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In This Lesson: Unit 2 Electrons, Orbitals, and Atomic Model History (Lesson 1 of 4) Today is Wednesday, September 27 th, 2017 Stuff You Need: Periodic Table Pre-Class: In your notebooks, draw a picture of electrons moving around the atom s nucleus. Include arrows to show direction.

Today s Agenda A little history review Electron Configuration Also known as Where the electrons at? Electron Orbitals and Quantum Numbers Heisenberg Uncertainty Principle Aufbau Principle Pauli Exclusion Principle Hund s Rule And coloring! Where is this in my book? P. 128 and following Oh, by the way, quantum numbers aren t in there. You heard me.

By the end of this lesson You should be able to describe the Quantum Mechanical Model of the atom. You should be able to indicate the arrangement and locations of electrons in multiple formats.

Guiding Video TED: George Zaidan and Charles Morton The Uncertain Location of Electrons

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.

Discovery of the Electron In 1897, JJ 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.

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.

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

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

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.

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

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

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/ato mic_structure/picture/bild_goldstein.jpg Eugen Goldstein

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

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

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

Erwin Schrödinger Louis de Broglie Schrödinger s Atom (1926) 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 threedimensional 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. 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

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 James Chadwick

Modern Atomic Theory All matter is composed of atoms. Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes can occur 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.

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.

Heisenberg Uncertainty Principle Werner Heisenberg discovered 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. http://www.wired.com/images_blogs/underwire/2012/09/heisenberg_660.jpg One cannot simultaneously determine both the position and momentum of an electron.

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

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!

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.

The Hog Hotel Analogy On your Hog Hotel worksheets, try #1-2. Then we ll go over it. Then we ll do the Classwork section.

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

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

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.

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.

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. I ll explain what I mean with a look back at two of my dorm rooms from college.

My Freshman Year of College e - e - I had the basic two bed/one roommate set up. Also, my roommate was awful but that s besides the point.

My Sophomore Year of College e - e - e - e - e - We had what our school called a suite, which was an arrangement of mini-rooms. Let s compare this to the atom and its rooms.

Orbital s Sublevel e - e -

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

Orbital p Sublevel e - e - e - e - e - e -

p Sublevels Shape: Dumbbell (3) Appears: n=2 and above. # of Orbitals: 3 (x, y, z) Capacity: 6 e -

Orbital d Sublevel e - e - e - e - e - e - e - e - e - e -

d Sublevels Shape: Double Dumbbells (4) and Dumbbell Doughnut Appears: n=3 through n=6. # of Orbitals: 5 Capacity: 10 e -

Orbital f Sublevel e - e - e - e - e - e - e - e - e - e - e - e - e - e -

f Sublevels Shape: Flowers and stuff. Appears: n=4 through n=5. # of Orbitals: 7 Capacity: 14 e -

And the hotel as a whole? 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

After f? Right now there are no elements in existence that have electrons at energy levels higher than 7. There are also no sublevels beyond f. However, if somehow we were to create an atom that had so many electrons we filled the f sublevel on the n=5 energy level, what would be next? g, then h and so on in alphabetical order.

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 farther 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.

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

Summary Table Energy Level (n) Sublevels Orbitals Per Sublevel Electrons Per Sublevel Electrons Per Energy Level (2n 2 ) 1 s 1 2 2 2 s p 1 3 2 6 8 3 s p d 1 3 5 2 6 10 18 4 s p d f 1 3 5 7 2 6 10 14 32 Floor Number Type of Rooms/Suites on Floor Rooms per Type of Room/Suite Capacity of Each Type of Room/Suite Capacity of Each Floor

Putting It All Together Let s try the back page of the Hog Hotel. It s the same thing we ve been doing, only using up arrows and down arrows instead of forward and backward letters.

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 electrons are placed. Mendeleev was on to something! Let s do some color-coding so we can predict what orbitals to write.

