Download The Building Blocks of Matter

Chapter 3
Atoms: The Building Blocks of
Matter
3.1 The Atom: From Philosophical
Idea to Scientific Theory
–
4th century B.C. Greek
philosopher Democritus
stated the universe was made
of invisible units called atoms
(atom “unable to be divided”).
–
Believed movements in atoms
caused changes observed in matter.
What held these atoms together?
Democritus did not know…. And no
one did until…
• Late 1700’s
– French Chemist, Antoine Lavoisier
established the Law of Conservation
of Mass.
Mass can NOT be created or destroyed; only
change forms.
– Joseph Proust, later established the
Law of definite proportions
the compound always contains the same
elements in the same proportions by mass.
• Law of multiple proportions- if 2 or more
different compounds are composed of the same
2 elements, then the ratios of the masses of the
2nd element combined with a certain mass of the
1st element is always a ratio of small whole #s.
Dalton’s Atomic Theory
– In 1803 John Dalton proposed an
atomic theory:
• Each element is composed of
extremely small particles called
atoms .
• Atoms of the same element are
exactly alike .
• Every compound always has
the same ratio and kinds of
atoms .
• A chemical rxn is a
rearrangement of atoms; they
are not created or destroyed .
• Remember: an atom is the smallest
part of an element that still has the
element’s properties.
Today…
• There are some exceptions to Dalton’s
theory; however they still are the basis for
understanding Chemistry.
– 1. all matter is composed of atoms
– 2. atoms of any one element differ in properties
from atoms of another element
• BUT… how can we be sure they exist?!?
3.2 The Structure of the Atoms
• Atom- “ the smallest particle of an
element that retains the chemical
properties of that element.”
The Structure of the Atom
• We now know that atoms can be divided into
many different subatomic particles. For
example:
– Nucleus- the center of the atom
– Protons- positively (+) charged subatomic particle.
– Neutrons- neutral (not charged) subatomic
particle.
– Electrons- negatively (-) charged subatomic
particle
Discovery of the Electron
• The English Chemist, Michael
Faraday, in 1839, showed
atoms contain electrical charge.
• American, Benjamin Franklin,
experimented with electricity
(think kite and key experiment).
Franklin was able to determine
the 2 charges an object has and
it was he who coined the names
“positive” and “negative”
charge. Opposite charges
attract and like charges repel.
• But where do these charges
come from? What are they?
The Discovery of the Electron
Cathode Rays and Electrons
• Cathode Ray Tube (CRT) - A battery is
connected to a tubing of partially evacuated
glass. The glass is lined with fluorescent
material, current flows to the ends of the tube.
The end connected to the (-) terminal of the
battery is called the cathode and the other is the
anode (+). A stream of radiation flows from the
cathode to the anode when the battery is turned
on.
http://www.youtube.com/watch?v=XU8nMKkzbT8
• English physicist, J.J. Thompson
(1856-1940), determined the ray
WAS made (-) particles by allowing
the ray to pass through a hole in the
anode and then through a magnetic
field. He determined that the
negative particles emanate from the
cathode, they had structure and he
named them electrons.
Charge and Mass of the
Electron
• American physicist, Robert Millikan
(1868-1953), was able to measure the
charge of an electron. He sprayed oil
and used X-rays to give the oil a
negative charged. Then measured how
different magnetic charges changed the
rate the oil fell.
He calculated the mass of the e- to be
9.11 X 10-19 grams.
Discovery of the Atomic
Nucleus
• New Zealander, Ernest Rutherford passed the
cathode ray of a radioactive substance between
two charged plates. The ray split- part of the
beam was deflected towards the (-) plate (alpha
radiation), another was deflected towards the
(+) plate (beta radiation) and a third passed
straight through undisturbed (gamma radiation).
• Alpha particles 2+ charge
• Beta particles 1- charge
• Gamma rays have no charge
• This experiment demonstrated that the atom
was much more complex than previously
thought.
Composition of the Atomic
Nucleus
• Thompson showed us
that atoms had electrons,
but that doesn’t explain
why atoms are electrically
neutral (they don’t have a
charge). If the have
electrons they must have
some type of (+) charges,
too.
• Thompson had suggested,
earlier, that the atom was
like plum pudding. The e-s
were spaced evenly
throughout the atoms (+)
interior, the way the plums
were distributed through the
pudding. (You can also think of
chocolate chip cookie dough- the dough is
the positive interior while the chocolate
chips are the e-s).
Alpha- Scattering Experiment
• 1909 The Alpha-scattering Experiment by
Rutherford- A beam of high-speed alpha particles
bombarded a thin sheet of Au foil. Most of the
alpha particles went straight through the foil but a
small portion DID deflect. And they would scatter in
every direction possible. Why??
http://www.youtube.com/watch?v=XBqHkraf8iE
• Rutherford’s data suggested the “plum
pudding model” was not true. He determined
that there must be an (+) core in the atom.
He named this the nucleus. The particles
which went straight through suggested the
atom was mostly empty space.
3.3 Counting Atoms
• Atomic Numbers
• Englishman, Henry Mosley (1887-1915)
discovered that each element had a
unique amount of positive charge.
This concept leads to the
understanding of why atoms of
different elements are unique. The
identity of the atom comes from the
number of protons contained in the
nucleus.
