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Chapter Six Electronic structure

Chapter Six Electronic structure. Jay Adams—Fall 2007. What this chapter is about:. The makeup of the atom. We know that the differences between atoms are based on the number of protons. We also know that the # of electrons is equal to the # of protons in a neutral atom.

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Chapter Six Electronic structure

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  1. Chapter SixElectronic structure Jay Adams—Fall 2007

  2. What this chapter is about: • The makeup of the atom. • We know that the differences between atoms are based on the number of protons. • We also know that the # of electrons is equal to the # of protons in a neutral atom. • This chapter considers how those electrons behave, and where in the atom they are located.

  3. What you’ll learn:

  4. Background-- • Light exists as a wave: • Wavelength (λ)is defined as the distance between two peaks or troughs. • Frequency (v) is defined as the number of cycles of a wave in a given time (usually cycles / second, or Hertz.) • The speed of a wave can be found using these two quantities. (λ) (v) = c

  5. More background: • The speed of light (c) is found to be 2.998 x 108 m/s. • The light that you see is only a part of the entire spectrum of light. You can see charcoal glow red, but you can’t see the heat that comes off of it.

  6. The Light Spectrum

  7. What we now know • We used to think that light was a continuous wave. • Max Planck proved (you don’t want to know how) that light was actually a wave of photons. • These photons are “spit out” at varying energy levels [explain] . The energy of a photon can be determined by the equation E = (h)(v) or E = (h)(c/λ) • H is known as Planck’s constant, and it is 6.626 x 10-34 J * s • The Joule is the standard unit of energy we will use in this chapter • Notice that energy is inversely related to wavelength.

  8. Atomic spectra • You’ve seen a prism. It was Newton who confirmed that light emanating from the sun could be broken down into its components. • When broken down, it gives a continuous spectrum (see previous slide or page 133) • Light given off by gaseous elements is different. Rather than being continuous, it is given off in discrete lines.

  9. What the????

  10. What do line spectra mean? • First of all, they mean that the elemental gases only give off light at specific, rather than continuous, wavelengths. • This means further that they only give off light at specific energies (Planck’s equation) • Since we know that photons are “spit out” only at specific energies, we say that the energy levels in an atom are quantized

  11. Quantized? • Imagine a 10 ft tall ramp: • Now imagine a set of 10 one-foot stairs: • On the ramp, you can slide from one level to the next. You can get to 6.5 feet. But on the stairs, you have to pick. You can be at five feet, or six feet. This is because the stairs are quantized—each stair is associated with a specific quantity of height.

  12. The quantum model • The problem with Bohr’s model has to do with uncertainty. • Heisenberg states that we can know either a particle’s energy or position, but we cannot know both simultaneously. • Think: if we stop a particle to look at it, we don’t know how it got there. But if we watch a particle move, we can’t measure it.

  13. Uncertainty in real life: • Driving directions to Wal-Mart • 20.2 mi – about 27 mins • Driving directions to Disney World • 534 mi – about 9 hours 13 mins • Consider the time: what could account for the difference? • What about the distance? Does that change?

  14. continued • Since we don’t know exactly how the electrons move, we can’t state with certainty that they move along a track. • We know say that electrons exist in orbitals, 4-d areas around the nucleus where we are most likely to find them. The four dimensions are x, y, z, and the probability that an electron inhabits that region.

  15. So where are the electrons? • Energy levels>sublevels (s,p,d,f)>orbitals (x,y,z) > electrons • There are never more than 2 electrons in any orbital. • S orbitals, the simplest, are spherical. Therefore, any s level can only hold 2 electrons • See page 80

  16. P orbitals • The P sublevel contains three separate dumbbell-shaped orbitals: x, y, and z • Because it contains three orbitals, the entire sublevel can hold a total of 6 electrons The sum of these orbitals can be seen on page 80 of your book

  17. Sublevel d • D sublevels contain five separate clover-leaf orbitals: • Because of five orbitals, the d level can hold 10 electrons. • The sum of these orbitals can be seen on page 80 of your book

  18. Sublevel f • The f sublevel is basically too complex to depict pictorially. • It contains seven orbitals, and, therefore, how many electrons? • 14…good if you got it. • Remember…these levels and orbitals are NOT something we’ve observed. They are models we’ve guessed at to explain behavior.

