Download Chapter 1 Student Notes

Chapter 1
The Science of Chemistry
VIDEO: “The World of Chemistry”
1-1 What is Chemistry?
A. Working with the Properties and Changes of Matter
Chemistry plays a vital role in your daily life and in the complex workings of your world.
Everything you see is made up of chemicals, EVERYTHING!
Chemistry the study of the composition and properties of chemicals and with the changes
chemicals can undergo.



Chemical any substance that has a definite composition (same stuff no matter where the
chemicals come from)
Natural chemicals gold iron, copper, water, etc…
Synthetic chemicals plastics, prescription drugs, etc…
Without chemicals nothing (soap, clothes, drugs, etc…) would exist. You depend on chemicals
everyday!
Chemical reactions happen all around you. DEMO on “Changes in Matter”
Chemical reaction the process by which one or more substances change to produce one or
more different substances, also called a chemical change
Ex) cooking food, rusting nail, striking as match, gasoline engine, etc…
The manufacture of chemicals is a big industry see page 5.
B. Physical States of Matter
Matter is anything that has mass and volume (occupies space).
Five States of Matter
1) Plasma: high energy gases, example is a star (Sun)
density depends on pressure
2) Gas: an indefinite shape and indefinite volume (takes the shape and volume of its container)
no free surfaces, molecules are far apart and free moving
density depends on pressure
3) Liquid: a definite volume but an indefinite shape (takes the shape of its container)
one free surface, molecules are relatively close together and free moving
density little affected by pressure
4) Solid: rigid, a definite shape and definite volume, many free surfaces
molecules are in fixed vibratory positions
density little affected by pressure
5) Bose-Einstein Matter (super atom): exists only at very low temperatures (0 K)
atoms move in single wavelike motion
cloudy blob
Changes in State
freezing, melting, condensation, evaporation, sublimation, deposition
Draw diagram.
C. Changes of Matter
Physical Property: anything you can observe without destroying or changing the
composition of a substance
ex) odor, color, volume, density, MP, BP
Chemical Property: describes how a substance(s) can change to form a new substance
Physical Change: identifying properties remain unchanged retain original properties,
characteristics, etc… appearance may change but not the properties
Changes in
1) amount of energy
2) physical state
3) form
4) particle size
ex) melting ice, tearing paper, splitting wood, mixing a cake
Chemical Change: also called chemical reactions, identifying properties of original
substance disappear as new substances with different properties are
formed. The new substance can longer be returned to its original form
ex) Na + Cl ------NaCl
ex) rusting iron, burning paper, wood, cooking an egg, baking a cake
Identify whether each of the following changes is a physical change or a chemical change.
1) Water boiling
6) Glass breaking
2) Iron rusting
7) Mowing the lawn
3) Butter melting
8) Magnetizing a nail
4) Alcohol evaporating
9) Baking a cake
5) Wood rotting
A chemical reaction is a process in which one or more substances are converted into new
substances with different physical and chemical properties. A chemical reaction is a
rearrangement of the atoms that make up the reactant(s). After rearrangement, those same
atoms are present in the product(s). Atoms are not created or destroyed, so mass does not
change during a chemical reaction.
Chemical reactions are represented by sentences known as chemical equations. A chemical
equation describes exactly what happens in a chemical reaction. A chemical equation shows
compounds before a chemical reaction takes place on the left (reactants) and compounds
formed from the chemical reaction on the right (products).
Reactants  Products
The arrow in the equation is read as “yields” or “produces”
Clues to a Chemical Reaction
 Color change
 Formation of a precipitate
 Formation of a gas (bubbles)
 Heat and/or flame is produced or absorbed
Exothermic: refers to a rxn where energy (as heat) flows from the system
(HOT)
Endothermic: refers to a rxn where energy (as heat) flows into system
(COLD)
DEMO
1.
2.
3.
4.