Electron Configuration Table s 1 s 2 s 2 1s 2s 3s s 4s 5s 6s 7s d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 3d d 4d 5d 6d p 1 p 2 p 3 p 4 p 5 p 6 2p 3p 4p 5p 6p p d 1 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 5d 4f f 6d 5f

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:

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 according to the Aufbau Principle. Stick with me here I ll teach you an easy way to remember that.

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. 1s 2 2s 2 2p 6 (electron configuration) (orbital notation) 1s 2s 2p

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.

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).

Electron Configuration for Ne s 1 s 2 s 2 1s 2s 3s s 4s 5s 6s 7s Ne: 1s 2 2s 2 2p 6 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 3d d 4d 5d 6d p 1 p 2 p 3 p 4 p 5 p 6 2p 3p 4p 5p 6p p d 1 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 5d 4f f 6d 5f

Electron Configuration Element Electron Configuration Hydrogen 1s 1 Helium 1s 2 Lithium 1s 2 2s 1 Beryllium 1s 2 2s 2 Boron 1s 2 2s 2 2p 1 Carbon 1s 2 2s 2 2p 2 Nitrogen 1s 2 2s 2 2p 3 Oxygen 1s 2 2s 2 2p 4 Fluorine 1s 2 2s 2 2p 5 Neon 1s 2 2s 2 2p 6

Use this for the next slide s questions s 1 s 2 s 2 1s 2s 3s s 4s 5s 6s 7s d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 3d d 4d 5d 6d p 1 p 2 p 3 p 4 p 5 p 6 2p 3p 4p 5p 6p p d 1 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 5d 4f f 6d 5f

Let s try a few practice elements Cobalt (Co): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Europium (Eu): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 6 Tungsten (W): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 4 Notice how we had to do a little rearranging at the end of the electron configuration for Tungsten.

Electron Configuration for W s 2 s 1 s 2 1s W: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 p 1 p 2 p 3 p 4 p 5 p 6 2s 5p 6 6s 2 5d 1 4f 14 5d 4 2p 3s s d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 3p 4s 3d 4p d 5s 4d 5p 6s 5d 6p p 7s 6d d 1 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 5d 4f f 6d 5f

On your worksheets Try the first page of the worksheet labeled, Electron Configurations Orbital Notation. We ll do the first one (Mg) together.

Things to Check Suppose you ve just written Magnesium s electron configuration: 1s 2 2s 2 2p 6 3s 2 To make sure you re right, check how many electrons magnesium has. 12 Do the exponents in the configuration add up to the same amount? 1s 2 2s 2 2p 6 3s 2 = 12

Ions Writing electron configurations of ions is easy: Step 1: Figure out how many electrons the ion has. [remember, protons electrons = charge] Step 2: Make that number the new atomic number. Step 3: Target an element with that new atomic number. Example: Oxygen with a charge of -2 (O 2- ) has two extra electrons. It s basically like diagramming a Neon atom. O 2- = 10 e - = 1s 2 2s 2 2p 6

Noble Gas Notation Try this: Write the electron configuration for Neon in your notebooks. 1s 2 2s 2 2p 6 Now try this: Write the electron configuration for Sodium underneath. 1s 2 2s 2 2p 6 3s 1 Notice anything?

Shorthand Notation Notice that the configurations build on one another. To save time, scientists use Shorthand Notation (or Noble Gas Notation) to condense the writing. Start from the last noble gas (Key: right-most column of elements) prior to your element and put it in brackets. Then, simply write the new configuration after it. Example: Sodium is [Ne] 3s 1 NOTE: Noble gases themselves can still be written in shorthand. Just use the previous noble gas and go from there. Helium does NOT have a shorthand configuration.

