Transcript Chapter One Chemistry: The Science of Change

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An organized process used by scientists to do
research and to verify the work of others
 Designed to produce a solution that can be tested
and supported by experimentation
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Observations:
 Qualitative: Descriptive
▪ Ex: color, texture, smell
 Quantitative: Numerical
▪ Ex: 2.5 grams, 55.8 mL
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Ex: My phone is running slower than my
friend’s phone
Ex: My phone has 50 more apps than my
friend’s phone
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A tentative explanation of observations
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Ex: As more apps are added onto a phone,
the phone will operate at a slower rate.
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Independent variable: what YOU control
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Dependent variable: what is MEASURED as
a result of changing the independent.
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What about our phone example?
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A set of controlled observations that test a
hypothesis
Contains independent & dependent variable
At least 3 trials
Control: In an experiment, this is the
standard for comparison.
Constant: factors that must remain
unchanging in the experiment to ensure
results are meaningful
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Revision of hypothesis
Evolution of a theory
 Explanation of a body of experiments and
observations
 Used to predict
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Scientific Law
 Summary of accepted facts of nature
 Concise verbal or mathematical statement
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Solid: definite shape & volume, low
energy, geometric structure
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Liquid: no definite shape, definite
volume, moderate energy, molecules
move past each other
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Gas: no definite shape or volume, high
energy, molecules move fast and
randomly
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Matter: anything with mass and volume
 Either a pure substance or mixture of substances
Determine if the following is an element, compound,
heterogeneous mixture, or homogeneous mixture
Distilled water
Aluminum cans
2012 penny
Silver
Carbon dioxide
Salt water
Orange Juice w/ pulp
Brass
Vanilla milkshake
Table salt
Mercury
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Physical Property: Something that can be
observed/measured without changing
composition of a substance
 i.e. color, malleable, ductile, density,
melting/boiling point
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Chemical Property: The ability of a
substance to combine with another
substance
 i.e. flammability, reactivity.
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Physical Change: A change that happens
without the chemical identity being altered
 i.e. cutting paper, phase changes
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Chemical Change: A change that alters the
chemical identity of a substance
 i.e. rusting, corrosion, combustion, baking
Identify if the following are physical properties (PP),
chemical properties (CP), physical changes (PC), or
chemical changes (CC)
1) Luster/Shiny
2) Flammable
3) Melting ice
4) Plants growing
5) Melting point of 15°C
6) Baking cookies
7) Reactivity
8) Salt forming on rocks
from the ocean
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Extensive Property: a property that depends
on amount of matter
 i.e. mass, volume
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Intensive Property: a property that DOES
NOT depend on amount of matter
 i.e. density, temperature
M x 10n
M = Base (a number between 1 and 10)
n = exponent (positive=bigger than 1;
negative=less than 1)
 Use scientific notation if number is in the
thousands and larger or in thousandths and
smaller
Ex:
0.00358
29,403
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English System: foot,
pound, gallon, etc.
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Metric System: meter,
kilogram, liter, etc.
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SI System: Revised
metric system (System
Internationale)
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Mean King Henry Died Unexpectedly
Drinking Chocolate Milk Monday
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Mega Kilo Hecta Deka Unit
Deci Centi Milli Micro
M _ _ K H D U D C M _ _ M
M _ _ k h da U d c m _ _ µ
106
103 102 101 100 10-1 10-2 10-3
10-6
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489.3 dg  kg
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0.0293 L  μL
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2.04 x 103 pm  cm
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How many nanograms are in one gram?
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How many grams are in one nanogram?
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A unit that is a combination of SI base units
 Ex: Volume & Density
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How many dm3 are in 1 m3?
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How many mm2 are in 1 m2?
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Celsius Scale: defined using the freezing
point (O°C) and boiling point (100°C) of pure
water at sea level
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Kelvin Scale: defined off of absolute zero (at
0 K, all molecular motion stops)
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Fahrenheit Scale: many varying stories of
how this scale was derived
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Celsius to Kelvin:
 K = °C + 273.15
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Kelvin to Celsius:
 °C = K – 273.15
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Celsius to Fahrenheit:
 °F = (1.8)(°C) + 32
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93.0 °C  K
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214 K  °C
D = m/V
 D = density, m = mass, V = volume
In Chemistry, density measured in g/mL
(for liquids) or g/cm3 (for solids)
 Since gases are very light, we measure
their density in g/L
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A sample of pure aluminum has a density of
2.70 g/mL. If a cube of aluminum has a length
of 2.0 cm on each side, how much mass will it
have?
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All non zero numbers are significant
 1.932
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All zeroes in between non zero numbers are
ALWAYS significant
 1.001
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3004
“Beginning zeros ” (All zeroes to the LEFT of the
first non zero number) are NEVER significant
 0.000234
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“Ending zeros” (All zeroes to the RIGHT of the last
non zero number) are significant ONLY IF there is a
decimal in the number
 400
5.00
400.
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1 003
209.50
0.000243
1.0040
500
303
1 000 000.
0.005080
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Exact numbers come from counting, not
measuring
 Example: 12 donuts = 1 dozen, 1000 mm = 1 m
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Sig figs are never based on exact numbers
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Adding/Subtracting
 Express the answer with the LOWEST # OF
DECIMAL PLACES
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Multiplying/Dividing
 Express the answer with the LOWEST # OF SIG
FIGS
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Perform the following calculations and
express the answer in the correct # of sig figs.
 89.332 m + 1.1 m =
 23.4 cm x 16.00 cm =
 54.78 m2 ÷ 2.3 m =
 128.2 mL – 43.98 mL =
90.4 m
374 cm2
24 m
84.2 mL
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When multiple operations occur, be sure to
complete one set of sig fig rules before
moving on to the other set.
 Avoid Rounding Errors!
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Example: Calculating percent error
(Accepted – Experimental)
Accepted
(8.641 g/mL – 8.43 g/mL)
8.641 g/mL
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Accuracy: How close you are to the TRUE
value
Precision: How close a group of
measurements are to each other
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A fraction set up with measurements in the
numerator and denominator that EQUAL
each other
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Ex:
1 foot
12 inches
12 inches
1 foot
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Using conversion factors to convert from one
unit to another
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Perform the following conversions:
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How many inches are in 2.3 meters?
Convert 54.5 mi/hr  m/s (1 mile = 1.6093 km)
How many 375 mg tables can be produced
from 2.60 kg of powdered aspirin?
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