5.
color change (sugar + H2SO4)
gas (chalk + H2SO4) (Zn + H2SO4)
ppt (HgNO3 + HCl)
Heat produced (Mg + flame) (Na + H2O)
Heat absorbed (Ba(OH)2 + NH4SCN)
Homework page 9 Que 1-10 page 31 Que. 12,13,14
Quiz to Follow
Working With Numbers
Measurements are often combined by adding, subtracting, multiplying, or dividing.
Several mathematical tools will help you in this process.
A. Significant Digits (Figures)
 Counting significant digits
Significant Figures  all of the certain digits plus the first uncertain or estimated digit in a
measured quantity.
Rules for Counting Significant Figures
1. Nonzero integers. All nonzero integers are significant.
Ex) 127.34 contains 5 s.f.
2. Zeros.
a) Leading zeros are zeros that precede all of the nonzero digits. They never count as
significant figures.
Ex) 0.00476 contains 3 s.f.
b) Captive Zeros are zeros that fall between nonzero digits. They always count as
significant figures.
Ex) 120.007 contains 6 s.f.
c) Trailing zeros are zeros at the right end of the number. They are significant only if the
number contains a decimal point.
Ex) 109,000 contains 3 s.f.
Ex) 25,100.00 contains 7 s.f.
Ex) 0.1000 contains 4 s.f.
3. Exact Numbers. Any number that represents a numerical count or an exact definition can
be assumed to have an unlimited number of significant figures.
Ex) 1 in = 2.54 cm, neither 2.54 nor 1 limits the number of sig. figs.
Determine the number of significant figures:
1.
2.
3.
4.
29.625
0.02006
3.017
10.082
5.
6.
7.
8.
0.00011
0.000009
3.000009
2,690
9. 6.50
10. 900.00
11. 60,013
12. 21.040


Significant Digits in Calculations
Calculators Do Not Identify Significant Figures, YOU MUST
Rules
1) Multiplication/Division
 The measurement with the smallest number of significant digits
determines how many digits are allowed in the final answer.
 Your answer cannot be more precise than your least precise measurement.
Ex) V = lwh
Additional examples
= (3.05 cm)(2.10 cm) (0.75 cm)
see page 58
3
= 4.80375 cm
= 4.8 cm3
2) Addition/Subtraction
 When measured quantities are added or subtracted, the number of decimal
places in the result is the same as that in the quantity with the smallest
number of decimal places.
Ex) 10.21 g + 0.2 g + 256 g = 266.41 g = 266 g
3) When a calculation is performed in several steps (a problem has both addition (or
subtraction) and multiplication (or division)), round to the proper number of
significant digits after each operation.
Suppose more than one mathematical operation is involved in the calculation?
Such a calculation may be "deceptive" as to how many significant figures are
actually involved. For instance:
The subtraction in the numerator must be performed first to establish the number
of significant figures in the numerator. The subtraction results in 0.50
Since the subtraction in the numerator resulted in a number to two significant
figures (rounding to two decimal places), and the least number of significant
figures in the resulting expression involving multiplication and division is now
two significant figures, the final result must be rounded to two significant figures.
4) Rounding
≥ 5 round up
< 5 leave alone
Ex) see page 59 Practice Exercises
Perform the following calculations, and express the results in the correct units and with the
proper number of significant figures.
a. (0.054 kg +1.33 kg) x 5.4 m2
b. 67.35 cm2 / (1.401 cm - 0.399 cm)
c. 4.198 kg x (1019 m2 - 40 m2 ) / (54.0 s X 31.3 s)
d. 3.14159 m x (4.17 m + 2.150 m)
e. 690 000 m / (5.022 h - 4.31 h)
f. (6.23 cm + 3.111 cm - 0.05 cm) x 14.99 cm
Hwk page 67 Que 20-26
7.5 kg/m2
67.22 cm
2.43 kg· m2/s2
19.9 m2
970 000 m/h
139 cm2
SCIENTIFIC NOTATION
In science it is common to work with very large and very small numbers. These numbers can be
expressed in scientific notation. M x 10n
Scientific notation simply expresses a number as a product of a number between 1 and 10 and
the appropriate power of 10.
Expressed as M * 10 n
Where M is a number between 1 and 10 and n is an integer.
