1

Conceptual Physics Study Notes & Questions: Week 14—The Atom-II (Chap. 21)

1) Niels Bohr developed a model of the atom based on the idea that the electrons’ orbits have fixed energy levels, so when the atom absorbs or emits photons, their energy must exactly match transitions between these discrete energy levels (p451).• In later years, scientists realized that these electron orbital levels were

determined by the electrons’ wave functions.

2) Photons are quanta of EM energy. (p453). When the atom absorbs a photon, an electron make a quantum jump to a higher energy orbital. When the atom emits a photon, an electron makes a quantum jump to a lower energy orbital. Because the orbitals have discrete energy levels, only certain photon frequencies are associated with atomic transitions. This set of photon frequencies define the atom’s absorption and emission spectra (p456), that is, the frequencies of photons absorbed or emitted, respectively, by the atom (see Figure 21-5, p454).

3) The transition of an electron from higher energy orbital to a lower orbital is called electron relaxation—the electron is “relaxing back to its ground state.” (Below, we will see this ground state is determined by the Pauli Exclusion Principle.) Electron relaxation has two basic paths: one involving photon emission, the other involves a series of radiationless vibrational relaxation steps where the transition energy is release by the atom imparting kinetic energy kicks to its neighboring atoms. This latter process can only occur when an atom is in close contact with other atoms.

4) When an excited atom emits a photon on its own, the process is called spontaneous emission or fluorescence (p411). Excited atoms can also be triggered to emit radiation—stimulated emission (p460)—by a passing EM wave of the appropriate frequency. The newly emitted photon “joins the parade.” This is the basis of a laser (p460). How does a laser work?

5) Subatomic particles fall into two categories: Fermions, such as electrons, protons and neutrons, which have a “½ spin” quantum mechanical (QM) state, and Bosons, such as photons, which have “an integer spin” QM state. These spin states are part of the fundamental QM properties of subatomic particles.

6) The Pauli Exclusion Principle is the observation that fermions can not occupy the same QM state. This is an extremely important characteristic of matter—it makes our world possible. On the other hand, bosons can occupy the same QM state, which makes laser beams possible.

2

7) Electrons in orbit around an atomic nucleus occupy a specific energy shell. The shell level corresponds to the number of electron wavelengths in the standing wave pattern that defines the orbital. Further, a given energy shell is composed of suborbital states, designated s, p, d, f… which correspond to the different ways the electron wave functions can wrap themselves around the atomic nucleus. These suborbitals differ in their energy levels and shapes: s being spherical and lowest energy, p being “door knob” shaped and next higher in energy, and so on. The suborbitals are shown below* *http://en.wikipedia.org/wiki/Atomic_orbital

8) Beside occupying a specific atomic shell and suborbital, an electron has two possible spin states: spin-up and spin-down, which indicates the orientation of its intrinsic magnetic field. These three parameters defined the quantum mechanical state of an electron in orbit around a nucleus. Since the Pauli exclusion principle forbids any two electrons from being in the same QM state, only two electrons (one spin-up, the other, spin-down) can occupy a given shell and suborbital.

. . .. . .. . .        n=7

. . .. . .                             n=6

. . .                                                               n=5

                                                                                                             n=4

                                                              n=3

                            n=2

       n=1

f (l=3)d (l=2)p (l=1)s (l=0)

3

10) The types of suborbtials available for a given electron shell corresponds to the shell number. The lowest shell only has the s suborbital, the second shell has the s & p suborbital, the third shell has the s, p & d suborbitals, etc.

11) For a given atomic nucleus (that is, for a given number of protons in a nucleus) electrons fill the available QM states in order of ascending energy. This is the ground state configuration. The first two electrons fill the first level, s suborbital; the next two electrons fill the second level, s suborbital, the next 6 electrons go into the second level, p suborbitals, then the next two electrons go in the third shell, s suborbital. However, higher level suborbits in a lower shell may have greater energy than low level suborbits in a higher shell, so not all suborbits in a given shell get filled before the next shell is occupied. This makes it complicated to know how many electrons are in the outermost, valence shell of the atom. The order in which electrons fill atomic orbits is shown below* and specifies how elements are arranged in the Periodic Table (p462) *http://en.wikipedia.org/wiki/Atomic_orbital

12) When an atomic shell’s s & p suborbitals are filled, the valence electrons are in a low energy configuration. Thus, having 8 valence electrons (2 in s, 6 in p suborbitals) is an energetically favorable state for an atom. This ‘drive’ to possess 8 valence electrons governs how atoms combine to form molecules.

