Transcript General Chemistry I

GENERAL CHEMISTRY 1
Chapter 1
Chemistry
Definition – Study of structure and interaction of
matter, including energy changes. Will discuss energy
in a later chapter. Matter is anything that has mass
and occupies space.
Scientific Method: Systematic approach to
scientific work
Observation
Hypothesis – Attempt to explain why
Law – Summary of observation
without trying to explain why
Experimentation – Test hypotheses
Theory – Hypothesis that stands the test
of time
Matter:
3 Types:
Elements – Basic substances – Can’t
be decomposed to a simpler substance. Only 115
known elements.
Compounds – 2 or more elements
chemically combined together in a fixed, definite
proportion by weight. Can only be decomposed
into its component elements by a chemical
reaction.
Elements and compounds are called
Pure Substances.
Mixtures – 2 or more pure substances physically
mixed together in any proportions whatsoever.
Can be decomposed, frequently fairly easily,
using physical methods, into its components.
Each component retains its own properties.
2 Types:
Homogeneous – All parts identical
Heterogeneous – Non-uniform
composition
One other quick definition is for an ion – a
charged particle. Positively charged particles are
specifically called cations, while negatively
charged ones are specifically called anions. We
will discuss these in much more detail later on.
3 States of Matter:
Solid – Keeps its shape and size
Liquid – Keeps its size but takes the
shape of its container
Gas – Takes the size and shape of
its container.
Any substance can be a solid, liquid or gas,
depending on the conditions.
All substances can be described by their
characteristics or properties.
Physical Property – Can be studied without
changing the substance into a new substance.
Chemical Property- Can only be studied by
changing (or attempting to change) the
substance into a new substance.
We study these properties by observing physical
& chemical changes.
Physical Changes are changes in the
appearance of a substance, but not in its
identity.
Chemical changes are actual changes in the
identity of the substance. A new substance or
substances is formed.
Physical Sciences require frequent measurements.
The official measurement system is called the
International System or SI, which defines basic
units for measuring various quantities and derived
units from them for other quantities. You can check
in your book for these on page 16.
The SI system replaced the older Metric System,
which is very similar and only differs in minor areas
of definition. I will explain the metric system.
centi
c
one-hundredth ( 1/100 )
milli
m
one-thousandth ( 1/ 1000 )
micro

one-millionth ( 1 / 10 6 )
nano
n
one-billionth ( 1 / 10 9 )
pico
p
one-trillionth ( 1 / 10 12 )
Basic units & abbreviations:
BASIC UNITS
• Name
Abbrev
• meter
m
• liter
• gram
L
g
Type of
Measurement
length
volume
mass
Approximate
Eng Equivalent
39.36 inches (about 1 yard)
1.06 quarts
0.035 oz (about 1/30 oz)
The scientific temperature scale is called the
Celcius (C) scale as opposed to the American
scale of Fahrenheit (F). A third scale, also
commonly used in science is called the Absolute or
Kelvin scale (F). They all measure the same thing,
intensity of heat, just using different units.
They can be converted into each other. The most
common formulas are:
C = (F – 32) / 1.8
F = C x 1.8 + 32
K = C + 273.15
Another method:
C = (F + 40) / 1.8 – 40
F = (C + 40)1.8 – 40
Accuracy vs. Precision
Accuracy describes how close a
measurement is to the correct value.
We usually don’t know how accurate
we are (otherwise we wouldn’t be taking the
measurement)
Precision describes how close measurements
are to each other. This depends on the quality of
the measuring instrument. This gives a
reasonable idea as to how accurate our
measurement is. Good precision usually, but not
always, indicates good accuracy. If we keep
getting the same measurement, then it most
likely is accurate.
Significant Figures: Help indicate the precision of
the measuring instrument. Any measurement
cannot be more precise than the measuring
instrument allows.
Assume the above picture is a ruler measuring in
cm. What is the measurement at the arrow? More
than 13 but less than 14. Not accurate to say 13 cm
or 14 cm. We can do better. Mentally, we can divide
the space between the smallest marks into 10 parts.
In our case, these mental marks are tenths of cm.
We then estimate how far along these mental marks
our arrow is. What do you think here?
Our final measurement will be reported with one
digit left of the decimal point (in the tenths position).