Rules for Scientific Notation



M is determined by moving the decimal point left or right so that only one nonzero digit is to
the left of it
The number of places the decimal point was moved determines n. If the decimal point is
moved to the left, n is positive. If the decimal point is moved to the right, n is negative.
All digits are significant when expressed as M.
OPTIONAL
If a calculator is not used to add, subtract, multiply, or divide numbers expressed in scientific
notation, then the following rules are needed.




When adding, exponents must be the same power to add.
When subtracting, exponents must be the same power to subtract.
When multiplying, multiply the numbers and add the exponents algebraically.
When dividing, divide the numbers and subtract the exponents algebraically.
Sample Problems
Express in Scientific Notation
1. 54,091
2. 98,000,000,000
3. 0.000589
4. 0.00000067
Sample Problems
1.
2.
3.
4.
5.
4.68 * 10 –8 + 1.23 * 10 –9
7.62 * 10 10 - 9.80 * 10 9
( 2.3 * 10 –6 ) ( 4.6 * 10 –7 ) ( 5.2 * 10 –3 )
9.72 * 10 10  3.45 * 10 –4
( 3.5 * 10 8 ) ( 1.72 * 10 –2 )  (2.3 * 10 4 )
Change to “ordinary” decimal number
1. 1.25 * 10 4
2. 1.67 * 10 –4
3. 9.347 * 10 7
4. 1.9444 * 10 –6
SCIENTIFIC NOTATION WORKSHEET
Express in Scientific Notation
1. 262, 000
2. 93, 100, 000, 000
3. 0.0195
4. 0.00000134
5. 65,300
6. 125,000,000,000,000,000
7. 0.000285
8. 7,790
Change to “ordinary” decimal number
1. 1.25 * 10 –9
2. 3.67 * 10 6
3. 8.91 * 10 15
4. 1.76 * 10 –3
5. 5.62 * 10 –1
6. 5.198 * 10 1
7. 2.76 * 10 15
8. 4.56 * 10 –10
Add
1. 4.72 * 10 10 + 3.21 * 10 10
2. 5.81 * 10 9 + 8.72 * 10 8
3. 2.79 * 10 –15 + 1.21 * 10 –14
Subtract
1. 9.74 * 10 10 - 2.89 * 10 10
2. 1.98 * 10 15 - 2.90 * 10 14
3. 3.45 * 10 –20 - 2.67 * 10 –21
Multiply
1. 6.7 * 10 10 X 2.5 * 10 15
2. 7.12 * 10 –10 X 3.05 * 10 15
3. 9.12 * 10 –20 X 8.71 * 10 –15
Divide
1. 8.5* 10 10  2.5 * 10 5
2. 9.3 * 10 –10  3.10 * 10 4
3. 8.4 * 10 –25  1.4 * 10 –15
Solve
( 1.5 * 10 15 ) ( 7.86 * 10 11 )  [ (6.2 * 10 –5 ) ( 8.10 * 10 6 )]
see page 67 Problems 35-40 for additional practice
Problem Solving
In Chemistry and every-day life you will often need to express a
measurement using a different unit than the one that was initially used.
To convert between units we can use dimensional analysis.
Dimensional analysis is a way to analyze and solve problems by using the
units, or dimensions, of the measurements.
Dimensional analysis also called factor-label method  technique or
problem solving method for converting between units via a conversion
factor.
“GIVEN”

CONVERSION
   DESIRED ANSWER
FACTOR
To convert units, conversion factors are used.
Conversion factors are equalities arranged as a ratio with different units.
For every equality, you can write two conversion factors
Ex.
1 dollar = 4 quarters
Correctly setting up dimensional analysis problems is
essential to being able to solve them.
Steps of Dimensional Analysis
1. Write down the conversion factors you can use
2. Write what is unknown equal to the known.
3. Multiply by the correct conversion factor
4. Cancel the units and do the math.
Fruit Exchange – Converting Units Practice
Use the following conversion factors to convert the units.
1 coconut = 3 bananas
1 banana = 4 peaches
2 peaches = 5 lemons
12 lemons = 1 orange
1) Convert 30 coconuts to kiwis.
