Transcript Chemistry 1A Mr. Kimball [email protected] http://www2.bakersfieldcollege.edu/dkimball Welcome to Chemistry 2A • Podcasts • A little about myself • A little about you – New? Major? ESL? International? –

Chemistry 1A
Mr. Kimball
[email protected]
http://www2.bakersfieldcollege.edu/dkimball
Welcome to Chemistry 2A
• Podcasts
• A little about myself
• A little about you
– New? Major? ESL? International?
– Learning Disorders
• Sign roll sheet
• Get phone numbers of others in class
The Class Syllabus
http://www2.bakersfieldcollege.edu/dkimball
Some students prefer to skim
through a course.
If you really
want to succeed
you need to go
deep!
Learning Skills
Take Responsibility (it’s your education)
Have Confidence (you can do it!)
Don‘t Procrastinate (study for final now)
Read/Listen Precisely (ignore things that just aren’t
there)
5. Practice (use practice tests)
6. Persistence (get up one more time than you fall down)
7. Recognize Patterns (most things are done the same
way)
8. Use Pictures (outline problem)
9. Think Sequentially (one step at a time)
10. Do Neat work (so you can check it)
11. Group Study (explain things to each other)
12. Try Something New (don’t keep repeating failures)
13. Get Help
Learning Skills Power Point
1.
2.
3.
4.
R
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P
P
P
P
P
P
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G
N
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Using This Book
• Concept check problems are found within the
chapter with the solutions right there with the
problem.
• Exercise problems are found within the chapter
with select answers in the back of the book.
• Homework is assigned from the Internet. You
should check the Eduspace link from my Main
Web site under Chemistry 1a for instructions.
Using This Book
• A CD comes with the book that has tutorials,
practice quizzes and other aids. Those not doing
well in the class should consider using some of
these aids.
• There is also a web site. It has practice tests,
flash cards, animations, etc.
(http://college.hmco.com/chemistry/general/ebbing/general_chem/8e/students/index.html )
Chemistry
and
Measurement
1.1 Modern Chemistry
1.2 Experiment and Explanation
1.3 Law of Conservation of Mass
1.4 Matter
Aristotle
Earth
Air
Fire
Water
Democritus (460-370 BC)
Greek Philosopher who first coined the word “atomos”.
What Is Chemistry?
• Chemistry is the study of the composition,
structure, and properties of matter and
energy and changes that matter undergoes.
– Matter is anything that occupies space and has
mass.
– Energy is the “ability to do work.”
Archimedes
Archimedes lived in Syracuse on the island of Sicily.
Archimedes
A comparison of Archimedes’ Pulleys and Study!
Big
Study!
Little
Study!
Galileo Galilei
Father of the Scientific Method
Heavy things
fall faster
than light
things????
Aristotle
Experiment and Explanation
• Experiment and explanation are the heart of
chemical research.
– An experiment is an observation of natural
phenomena carried out in a controlled manner
so that the results can be duplicated and rational
conclusions obtained.
– After a series of experiments, a researcher may
See some relationship or regularity in the
results.
Experiment and Explanation
• If the regularity or relationship is
fundamental and we can state it simply, we
call it a law.
– A law is a concise statement or mathematical
equation about a fundamental relationship or
regularity of nature.
– An example is the law of conservation of mass,
which says that mass, or quantity of matter,
remains constant during any chemical change.
Experiment and Explanation
• Explanations help us organize knowledge
and predict future events.
– A hypothesis is a tentative explanation of some
regularity of nature.
– If a hypothesis successfully passes many tests,
it becomes known as a theory.
– A theory is a tested explanation of basic natural
phenomena.
Experiment and
Explanation
• The general process
of advancing
scientific knowledge
through observation,
laws, hypotheses, or
theories is called the
scientific method.
The Scientific Method
Examples:
1. Pons and Fleishman, Univ. of Utah.
2. Horoscope
3. Weather.
Your Assignment:
1. Formulate a Problem.
2. Observe and collect Data.
3. Interpret Data.
4. Test your Interpretation.
Matter: Physical State and
Chemical Constitution
• There are two principal ways of classifying
matter:
– By its physical state as a solid, liquid, or gas.
– By its chemical constitution as an element,
compound, or mixture.
Solids, Liquids, and Gases
• Solid: the form of matter characterized
by rigidity; a solid is relatively
incompressible and has a fixed shape
and volume.
• Liquid: the form of matter that is a
relatively incompressible fluid; liquid
has a fixed volume but no fixed shape.
• Gas: the form of matter that is an easily
compressible fluid; a given quantity of
gas will fit into a container of almost
any size in shape.
Elements, Compounds, and Mixtures
• To understand how matter is classified by
its chemical constitution we must first look
at physical and chemical changes.
