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