2) Convert 20 lemons to bananas.
3) Convert 1 mango to lemons.
4) Convert 18 grapefruit to peaches.
5) Convert 10 oranges to bananas.
6) Convert 27 grapefruit to peaches.
7) Convert 36 kiwis to peaches.
1 mango = 6 oranges
1 coconut = 9 grapefruit
3 grapefruit = 2 kiwis
Units of Measurement
A measurement must contain a number and a unit.
In science the Metric system or SI system is used.
Scientists worldwide use a set of units called the Système Internationale d’Unités or SI.
7 base units all other units are derived units
All other units are derived by multiplying and dividing the base units.
Speed = distance/time = meters/secs
Length  meter
 meter stick
V=lwh = m*m*m = m3
V=lwh = cm*cm*cm = cm3
1 cm3 = 1 cc = 1mL
Volume  m3 (approx. 1 ton) = 1000 L
 graduated cylinder
Other Units (non SI units)
Liter
ºC
atm (atmospheres)
mm Hg
Mass  kilogram
 balance
calories
The base SI units are not always convenient to use. To write the base units in a more
convenient way metric prefixes are used.
Metric Prefix Scale
Prefix
Symbol
Meaning
giga -
G
1,000,000,000
1 x 109
mega -
M
1,000,000
1 x 106
kilo -
k
1,000
1 x 103
hecto -
h
100
1 x 102
deka -
da
10
1 x 101
Base Units: meter (m), gram (g), liter (s), second (s)
deci -
d
0.1
1/10
centi -
c
0.01
1/100
milli -
m
0.001
1 x 10-3
micro -
µ
0.000001
1 x 10-6
nano -
n
0.000000001
1 x 10-9
Conversion Factors  a ratio that is derived from the equality of two different units and that
can be used to convert from one unit to the other.
Conversion Factors
1 km = 1000 m
1 m = 1000 mm
1 m = 100 cm
1 cm = 10 mm
1kg = 1000g
1 g = 1000mg
1 kL = 1000 L
1 L = 1000mL
When moving up the prefix scale to a larger prefixes (units), move the decimal point one
place to the left () for each step up the scale.
When moving down the prefix scale to smaller prefixes (units), move the decimal point one
place right () for each step down the scale.
Metric Worksheet
1) 125 mm = ________ m
11)
152 cm = __________ mm
2) 29.4 L = _________ ml
12)
1.46 ml = __________L
3) 152,000 m = _______ km
13)
982 mg = __________ g
4) 0.00156 kg = ________ g
14)
16.5 L = __________ ml
5) 892 g = __________ kg
15)
6.98 mm = __________ m
6) 1.694 kl = _________ L
16)
435.68 cm = __________ m
7) 0.0986 mm = __________cm
17)
4.9 L = __________ ml
8) 3.49 cm = __________ m
18)
8.95 km = __________cm
9) 12,498 mg = __________ g
19)
65.7 g = __________kg
10)
20)
0.00078 km = __________mm
26.49 km = __________ mm
Problem Solving
Dimensional analysis also called factor-label method  technique or problem solving
method for converting between units via a conversion factor.
“GIVEN” 
CONVERSION
   DESIRED ANSWER
FACTOR
Ex) 250 gallons = ____________L
250 gal 
3.785 L
1 gal
 946 L
learn the unit equalities (an equation showing how different units are related)
1 kg = 2.2 lb
1 lb = 454 g
1 m = 39.37 in
1 in = 2.54 cm
1 L = 1.056 qts
1 qt = 0.946 L
conversion factors (an equation equal to 1) must be written. Two conversion factors can
be written for each equality.
1 inch = 2.54 cm
1 inch
2.54 cm
or
2.54 cm
1 inch
set-up the conversion factor so the units will cancel.
Do an example!!!
