CHM135H1 Lecture Notes - Lecture 4: Photon, Principal Quantum Number, Angular Momentum

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14 Sep 2018

School

UTSG

Department

Chemistry

Course

CHM135H1

Professor

Gilbert Walker, Eugenia Kumacheva, Kristine Quinlan

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CHM135H1 Lecture Notes - Lecture 3: Waves, Execution Unit

CHM135H1 Lecture Notes - Lecture 4: Photon, Principal Quantum Number, Angular Momentum

CHM135H1 Lecture Notes - Lecture 5: Pauli Exclusion Principle, Atom, Aufbau Principle

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Chm135 lecture 4: electronic structure of an atom 13th september 2018. A sample of gas is present inside the gas discharge tube. The gas is excited, and as it goes back down to a lower energy state, a frequency/ wavelength of light is emitted. This line spectra is unique for different atoms. This provides evidence that atoms have certain allowed energy levels. Certain allowed energy levels which have specific energies attached to them. M and n are integers of shell where n must be larger than m. This is empirical, it does not explain why it happened and fits only for hydrogen. Bohr predicted a solar system model of the atom, with the nucleus in the center and electrons orbiting around it. The model does not predict the precise location of the electron. The wavelength of an electron is on the order of its size and hence, is significant as compared to calculating the wavelength of a baseball.

Related Questions

The Bohr model of the hydrogen atom served as a major step toward describing the structure of atoms. It is essentially a classical model with one quantum modification, namely the orbital angular momentum of the electron is an integer multiple, n, of Planck's constant, h, divided by 2*pi. That is L=rmv=n(h/(2*pi)).

An electron in a hydrogen atom transitions from state n to state n-2. Find a general formula for the frequency of the emitted light for this transition. Does your formula predict the same energy regardless of the initial value of n? What does this mean regarding the

CHEM 1430 LIGHT LAB SPRING 2018 A Danish physicist, Neils Bohr, a post-doctoral student in J. J. Thomson's lab and friend of Rutherford, took up the problem of electron arrangement within the atom. Bohr's breakthrough was recognizing Balmer's fixed emission wavelengths, when taken with Planck's Relationship, led to the conclusion the shells, like the radiation emitted, must be quantized (have set, specific energles). Bohr proposed the electrons circulated about the madeus, much like planets cirde surãin seres of concentric shells of increasing size. Each shell stood at a fixed (set) distance from the nucleus with the shell nearest the nucleus taking the m and a shell further away taking the n in the B-R equation. The integral values of m and n taken with the discrete wavelength led Bohr to propose the shells were "quantized," meaning electrons could only exist at the surface of the shells. When an electron moved from a high shell, n, to a low shell, m, light of a foxed wavelength was emitted Bohr reasoned the shells were special due to this quantized nature. An electron in a shell occupied what Bohr described as a "steady state," meaning the electron could circulate about the nucleus without loss or gain of energy. This steady state proposition, which conflicted with classical physics, could not be explained by then available data but was required for Bohr's proposed structure to be stable. Later work revealed the wave-particle duality of the electron which lead to acceptance of the Steady State Proposition. The energy of an electron in any shell, represented by n, can be calculated using an equation derived from combining the Balmer-Rydberg and Planck equations. The energy of an electron in the no shell is given by é¦¯.h RL where hcR,-2.18x10-*) where Ru, is the Rydberg's Constant for the hydrogen atom, h is Planck's Constant and c is the speed of light NOTE: The energy is given a nexative value to refect the stabillting, attractive force between the negatively charged electron and the positively charged nucleus. As the shell number n increases, the attractive force decreases leading to a rising shell energies (E, increases, values less negative). ACTIVITY Calculate the energles for an electron in the first 7 shells of the hydrogen atom using the above relationship, recording the values in the center column of the Hydrogen the right-hand column in the table, convert the shell with the exponent of 10 Improper sclentific notation is typically used to plot data covering several orders of magnitude (powers of ten). Shell Energy Table. For energies to improper scientific notation Hydrogen Shell Energy Table Shell Value, n Calculated Shell Energy.ã Express Shell Energy in improper Scientific Notation, x 10-1 Page 9 of 12

Shown below is a qualitative diagram of the atomic orbital energies for an Na atom. The number of orbitals in each subshell is not shown.

(a) Are all of the subshells for n = 1, n = 2, and n = 3 shown? If not, what is missing?

(b) The 2s and 2p energy levels are shown as different. Which of the following is the best explanation for why this is the case? (i) The 2s and 2p energy levels have different energies in the hydrogen atom, so of course they will have different energies in the sodium atom, (ii) The
energy of the 2p orbital is higher than that of the 2s in all many-electron atoms. (iii) The 2s level in Na has electrons in it, whereas the 2p does not.
(c) Which of the energy levels holds the highest-energy electron in a sodium atom?
(d) A sodium vapor lamp (Figure 7.23) operates by using electricity to excite the highest-energy electron to the next highest-energy level. Light is produced when the excited electron drops back to the lower level. Which two energy levels are involved in this process for the Na atom?

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CHM135H1 Lecture Notes - Lecture 4: Photon, Principal Quantum Number, Angular Momentum

CHM135H1 Lecture Notes - Lecture 3: Waves, Execution Unit

CHM135H1 Lecture Notes - Lecture 4: Photon, Principal Quantum Number, Angular Momentum

CHM135H1 Lecture Notes - Lecture 5: Pauli Exclusion Principle, Atom, Aufbau Principle

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