Part II
   
Demonstrations
Demonstration 1: Spectra in the Real World
Give each student a small square of plastic diffraction grating to view H 2 , He, Ne,
etc. emission from spectrum tubes. Let students take the grating home to view
sodium street lamps and mercury vapor lamps. WARNING: Do not look at the sun!
 
Demonstration 2: Getting a Charge out of Things
If students havenít seen the demonstration of an amber or glass rod (or a length of
PVC pipe works very well) rubbed with fur, silk, or hair, and then used to attract or
repel pith balls or other objects (such as suspended balloons), this is a good place to
perform it.
Demonstration 3: Standing Waves
An old-style manís electric razor, with the cutting top removed, has an exposed
vibrating post. By attaching a string to this post, you can show quantized standing
waves when the razor is turned on.
 
  
Demonstraion 4 modeling orbitals.
Spherical balloons can be used to model s-orbitals. Two spherical or oblong balloons tied together
can be used to model p-orbitals. 
 Materials 
 6 Balloons (2 each of 3 differentcolors).Tiepairs
of balloons (2 the same color) together. Bunch. 
Demonstration 5: Wavelength and Color
To demonstrate the relationship between wavelength and color in the visible range
from red to violet, use a spectrophotometer. Place a piece of white chalk in a
spectrophotometer cell (see Figure 7). Cut (or sandpaper) the chalk at an angle that
will reflect light upward to the studentís eye looking down into the cell holder. Open
the slit width. Rotate the wavelength knob to compare colors and wavelengths.
 
 
 
Demonstration 6: Simple Mass Spectroscopy
Materials
Overhead projector
Sheet of Plexiglas, 8-1/2 x 11 in
Magnet, disk shape, 3-4 in diameter
3-4 Steel ball bearings of various sizes
Procedure
1. Form a shallow inclined plane with the Plexiglas sheet on top of the overhead
projector, by using a small block of wood or other material to elevate one end of
the sheet.
2. Insertthe magnetunderneaththecenter of the inclined plane.
3. Roll the steel balls down the incline. If they are rolled near but not over the
magnet(theywill stop)theirdirectionof motion will be altered. The smaller ball
bearings will be deflected through a greater angle than the larger ones.
Demonstration 7: Flame Tests
Absorptionandemissionspectra areusedtoidentify manyelements.Whenelements
are heated to high temperatures they may be placed in an excited state; in this
excited state valence electrons move to higher energy levels. When the electrons
return to their ground state, they may emit visible light of characteristic colors that
can be used to identify the element.  In this demonstration, students will identify the colors of the emission spectra of
some metallic ions, using a burner and wood splints.  Place the substance in the hottest part of the flame Substance to be tested Wood splint a a a a
Materials
Burner
Wood splints
Copper(II) nitrate, Cu(NO 3 ) 2 , solid crystals
Barium nitrate Ba(NO 3 ) 2 , solid crystals
Sodium nitrate, NaNO 3 , solid crystals
Strontium nitrate, Sr(NO 3 ) 2 , solid crystals
Procedure
Place a small amount of one of the solid ionic compounds (about the size of
a rice grain) on the tip of a wooden splint. Place the splint at the edge of the
hottest part of the flame (top of inner blue cone) and observe the color. (Some students, particularly some
males, may be color blind. Consider allowing students to compare their observations in small groups, alerting
them to the possibility of color blindness.)  (Note: Good results have been obtained with
concentrated syrups of the nitrate salts.) 
Results
Key Questions
1. What are the principal parts of an atom? [The nucleus and electron cloud. The
nucleus contains protons, neutrons and many other subnuclear particles (see
Nuclear Chemistry module). The electrons can be divided into core and
valence electrons.]
2. What forces hold atoms together? [Coulombic (electrostatic) attractions between
oppositely charged protons and electrons. Coulombís law is ]
3. How are electron orbitals different from orbits?
[Orbitals refer to the probable location or region of space (a volume) where an
electron is likely to be found with a specified degree of certainty. It arises from
application of wave and uncertainty considerations. Orbits are Niels Bohrís
representation of rigid circular or elliptical paths followed by particulate
electrons.]
4. What is an energy level diagram? [A graphic representation of the relative
energies possessed by electrons in various orbitals as described by their
quantum numbers.]
Ionic salt Metal ion Color
Copper(II) nitrate, Cu(NO 3 ) 2 Cu 2+ Blue-green
Barium nitrate, Ba(NO 3 ) 2 Ba 2+ Green
Sodium nitrate, NaNO 3 Na + Yellow-orange
Strontium nitrate, Sr(NO 3 ) 2 Sr 2+ Bright red
F = k q1 q2
 