Ex) 250. cm = ________in
Ex) 39 L = ________ gal
Ex) 16 in = _______cm
Ex) 5.0 hrs = ________sec
Ex) 2.3 gal =_______cm3
Ex) 3.5 mi =_________m
Ex) 86 cm = _________ft
Ex) The drain in a 220 L tub leaks at a rate of one drop of water every three seconds. Twenty
drops of water is equivalent to one milliliter. How many days will it take for the filled bathtub to
empty (hint the desired unit is days/tubful)?
It is very important to use a systematic approach such as the factor-label method with these
easier problems. It is in your interest to master the problem solving method on less demanding
problems. When more difficult problems are encountered (later in the year), the method of
approach itself does not add to the problem’s difficulty.
METRIC WORKSHEET
1) 15 in = __________cm
2) 1.89 qt = __________ L
3) 340 lbs = _________ kg
4) 16.5 ft = __________ m
5) 0.96 gallons = _________ ml
6) 125 oz = __________ g
7) 156 yds = __________ km
8) 0.621 lbs = __________ g
9) 6 ft 2 in = __________ cm
10)
25 qts = __________ ml
11)
125 mm = __________ in
12)
15 ml = __________ qts
13)
622 g = __________ lbs
14)
5200 km = __________ yds
15)
9.62 cm = __________ ft
16)
0.0194 kg = __________oz
17)
15,000 m = __________ miles
18)
0.88 L = __________ gallons
19)
75,000mg = __________ lbs
20)
215,000 mm = __________ ft
Properties of Matter
Physical Property: anything you can observe without destroying or changing the
composition of a substance
ex) odor, color, volume, density, MP, BP
Density = the amount of matter present in a given volume of a substance
= mass per unit volume
= units are g / ml or g / cm3
density = _mass
volume
Determining Density
1) Density of gas will be discussed later.
2) Density of a liquid can be determined by weighing a known volume of the substance.
3) Density of an insoluble solid is determined by a method called water displacement.
Finding the Different Variables
1) Finding d
d=
m
v
2) Finding m
m=d*v
3) Finding v
v=
m
d
Page 16 Determine slope of graph. Pick two points on the line that are not part of the data.
Slope =
(Y2  Y1 )
rise Y


run X ( X 2  X 1 )
Substance
H2(g)
CO2(g)
Air
C2H5OH
Ice
Water
C12H22O11
NaCl
Al
Fe
Cu
Au
Os
Density of Some Common Substances
Density (g/cm3)
0.0000824
0.00180
0.0013
0.789
0.917
0.997
1.587
2.164
2.70
7.86
8.94
19.3
22.48
Ex) Using the above table, calculate the mass of 200 cm3 of air.
Ex) An unknown liquid has a mass of 30.6 g and a volume of 53.3 mL. What is the density of the liquid?
Ex) Could a block of metal with a mass of 18.2 g and a volume of 2.56 cm3 be iron?
Ex) Using the above table, calculate the volume occupied by 160. g of iron.
Homework
1. What is the volume, in cubic centimeters, of a sample of cough syrup that has a mass of
50.0 g? The density of cough syrup is 0.950 g/cm3.
2. A student finds a shiny piece of metal that he thinks is aluminum. In the lab he
determines that the metal has a volume of 245 cm3 and a mass of 612 g. Is the metal
aluminum?
3. What is the density of an object having a mass of 4.0 g and a volume of 30 cm3 ?
4. Using the above table, find the mass of a 35 cm3 sample of aluminum.
Chemical Property:  a property of matter that describes a substance’s ability to participate in
chemical reactions
Ex) iron rusts, silver tarnishes, gold does not rust
Ex) Mercuric(II) oxide  oxygen + mercury (pg 19)
Clues to a Chemical Reaction
 Color change
 Formation of a precipitate
 Formation of a gas (bubbles)
 Heat and/or flame is produced or absorbed
Exothermic: refers to a rxn where energy (as heat) flows from the system (HOT)
Endothermic: refers to a rxn where energy (as heat) flows into system (COLD)
1-3 HOW IS MATTER CLASSIFIED?
All matter is composed of about 118 different kinds of atoms. These atoms can be physically
mixed or chemically joined together to make up all kinds of matter.
Atom  the smallest unit of an element that maintains the properties of that element.