We know we are making a guess, but it is a
reasonable guess, but there still is some
uncertainty, probably + or – 0.1cm. It is not
considered reasonable to go any further ( our
minds are not capable of dividing a small space
into 100 parts, only into 10 parts). This last digit is
called the last significant digit (digit means
numeral).
We always measure until we have to make a
guess, then we stop. This indicates to anybody
reading our measurement without having the
instrument in front of them, to know how precise it
is. Or, in other words, what the smallest division on
the device is.
In many calculations, need to know how many
significant digits (or significant figures) are present
in a measurement. All non-zero digits are always
significant. The possible confusion lies with
zeroes.
Zeroes in Sig Figs
Leading zeroes - To the left of first non-zero
digit. All are not significant.
Middle zeroes - In between 2 significant
digits. All are significant.
Trailing zeroes - To the right of the last nonzero digit. Trailing zeroes to the right of the
decimal point are significant. Trailing
zeroes in whole numbers may or may not
be significant. Only the measurer knows.
The ambiguity can be removed by using
scientific notation. We will discuss this
shortly.
When calculations are done involving
measurements, there are specific rules to follow for
handling significant figures (From now on we will
use S.F. to refer to significant figures)
S.F. in Calculations
Multiplication and/or Division - The final
answer will have the same number of significant
digits (figures) as the measurement with the least
number of significant digits. Exact values and
counting units are considered to have an infinite
number of significant digits.
Addition and/or Subtraction - The final answer
will have the same number of decimal positions
as the measurement with the least number of
decimal positions. Again, exact values and
counting units are considered to have an infinite
number of significant digits.
Mixed Calculations - Round off when you switch
from one type of calculation to the other.
In all cases the final answer is rounded off to
proper significant figures.
Scientific Notation.
Makes use of the exponential powers of 10:
for example:
2000 = 2 x 1000 = 2 x 103
We need to be able to write a number in
normal decimal notation (1045.23) or in scientific
notation ( 1.04523 x 103). We need to be able to go
back and forth. My approach is to realize that for
every unit of exponent we have to move the decimal
point one place. If the exponent is decreasing, then
the decimal point is moved to the right, while it is
moved to the left if the exponent is increasing. For
this approach remember that a normal decimal
number, such as 156.2, is the same as 156.2 x 100
There is a difference between exponential notation
and scientific notation. Exponential notation
(frequently used in engineering applications, allows
any # of digits to the left of the decimal point, while
scientific notation allows only one non-zero digit
to the left of the decimal point.
Scientific notation avoids S.F. ambiguity, because
the non-exponential part of the number will
contain only the S.F. of that number.
(Remember: a number in scientific notation is still
only one number, even though it looks like a
multiplication of 2 numbers.
Modern scientific calculators all can easily
make use of scientific notation. Get familiar with
your particular calculator.
NOTE: When doing calculations involving
measurements, remember that almost all
measurements in science have units, such as g, mL,
km etc. Units are treated like numbers in
multiplication and division, while in addition and
subtraction, only like units can be added or
subtracted. All answers must have correct units or it
is not a correct answer.
Two important physical properties that you will be
studying in the third experiment this semester are
density and specific heat.
Density measures the mass of a specific volume of
the substance. It is essentially constant for the
solid state and liquid state (slight variation in liquid
state) and can be used to help identify a substance.
Mass
Density 
Volume
M
or D 
V
Specific Heat (SH) measures the amount of heat
needed to raise the temperature of 1 gram of a
substance by 1 C. Therefore the total heat
involved in a temperature change can be calculated
if we know the specific heat of the substance.
Determining the SH is also useful in helping to
identify a substance. The amount of heat is usually
symbolized by the letter q.
q = (SH)(mass)(tfinal – tinitial)
Dimensional Analysis
Conversion factors
Ratio of 2 measurements = 1.00000
Choose the conversion factor that
cancels the appropriate unit or units. Also, if
possible, make sure the conversion factors using
book values has enough significant figures so that
the number of sig. figs. from measured values is not
decreased.
Permanent & temporary
Do examples 1.6, 1.7 and 1.8 in text (pages 29-30).
In most cases we will solve problems using this
method, although at times I will use algebra
instead. Either method is acceptable.