  19. Capacities of energy levels

  20. Energies of sublevels • In general…the farther an orbital or level is from the nucleus, the higher its energy. • The 1s sublevel has the least energy. • Everything makes sense until you get to 3p…for some reason, you don’t go to 3d next…4s comes first.

  21. Sublevel ordering If you’ll draw them out like this, you’ll never ever get them wrong. It’s impossible, unless you do something boneheaded 7s 7p 6s 6p 6d 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

  22. How to build an atom • The Aufbau principle states that the electron arrangement of an atom is determined by adding electrons to a smaller atom • Therefore, if you know how many electrons there are in a mystery atom, you can draw the electron configuration perfectly …and if you know the charge, you can identify the atom.

  23. Hydrogen and Carbon • Pauli’s exclusion principle states that each orbital can only contain two electronw with opposing “spins” • Hydrogen has 1 electron. Therefore, hydrogen’s electron configuration is simple: 1s1 • Carbon, on the other hand, has 6 electrons. Using my diagonal rule, though, it’s easy to determine its electron configuration: 1s22s22p2

  24. Filling multiple orbitals • The p-block has three orbitals: x, y, and z. How do we know which orbitals get which electrons? • Hund’s rule states that, as orbitals fill, each orbital gets one electron before any get two. This is because electrons repel each other and energy must be held to a minimum. • An atom with four valence electrons in 2p-orbitals looks like this: 2px22py12pz1 , not this: 2px22py22pz0

  25. Do you get it? • Give the electron configuration of manganese (25 electrons). • 1s22s22p63s23p64s23d5

  26. Quantum numbers: electron description • Electron configuration is a description of an atom. Each type of atom has a unique electron configuration. • Quantum numbers work the same way for electrons. They are a means to quantify and describe particular electrons. • N is the first quantum number. It describes the principal energy level of the electron in question. As far as you’re concerned, it can go from 1-7.

  27. More quantum numbers • L, the second number, identifies the type of sublevel in which an electron exists—0(s), 1(p), 2 (d), or 3 (f) • M, the third number, identifies the specific orbital in which an electron exists. It can range from –m to +m. • Ms, the fourth number, identifies the “spin” of the electron: +1/2 or -1/2 • Remember Pauli: no two electrons can have the same four quantum numbers in an atom.

  28. Identifying atoms • Atomic number=number of protons in an atom. This is the principal way we identify atoms, because it never changes. • Mass #=number of protons + number of neutrons. The number of neutrons is variable for any atom. More neutrons = more mass. • Atoms with numbers of neutrons differing from the most common are called isotopes

  29. Isotopes and averages • Example: 92.58 % of naturally occurring lithium is Li-7. • However, that means there’s 7.42 % out there with some other mass (Li-6). • How do we know how much lithium weighs? • We use a weighted average: (.0742)(6.015)+(.9258)(7.016)=6.941 amu • This is similar to grading your work.

  30. Valence electrons • Valence is a special, fancy, scientific name that means “outside” • Valence electrons are important because, as the electrons feeling the least pull from the positively charged nucleus, they are the most likely to be involved in bonding with other molecules. • How many valence electrons? • Hydrogen • Carbon • Oxygen

  31. Electron Dot Symbols • Also called Lewis structures • A way to visualize an atom’s valence electrons • The first two electrons go on as a pair; any remaining go one at a time: Mg S Ar

  32. Ions: charged atoms • Anion= A Negative ION • Cation= (I know it’s cheesy) C A + ION • The important thing: adding electrons makes an atom negative; taking them away makes an atom positive. • How many electrons would C+2 have? • What about O-3

  33. HOMEWORK • Terms: atom, model, electron-neutron, quantized, orbital, sublevel, electron configuration, mass number, atomic number, isotope, valence electron, ion, cation, anion • Questions: 2, 5-10, 12-18, 20-24

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