Shorthand Notation Element Electron Configuration Shorthand Notation Hydrogen 1s 1 -- Helium 1s 2 -- Lithium 1s 2 2s 1 [He]2s 1 Beryllium 1s 2 2s 2 [He]2s 2 Boron 1s 2 2s 2 2p1 [He]2s 2 2p 1 Carbon 1s 2 2s 2 2p 2 [He]2s 2 2p 2 Nitrogen 1s 2 2s 2 2p 3 [He]2s 2 2p 3 Oxygen 1s 2 2s 2 2p 4 [He]2s 2 2p 4 Fluorine 1s 2 2s 2 2p 5 [He]2s 2 2p 5 Neon 1s 2 2s 2 2p 6 [He]2s 2 2p 6

Exceptions Unfortunately, there are some exceptions to the electron configuration rule. Copper and Chromium are two good examples of this. Try diagramming them. Cr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 Contrary to what you may have come up with, in reality their configurations are: Cr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 The reason for this is that filled sublevels are the most stable. Half-filled sublevels are not as stable as filled, but more stable than others.

The Full Hotel 7s 7p 6s 6p 6d 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

Quantum Numbers Because the current model atom is three dimensional and based on mathematics, we use a series of descriptions (numbers) to denote electrons. This system allows us to combine Electron Configuration and Orbital Notation into one. The descriptions are called quantum numbers, and they include the principal quantum number (n). KEY: Think of these as mathematical code language for stuff like 3d 10. You already know this! http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html AND FOLLOWING

Quantum Numbers n = Principal Quantum Number Indicates energy. l = Angular Quantum Number Indicates sublevel: 0 = s 1 = p 2 = d 3 = f m l = Magnetic Quantum Number Indicates orbital. m s = Spin Quantum Number Indicates particular electron by its spin (more to come).

Quantum Number Rules n is from 1-7 (you knew that already). l is from 0 to n-1. This should make sense to you because: On n=1, only s (0) sublevels appear. On n=4, s (0), p (1), d (2), and f (3) sublevels appear. m l is from l to +l. Each m l value represents a different orbital. When l = 1, we re talking about a p sublevel. In that case, m l can be either -1, 0, or 1, each representing one of the three dumbbells in space. m s (spin) is either -½ or +½. In short, one direction or another. This indicates a single electron.

n = 2 l = 1 m l = -1 m s = ½ z Breaking Down The Code So you could be talking about 2s and its single orbital y z or you could be talking about 2p and 2p its three orbitals. y x e - e - x e - e - - Even That So Suppose then identifies after I providing add I were just that If Finally, one to So n mtwo = give l then = 2, electron -1. quantum then I you add I tell That just l that = within you identifies 0numbers, one or ml = s l one quantum = 1. +½. 1. orbital just there one number: within are orbital, still one up sublevel to on six one which electrons energy still contains level. three n = Thus, 2. orbitals two four electrons. remaining. quantum numbers always indicate just one electron.

Breaking Down The Code If I described something as having these quantum numbers, what am I really saying? n = 3 l = 2 m l = 2 m s = +½ Translated: n = 3 (so third energy level) l = 2 (so it s a d sublevel we re talking about 3d) m l = 2 (so one particular 3d orbital) m s = +½ (so one electron in one orbital in 3d)

Putting It Into Code Alternatively, what if I wanted to refer to two electrons in the 2p x orbital? How would it be written in quantum numbers? n = 2 (that s an easy one) l = 1 (because when l = 1, that s code for p) m l = -1 (because we just want one p orbital/dumbbell) For our purposes, we could have also picked 0 or 1. m s is not needed because we re talking about two electrons.

Quantum Number Practice What combinations of l and m l can there be when n = 3? l can be 0, 1, or 2 (reflecting s, p, or d sublevels). m l can be -2, -1, 0, 1, or 2 (reflecting the orientation of either one s orbital, three p orbitals, or all five d orbitals). Describe the 3p sublevel using quantum numbers. n=3, l=1, m l =-1, 0, 1 How many electrons am I describing if I indicate quantum numbers of n=4, l=2, m l =2? n indicates a set of 2n 2 electrons (32). l indicates a d sublevel, so that cuts us down to 10 electrons. m l indicates the orientation of one of the d orbitals, so 2 e -.

Cracking the Code Another way to look at it Quantum Numbers n = 3 l = 2 m l = -2 m s = +½ 3d Orbital Notation

Quantum Numbers Summary Image

Quantum Number Practice Quantum Number Practice Worksheet 13 is a CHALLENGE.

