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  1. In This Lesson: Unit 2 Electrons, Orbitals, and Atomic Model History (Lesson 1 of 4) • Stuff You Need: • Periodic Table • Paper Towel Today is Wednesday,February 26th, 2014 Pre-Class: In your notebooks, draw a picture of electrons moving around the atom’s nucleus. Include arrows to show direction. You’re going to put this on a whiteboard shortly, so grab a SMALL paper towel.

  2. Today’s Agenda • A little history review… • Electron Configuration • Also known as “Where the electrons at?” • Electron Orbitals • Heisenberg Uncertainty Principle • Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule • Where is this in my book? • P. 127 and following…

  3. Guiding Video • TED: George Zaidan and Charles Morton – The Uncertain Location of Electrons

  4. In the beginning… • There was Democritus, a Greek professor (460 BC - 370 BC). • He came up with the term “atom” to describe the tiny particles he suggested. • Then there was John Dalton (1803). • He studied combinations of elements in chemical reactions. • His atomic model was just a solid ball.

  5. Discovery of the Electron • In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. • Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

  6. Conclusions from Studying Electrons • Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. • Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons. • Electrons have so little mass that atoms must contain other particles that account for most of the mass.

  7. In shorter terms… • Electrons are important because: • They create ions. • They lead to bonding. • They determine how atoms behave.

  8. Thomson’s Atom (1897) • Called the Plum Pudding Model, as Thomson thought that electrons were like plums sitting in a positive pudding. JJ Thomson

  9. Rutherford and the Nucleus • Ernest Rutherford fired α particles (helium nuclei) at an extremely thin sheet of gold foil. • He recorded where the particles “landed” after striking (or passing through) the gold. Ernest Rutherford “Like Howitzer shells bouncing off of tissue paper.”

  10. Rutherford’s Findings • Because most particles passed through and only a very few were significantly deflected, Rutherford concluded that the nucleus: • Is small • Is dense • Is positively charged

  11. Rutherford’s Atom (1913) • After the Rutherford experiment, the atom model looked like this: • Looked like the Infinity Ward logo, but it’s wrong. http://www.epa.gov/radiation/images/ruthbohr.jpg http://www.thatvideogameblog.com/wp-content/uploads/2010/08/infinity-ward-logo.jpg

  12. Eugen Goldstein and the Proton • Eugen Goldstein is sometimes credited with the discovery of the proton. • Other times it goes to Wilhelm Wien who performed other critical measurements of the proton using an anode ray (somewhat like Thomson’s cathode ray). http://www.pkc.ac.th/kobori/Assets/ChemistryMahidol1/www.il.mahidol.ac.th/course/ap_chemistry/atomic_structure/picture/bild_goldstein.jpg

  13. Jimmy Neutron and the Rutherford Atom? • Even Jimmy Neutron has an image of the Rutherford Model on his shirt! • Not so “boy genius” after all…

  14. Bohr’s Atom (1913) • Bohr thought of electrons moving around the nucleus like planets around the Sun. • His was a flat model of the atom. • In reality, the electrons actually move around the nucleus like bees around a hive. Niels Bohr

  15. The Bohr Model • Niels Bohr, among other things, proposed the Bohr Model. • Unlike Rutherford’s atom, which had electrons all at approximately the same distance from the nucleus, Bohr’s model showed them orbiting in a flat space but at different, fixed distances: http://www.thephysicsmill.com/blog/wp-content/uploads/bohr_model_no_emission.png

  16. Schrödinger’s Atom (1926) Louis de Broglie • In 1923, Louis de Broglie discovered that particles as small as electrons have some wave-like properties (as opposed to strictly particle-like). • More on this in our next lesson. • In 1926, Erwin Schrödinger develops equations that lead to the electron cloud model of the atom. • Electrons around found in a three-dimensional space around the nucleus and are more likely to be found closer-in. • Combined, these two discoveries do away with the Bohr model but require a more complex model of the atom. Erwin Schrödinger http://upload.wikimedia.org/wikipedia/commons/thumb/2/26/Erwin_Schrödinger.jpg/220px-Erwin_Schrödinger.jpg http://1.bp.blogspot.com/_GVA115I1I8Y/TT6_AHLks3I/AAAAAAAABWo/uvD4LGMKRgY/s1600/Broglie_Big.jpg

