Chapter 5 Review CHM 151

Periodic Table
1. Created in 1869 by Mendeleev
2. Organized so that columns (groups) have similar properties
Electromagnetic (EM) Spectrum
1. Radiant energy (light, rays, waves) travel in waves
2. Parts of a wave
  1. wavelength (l) distance from peak to peak (units = m)
  2. frequency (n) - how many peaks pass per second (units =1/s = s^-1 = Hz = Hertz)
  3. energy (E) - higher frequency means more energy (units = Joules = J = kg m² / s²)
3. All the frequencies put together is the EM Spectrum (important)
4. Also see Fig 5.3
5. All of these waves are radiation. Radiation is mostly good - sun, TV, cell phones, microwaves, medical applications. Some is bad is exposed too much - X rays, gamma rays. Click here and go to Activities then EM spectrum activity.
6. Speed of light (c) = 3.00 x 10⁸ meters per second (m/s). Which travels faster? UV light or microwaves? Well guess what all EM radiation travels at the same speed - light speed.
  1. As l increases, E decreases and n decreases. Longer waves have less energy and lower frequency.
  2. c = ln (important) Look at Example 5.1. Try problem 5.1, 5.2, 5.3.
Atomic Spectra (no math)
1. The sun, old style bulbs give off white light - all visible light together. If put white light through a prism it splits into all the colors of the rainbow and is continuous.
2. If we heat up pure elements or add electricity to them they emit different colors. We use this principle in neon signs. Take glass tubes shaped however you want and fill with different gases. Attach to electricity and the "excited" gases emit specific colors of light. Many gases are used besides Neon. Fluorescent bulbs use mercury gas! Electrical energy is converted into light energy.
3. If you put the light from these heated or electrified elements through a prism you can see the exact wavelengths involved - just a few lines show up - not all of them like a rainbow since the light is not white. We call this line spectra. Examples:
  5. More are at the bottom of this review. Line spectra are unique for each element. Can be used to identify unknown gases. Note there may be some wavelengths in the IR or UV regions - we just can't see them.
  6. Examples of burning metals in flame tests are at the end of this review. Click here go to Activities then flame test movie.
  7. Later we'll learn WHY we get distinct or discrete lines for elements instead of a continuum.
Plank
1. Planck (~1900) discovered E = hn where h = 6.626 x 10^-34 Js (Joules x seconds) (important)
  1. often used to solve problems with c = ln
  2. What radio station (MHz) uses a 3.05 m carrier wave? What is the energy of this wave? Well plug in: 3.00 x 10⁸ m/s = (3.05m)n so solve for n = 9.84 x 10⁷s^-1 or Hz. Now change to MHz. 9.84 x 10⁷Hz ( 1 MHz / 10⁶ Hz) = 98.4 MHz. The E = (6.626 x 10^-34 Js)(9.84 x 10⁷ s^-1) = 6.52 x 10^-26 J.
2. We see that radiation (light, waves, rays) are wavelike. Well radiation is also particle like. Einstein (1905) finally explained the photoelectric effect (when you hit metals with radiation above a certain minimum energy it causes electrons to be ejected off the surface of the metal.) Einstein said this is because light is a beam composed of what he called photon particles. Particles have mass, does light? Well yes since black holes suck in light by gravity. Light and all the other EM radiation is both wave and particle like. (important)
3. EM radiation exists in very small discrete units called photons. Just like matter exists in very small units called atoms.
4. Quantized means exists in discrete units like steps, stairs, shelves, and now light and EM radiation. We say a photon is a quantum (discrete unit) of light energy. (important)
de Broglie
1. In 1924 de Broglie derived this equation l = h / mc. If something other than EM radiation is traveling then light speed just becomes velocity: l = h / mv. Because of h being in Joules ( J = kg m² / s²), the mass must be in what units? yep, kg
2. The mass of an electron is 9.11 x 10^-31 kg and the velocity is 2.2 x 10⁶ m/s. Check yourself that you get l = 3.3 x 10^-10 meters. You mean an electron has a wavelength? Yes, because an electron is both wave like and particle like.