Summary Table Quantum Numbers Principal Quantum Number (n) Possible Angular Quantum Numbers (l) Possible Magnetic Quantum Numbers (m) 1 0 (s) 2 0, 1 (s, p) 3 0, 1, 2 (s, p, d) 4 0, 1, 2, 3 (s, p, d, f) 0 (up to 1 orientation for s) -1, 0, 1 (up to 3 orientations for p) -2, -1, 0, 1, 2 (up to 5 orientations for d) -3, -2, -1, 0, 1, 2, 3 (up to 7 orientations for f)

The Rules We ve already learned one rule: Aufbau Principle non-excited electrons fill energy levels from the lowest level up. Now let s learn the other two: Pauli Exclusion Principle Hund s Rule

Pauli Exclusion Principle No more than two electrons can occupy the same orbital (not sublevel, though). Each must have opposite spins within a magnetic field. This is the fourth quantum number m s. +½ ½ Wolfgang Pauli

PEP and Orbital Notation In electron configuration, there is no indication of spin. In the Hog Hotel, electrons in the same orbital were illustrated by opposite-facing hogs. In orbital notation, scientists use up and down arrows to describe electrons opposite spins.

Chemistry versus Hogs Hog Hotel Fill floors from the ground up. Hogs hate to go up stairs if they can avoid it. Chemistry Aufbau Principle Fill energy levels from lowest to highest. Only two hogs per room. They face opposite ways. One hog per room until forced to put two in. Hogs hate to go up stairs.

Chemistry versus Hogs Hog Hotel Fill floors from the ground up. Hogs hate to go up stairs if they can avoid it. Only two hogs per room. They face opposite ways. One hog per room until forced to put two in. Hogs hate to go up stairs. Chemistry Aufbau Principle Fill energy levels from lowest to highest. Pauli Exclusion Principle Only two electrons per orbital. Electrons spin opposite ways.

Hund s Rule Two electrons can occupy a given orbital only after all other orbitals have been filled with one. In the Hog Hotel, Hund s rule was illustrated by the singles only concept. Friedrich Hund

Hund s Rule You can also think of it with a plain English example: Imagine a school bus being filled with students who all dislike each other. http://www.instructables.com/image/fupuhd6fgh3ufiv/removing-school-bus-seats.jpg

Hund s Rule Each student will take a seat by himself until there are no free seats left. Only then will they pair.

Hund s Rule Let s explain Hund s Rule with an example: Oxygen. Oxygen is atomic number 8, so it has 8 electrons. Electrons O Left: 1s 2s 2p 87 63 21 54 0 First, fill the 1s shell with electrons. Then, fill the 2s shell with electrons. Then, begin filling the 2p shell, but only put one electron in each orbital (keep em all spinning the same way). Finally, place a second electron in each shell.

Putting It All Together Using the three rules (Aufbau Principle, Pauli Exclusion Principle, Hund s Rule), let s draw some electron diagrams! Let s start with Helium: He 1s 1s 2 Notice that Helium has a full 1s shell (like a full first floor), with no other electrons occupying any other energy level. This comes into play on the next slide.

Element Electron Configuration Orbital Notation Shorthand Notation Li 1s 2 2s 1 1s 2s 2p Be 1s 2 2s 2 1s 2s 2p B 1s 2 2s 2 2p 1 1s 2s 2p C 1s 2 2s 2 2p 2 1s 2s 2p N 1s 2 2s 2 2p 3 1s 2s 2p O 1s 2 2s 2 2p 4 1s 2s 2p F 1s 2 2s 2 2p 5 1s 2s 2p Ne 1s 2 2s 2 2p 6 1s 2s 2p [He]2s 1 [He]2s 2 [He]2s 2 2p 1 [He]2s 2 2p 2 [He]2s 2 2p 3 [He]2s 2 2p 4 [He]2s 2 2p 5 [He]2s 2 2p 6

Putting It All Together Finally, using all that we ve learned, let s do the following: Complete the Electron Configurations and Orbital Notation sheet. Complete the Electron Configuration Evaluation Worksheet If you can do all this, you re ready.