Transcript CHEM 101 - Virginia State University

CHEM 101
Chapter 1
Matter and Measurement
Virginia State University
Summer 2008
Dr. Victor Vilchiz
5/19/08
Physical Properties
• There are two types of physical properties:
– Extensive Properties
– Intensive Properties
• Extensive Properties are those that
depend on the amount of substance
present (mass, volume).
• Intensive Properties are independent of
substance amount (color, taste, density).
Density
• It has been mentioned several times… but
what is density?
– Density is the distribution of something over a
given amount of space.
• Mass density: given amount of mass in a given
volume.
• Charge density: amount of charge in a given space.
• Particle density: amount of ………….
Observations and Measurement
• Last time we said that if you analyze
something w/o any measurements all you
will get is qualitative information.
– Qualitative: tells you the quality (what is
present) but it does not tell you %
• If you start making measurements then the
information obtained is quantitative.
– Quantitative: tells you the quantity of
components in the sample.
Measurements
• Before we continue talking about
measurements we must agree on several
things…
– Units
– Standards
– References
• In reality all three are the same thing.
• There are several unit systems in use but
we will concentrate in one and only one (SI)
SI Units
• First things first… it is not the SI system…
that’s redundant.
– SI = Systeme Internationale
• Established in the 1960’s by the world’s scientific
community.
• Based on the metric system
• Until the late 1990’s two countries did not use this
system
– US and Liberia
– We are now the only ones, but we are slowly changing.
The Fundamental SI Units
Common Derived Units
Temperature Scales
• There are many temperature scales but 3
of them are the most important.
– Fahrenheit
– Celcius/Centigrade
– Kelvin
• We will discuss their origins and points of
reference.
Fahrenheit
• Proposed in the early 1700’s by Gabriel
Fahrenheit
– It is believed that he used 4 reference points
to build his scale
• 0°F Lowest temperature attained by a mixture of
salt/ice/water
• 30°F freezing point of water
• 100°F human body temperature
• 240°F boiling temperature of water
Fahrenheit
• Since his proposal all these temperatures
have changed.
• 0° was discarded since it did not specify
how to get the mixture (bad science);
some sources claim (1:1:1 ammonium
chloride, ice, water)
• 30°32° we now know water freezes at
32°F (not bad)
Fahrenheit
• 100°  98.6°F (it is believed that he did
not really reported a human’s temperature
but the temperature of a pig’s rectum.)
• 240°212°F (big difference) however,
considering that this had to be an
extrapolation since during his time
thermometers were made using alcohol
which boils before water it was a good
estimate.
Centigrade/Celsius
• First proposed by Anders Celsius
• Developed by using a rudimentary
thermometer (column filled with mercury)
• The column is brought into contact with a
mixture of water and ice and the height of
the liquid mercury is marked.
• Column brought into contact with boiling
water and height of mercury column is
marked.
Celsius
• Lower mark is set to be 0 and higher mark
100.
• The column is divided in 100 degrees
(grades)
– Hence the scale was originally referred to as
CentiGrade (100 degrees)
– Named changed to avoid confusion with other
centigrade (100th of a degree).
Kelvin
• Also known as the absolute temperature scaled
• Proposed by Lord Kelvin.
• The scale has “absolute zero” as its zero point.
– Absolute zero: in this scale there can’t be negative
temperatures.
• The scale was derived from P, T work done in
gases.
• Kelvin observed that if he extrapolated his
results to P=0 he always got the same
temperature. This is now defined as absolute
zero and/or -273.15°C.
Comparison of temperature scales.
Temperature Conversions
• The conversion from Celsius to Kelvin is simple
since the two scales are simply offset by
273.15°
o
K  C  273.15
• The conversion of Fahrenheit to Celsius, and
vice versa, can be accomplished with the
following formulas.
o
F  32 o F  1.8 (o C)  32
o
C
1.8
Temperature vs Heat
• Temperature is a relative measurement,
remember how the Centigrade scale was
developed. An object is hotter/colder than
a reference point.
• Heat is energy available to flow from
system to surroundings or vice versa
Metric System
• The system is based on factors of 10
• The basic unit (besides the kilogram) has
a “factor” of 100 =1
• You need to become familiar with the
following prefixes:
– Giga, Mega, kilo, deci, centi, milli, micro,
nano, and pico.
Metric Prefixes
Common Non-SI Units
• There are three common non-SI units used in
scientific measurements
• Mass: while the SI Unit is the kg… we usually
report mass in grams.