3. I throw a 16 pound bowling ball down an alley at 7.55 m/s. What is the wavelength of the ball? Try it first. Answer below.
4. Look at example 5.5 in your text and try problem 5.9.
5. 16 lbs (1 kg / 2.20 lbs) = 7.3 kg bowling ball. l = (6.626 x 10^-34 kg m² s/ s²) / (7.3kg)(7.55m/s) = 1.2 x 10^-35 m Does this really mean anything? Not much for such a large object. It is more important with tiny objects.
Quantum Mechanics and the Heisenberg Uncertainty Principle
1. Heisenberg (1927) said it is impossible to know both where an electron is AND its velocity at the same time. Kind of like taking a picture of a moving car - you can't tell how fast it is going.
2. no math
Quantum Numbers
1. Solving quantum mechanics we get solutions that tell us where electrons are most likely to exist = areas of probable electron density = wave functions and we call these shapes of electron density orbitals. See electrons can't be just anywhere, they can exist only in certain areas or levels. Orbitals are quantized like steps. We call these areas shells.
2. Quantum Numbers (only cover first 2, not all 4) (important)
  1. Principle Quantum Number (n) = shell number. n = 1, 2, 3, 4, 5... and represents the distance from the nucleus.
  2. Second quantum number (l) = subshell number. Shells are divided into one or more subshells. l=0 for s subshell. l=1 for p subshell. l=2 for d subshell. l=3 for f subshell. Rule: n>l . So on the first shell where n=1, l must =0. On the second shell where n=2, l can be =0 or 1. On the third shell where n=3, l can be =0, 1 or 2. And so forth. See Figure 5.9
  4. Examples: Give the first two quantum numbers for the following subshells:
    1. 2nd shell, s subshell-------n=2 and l=0
    2. 7th shell f subshell--------n=7 and l=3
    3. 5th shell p subshell----------n=5 and l=1
    4. 4th shell d subshell----------n=4 and l=2
  5. What subshell is this?
    1. n=3 and l=0 -----------3s
    2. n=1 and l=1-----------no such thing as 1p
    3. n=4 and l=3-----------4f
    4. n=7 and l=2-----------7d
3. Subshells contain orbitals. Each orbital can contain a maximum of two electrons. (important)
  1. s subshell has one orbital so a max of 2 electrons
  2. p subshell has 3 orbitals so a max of 6 electrons
  3. d subshell has 5 orbitals so a max of 10 electrons
  4. f subshell has 7 orbitals so a max of 14 electrons
4. Questions go to 151 Help and click on quantum numbers (important)
Shapes of Orbitals
1. s subshell has 1 orbital (l=0) and it is spherical. The 1s is a sphere. The 2s is a sphere within a sphere separated by a node (area of 0% electron density). The 3s is three spheres separated by 2 nodes. And so forth. Fig 5.10 and 5.11
2. p subshell has 3 orbitals (l=1) and they are shaped like figure 8's or dumbbells. Each follows an axis in space - x, y and z. Fig 5.12
3. d subshell has 5 orbitals (l=2) and four of them are clover like with 4 lobes, the other is like a p orbital with a donut. Fig 5.13
4. f subshell has 7 orbitals (l=3) and most of them are eight lobed. Try problem 5.15.
5. Best pictures ever of orbitals
6. Back to line spectra introduced above. So why do the elements when heated up emit only certain frequencies of light giving us the distinct spectra instead of a continuous rainbow? View a video here about emission lamps like neon signs. Because atoms can only absorb wavelengths that make electrons jump from one shell to another shell. Electrons can't exist between shells so only a few energies are absorbed - the ones that exactly match the energy needed to go between shells. Figure 5.14 (important)
7. So an absorption spectra is all the colors of the rainbow except those that were absorbed - they are missing or black. An emission spectra (pictured earlier) is when the atoms are given energy by heat or electricity and then they emit the wavelengths of energy so just a few lines show up. They are opposites of each other.
