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A Few Laws
• Conservation of Matter-For any chemical or
physical change , the mass of system must be
conserved.
• Definite/Constant Proportions-In a pure
chemical compound, the mass ratio of any two
elements (and the fraction that the mass of a given
element is of the total) is an intrinsic property.
• Multiple Proportions-If an element exhibits more
than one equivalent weight, the ratio of the
equivalent weights will be a small whole number
ratio. (2/3 or 1/2, never 5.7/4.8)
Dalton’s Atomic Theory
• Nowadays it is common to begin discussions of elements
and their behavior from the atomic viewpoint. No one
questions the existence of the atom or its centrality in
chemistry
• That was not always the case. Modern atomic theory is
new and was being challenged by eminent scientists into
the beginning of the 20th century.
• The presentation of the atomic theory stated by John
Dalton varies enormously. For our purposes, we will
present three major points.
– Elements consist of indestructible particles called atoms
– There is a unique atom for each element and all atoms of the same
element are identical (not surprisingly, Dalton used mass as the
defining characteristic of an atom.)
– The atom is the unit of chemical change
• From a general perspective, Dalton’s theory is a perfect
explanation for how elements behave. Its weaknesses,
which are always obvious in hindsight, have to do with the
fact that he could not possibly have obtained the data
needed for the modern description of the atom.
• We know today that mass is not the defining characteristic
and that the atom has three subatomic particles. Further,
the physicists have gotten even deeper into the atom and
may eventually find that matter isn’t really.
• the electron-discovered in 1897 by JJ Thomsom. Thomas
Edison received his patents on the electric light two
decades earlier. Edison and his contemporaries thought
that electricity involved molecules
• the proton-discovered in 1920 by Ernest Rutherford
• the neutron-discovered in 1932 by a former student of
Rutherford-James Chadwick
Elements and the Periodic Table
• It was found “early on” that certain materials could not be
broken down further by physical or chemical means.
• Over the centuries, the number of such materials, called
elements, grew steadily
• In the mid 19th century, several scientists, chief among
them being Mendeleev, organized the elements based upon
their physical and chemical properties. This grouping was
very logical and you could easily duplicate much of it
today with sufficient data.
• The marvelous aspect of Mendeleev’s work is that it fits
our modern view of the elements-one based on atomic
structure.
• Three interesting web sites for simple exploring:
– http://www.webelements.com
– http://www.chemicalelements.com
– http://chemlab.pc.maricopa.edu/periodic/default.html
What you need to know about the Elements and the
Periodic Table-deadline Oct. 21,2003
Understand the general separation of the periodic table into
metals and nonmetals and be able to list and define the
characteristics of a metal
Identify the positions of the alkali metals, alkaline earths and
the halogens. Know how the major chemistry of these
groups is predictable.
Using periodic position-be able to predict whether a
combination of two elements is likely to be ionic
Know the elements which are gases and liquids in their
elemental forms
Know which elements are monoatomic, diatomic and
polyatomic
Know what is meant by allotrope and isotope
As much as possible, once you begin doing problems required
data about the elements, you should use a periodic table.
This will reinforce the positions of the elements.
Some elemental facts
• Mendeleev had 63 elements in his periodic table. The
current element count is about 114 with some disputes
ongoing
• Many elements have been known since antiquityparticularly the coinage metals. No date of discovery is
claimed for iron, copper, silver, tin,gold or lead.
• nitrogen and oxygen were discovered in 1772 and 1774
respectively
• By 1900, most of the naturally occurring elements had
been discovered with the last group consisting largely of
the inert gases, all discovered between 1895 and 1898.
