Chapter 6 - Atomic Structure
- The Wave Nature of Light
- Electromagnetic radiation
- Wavelength () and Frequency ()
- * = c
- c = 3.00 x 108 m/s
- speed of light
- visible light
- 400 - 700 nm
- chemical reaction in eyes
Problems 1-7 odd
- Quantum Effects and Photons
- Planck - minimum energy packet
- related to frequency
- E = h
- h = 6.626 x 10-34 J s (Planck's constant)
- E= energy of photon
- photon is quanta of energy discussed by Planck.
- baritone - quantized notes vs trombone - continuous notes.
- Photoelectric effect
- only light above a certain frequency causes electrons to be ejected from
substance.
- explain, using photon theory.
- IF (photon energy > energy holding electron) electron is ejected.
- shorter wavelength - more energy
- X-rays, gamma rays cause cell disruption (very high energy)
- radio waves - no effect.
- Wave - Particle duality.
Problems 9-19 odd
- Bohr's Model of the Hydrogen Atom
- Continuous spectra
- black body radiation
- color indicative of temperature
- Line spectra
- energy relates to electron transfer within an atom
- most notably Ne, Hg, and H
- Balmer series nu ~=~ C``(``{{1} over {2 ^2}} ~-~ {{1} over {n ^2}}``)
- C = 3.29 x 1015 s-1
- n = integer > 2
- Bohr - energy for electron movement between orbits
- orbit energies quantized
- E ~=~ R SUB H ``(``{{1} over {{n SUB 1} ^2}} ~-~ {{1} over {{n SUB 2} ^2}}``)
- n2 > n1
- RH = 2.18 x 10-18 J
- C = RH / h
- Bohr = Balmer when final orbit = 2
- empirical vs theoretical equations match!
Problems 21-27 odd
- The Dual Nature of the Electron
- matter also exhibits wave/particle duality
- effect minimal except for the electron
- Uncertainty Principle - Heisenberg
- Quantum Mechanics
- Since cannot know both position and momentum of electron, describe as 'probability
function'.
- from wave mechanics - called 'wave function' (Schrodinger)
- calculations match those for hydrogen and orbits, but also explain line spectra
from other elements that cannot be explained by orbits.
- wave functions called orbitals
- 4 quantum numbers
- Principal (n) - integral values from 1 - [shell]
- Azimuthal (l) - integral values from 0 - (n-1) [subshell]
- also represented by letters s p d f for l = 0 1 2 3 respectively.
- Magnetic (ml) - integral values in the range l [orbital]
- Spin (ms) -
- Set of 4 qn is unique for each electron in an atom
- Representation of Orbitals
- s - sphere (1 lobe)
- p (3) - dumbbell (2 lobes)
- d (5) - 4 with 4 lobes; 1 with 2 lobes and 'donut'
- f (7) - 6 with 8 lobes; 2 with 2 lobes and 'double donut'
- for multiple orbitals of the same type (p, d, f) the orbitals are degenerate
Problems 33-45 odd
- Orbitals in Many-Electron Atoms
- Aufbau Principle - occupy lowest energy orbital available
- for subshells - the higher value of the shell the higher the energy
- for shells - the higher the value of the subshell the higher the energy
- more complex shape - higher energy.
- Pauli Exclusion Principle - no two electrons of the same spin can occupy the same
orbital.
Problems 47-55 odd
- Electron Configuration
- Representation with 1s2 2s2 2p6 ...
- use diagram to predict less obvious choices.
- Using arrows to represent electrons.
- Hunds rule - stability with filled or completely filled orbitals with electrons of the
same spin.
- Outer electrons - valence electrons
- Inner electrons - core electrons.
- Electron Configuration and the Periodic Table
- Describe electron addition vs type of element (rep, transition, inner tx)
- Noble gas configuration.
- Explain anomalies in electron configuration.
Problems 57-61 odd
Learning Goals
1. Describe the wave properties and characteristic speed of propagation of radiant energy
(electromagnetic radiation.)
2. Use the relationship = c which relates the wavelength and the frequency of radiant
energy to it's speed.
3. Explain the essential feature of Planck's quantum theory, namely, the smallest increment,
or quantum, of radiant energy of frequency that can be emitted or absorbed is h, where h
is Planck's constant.
4. Explain how Einstein accounted for the photoelectric effect by considering the radiant
energy to be a stream of particle-like photons striking a metal surface.
5. Explain the origin of the expression line spectrum.
6. List the assumptions made by Bohr in his model of the hydrogen atom.
7. Explain the concept of an allowed energy state and how this concept is related to quantum
theory.
8. Calculate the energy difference between any two allowed energy states in hydrogen.
9. Describe the uncertainty principle, and explain the limitations it places on our ability to
define simultaneously the momentum and location of a subatomic particle, particularly an
electron.
10. Explain the concept of orbital, electron density, and probability as used in the quantum
mechanical atom.
11. Describe the four quantum numbers used to define electrons in an atom, and list the
limitations placed on each of them.
12. Describe the shapes of the s, p, and d orbitals.
13. Explain why electrons with the same value for n but different values for l have different
energies.
14. Write the electron configuration for any element.
15. State the Pauli exclusion principle and Hund's rule, and describe how they are used in
writing the electronic structure of the elements.
16. Describe what we mean by the s, p, d, and f blocks of elements.
17. Write the valence electron configuration for any element once you know it's position on the
periodic table.
18. Write the orbital diagram representations for electron configuration of atoms.
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