– A physical change is a change in the form of
matter but not in its chemical identity.
– Physical changes are usually reversible.
– No new compounds are formed during a
physical change.
– Melting ice is an example of a physical change.
Elements, Compounds, and Mixtures
(cont’d)
• A chemical change, or chemical reaction, is
a change in which one or more kinds of
matter are transformed into a new kind of
matter or several new kinds of matter.
– Chemical changes are usually irreversible.
– New compounds are formed during a chemical
change.
– The rusting of iron is an example of a chemical
change.
Elements, Compounds, and Mixtures
(cont’d)
• A physical property is a characteristic that
can be observed for material without
changing its chemical identity.
• Examples are physical state (solid, liquid,or
gas), melting point, and color.
• A chemical property is a characteristic of a
material involving its chemical change.
– A chemical property of iron is its ability to react
with oxygen to produce rust.
Separate by
Chemical
Processes
-burning
-fermentation
-rusting
Matter
Pure Substances
Elements
(atoms)
Hydrogen
Oxygen
Copper
Zinc
Separate by
Physical
Processes
-filtering
-distillation
-centrifuging
Compounds
(molecules)
Water
Alcohol
Sugar
Salt
Mixtures
Homogeneous
(solutions)
Air
Sodas
Ocean Water
Alcoholic drinks
Heterogeneous
(most things)
Granite
Sand
Wood
Orange Juice
Separation by
distillation.
Elements:
sulfur,
arsenic,
iodine,
magnesium
, bismuth,
mercury.
Photo
courtesy of
American
Color.
A mixture of potassium dichromate and iron
fillings. Photo courtesy of James Scherer.
Return to slide 15.
A magnet
separates the
iron filling
from the
mixture.
Photo courtesy
of James
Scherer.
Return to slide 15.
Chemistry and
Measurement
1.5 Measurement and
Significant Figures
1.6 SI Units
1.7 Derived Units
1.8 Units and Dimensional
Analysis
Measurement and Significant Figures
• Measurement is the comparison of a
physical quantity to be measured with a unit
of measurement -- that is, with a fixed
standard of measurement.
– The term precision refers to the closeness of the
set of values obtained from identical
measurements of a quantity.
– Accuracy is a related term; it refers to the
closeness of a single measurements to its true
value.
Precision vs. Accuracy
Measurement and Significant Figures
(cont’d)
• To indicate the precision of a measured
number (or result of calculations on
measured numbers), we often use the
concept of significant figures.
– Significant figures are those digits in a
measured number (or result of the calculation
with a measured number) that include all
certain digits plus a final one having some
uncertainty.
Scientific Notation
• Useful with very large and very small numbers.
• Decimal always after first digit.
• Use x 10n where n is the number of decimal places
you must move the decimal to get it just after the
first digit.
• Positive exponents represent large numbers.
2,340,000,000,000,000 = 2.34 x 1015
• Negative exponents represent small numbers.
0.00000000000000234 = 2.34 x 10-15
Measurement Accuracy
How long is this steel rod?
There is no such thing as a totally accurate measurement!
Significant Figures
• Numbers that measure or contribute to our
accuracy.
• The more significant figures we have the
more accurate our measurement.
• Significant figures are determined by our
measurement device or technique.
Rules of Determining the
Number of Significant Figures
1. All non-zero digits are significant.
234 = 3 sig figs
1.333 = 4 sig figs
1,234.2 = 5 sig figs
2. All zeros between non-zero digits are
significant.
203 = 3 sig figs
1.003 = 4 sig figs
1,030.2 = 5 sig figs
Rules of Determining the
Number of Significant Figures
3. All zeros to the right of the decimal and to the
right of the last non-zero digit are significant.
2.30 = 3 sig figs
1.000 = 4 sig figs
3.4500 = 5 sig figs
4. All zeros to the left of the first non-zero digit are
NOT significant.
0.0200 = 3 sig figs
0.1220 = 4 sig figs
0.000000012210 = 5 sig figs
Rules of Determining the
Number of Significant Figures
5. Zeros to the right of the first non-zero
digit and to the left of the decimal may
or may not be significant. They must
be written in scientific notation.
2300 = 2.3 x 103 or 2.30 x 103 or 2.300 x 103
2 sig figs
3 sig figs
4 sig figs
Rules of Determining the
Number of Significant Figures
6. Some numbers have infinite significant
figures or are exact numbers.
233 people 14 cats (unless in biology lab)
7 cars on the highway 36 schools in town
How many significant figures are
in each of the following?
