AP Chem- Atomic Structure and Periodicity, Bohr model?

Question

What are Neils Bohr's theoretical connections between the energy level transitions of an electron and the specific color of a photon in the hydrogen line spectrum?

Please explain the advantages and disadvantages in the modern quantum mechanical model over the Bohr model. Use the following concepts in your answer: Heisenberg Uncertainty Principle, quantum numbers, orbital, sublevel, Hund's Rule, Pauli Exclusion Principle.

There is a greater variation between the properties (both chemical and physical) of the first and second of a group or family in the periodic table than between the properties of the second and third members of the group. Consider as examples either the group containing nitrogen or the one containing oxygen. Select three properties and discuss the variation of these properties to illustrate the generalization expressed in the first sentence of the question.

Answer 1

I will give concepts in outline form. Anything more would be doing your homework (or taking an exam) for you.

Dalton's model had atoms like little balls. Every element had a different kind of ball (atom). Neils Bohr's experiments with cathode rays showed that atoms had parts. He used the planetary system as a model. Heavy nucleus in the center and light electrons in orbits around this nucleus. We still call the energy levels orbits (although some use the term shell instead).

For Hydrogen, a single proton is the nucleus, orbited by a single electron. If the Hydrogen atom is not interacting with the environment, the electron is in the lowest possible energy level, called the ground state.

The Hydrogen atom can only absorb specific amounts of energy. This energy excites the electron into a higher or more energetic level. This is called the excited state. The excited electron can only emit specific amounts of energy to return to the ground state.

The energy of a photon is related to the wavelength of light:
E = hc / lambda
where lambda is the wavelength.

Different wavelengths of light appear as different colors. Thus specific amounts of energy means specific wavelengths of light which means specific colors in the spectrum.

Your second question requires an entire text book chapter to answer. The Bohr model is too simplistic (unrealistic) to be able to predict the behavior of subatomic pieces (they are not classical particles). Electrons are not like little round balls. They are clouds of location probability (the Heisenberg Uncertainty fits in here). The shapes of the clouds are determined mathematically by considering the quantum numbers associated with the specific electrons involved.

The advantages include a better predictive model and a clean transition into chemical bonding based upon electron behavior. The disadvantages include mathematics beyond most humans, even with assumptions, The model also leaves common sense behind (how can a p-orbital electron be on one side of the nucleus or the other without ever passing through the nucleus?).

The principal quantum number determines the shell (energy level) and limits the number of subshells (orbitals) involved. How these subshells are populated is called the Aufbau rules which includes Hund's Rule and the Pauli Exclusion Principle (see the second site below).

The last part of your question is beyond the scope of this format. What you need to do is work on your answer and then post more specific questions on the details. I can tell you that part of the answer is determined by relative size and energy of the shells involved.

In your examples, Nitrogen and Oxygen are diatomic gases, while Phosphorus & Arsenic and Sulphur & Selenium are all solids.

Answer 2

Actually, there are no theoretical connections. After formulating the atom as the electron attracted by electrostatic force to the proton, balanced by the centrifugal repulsion of the electron from the proton, Bohr could only postulate arbitrary fixed enegy levels for the electrons to reside. Then, when electrons descended from upper to lower states, they emitted characteristic frequencies of radiation.

Actually, the Bohr model of hydrogen led to the model of all other atoms, which led Gilbert N. Lewis to formulate his theory of the octet rule and the two-electron bond. The Lewis theory has survived useful to this day. When you do your Ph.D. in organic chemistry, you will find yourself doing research with Lewis theory on Bohr atoms on the back of an envelope while sitting in an airport, because you don't have the computing power to use quantum mechanics.

The third and fourth members of any series have the 3d10 and 4d10 electrons to use. They can invoke amazing hybridized orbitals. Also, your teacher is wrong to suppose that boron is "pretty much like" aluminum metal. Or that carbon is like silicon. Oxygen gas and oligomeric sulfur? Inert nitrogen gas and flammable phosphorus? C'mon teach: Gimme a break!