2. Rutherford Model While the Rutherford Model was an improvement over previous models, it still did not explain how the electrons were distributed around the nucleus. What prevented the electrons from being pulled into the nucleus?
3. Bohr Model Niels Bohr proposed that electrons can circle the nucleus only in allowed paths, or orbits. Electrons in these orbits give the atoms a fixed energy. Electrons are in the lowest energy state closest to the nucleus. Orbits are separated by areas where electrons can’t exist. Electron energy is higher when the electron is in orbits farther from the nucleus.
4. The Quantum Mechanical Model Energy is quantized. It comes in chunks. Quanta – the amount of energy needed to move from one energy level to another. Quantum leap in energy Schrödinger derives an equation that described the energy and position of the electrons in an atom.
5. A mathematical solution It is not like anything you can see. It does have energy levels for electrons Orbits are not circular It can only tell us the probability of finding an electron a certain distance from the nucleus. The electron is found inside a blurry “electron cloud.”
6. Atomic Orbitals Principal Quantum Number (n) = the energy level of the electron Within each energy level the complex math of Schrödinger’s equation describes several shapes. These are called atomic orbitals. Regions where there is a high probability of finding an electron.
7. S Orbitals 1 s orbital for every energy level Spherical shaped Each s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc. orbitals S Orbital
8. P Orbitals Start at the second energy level 3 different directions 3 different shapes (dumbell) Each can hold 2 electrons P Orbital
9. D Orbitals Start with the third energy level 5 different shapes Each can hold 2 electrons D Orbital
10. F Orbitals Start at the 4th energy level Have seven different shapes 2 electrons per shape F Orbital
12. By Energy Level 1st Energy Level Only s orbital Only 2 electrons 1s2 2nd Energy Level s & p orbitals 2 in s, 6 in p 2s22p6 8 total electrons
13. Filling Order Lowest energy fills first The energy levels overlap The orbitals do not fill up order of energy level Counting system Each “___” is an orbital shape Room for 2 electrons
15. Electron Configurations The way electrons are arranged in atoms Aufbau principle – electrons enter the lowest energy level first This causes difficulties because of the overall of orbitals of different energies Paulie Exclusion Principle – at most 2 electrons per orbital – different spins Hund’s Rule – when electrons occupy orbital of equal energy, they don’t pair up until they have to
29. Rewrite when done.. 1s22s22p63s23p64s23d104p65s24d105p66s2 4f145d106p67s25f146d107p66f147d107f14 Group the energy levels together 1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p66d106f147s27p67d107f14
31. Orbitals fill in order… Lowest energy to higher energy Adding electrons can change the energy of the orbital Filled and half-filled orbitals have a lower energy Makes them more stable Changes the filling order of d orbitals
32. Practice Vanadium 1s22s22p63s23p63d34s2 Titanium 1s22s22p63s23p63d24s2 Chromium (expected)1s22s22p63s23p63d44s2 (correct) 1s22s22p63s23p63d54s1 Why? It more stable – gives us two half-filled orbitals!
35. Table 3(p. 116) Notice that as you move across the periodic table, the elements increase in the number of electrons. Notice the electron-configuration notation Aufbau Principle
36. Noble Gas Notation The elements all the way to the right of the periodic table make up Group 18. Their sublevels are completely filled, and they are said to have an octet of electrons. These elements (helium, neon, argon, krypton, xenon, and radon) are called Noble Gases. We can use their electron configurations to simplify the electron configurations we write for other elements!
37. Example Neon Ne = 1s22s22p6 Sodium Na = [Ne]3s1 Tables 4 - 6 (pp. 117-120) Noble-Gas Configuration Refers to an outer main energy level occupied, in most cases, by eight electrons.