4

Conceptual Physics Study Notes & Questions: Week 14—Chemistry (Chap. 23)

1) The notes for Chapters 22 and 21 explain much of what is presented in this chapter. Please review that material before proceeding.

2) Atoms combine with other atoms to form molecules, driven by the low energy configuration that 8 valence electrons provide each atom. Thus, chemistry is governed by the fact atoms ‘want’ to have 8 electrons in their valence shells. Elements have different electronegativities—the strength of their pull on valence electrons.

3) There are three types of molecular (chemical) bonds a) Ionic—electrons are ripped from one atom by another (p494)b) Covalent—valence electrons are shared (p497)c) Metallic—valence electrons are donated to an electron sea, shared

by all neighboring metal atoms (p496).

4) As reflected in the Periodic Table of Elements (p493), elements have different numbers of electrons in their valence shells. This governs which atoms combine together and how. Whether ionic, covalent or metallic bonds form depend on who’s mixed with what, and at what pressure, density & temperature. Generally…

a) Atoms with 1 or 2 valence electrons will tend to give up their electrons (if they can) to uncover the 8 electrons in the next lower shell. ( ionic bond)

b) Atoms with 6 or 7 valence electrons will tend to grab electrons from their neighbors (if they can) in order to obtain a full complement of 8. ( ionic bond)

c) Atoms with 2—6 valence electrons will tend to share their electrons with other atoms to obtain 8 via the sharing arrangement. ( covalent bond)

d) Atoms with 1 or 2 valence electrons that are surrounded with similar atoms will tend to release their valence electrons into a shared ‘cloud’ shared by their neighbors. ( metallic bond)

e) Atoms with 8 valence electrons will not combine with anything; these comprise the noble gases.

5) Covalent molecules composed of atoms with different electronegativities can exhibit polarization—some parts of the molecule will tend to be more positively or negatively charged, even though the overall molecule is electrically neutral. These are called polar molecules (p500). They will be attracted toward ions or other polar molecules. This is the primary cause of intermolecular attractive forces.

5

6) The strongest type of intermolecular force is the hydrogen bond. A polar molecule like water leaves its constituent hydrogen atoms partially uncovered of its electron cloud. For hydrogen, there is no inner electron shell to cover its nucleus—the bare proton is nearly ‘hanging out in the wind.’ Any nearby polar molecule can thus get very close to the partially exposed proton, forming a strong intermolecular bond.

7) Other polar molecules are mutually attracted to one another. Water is a polar molecule. Some large extended molecules like proteins will contain some segments that are polar, and other segments that are non-polar. Polar molecules or their polar segments are often called hydrophilic, because they readily mix with water molecules. Molecules or molecular segments that are not polar will tend to avoid mixing with water. These are termed hydrophobic.

• Lipids are long molecules that are hydrophilic at one end and hydrophobic on the other end. In aqueous solution, lipids will crowd together with their hydrophobic heads close together, leaving their hydrophilic tails in contact with water. The result is an extended two dimensional membrane—called a lipid membrane—that can form an enclosing vesicle. These are the basis of biological cell membranes, which are a key necessity for life.

8) Van der Waals forces are weak intermolecular bonds that arise from momentarily polarization in non-polar molecules (p501), that arise from statistical fluctuations in the enclosing electron cloud of the molecular orbitals. This weak attraction is easily broken.

9) There are three aspects to the strength of materials: (i) compressive, (ii) tensile, and (iii) shear (p503). To what types of push/pull/twisting forces do these apply?

10) Composite material are composed of two or more different types of molecules (p505). These can be quite exotic and exhibit surprising quantum mechanical properties. This is a fertile area of current scientific research.

11) Chemical reactions are reconfigurations of atoms & molecules. Reactions occur if there is sufficient thermal energy to overcome existing stable molecular bonds—for example, wood and oxygen will not burn unless the temperature is sufficiently high.

a) Reactions that release more energy than that required to initiate them are called exothermic. (p507)

b) Reactions that require a net input of energy to form stable compounds are called endothermic.