General thoughts on the modern atom
• Consists of a nucleus which contains protons and,usually,
neutrons
• Outside of the nucleus we find the electrons-the volume
they occupy determines the effective size of the atom
• The defining characteristic of an element is the number of
protons which gives the element its atomic number
• Neutral atoms have equal numbers of electrons and protons
• An atom has a mass number equal to the sum of protons
and neutrons
• The chemistry of an element can be described in terms of
the behavior of its electrons
• ion formation is through by electron loss or gain
– cations (+ charge) electrons have been lost
– anions (- charge) electrons have been gained
Representing everything
N=atomic number
M=mass number
C=charge
e=# of electrons
p=N=#of protons
n=#of neutrons
p-e=C(this carries a sign!)
p-C=e
M-N=n
X=atomic symbol
M
N
X
C
Atomic weight/mass
It is assumed that everyone knows the difference
between mass and weight. However, use of terms
like “atomic weight” and “molecular weight” is
common, even if not purely correct. You find such
usage throughout the chemical literature and you
should not be bothered by it or troubled if you use
it.
Before beginning discussion of the mass/weight of
the atom, some attention should be given to
another measure of an element’s properties for
which mass is the basis-the equivalent weight.
Equivalent Weight
• It has already been noted that elements have a restricted combining
capacity with regards to the formation of compounds. This behavior
was tracked by analysis of mass proportions.
• One can readily take it to another level by choosing an element which
forms a large number of chemical compounds and defining a certain
mass as its standard combining mass. Oxygen is a logical choice and
lets define its reacting mass as 8.0grams(the actual value of the number
isn’t important). We then find the following
– In a compound with hydrogen, there are 2g of H per 16g of O-therefore
the mass of H equivalent to 8g of C is 1 which becomes its equivalent
weight.
– In 2 different compounds of C and O, we find equivalent weights for C of
4 and 6.
– If this all works, we should be able to find a simple compound of C and H
with the mass proportions of C to H equal to the ratio of their equiv
weights(1 to 4 (or 6). There are lots of these.
– Note also that the ratio of the two equivalent weights for C is a small
whole number ratio
Back to Atomic Weights/Masses
• The key ideas in the atomic weight scale
are:
– It’s relative
– It’s arbitrary and may be changed
• Currently we set the mass of 12C=12amu
and the remaining atoms are scaled relative
to that. An atom which had a mass
1.45times that of 12C would have a mass of
17.4amu.
Isotopes and mean atomic masses
• As has already been noted, atoms of the same element can
differ with regards to the number of neutrons present and,
thus, also with regards to the atomic mass
• Atomic masses noted in the text and on periodic tables are
mean atomic masses based upon the natural isotopic
abundances of the element.
• The calculation of a mean atomic mass is straightforward
and illustrated in your text several times. This calculation
will not be asked on quizzes or exams, but students should
be able to answer questions like: Element X has two
isotopes with masses of 45 and 47. If the mean atomic
mass is 46.3, which isotope is predominant?
Practical Atomic Masses-what’s a mole
• In chemistry, we deal with matter on the large scale
• Dalton’s Law tells us we need to measure elements in some manner
that can be looked at as an “atomic measure”
• If all apples weighed 23.0ounces and all oranges weighed 13.2ounces,
I can readily scale these up to any other weight and be assured that I
have the same number of apples and oranges, so long as the ratio of
23.0 to 13.2 is maintained. Thus 23.0tons of apples contain the same
number of apples as oranges in 13.2tons of oranges.
• Thus, I simply convert the mass unit used for atoms (the amu) into a
macro mass unit (grams) and call quantity the mole. The mass of one
mole of any pure substance is called its molar mass. For atoms, one
often hears gram atomic weight.
• For chlorine-atomic weight = 35.453amu, gaw=35.453g
• n(moles)=mass/gaw (single most important equation in all of general
chemistry)
• Useful thought: gaw is the minimum quantity needed to have a mole. If
n<1.0 mass<gaw
• What is the gaw of chlorine expressed as mg/mmole?
Now for something completely different-molar volume
• It should be apparent that knowing the volume of
one mole of an element could be useful in a
number of ways. At a minimum, such a quantity
will make a mole more “real”. Molar volumes are
readily determined from the ratio of the gaw to the
density.
• MV=V/n (gee-a three variable equation-who
would have ever guessed)
• dimensions are usually cm 3/mole
• Typical values(non-gaseous elements)
– Al:10.0, Zn:9.2,Ag:10.3,Fe:7.1,Pb:18.3
– roughly how large is a mole?