Lecture 1 Notes, Immaculata Week, July 2011, Charles H. Mahler, Lycoming College
My email address is mahler@lycoming.edu if you have questions. Please note I also built
the 36 air powered rocket launchers for the Summer of Innovation Camps and would be glad to
try and answer questions on those.
I opened by playing “The Elements” song by Tom Lehrer from the e-book “The
Elements” by Theodore Gray for the iPad. Showed some features of the e-book (about $13 for
the iPad), as well as the corresponding free website (www.periodictable.com). Noted that Geay
says he will freely license his images for educational (though not commercial) purposes, if you
(or your students) contact him.
Moved on to the NIST Periodic Table (www.nist.gov/pml/data/periodic.cfm) and opened
discussion on “How do you think about the periodic table, how do you talk to your students
about the periodic table?”.
One way I think about the periodic table is as a way to summarize a lot of information in
as concise a form as possible: atomic mass, atomic number, electron configuration, ionization
energy, etc. can all be shown on a periodic table. Another example of this idea is the ideal gas
law, which takes many observations and experiments on pressure, volume, number of moles, and
temperature and summarizes their relations (PV = nRT).
Next we looked at Mendeleev and the historical development of the periodic table.
Pointed out how Mendeleev organized the elements by their chemical properties and by
increasing atomic mass. He was not afraid to ignore the apparent order given by atomic masses if
the chemical properties were more consistent that way (so Tellurium has a slightly greater mass
than Iodine, but he put Te before I – we now know the difference is due to their isotopic
compositions).
Mendeleev was smart enough to realize that there were missing elements which had not
yet been discovered. He used the periodic table to predict the properties of what he called “ekasilicon” (Germanium). “Eka” is the Sanskrit word for “one” so the name means the element one
beyond Silicon. Mendeleev predicted many properties of the missing element, and when
Winkler discovered / isolated Germanium in 1886, Mendeleev’s prediction were found to be
pretty accurate. These included the as atomic mass (predicted 72 amu, actual 72.6 amu), density
(predicted 5.5 g/cm3, actual 5.35), the correct chemical formulas for the oxide and chloride
(GeO2 and GeCl4), and others such as melting point of the element and its compounds.
The use of a theory to make and test predictions is very much part of the scientific
method, and Mendeleev’s successful predictions of Germanium and Gallium and their properties
helped convince chemists of his day that his proposed periodic table was correct.
Other ways of looking at the periodic table include dividing it into the metals on the left
and center, non-metals on the right, and the semimetals (or metalloids) on the “stair step line”
between them. Discussed the properties of Si, a semimetal, and how its electrical conductivity is
between the high conductivity of metals and the non-metals, which are insulators. Silicon is a
semiconductor and this property has allowed the development of integrated circuits, microchips,
and computers and devices that use such chips.
Also noted that the NIST periodic table shows which elements are solids, liquids, and
gases, as well as which are naturally occurring and synthetic (radioactive). Talked about
naturally occurring radioactive elements found in the Earth’s crust, and synthetic ones made in
particle colliders.
We looked at how the periodic table reflects atomic structure. The nucleus is very small
and heavy and contains positive protons and neutral neutrons. The atomic number, Z, is just the
number of protons in the nucleus and is the same for all atoms of a given element (so all
Uranium atoms have 92 protons and Z = 92). Moseley discovered the atomic number Z (and was
killed in WWI at Gallipoli). Knowing about and being able to measure Z meant that the order of
elements in the periodic table could be checked (Mendeleev got it right). Knowing Z also
allowed any missing elements (or “holes” in the periodic table) could be identified, and showed
where there were no “holes” possible.
We briefly discussed isotopes, which are atoms with the same Z but different numbers of
neutrons and considered the three isotopes of hydrogen: 1H (one proton in the nucleus, well over
99% of hydrogen), 2H (deuterium, one proton and one neutron about 0.01%), and 3H (tritium,
one proton and two neutrons, this is radioactive and used in H-bombs). As the number of
neutrons in a nucleus increase, the nucleus becomes more likely to be unstable (radioactive).