  17. Chadwick and the Neutron • Chadwick discovered the neutron in 1932 and won the Nobel Prize three years later for it. http://www.dnahelix.com/jimmy/jnmov_jn_ext_shrinkray.jpg http://www.nobelprize.org/nobel_prizes/physics/laureates/1935/chadwick.jpg

  18. Modern Atomic Theory • All matter is composed of atoms. • Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changescanoccur in nuclear reactions! • Atoms of an element have a characteristic average mass which is unique to that element. • Atoms of any one element differ in properties from atoms of another element.

  19. The Quantum Mechanical Model • The currently-accepted model is the Quantum Mechanical Model of the atom. • In it, mathematical models determine the most likely positions of electrons around the nucleus. • Sound complicated? It is. • Instead of exploring the laws, we’re going to look at some of the “results” of them. • But first, an actual look at atoms on camera. • NOVA video.

  20. Heisenberg Uncertainty Principle • Werner Heisenbergdiscovered that you can find out where an electron is, but not where it’s going. • Alternatively, you can find out where it’s going but not where it is. • Not both. “One cannot simultaneously determine both the position and momentum of an electron.” http://www.wired.com/images_blogs/underwire/2012/09/heisenberg_660.jpg

  21. Heisenberg Uncertainty Principle • To be able to see things, light must strike an object and then bounce off of it, returning to your eye. • For objects like, say, bowling balls, light strikes it and the bowling ball just sits there. • For electrons, however, they have so little mass that when light strikes them, they move in a different direction. http://cdn2-b.examiner.com/sites/default/files/styles/image_full_width/hash/0d/2e/0d2e398879c6b94255370961648165a2.jpg

  22. Guiding Example • Now, before we dive face-first into electron orbitals, we’re going to use a “guiding example” from something not-so-scientific to understand the concepts behind them. • The Hog Hotel! • Remember, as we explore this analogy, the goal of this entire lesson is to learn how electrons configure themselves around the nucleus. • It’s a big game of hide and seek with electrons!

  23. The Hog Hotel Analogy • Imagine you’re the manager of a towering hotel (for pigs) and you have a list of pigs that want to stay there. • Here are the rules you need to follow: • Rooms must be filled from the ground up. • Only singles first. No pig gets a roommate until all rooms on one floor are filled. • If two pigs are staying in the same room, they will face opposite directions. Weird.

  24. The Hog Hotel Analogy • On your Hog Hotel worksheets, try the first page and #2 on the back of the first page. • Then we’ll go over it. • Then we’ll do the rest of the back page.

  25. Electron Energy Levels (Shells) • Rising up from the lobby of the hotel are the various floors hogs might occupy. • Moving away from the nucleus are the various energy levels electrons might occupy. • These energy levels are symbolized by n. • Energy Level 1 n=1 • Energy Level 2 n=2

  26. n • n is the Principal Quantum Number. • To determine how many electrons fit into a given energy level, use this formula: Electrons = 2n2 • Energy Level 1 n=1 • Energy Level 2 n=2

  27. Aufbau Principle • In German, aufbau means “building up.” • The Aufbau Principle states that electrons, when not excited, will fill energy levels starting at the lowest energy. • In the Hog Hotel, this was the rule that the hogs are lazy and prefer rooms on the lowest floors possible.

  28. Orbital Shapes • Imagine that each room in the hotel, even on the same floor, has a different shape. • In the atom, on the energy level are sublevels consisting of orbitals where there is a 90% probability of finding an electron. • An orbital is like a specific room (indicated sometimes by a direction). • Orbitals can hold up to 2 electrons. • A sublevel is like a group of rooms or a suite (indicated by a letter – also called subshells). • Sublevels can hold 1, 3, 5, or 7 orbitals.

  29. Orbital Hotel Rooms? • For the next few slides, I’m going to show you pictures of orbitals. • Think of these as rooms in a weird atomic hotel. • Some are basic rooms, holding only two electrons. • Some are like suites, with individual rooms comprising a larger room. • They don’t all appear on every floor, however.