Skip section 9
Electron Spin
1. Electrons spin clock or counterclockwise creating tiny magnetic fields.
2. Pauli exclusion principle (1925) states that the two electrons in an orbital must have opposite spins. (important)
Skip section 11
Electron Configurations (important)
1. The closer to the nucleus the lower in energy (more stable). s subshells are closer to the nucleus than p subshells so the s is more attracted to the protons in the nucleus and has lower energy. Similar for d and f. E of s < E of p < E of d < E of f.
2. Ready for filling electrons into their orbitals in the shells. Aufbau rules:
  1. Fill in lower energy subshells first (we don't fill in by shell necessarily)
  2. Put 2 electrons per orbital and they must have opposite spin.
  3. Put 1 electron per orbital in subshells before pairing them up = Hund's rule. For example you must put one electron in each of the 5 d orbitals before pairing them up. Remember the orbitals in a subshell are degenerate (equal in energy) This means the 5 orbitals in a d subshell are all equal in energy to each other.
3. The actual order we fill in is 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
4. How can you remember that? Well just follow the Periodic Table row by row. Figure 5.18 and 5.17
5. Examples: (superscript refers to how many electrons in that subshell, can't exceed the max)
  1. H - 1s¹
  2. F - 1s²2s²2p⁵
  3. P - 1s²2s²2p⁶3s²3p³ or [Ne]3s²3p³
  4. I^-1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁵ or [Kr]5s²4d¹⁰5p⁵
  5. You try U, Ag, O, Li, and Fe. Don't look yet!
  6. U - 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s²4f¹⁴5d¹⁰6p⁶7s²5f⁴ or [Rn]7s²5f⁴, Ag - 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹4d¹⁰or [Kr]5s¹4d¹⁰, O - 1s²2s²2p⁴ or [He]s²2p⁴, Li - 1s²2s¹, Fe - 1s²2s²2p⁶3s²3p⁶4s²3d⁶
  7. Practice Click here go to Activities then electron configuration activity and the movie.
  8. Another great movie
Electron Configuration Trends
1. All Alkali metals end with s¹
2. All Alkaline Earth metals end with s²
3. All Halogens end with s²p⁵
4. All Noble Gases end with s²p⁶
5. Valence electrons - those in the outermost shell. See Table 5.3. Example 5.10 and 5.11. Try Problems 5.17, 5.18, 5.19 (important)
Electron Configuration Anomalies
1. When we did Ag above did you think I made a typo? Well I did not! Some transition metal configurations are not as expected. This is because it is lowering in energy to have a half full or full d subshell. Atoms would like to have 5 or 10 electrons in the d subshell if possible. So they will steal one electron from the s to make this possible.
2. Expect Mo to be [Kr] 5s²4d⁴ but it is really [Kr] 5s¹4d⁵
3. So that is why Ag is not [Kr]5s²4d⁹but really [Kr]5s¹4d¹⁰
4. What is the electron configuration for Cu??? Expect [Ar]4s²3d⁹ but it is really [Ar]4s¹3d¹⁰
Atomic Radii
1. Size increases f and i in the periodic table. See Figures 5.19 and 5.1 (important)
2. It makes sense as you down a group the atoms get bigger - we started filling in another shell. But as you go across a row we are filling in the same shell and as more electrons are added to that shell they are more attracted to the increasing number of positive protons in the nucleus. So that attraction pulls them a bit closer and the atoms get a tiny bit smaller in general across a row. There is a big jump in size when we start filling a new shell - Alkali metals are the biggest atoms.
3. Z_eff is the symbol for effective nuclear charge which is the "feel" or attraction the electrons feel for the protons in the nucleus. The electrons in the 1st shell feel more pull than the electrons in the 2nd shell etc. This is because the inner electrons being negative actually repel the outer electrons since they are also negative. So we say the inner electrons "shield" the outer electrons from feeling the full attraction of the protons. Z_eff increases g and h in the periodic table. Try Prob 5.21.