb = 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 = m/s = 31.3 m/s

Recall that 1 L = 1,000 cm3 and (1 cm)3 = (0.01 m)3.
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 0.808 g/cm3 = ? kg/m3 Liquid nitrogen
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 g/cm3. Convert the density to units of kg/m3. Liquid nitrogen 0.808 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:

1 Mm = 106 m 10-6 Mm = 1 m 1Km = 103 m 10-3 Km = 1 m meter (base)
Conversion factors can be written/used 2 ways 1 Mm = 106 m 1Km = 103 m meter (base) 1 cm = 10-2 m 1 mm = 10-3 m 1 mm = 10-6 m 10-6 Mm = 1 m 10-3 Km = 1 m meter (base) 102 cm = 1 m 103 mm = 1 m 106 mm = 1 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 mm 2.35 x 1012 inches to cm (1 ft = m) 3.50x104 mL to cL 42.0 km/h to ft/ms 0.55 Acres to m2 (247 acre = 1 km2) 106 g/mm3 to kg/m3 Review Unit Conversion & Significant Figures: Crash Course Chemistry #2

Chemistry: The Study of Matter

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