Transcript Scientific Method - Olympic High School Home Page

AP Chemistry
Olympic High School
Mr. Daniel
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The Scientific Method
Hypothesis: a tentative explanation or prediction
based on experimental observations
Law: a description of observed natural phenomena
Theory: a unifying principle that explains a body of
facts and the laws based on them. A good theory
can be modified to explain new information.
What is chemistry?
“Science that deals with the properties, composition, and
structure of substances (elements and compounds), the
reactions and transformations they undergo, and the
energy released or absorbed during those processes...
Often called the "central science," chemistry is concerned
with atoms as building blocks..., with everything in the
material world, and with all living things.” Encyclopaedia
Britannica, 2002.
“The study of the properties of materials and the
changes that materials undergo.” Chemistry: The Central
Science, Brown, LeMay, Bursten 7th ed.
Why should we study chemistry
/ How does chemistry affect us?
Food science
Forensics chemistry
Drugs
Environmental chemistry: geo- and atmospheric
Materials (construction materials, clothing, paints, computer
components, etc)
New energy sources
Et cetera...
Classification & Properties of Matter
Matter – Anything that occupies space and has mass.
Substance – Pure matter with unique properties; it
cannot be separated by physical processes.
Element – Cannot be subdivided by chemical or
physical processes into simpler substances.
Atom – The smallest particle of an element that retains
the characteristic chemical properties of that element.
Compounds
Compounds – Contains more than one type of
atom, but all molecules (or repeat units) are the
same, e.g., water (H2O), sodium chloride (NaCl),
and “fool’s gold” pyrite (FeS2).
Chemical formula – gives the ratio of atoms of
each element in a compound.
Molecule – smallest discrete units that retain the
composition and chemical characteristics of the
compound. These groups of atoms are held
together with covalent bonding
Mixtures
Mixture – Have variable composition and can be
separated into component parts by physical methods.
Mixtures contain more than one kind of molecule, and
their properties depend on the relative amount of each
component present in the mixture.
Homogeneous Mixture (solution) – Uniform
composition.
Air: principle components include O2, N2 & CO2
Brass: solid solution of Cu and Zn
Heterogeneous Mixture – Non-uniform composition.
Chocolate Chip Cookie: Chocolate, flour, water, etc
Blood– red & white blood cells, plasma, etc.
Ne
SO3
N2
Homogeneous
Mixture
Physical Properties
Physical Properties: Properties that can be observed and
measured without changing the composition of a substance.
color, state of matter, melting and boiling points
density, solubility, electric conductivity,
malleability and ductility, viscosity, etc...
Extensive Properties: Depends on the amount of
a substance present, e.g., mass and volume.
Intensive Properties: Does not depend on the
amount of a substance present, e.g., boiling
point and density.
States of Matter
• Solid: fixed volume; rigid shape;
particles are packed closely together.
• Liquid: fixed volume; no definite shape;
atoms or molecules are arranged randomly
• Gas: volume
varies with
temperature &
pressure; no
definite shape;
gas molecules
are far apart.
Chemical Properties
Chemical Properties: Describe the way a substance may
change or react to form other substances. Examples include:
Ethanol burns in air (reacts with oxygen in the air)
Sodium reacts vigorously with water
Corrosion of metal parts (rust)
Etc.
Chemical Change
2H2(g) + O2(g)  2H2O(g)
SI (Metric) Units
Table 1.2 Some SI Base Units
Measured Property
Name of Unit
Abbreviation
Mass
kilogram (not gram!)
kg
Length
meter
m
Time
second
s
Temperature
kelvin
K
Amount of substance mole
mol
Electric current
A
ampere
Metric Prefixes (Table 1.3)
Prefix
Abbreviation
Meaning
Example
mega-
M
106
1 megaton = 1×106 tons
kilo-
k
103
1 kilogram (kg)= 1×103 grams
deci-
d
10-1
1 decimeter (dm) = 1×10-1 m
centi-
c
10-2
1 centimeter (cm) = 1×10-2 m
milli-
m
10-3
1 millimeter (mm) = 1×10-3 m
micro-

10-6
1 micrometer (m) = 1×10-6 m
nano-
n
10-9
1 nanometer (nm) = 1×10-9 m
pico-
p
10-12
1 picometer (pm) = 1×10-12 m
femto-
f
10-15
1 femtometer (fm) = 1×10-15 m
Temperature
Fahrenheit scale
Celsius scale
water freezes
32 °F
0 °C
body temperature
98.6 °F
37 °C
water boils
212 °F
100 °C
T (°C) = 5 °C / 9 °F [ T(°F) – 32]
Kelvin scale = same size unit as Celsius, but
the lowest temperature that can be achieved is
absolute zero
K = °C +273
°C = K - 273
What about Volume?
Liter (L) is not a base SI unit.
It is a derived unit.
1 m3 = 1000 dm3 (L)
= 1,000,000 cm3
1,000,000 cm3 = 1  106 cm3
Density
Density = Mass
Volume
Units used are : g/mL (liquids),
g/cm3 (solids),
g/L (gases)
Precision and Accuracy
Significant Figures!!
Significant Figure
• Significant figures are a scientist’s way of
reporting on how accurately a measurement
was made.
• Accuracy depend on two things:
• The device used to measure
• The care taken by the individual
Significant Figures
• A Significant figure is defined as including
all certain digits plus one estimated digit.
What would be the
measurement of the liquid in
the graduated cylinder?
Significant Figures
Which measurement is the best?
What is the
measured value?
What is the
measured value?
What is the
measured value?
Rules for Counting Significant
Figures
Non-zeros always count as significant
figures:
3456 km has
4 significant figures
Rules for Counting Significant
Figures
Zeros
Leading zeros never count as significant
figures:
0.0486 mL has
3 significant figures
Rules for Counting Significant
Figures
Zeros
Captive zeros always count as
significant figures:
16.07 cm has
4 significant figures
Rules for Counting Significant
Figures
Zeros
Trailing zeros are significant only if the number
contains a written decimal point:
9.300 g has 4 significant figures
9300 g has 2 significant figures
9300. g has 4 significant figures
Rules for Counting Significant
Figures
Two special situations have an unlimited number
of significant figures:
1. Counted items
a) 23 people, or 425 thumbtacks
2. Exactly defined quantities
b) 60 minutes = 1 hour
Significant Figures
1. Read the number from left to right and count all the
digits, starting with the first digit that is NOT zero if
a decimal is present.
1200
2 sig figs
1200.
4 sig figs
1200.0
5 sig figs
0.00120
3 sig figs
0.01002
4 sig figs
Significant Figures
1) In multiplication or division, the number of
significant figures in the answer should be the same
as that in the quantity with the fewest significant
figures.
0.080 × 125 =?
0.080 × 125 = 10 → 10. or 1.0 × 102
(2)
(3)
(answer should have 2 sig figs)
2) When the number is rounded off, the last digit to be
retained is increased by one only if the following
digit is 5 or greater.
Significant Figures
3) When adding or subtracting numbers, the number
of decimal places in the answer is equal to the
number of decimal places in the number with the
fewest digits after the decimal.
10.1 + 100.001 + 1010 =?
10.1
(10-1 place)
100.001
(10-3 place)
1010.
1120.101
(100 place)
→
1120. or 1.120×103
Dimensional Analysis
(factor label method)
Use units to guide you through calculations.
Conversion factors are fractions where the numerator and
denominator are the same quantity with different units. Many
conversion factors are EXACT.
1000 mm
1m
2.54 cm
1 ft
1 in
12 in
1 lb
453.6 g