• Volume: The derived unit of Volume is the m3 yet
in measurements we vary the unit depending on
the substance… liquids are usually measured in
mL, gases in L
– 1000L=1m3 and 1mL=1cm3
• Temperature: we report temperatures in °C
Making Measurements
• Now that reference points (units) have
been agreed on we can proceed to actual
measurements.
• So you are making a measurement… what
should you worry about...
– There are two things that you should pay
close attention to.
• Reliability
• Reproducibility
Reliability
• Reliability
– Can we trust you to make good
measurements?
• Are your values close to the true values?
– Can you hit the Bull’s Eye?
• AKA Accuracy
Reproducibility
• Reproducibility
– Can we trust you to get the same result every
time?
• Did you get lucky?
– Are your values close to the average?
• Can you hit the same spot over and over?
• AKA Precision
•
Note: some books will tell you that you cannot be accurate if you are not
precise… that is false!!! It is possible to hit the 4 corners of a square and
you average will be the center (accurate)
Measurement and Significant
Figures
The first panel has no
precision but may be
accurate
The second panel shows
precision but no
accuracy
The third panel shows
both precision and
accuracy
Errors
• Since you are not 100% precise there will
always be an error.
% error = (mv - tv)/tv *100
mv = measured value
tv = true value
• No matter how careful you are there is
always some error.
– There are two types of errors
• Systematic
• Random
Uncertainty
• Uncertainty: this is the degree of
“unknown” or unreliability of your
measurement.
– To minimize this unreliability you need two
things…
• Practice in making measurements
• Precise equipment… the more precise the more
expensive the instrument will be.
Systematic Error
• When something goes wrong we tend to
blame the instrument we are using.
– At times it is the machine that is wrong
• The error can be fixed by recalibrating the
instrument.
– Example: You use a mass balance and when you check
the mass of your standards no matter which one you use
you seem to be off by 10grams on the heavy side.
• You recalibrate by adjusting your machine by 10
grams or subtracting 10 from the measured value.
Random Error
• This is just part of life and no matter what
you do you cannot remove it from your
experiment.
– The only thing you can do is to perform many
trials and average out the error.
• While in systematic error the measure value is
always either too high or too low… on random
error at times it is low and at times it is high and
the value will vary.
Reporting Values
• Ok so you are careful, reliable, and can
reproduce measurements… NOW WHAT?
• You need to know how to report numbers.
– At times it is best to report numbers in
scientific notation… a good rule of thumb is.. If
its bigger than 16 use Sci Notation… if it is
smaller than 0.1 use Sci notation.
Measurement and Significant
Figures
• 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.
Measurement and Significant
Figures
All measurements,
no matter how
careful they are
taken, have an
error, uncertainty,
associated with
them.
Measurement and Significant
Figures
• Number of significant figures
refers to the number of digits
reported for the value of a measured
or calculated quantity, indicating the
precision of the value.
• To count the number of significant figures in
a measurement, observe the following
rules:
Measurement and Significant
Figures
All nonzero digits are significant.
Zeros between significant figures are significant.
Zeros preceding the first nonzero digit are not
significant.
Zeros to the right of the decimal after a nonzero
digit are significant.
Zeros at the end of a nondecimal number may or
may not be significant. (Use scientific notation.)
Significant Figures
• 3456 has
• 4 sig figs.
0.0486 has
3 sig figs.
• 16.07 has
• 4 sig figs.
• 1.300 has
• 4 sig figs.
• 1310 has
• 3 sig figs.
• 1310. has
• 4 sig figs.
Measurement and Significant
Figures
When multiplying and dividing measured
quantities, give as many significant figures as the
least found in the measurements used.
When adding or subtracting measured
quantities, give the same number of decimals as
the least found in the measurements used.
Rounding
• It is good practice to carry 2 extra
significant figures through the
calculation and only round at the end
• 1st non-significant figure >5 round up
• 1st non-significant figure <5 round down
• 1st non-significant figure =5 then:
– If last significant figure is odd round up
– If last significant figure is even round down
Measurement and Significant
Figures
• 14.0 g /102.4 mL = 0.136718
g/mL
Measurement and Significant
Figures
• 14.0 g /102.4 mL = 0.136718
g/mL
only three significant figures
Measurement and Significant
Figures
• 14.0 g /102.4 mL = 0.137 g/mL
only three significant figures
Measurement and Significant
Figures
• An exact number is a number
that arises when you count items
or when you define a unit.
For example, when you say you have nine
coins in a bottle, you mean exactly nine.
When you say there are twelve inches in a
foot, you mean exactly twelve.