  30. s Sublevel e- e- Orbital

  31. s Sublevels • Shape: Sphere • Appears: n=1 and above. • # of Orbitals: 1 • Capacity: 2 e-

  32. p Sublevel e- e- e- e- e- e- Orbital

  33. p Sublevels • Shape: Dumbbell (3) • Appears: n=2 and above. • # of Orbitals:3 • Capacity: 6 e-

  34. d Sublevel e- e- e- e- e- e- e- e- e- e- Orbital

  35. d Sublevels • Shape: Double Dumbbells (4) and Dumbbell Doughnut • Appears: n=3 and above. • # of Orbitals: 5 • Capacity: 10 e-

  36. f Sublevel e- e- e- e- e- e- e- e- e- e- e- e- e- e- Orbital

  37. f Sublevels • Shape: Flowers…and stuff. • Appears: n=4 and above. • # of Orbitals: 7 • Capacity: 14 e-

  38. And the “hotel” as a whole? 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

  39. Quick Review • How many electrons can fit into that s sublevel? • 2 • Which energy level is further from the nucleus, n=2 or n=5? • 5 • How many electrons can fit at the 2nd energy level? (n=2) • 8 (remember 2n2 ?) • In which energy level does the f orbital start to appear? • n=4

  40. You Should Know… • You may be feeling a little overwhelmed. • If you understand this, you’re in good shape: • Around the atom are energy levels, like floors in a hotel room. The further out, the higher energy. • Each energy level has sublevels, like “types of rooms” in a hotel. • Each sublevel has one or more orbitals, which are like individual rooms. For example, s sublevels have one orbital, whereas p sublevels have three orbitals. • These orbitals each can hold two electrons and show the 90% likely location of those two electrons at any time.

  41. Summary Table

  42. Putting It All Together • Let’s try the third and fourth pages of the hog hotel worksheet. • It’s the same thing we’ve been doing, only using “up arrows” and “down arrows” instead of forward and backward letters.

  43. Orbital Notation • What you have just learned (the arrow way of writing electrons) is called orbital notation. • As it turns out, there’s a pattern to finding the orbitals in which the last electrons are placed. • Mendeleev was on to something! • Let’s do some color-coding so we can predict the ending orbital.

  44. First shade the blocks…

  45. Inner Transition Metals • Below the table are the inner transition metals (f block). • They look disconnected, but really they are “within” the transition elements (d block). • Expanded, the table would look like this.

  46. d and f Sublevels • Uh, wait a second… • It looks like according to the table we just shaded, d and f sublevels are going out of order. • In the n=6 row, it’s 5d and 4f. • What’s the deal? • d and f sublevels exist at lower energy levels than p sublevels (starting at n=4), so they’ll be filled first. • Stick with me here – I’ll teach you an easy way to remember that.

  47. Writing Configurations • Chemists need to be able to effectively record the electron configurations of various atoms. Consider Neon, the first element on the last page of the Hog Hotel. • Neon is in the second row (n=2), so there are electrons in n=1 and n=2. • 1 2 • There are electrons in sublevels 1s, 2s, and 2p. • 1s 2s 2p • Finally, there are two electrons in sublevel 1s, two in subshell 2s, and 6 in subshell 2p. • 1s2 2s2 2p6 (electron configuration) • ↑↓↑↓↑↓↑↓↑↓(orbital notation) 1s 2s 2p

  48. Two Ways to Figure This Out… • It can be hard to remember the order of the various quantum numbers and subshells. • You can figure out the electron configuration of an element two ways. • The easy way and the hard way. • Just kidding. They’re just different. • One way is the diagonal rule. • This: • The other way is hard to explain in writing, but I like it better.

  49. Directions for Using the Cheat Sheet • Target your element. • Starting with hydrogen, move left to right across the rows, moving down one each time you reach the end. • Every time you either A) reach the end of a row or B) change blocks, write down the “address” of the last element in that section. • Stop when you get to your element. • Check your work! You should be able to count the same number of electrons (more on that in a bit).