1) 23.34
4 significant figures
2) 21.003
5 significant figures
3) .0003030
4 significant figures
4) 210
2 or 3 significant figures
5) 200 students
infinite significant figures
6) 3000
1, 2, 3, or 4 significant figures
Using Significant Figures in
Calculations
Addition and Subtraction
1. Line up the decimals.
2. Add or subtract.
3. Round of to first full column.
23.345 +14.5 + 0.523 = ?
23.345
14.5
+ 0.523
38.368
= 38.4 or three significant fingures
Using Significant Figures in
Calculations
Multiplication and Division
1. Do the multiplication or division.
2. Round answer off to the same number
of significant figures as the least
number in the data.
(23.345)(14.5)(0.523) = ? 177.0368075
= 177 or three significant figures
SI System
British
Length
Mass
Volume
Time
meter
gram
Liter
second
Km=1000m
Kg=1000g
KL=1000L
1min=60sec
100cm=1m 1000mg=1 g 1000mL=1L 60min=1hr
1000mm=1m
Foot
pound
gallon
12in=1ft
16oz=1 lb
4qt=1gal
3ft=1yd
2000 lb=1 ton 2pts=1qt
5280ft=1mile
second
(same)
Table 1.5 Relationships of Some
U.S. and Metric Units
Length
1 in = 2.54 cm
1 yd = 0.9144 m
1 mi = 1.609 km
1 mi = 5280 ft
Mass
Volume
1 lb = 0.4536 kg 1 qt = 0.9464 L
1 lb = 16 oz
4 qt = 1 gal
1 L = 1.06 qt
1 oz = 28.35 g
1 lb = 454 g
Table 1.3 SI Prefixes
Multiple
106
103
10-1
10-2
10-3
10-6
10-9
10-12
Prefix
mega
kilo
deci
centi
milli
micro
nano
pico
Symbol
M
k
D
C
m
m
n
p
Units: Dimensional Analysis
• In performing numerical calculations, it is
good practice to associate units with each
quantity.
– The advantage of this approach is that the units
for the answer will come out of the calculation.
– And, if you make an error in arranging factors
in the calculation, it will be apparent because
the final units will be nonsense.
Unit Conversion
• Sodium hydrogen carbonate (baking soda)
reacts with acidic materials such as vinegar
to release carbon dioxide gas. Given an
experiment calling for 0.348 kg of sodium
hydrogen carbonate, express this mass in
milligrams. 3
3
0.348 kg x
10 g
1 kg
x
10 mg
1g
= 3.48 x 105 mg
Units: Dimensional Analysis
• Dimensional analysis (or the factor-label
method) is the method of calculation in
which one carries along the units for
quantities.
– Suppose you simply wish to convert 20 yards to
feet.
3 feet
20 yards 
 60 feet
1 yard
– Note that the units have cancelled properly to give
the final unit of feet.
Units: Dimensional Analysis
• The ratio (3 feet/1 yard) is called a
conversion factor.
– The conversion-factor method may be used to
convert any unit to another, provided a
conversion equation exists.
– Relationships between certain U.S. units and
metric units are given in Table 1.5.
Unit Conversion
• Suppose you wish to convert 0.547 lb to
grams.
– From Table 1.5, note that 1 lb = 453.6 g, so the
conversion factor from pounds to grams is
453.6 g/1 lb. Therefore,
453.6 g
0.547 lb 
 248 g
1 lb
Temperature
• The Celsius scale (formerly the Centigrade
scale) is the temperature scale in general
scientific use.
– However, the SI base unit of temperature is the
kelvin (K), a unit based on the absolute
temperature scale.
– The conversion from Celsius to Kelvin is
simple since the two scales are simply offset by
o
273.15o.
K  C  273.15
Figure 1.23: Comparison of
Temperature Scales
Temperature
• The Fahrenheit scale is at present the
common temperature scale in the United
States.
– The conversion of Fahrenheit to Celsius, and
vice versa, can be accomplished with the
following formulas
o
F  32 o
o
o
C
1.8
F  1.8 ( C)  32
Derived Units
• The SI unit for speed is meters per second,
or m/s.
– This is an example of an SI derived unit,
created by combining SI base units.
– Volume is defined as length cubed and has an
SI unit of cubic meters (m3).
– Traditionally, chemists have used the liter (L),
which is a unit of volume equal to one cubic
3
3
decimeter. 1 L  1 dm and 1 mL  1 cm
Derived Units
• The density of an object is its mass per unit
volume,
m
d
V
where d is the density, m is the mass, and V is the
volume. Generally the unit of mass is the gram. The
unit of volume is the mL for liquids; cm3 for solids;
and L for gases.
A Density Example
• A sample of the mineral galena (lead
sulfide) weighs 12.4 g and has a volume of
1.64 cm3. What is the density of galena?
mass
12.4 g
3
Density =
=
=
7.5609
=
7.56
g/cm
volume
1.64 cm3