Every element about Lead and Bismuth is radioactive.
Next we looked at the third particle making up atoms, the electron, which has a negative
charge. If an electron has a mass of 1, a proton’s mass is about 1837 and a neutron is about 1838,
so the electron is very light compared to the nucleus and its particles. Rutherford’s experiment
where he bombarded thing Gold foil with alpha particles (Helium nuclei) and saw that some
bounced back allowed him to formulated current model of the atom nucleus as small, massive,
and positively charged. The nucleus is surrounded by a cloud of electrons and the atom is mostly
empty space. There are as many electrons as protons in a neutral atom (to maintain charge
balance).
Chemistry is all about how atoms can transfer and share electrons. If an atom loses
electrons through ionization, it forms a positive cation. If it gains electrons, it forms a negative
anion. The periodic table allows us to predict the charges of many simple ions, as they follow the
octet rule. The octet rule says that atoms will gain or lose electrons to form ions that have the
same number of electrons (or electron configuration) as a noble gas. So for example, Chlorine
forms a -1 anion, and Sodium forms a +1 cation. This only holds true for the columns of the
periodic table closest to the noble gases – changes greater than 3 or 4 electrons are not seen as
simple ions.
Atoms can also form covalent bonds and share electrons. Although the sharing is not
always equal, electrons are typically found in bonds in such a way that the atoms still share eight
electrons (or two in the case of Hydrogen, since the noble gas Helium only has two electrons).
We looked briefly at Lewis Dot structures as a way of looking at and predicting covalent bonds
in molecules, and went over the concept of valence electrons, which are the electrons above the
noble gas shell (so sodium has one valence electron, and carbon has four).
Quantum mechanics shows that in atoms the electrons are in orbitals around the nucleus.
We looked at the Orbitron, a website which shows these orbitals in many different ways,
including images, animations, graphs, and equations. The periodic table’s structure reflects these
orbitals, with the s block (alkali metals and alkaline earths), p block (last six columns, includes
all the non-metals), d block (10 columns of transition metals), and f block (14 columns of
lanthanoids and actinoids).
We looked briefly at the WebElements website, which has more information on the
elements. For fun, we looked at a website alerting its readers to the dangers of Dihydrogen
Monoxide (water), and noted that this could be a way to help students think critically about what
they find on the web.
We did a hands on demonstration on making density columns with sugar water of
different concentrations (see separate description). This led to discussions of density and the
suggested demonstration of putting two cans of soda pop into water to see if they sink or float. A
can of sugar-containing soda pop sinks (dissolved sugar makes it denser than water) and can of
diet soda pop floats (no sugar, less dense).
We talked about how to make solutions and how to dissolve things like sugar in water.
Water has a dipole moment and is a polar molecule that can form hydrogen bonds; this helps it to
dissolve sugar (which has lots of OH groups) as well as salts (cations go to the negative end of
the molecule, anions to the positive end).
We discussed how heating a solution will generally make more solids (like sugar)
dissolve in it, but for gases the opposite trend is seen, and cold water dissolves more gas than
does warm water. Both trends can be explained by the increased energy of the water molecules
in a warm solution – this helps more sugar and most solids dissolve, but for gases the extra
energy allows the gas molecules to escape from the surface of the solution more easily. Gave
examples (cold oceans support lots of aquatic life, warm creeks do not) and a suggested
demonstration, opening two identical cans of soda pop except for temperature and pouring them
into identical containers. The warm soda pop will have more bubbles than the cold, as it cannot
dissolve the gas as well.
Finally I did exploding Pringles can demonstration, burning hydrogen gas with oxygen
from the air to make water. The flame burns steadily at first above the lid, then as it enters the
can through the hole, a radical reaction leads to the loud explosion. Discussed how hydrogen
burning is one of the most energy dense fuels known, and how the space shuttle used this
reaction to blast off. Pointed out that hydrogen could become a major alternative fuel in the
future.