Note that exact numbers have no effect on
significant figures in a calculation. They have
infinite number of sig figs if you will.
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.
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 131m to
millimeters.
131meters 
1000mm
 131,000mm
1 meter
Note that the units have cancelled properly to
give the final unit of mm.
Units: Dimensional Analysis
• The ratio (1000 mm/1 m) 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 2.2
General Chemistry I
Virginia State University
Chapter 2
Dr. Vilchiz
Summer 2008
Alchemy
• Is in essence the ancestor of modern
chemistry
• Alchemy was more than a science
– It was a philosophy
– A way of life
• Alchemists strived to reach pureness and
perfection.
– Alchemists venerated gold as the symbol of
perfection
Alchemy
• It was believed that to posses gold will
make you rich and pure
• To drink gold meant to live forever
• Alchemy became the movement to find a
way to transform (transmute) matter into
gold.
• While these believes may seem silly
and/or far-fetched they were widely
accepted.
The search for the elixir of life
• It was known that it was possible to take
iron and make steel and that if you mixed
copper and zinc you will get brass.
• Why wouldn’t be possible to make gold?
• Needless to say the search for the elixir
was futile and eventually alchemy gave
way to new scientific questions and
approaches.
Alchemy’s Legacy
• While it might be true that alchemy
failed to produce an answer to its driving
force it was not by any means a waste.
– Many process we now use were discovered
or developed during the alchemists years.
• Distillation
• Fermentation
• Putrefaction
– Many elements were also discovered
• Bi, Zn, As, Co, and P
Alchemy to Chemistry
• Where does Chemistry come from?
– We are not 100% sure where the name
comes from but there are several
possibilities
• It could had come from Egypt Khem = turn
black
• It could had come from GreeceCheo=to cast
• It could had come from ChinaChin-I=gold
making juice.
Where to now?
• So making gold was not possible… now
what?
• The obvious question will be then… why
can’t we make gold?
• The quest to understand what was going
on began and thus modern chemistry was
born.
Ancient to Modern
• In the ancient times of alchemy we had
only “AFEW” Elements.
–
–
–
–
Air
Fire
Earth
Water
• Currently we know 116 elements of
which only 111 are recognized by the
IUPAC.
The study of AIR
• Gold was replaced by air as the primary
study subject.
– It is abundant and it behaves differently
under different circumstances.
• It was pointed out that at times when air
came in contact with lime water it will
produced a cloudy solution.
– This air was baptized as “Fixed Air”
• We now know it as CO2 (carbon dioxide).
Types of AIR
• There were other times when air led to
fiery explosions.
– This type of air is referred to as “explosive”
air
• It is now known as Hydrogen
• Air at times produce very noxious odors.
– Thus, it was referred to as “noxious air”
• This one is now known as Nitrogen
• We are missing one type of air that is
very important.
Where did the OO go?
• Oxygen was discovered while
experiments with mercury (I) oxide
were performed.
– As HgO is heated a separation of the
elements takes place resulting in liquid
mercury and gaseous oxygen.
• As the experiment was concluded a
smoldering piece of wood burst into
flames, hence Oxygen was known as
“flammable air,” as the just heated HgO
sample was placed close by.
Laws of Mass
• We are all familiar with the Law of
Conservation of Mass
– “Matter can not be created nor destroyed”
• While this notion is very familiar to us it
was not until the 1700’s that it was
actually stated:
– “Any mass gained by a substance in a
process comes from the surroundings”
Atomic Theory
• Atomic Theory is derived from the 3
mass laws:
– Law of Conservation of Mass
– Law of Definite Composition
– Law of Multiple Proportions
Atomic Theory of Matter
A chemical reaction consists of the
rearrangements of the atoms present in the
reacting substances to give new chemical
combinations present in the substances
formed by the reaction.
Atoms are not created, destroyed, or broken
into smaller particles by any chemical
reaction.
Atomic Theory of Matter
• The Atomic Theory was introduced by
Dalton
– The theory represented a new way of
thinking or chemical philosophy.
– It was based on four assumptions now
known as “The Atomic Theory
Postulates”
• Keep in mind that a lot of experiments have
been performed since the theory was first
introduced.
Postulates of Dalton’s Atomic
Theory
Postulates of Dalton’s Atomic Theory
Atoms are the smallest unit of matter. An atom is an
extremely small particle of matter that retains its identity
during chemical reactions.
Atoms of Element A cannot be converted to Atoms of
element B
Atoms of the same element are identical. Each atom of
an element has the same properties. Mass is one such
property. Thus the atoms of a given element have a
characteristic mass.