5. How are electron energies related to orbitals? [The principal quantum number
gives a general indication of an electronís energy and sets restrictions on its
orbital type. Higher energy orbitals are those with higher probability of being
far from the nucleus.]
6. How do cations form? [If an electron absorbs energy in excess of the amount
corresponding to the highest energy orbital in its electron cloud, it will be ionized,
that is, leave the atom. The resulting positive ion left behind is a cation.]
7. What determines the ionization energy of an element? [The nuclear charge (the
binding forceís source) and the energy possessed by the electron (its energy level)
combine (actually compete) to determine additional (ionization) energy required.]
8. What determines the commonly formed cation of an atom? [When a metal
and nonmetal combine, the overall process is exothermic, largely due to the
highly negative lattice energy involved; thus reaction takes place. One of the
processes that requires energy in forming the compound is the ionization of
the metal ion. It is the relative magnitudes of the ionization energies of a
particular metallic species that determines the charge on the cation. For
example, sodium has a first ionization energy (I 1 ) of 495.8 kJ/mol; its second
ionization energy (I 2 , the energy needed to remove a second electron from
sodium) is 4562.4 kJ/mol, almost 10 times I 1 ! No reaction involving the formation
of Na 2+ could ever be expected to be exothermic; hence, Na 2+ does not exist. A
second example is the formation of magnesium ions. I 1 , I 2 , and I 3 for
magnesium are 737.7, 1450.7, and 7732.6 kJ/mol, respectively. The large
energy requirement for the formation of Mg 3+ precludes its existence. Mg +
requires the least amount of energy for formation, but is unstable because of
its half-filled 3s orbital. Hence, the most stable magnesium ion is Mg 2+ ,
requiring a total energy expenditure of 2188.4 kJ/mol, which is certainly
within the realm of a possible overall exothermic process for the formation of
a magnesium salt.]
9. What is the process of electron excitation? [An electron is said to be excited
if it absorbs energy and moves to a higher energy orbital.]
10. Why do atoms emit light when excited? [Being in an excited state means having
excess energy. Like anything else, atoms tend to lose their excess energy, so the
electron drops to a lower energy orbital. The lost energy is given off as light (DE = hv).]
11. How do the physical sizes of atoms and nuclei compare? [Atoms tend to have
diameters of around 0.1 nm (1-10 Å). The nucleus is about 10 15 m (1 femtometer)
in diameter but contains 99.9% of the mass.]
12. What are the three dimensional shapes of s and p orbitals? [The 90% probability
boundary for an s-orbital is a sphere. For p-orbitals the boundary is dumbbell
shaped. Three p-orbitals, one dumbbell per axis, comprise a set.]
13. What is meant by wave-particle duality? [In some instances, the behavior of
electromagnetic radiation and electrons can best be understood in terms of wave
theory, and in other instances, their behavior can best be understood if we treat
them as discrete particles. Electromagnetic radiation is usually described in
terms of wavelength or frequency; the quantity most closely associated with
particles is mass or momentum, the mass-velocity product. Louis de Broglie in
1924 offered the startling proposition that light may sometimes display particle-like
characteristics, and that small particles may sometimes display wavelike
properties. He summarized his proposition in a now famous equation l = h/mv,
where l is the wavelength, h is the Planck constant, m is the mass, and v is the
velocity of the object in question. Wave-particle duality is only important when
the wavelength is of atomic dimensions, i.e., of the order of 10 1 - 10 3 picometers.]
14. What information is provided by quantum numbers? [All we can hope to
know about an atomic electron: its energy, angular momentum, magnetic
moment and spin.]
15. How do quantized and continuous processes differ? [Anything quantized can
be described in terms of separate discrete units. Those that are continuous are
without beginning or end and may be had in any quantity. Ice cubes are
quantized but water flows continuously.]
16. What is the electron configuration of (specify an element)?
 
Examples: H 1s 1
Li 1s 2 2s 1
Ga 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1
Counterintuitive Examples
1. Adjustable light dimmer switches (unless specifically designed) do not work on
fluorescent lamps. You can demonstrate this fact by first connecting a standard,
preferably clear, incandescent light bulb to a dimmer switch or laboratory
voltage controller such as a Powerstat. By slowly increasing the voltage,
you can brighten the incandescent bulb in a continuous manner. Repeat
using a fluorescent lamp. Since the mercury gas enclosed in the fluorescent
lampís envelope requires a minimum ionization energy, a quantized process,
the lamp will not go on until this minimum energy is applied.
2. A weight attached to a string and rotated around oneís head continues in
motion only as long as energy is applied. An electron, perceived to be in a
classic Bohr orbit, does not slow down. This behavioral discrepancy was one
of the counterintuitive problems that Bohr had to solve.
3. Energy changes do not always seem to be quantized. If you heat a piece of
metal, its temperature rises on a continuum; the metal does not seem to
acquire energy in quantum leaps. On the other hand, if you step halfway
between two ladder rungs, you find that you remain on the first step; you
must acquire potential energy in quanta dictated by the step-size of the
ladder. The manner of acquiring energy in these two instances is not
different. The energy step-sizes for a metal are so infinitesimally small that
they cannot be detected by a relatively crude probe such as a thermometer.
4. As electrons are added to atoms in the same period of the Periodic Table, the
atomic size decreases. Adding an electron also entails adding a proton to the
nucleus, thus increasing effective nuclear charge. The nuclear charge
density increase is proportionally greater than the electron energy increase,
leading to more compact (smaller) atoms. A Periodic Table showing atomic
sizes (such as the one included in the Periodicity module) will help illustrate this
point.
5. Light radiated by an excited atom consists of several colors, but the human
eye perceives only one composite color. This composite color can be resolved into
its components by viewing it through a diffraction grating. A spectral tube for
neon or helium, or even an orange night light, illustrates this fact very nicely.
6. Because electron probability functions are asymptotic, they do not become
zero even at a great distance from the atomís nucleus. Consequently, there
is a very small chance that some electron associated with an atom in your
body is at this moment very far away. Perhaps a part of you has been to the
moon and you just donít know it!
 
 
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