Since matter exists in so many different forms, having a way to classify matter is important for
studying it. In chemistry, it helps you predict what characteristics a sample will have based on
what you know about others like it.
Pure Substances  a homogeneous material consisting of one particular kind of matter and has
a definite chemical composition (same composition).
 a sample of matter, either a single element or a single compound, that has
definite chemical and physical properties.
ex) sugar, NaCl, water, Na, Mg, etc....
each element and compound is a pure substance
Elements
Molecule
Allotrope
 the simplest, least complex forms of matter that we encounter in the laboratory
or anywhere else, and they cannot be decomposed into simpler substances by
ordinary chemical reactions.
 approx. 113 elements to date
 some elements exist as single atoms, monatomic ex) hydrogen
 some elements exist as molecules consisting of as few as two or as many as
millions of atoms
 the smallest unit of a substance that keeps all of the physical and chemical
properties of that substance, consist of two or more atoms bonded together
Diatomic molecule  two atoms bonded together
Ex) H2, F2, Cl2, Br2, I2, N2, O2
 one of a number of different molecular forms of an element.
Ex) O2 and O3
Element symbols: some come from Latin names, others from scientists, etc…
first capitalized, if second never capitalized see page 22
Compound : a substance formed from two or more elements in which the elements are
always combined in the same fixed ratio by mass (constant comp.)
ex) H2O , C6H12O6 , (NH4)3PO4 , etc...
when elements react to form a compound, their individual properties are lost
and in their place we find the properties of the compound
ex) 2Na + Cl2 ---- 2NaCl
compounds are separated into elements by chemical means (elec., heat)
Compounds are represented by formulas.
Molecular formula  C9H8O4
Structural Formula  see page 24
Ball-and-Stick  see page 24
Space-Filling Model  see page 24
Identify as either an element or compound
1) carbon
2) water
3) aluminum foil
4) plastic
5) tin
6) silicon dioxide
7) carbon dioxide
8) helium
9) arsenic
10) sodium chloride
Mixture
 a combination of two or more pure substances that are not chemically combined
 a material consisting of two or more kinds of matter each retaining its own
characteristics, they have a variable composition (two or more pure
substances)
ex) sugar water, coffee (bean, water, milk, sugar, flavor)
Mixtures can be either homogeneous or heterogeneous.
Homogeneous mixtures : same properties throughout the sample, also called solutions
(need not be liquids)
ex) sugar water, steel, brass, air, milk, gasoline
Heterogeneous mixtures : consists of two or more regions called phases that differ in
properties
ex) vinegar & oil, gasoline & water, carbonated soda,
orange juice, tomato juice, chocolate chip cookies, granite
Separating the Components of a Mixture
Heterogeneous mixtures composed of liquids and solids are separated using a filtration
set up.
Homogeneous mixtures are separated by distillation, crystallization, and chromatography.
Distillation takes advantage in the difference of boiling points of two or more
liquids. It can also be used to separate solid impurities from liquids.
DEMO using CuSO4 · 5H2O
Crystallization when liquids are evaporated crystals will form.
Chromatography is a process where a solution can be separated by allowing it to
flow over a stationary substance.
FLOW CHART
See page 28
HOMEWORK
Identify the following substances as pure substances, heterogeneous mixtures, or homogeneous
mixtures
1) Alphabet soup
2) Salt
3) Concrete
4) Vegetable oil
5) Air
6) Paint
7) Sea water
8) Granite
9) Steel
10) Sugar
Completion
Word Bank
filtration, crystallization, chromatography, electrolysis, distillation
1) Heterogeneous mixtures are often separated by _____________.
2) Separating sand from water can be done by _______________.
3) The sugar in sugar water can be removed by _____________.
4) The separation technique that takes advantage of different boiling points is called
____________.
5) Removing chlorophyll pigment from leaves might be done by ____________.
6) The best way to decompose water into oxygen and hydrogen is by
_________________.
7) Crude oil is broken down by heat, vaporized, and allowed to condense into
various liquids such as gasoline. This process is called ____________.