• Atomic # (Z)- the # of protons in the nucleus of
an atom. (Always a whole #)
– Ex. O-8 = 8 protons in the nucleus
– C-6 = 6 protons in the nucleus
– Pb-82= 82 protons in the nucleus
• Now, individual atoms are electrically neutral… meaning
they have an equal # of protons as electrons.
– O has 8 protons (we know this by its atomic number) and 8 electrons.
– C has 6 protons and 6 electrons
– Pb has 82 protons and 82 electrons.
Ions:
When p+ and e- are not equal.
• The atom can gain or lose electrons (NOT PROTONS- if it change
the # of protons you change the element/atom.) This is an ion. If you
gain an electron you get a negative ion. If you lose an
electron you have a positive ion.
• Charge on Ion = # of protons - # of electrons
– Magnesium (Mg) can loose 2 electrons making it a positive ion- written
as:Mg2+ . O will gain 2 electrons making it a negative ion. It is written as
such: O2- .
• Atoms that lose electrons, like Mg, are called cations.
Atoms that gain electrons, like O, are called anions.
Isotopes
• Just as we have the atomic # (# of protons) we
have atomic mass #:
Atomic mass # (A) - the # of protons + neutrons in
an atom.
• Because the bulk of the atom’s mass is provided
by the protons and neutrons, we only consider
their masses when calculating the atomic mass #.
– Ex: O has 8 protons and 8 neutrons, so A= 16.
• Dalton said that every atom of the same
element is exactly alike… NOT SO. They
do have the same # of protons; however,
they do not necessarily have the same #
of neutrons!
• This means that different atoms of the same
element may/will have different masses.
• These are:
isotopes (atoms having the same # of protons
but different # of neutrons).
– Some isotopes are more common than others. H
is present on the earth and the sun. The most
common isotope of H is protium, then deuterium
and finally, tritium.
– Ex: there are 3 isotopes of H (why A = 1.00794).
• Protium (only has one proton in the nucleus) A=1
• Deuterium (1 proton + 1 neutron) A=2
• Tritium (1 proton + 2 neutrons) A=3
• Calculating the # of neutrons in an atom:
General Formula:
A – Z= # of neutrons.
– Ex: calculate the # of neutrons in the isotopes of Cl (Cl-35 and Cl-37)
Relative Atomic Masses
• The mass of a single element is extremely small
(in the one trillionth of a billionth range) and is very difficult
to work with. So instead we express the mass of
atoms in atomic mass units (amu).
– One amu is equal to 1/12th of the mass of a carbon12 atom. This isotope has exactly 6 protons and 6
neutrons so the mass of each has to be about 1.0
amu.
• 1amu = 1/12 (mass of 126C atom) = 1.66x10-24 g
Average Atomic Masses of
Elements
• The atomic mass of an element is often listed as
the average atomic mass as found in nature. This
is a weighted average of the isotopes for that
particular element. The more commonly found
isotopes have a greater effect on the averages
mass than the more rare isotopes.
– Table 3-4 on p.80 in your book
– Ex: Cl 24% Cl-37 and 76% Cl-35, thus the average
atomic mass (35.45 amu) is much closer to 35 than 37.
Relating Mass to Numbers of
Atoms
• If we can have an atomic mass measured in
amu…. we can also have FORMULA MASS:
• Formula Mass- the sum of the atomic
masses of all the atoms in the compound.
• Ex. CO2
C- 12.01
O- 16.00 X 2
44.01 amu – formula mass
YOU TRY!
1. H2O
2. CH4Cl2
So… What’s a MOLE?!?
• We have a problem… how do you
measure amu’s? How do you count
atoms? They are so small… that you
don’t. Instead, we use a number called
the mole to help us.
• A mole (of any element) – the number of
atoms equal to the number of atoms in
exactly 12.0 g of carbon-12.
How does it work?
• There is an actual # … it is 6.02 x 1023
atoms/mole.
– If you have 1.0 g of H, you have 6.02 x 1023
atoms of H
– If you have 16.0 g O, you have 6.02 x 1023 atoms
of O
How many grams of Pb would you need in order to
have 6.02 x 1023 atoms of Pb?
Notice the amu- gram link?
• “The mass in grams of 1 mole of a substance is
numerically equal to its atomic mass or formula
mass in atomic mass units.”
• “Mole” is just the name of the number we use in
chemistry to help count atoms.
– Like…
•
•
•
•
A coupleA dozen –
A baker’s dozenA year-
Avogadro’s Number
• The Mole is also known as Avogadro’s Number (N)
in honor of Amadeo Avogadro, the Italian chemist.
How Big is Avogadro’s number?
602,000,000,000,000,000,000,000!
if it was a stack of paper… it would reach beyond our
solar system!
…basketballs… it could create a planet the size of the
earth!
…grains of rice… would cover the earth to a depth of
246 feet!
Molar Mass
• “The mass in grams of 1 mole of a substance..” (M).
• Molar Mass of…
–
–
–
–
–
–
Ca 40 amu –> 40 g/mol
C 12 amu --> 12 g/mol
H2O2 34 amu --> 34 g/mol
NaCl
CaCl2
C6H12O6
Gram-mole-Particle
Conversions
• We can use the molar mass to convert moles to
grams or grams to moles.
• Likewise, moles can be converted to # of particles
via Avogadro’s #.