A compound is a type of matter composed of atoms of
two or more elements chemically combined in fixed
proportions.
Atomic Theory in Present Times
• Dalton’s theory has not been able to
withstand all the experiments
performed since it was introduced.
– The problem with the theory is that it is
too simple.
• Yet it was revolutionary in its own time.
• The Theory tells us about simple ratios of
elements in compounds but it does not tells
us why.
• The theory does not explain charged
particles
Radioactivity
• One of the pieces of evidence for the
fact that atoms are made of smaller
particles came from the work of
Marie Curie (1876-1934).
• She discovered radioactivity, the
spontaneous disintegration of some
elements into smaller pieces.
“New” Experiments and the
Atomic Theory
• Since Dalton introduced Atomic
Theory new experiments have been
performed:
– Alpha radiation which lead to the
discovery of the nucleus
• Nucleus is 10,000 x smaller than the atom
– Nuclear reactions have been performed
– Isotopes were discovered
Atomic Theory Revisited
• Postulate #1 atoms are the smallest
component of matter
– Not true, smallest are protons/electrons
and neutrons, but the atoms are the
smallest body to retain unique identity
• Postulate #2 Atoms of cannot be
converted to another element.
– Not true, nuclear reactions allows us to
do just that
Atomic Theory Revisited
• Postulate #3 atoms of the same
element are identical
– Not true, isotopes of elements have been
discovered where the number of
neutrons may vary.
• Postulate #4 Ratio of elements in a
compound is specific.
– Still true
• While the model was too simple it has
been a great starting point.
The Electron
• The electron was discovered by J.J. Thomson
between 1898 and 1903. NP 1906
• While an electrical discharge is applied to a
tube filled with gas he noticed a “ray” (cathode
ray tube). If a magnetic field was present near
the ray it bent the ray away from the negative
pole.
• The ray was made of negative particles
(electrons).
A cathode-ray tube. The fast-moving
electrons excite the gas in the tube, causing a
glow between the electrodes.
Deflection of cathode rays by an applied
electric field.
The plum pudding model of
the atom.
Charge of an Electron
• We all know that the electron is negatively
charged… but just how much charged does it
have?
• Robert Millikan devised an experiment that
answer this question.
Millikan’s Experiment
The Nucleus
• In 1911 Rutherford performed an experiment in
which he bombarded a thin piece of gold foil
with alpha particles.
• An alpha particle is a helium nuclei.
• He noticed that while most of the alpha
particles went through some were
reflected/deflected. The only explanation to
this was that there was a solid component in
the inside of the gold atoms.
Rutherford's experiment on
-particle bombardment of metal
foil.
(a) The expected results of the metal foil
experiment if Thomson's model were
correct. (b)Actual results.
Composition of an Atom
The nucleus of an atom is composed of
two different kinds of particles: protons
and neutrons.
Since they reside in the nucleus sometimes
they are referred as nucleons.
The nucleus is surrounded by electron
clouds which are much bigger than the
nucleus.
Electron clouds have no definite
shape and therefore are hard to
measure in size.
Atomic Theory of Matter
A proton is the nuclear particle having a
charge equal in magnitude to that of the
electron’s (a “unit” charge) but opposite
in sign. The proton’s mass is more than
1800 times that of the electron’s yet it is
10000 times smaller. (representation of an
atom)
The number of protons in the nucleus of
an atom is referred to as its atomic
number (Z).
Atomic Theory of Matter
An element is a substance whose atoms all
have the same atomic number. This means
that the number of protons dictates the type
of element.
The neutron is a nuclear particle having a
mass almost identical to that of a proton, but
no electric charge. Notice there is no
restriction as to the number of neutrons for a
given element.
Table 2.1 summarizes the masses and
charges of these three fundamental particles.
Atomic Theory of Matter
The mass number is the total number of
protons and neutrons in a nucleus.
A nuclide is an atom characterized by a
definite atomic number and mass number.
The shorthand notation for a nuclide consists
of its symbol with the atomic number as a
subscript on the left and its mass number as
a superscript on the left.
sodium  23
23
11 Na
Atomic Theory of Matter
Isotopes are atoms whose nuclei have the
same atomic number but different mass
numbers; that is, the nuclei have the same
number of protons but different numbers of
neutrons.
Sodium, for example, exists as two isotopes:
sodium-23 and sodium-24.
23
11
Na
24
11
Na
The fractional abundance is the fraction of
a sample of atoms that is composed of a
particular isotope.