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Fireworks! 1. Do you enjoy watching fireworks? Why or why not? 2. What are the colors and shapes and sounds of a firework that have most impressed you? Today’s assignment: Re-read the article and find 2-3 interesting facts. Research 2-3 amazing fireworks facts or stories that you would have included in the article. Resources are listed on the handout. Then in your group, design your own firework display. <google “fireworks graphic design”> Here’s one for you: The electron configuration of carbon is 1s22s22p2. How many electrons are present in carbon? What is the atomic number of carbon? Bellwork 1) The electron configuration of nitrogen is 1s2 2s2 2p3. What is the atomic number of nitrogen? How many electrons are in a nitrogen atom? Write the orbital notation for nitrogen. Bellwork – 1/27 The electron configuration of boron is 1s2 2s2 2p1. What is the atomic number for boron? How many electrons are present in an atom of boron? Which period of the periodic table is it in? Which elements are in the same period? At the end of class today, 1/26 you should be able to: Write the orbital diagram and electron configuration for any 1st or 2nd period element. Write the noble gas configuration for any 3rd or 4th period elements! Bellwork 1/13 1) What is the total number of orbitals (not shapes) in the 3rd main energy level? 2) What’s the difference between a shell, a subshell and an orbital? 3) How many electrons can a p sublevel hold? Bellwork 1/15 Which of the principles do each of the following electron energy diagrams violate? 1) 2) 3) At the end of class today you should be able to : State in your own words the Aufbau principle, the Pauli exclusion principle, and Hund’s rule. Target Practice! Show what cha know! Bellwork 1/28 Draw and label (with the arrows) the orbital diagram for fluorine. Identify these elements: [Kr] 5s1 [Ar] 4s2 3d5 1/13 Target – turn it in! In your own words: 1. State the Aufbau principle. 2. State the Pauli exclusion principle. 3. State Hund’s rule… Section 3: Electron Configurations Previously in Chapter 4… Electrons live in little houses called orbitals. Bohr thought that a whole bunch of electrons could live in one house at a time, but the quantum model made it clear that, like a studio apartment with an unfriendly landlord, only two can live in one at the same time. In this section: We will list the total number of electrons needed to fully occupy each main energy level. State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule. Describe the electron configurations for the atoms of any element using orbital notation, electronconfiguration notation and noble-gas notation. Finally! The quantum model of the atom improves on the Bohr model because it describes the arrangements of electrons in atoms – other than hydrogen… The arrangement of electrons in an atom is known as the atom’s electron configuration. Unique! Because atoms of different elements have different numbers of electrons, a unique electron configuration exists for the atoms of each element. Like all systems in nature, electrons want to be arranged with the lowest possible energy… just like high school students! Flashback!! What is an atom’s ground state??? Remember ground state? The element’s ground state configuration is the lowest-energy arrangement of electrons for that element. A few simple rules combined with quantum numbers allow us to determine the ground state configuration. This is carbon’s electron configuration: 2 1s 2 2s 2 2p The integers are n (main energy levels). The letters are l (shape/orbital). The superscripts are how many electrons are in each level. Electrons fill the lowest-energy orbitals first. Step 1: Figure out the energy levels of the orbitals. Step 2: Electrons are added to the orbitals, one by one, according to 3 basic rules. Rule #1: Aufbau principle Rule #1: According to the Aufbau principle, electrons fill up lower energy orbitals first, and then start filling up higher energy orbitals. In a ground-state H atom, the electron is in the 1s orbital. Interesting sidenote: The word “aufbau” isn’t somebody’s name. It means something in German, but it doesn’t really matter… And back to the lesson… Chapter 4 Configurations Relative Energies of Orbitals Rule #2: People who are remembered for only one thing… Wolfgang Pauli stated that all electrons in an element have unique sets of quantum numbers, an idea known as the Pauli Exclusion Principle. Because quantum numbers are the mathematical description of each electron, making sure all elements have unique sets of quantum numbers allows us to tell them apart. And that’s the one thing we remember Pauli for. Rule #2: Pauli Exclusion Principle Rule #2 reflects the importance of the spin quantum number. According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers. An orbital only holds 2 electrons of opposite spin states. Rule #3: Hund’s Rule According to Hund’s rule, electrons prefer to be unpaired whenever possible. This means they will always enter an empty orbital before they pair up. Hund’s rule also states that all single electrons in orbitals must have the same spin state. Hund’s Rule in Action… Show 2 e- filling the p orbitals Show 3 e- filling the p orbitals Show 4 e- filling the p orbitals Did you mind the rules?! With a partner, take turns to state Aufbau principle, Pauli Exclusion principle and Hund’s Rule. Use the 5 orbital boxes in your notes to explain how 3, then 5, then 8 then 10 electrons will fill these orbitals. Show how 3, 5, 8 and 10 electrons fill these orbitals. Try it yourself and let’s check it! All you need is right here! Find the period/row… which equals “n”. Then use what you know about quantum numbers to fill in the electrons. Orbital diagrams are easy! Orbital diagrams are pictures that show what electron configurations look like when you draw them. Because each term in an electron configuration is lower in energy than the one before it, we pretty much just draw a bunch of lines and fill ‘em full of electrons. Each electron is represented by an arrow, and the up and down arrows in a single orbital represent the two different electrons in this orbital. • In this method, each line (or box) represents a single orbital, which is why there’s one line for s-orbitals (there’s only one s-orbital for every energy level), • three lines for p-orbitals (there are 3 porbitals for every energy level – except n=1, right??), Let’s start at the start… H and He What is the atomic number? How many electrons? Draw the orbital diagram then the electron configuration. Let’s try some 2nd period notations Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Li Be B C N O F Ne Noble gas notation • If we want to make our lives A LOT simpler, we don’t have to write out all of these terms that are in normal electron configurations. • Instead of writing out a big string of orbitals, just abbreviate it so you start at the most recent noble gas to the element you’re interested in. For example: Ne is 1s2 2s2 2p6, right? In the 3rd period, Na is 1s2 2s2 2p6 3s1 Using noble configuration it becomes Na = [Ne] 3s1 Mg = [Ne] 3s2 This is sooooooo much simpler! Use noble gas notation for the 3rd period elements Na [Ne] 3s1 Mg [Ne] 3s2 Al Si P S Cl Ar Let’s look at 4th period elements K [Ar] 4s1 Ca [Ar]4s2 Sc [Ar] 4s2 3d1 Just keep following the row in the periodic table BUT REMEMBER, the d block is LOWER than the s block. From 4s1 and 4s2… it is 3d1, 3d2, 3d3…3d10 Ti [Ar] 4s2 3d2 V [Ar] 4s2 3d3 Whoa! Look! Cr [Ar] 4s1 3d5 Why is Cr: [Ar] 4s1 3d5?? Shouldn’t that be Cr [Ar] 4s2 3d4??? Let’s look at the orbital diagram of chromium and sort this out! Compare these: Cr: [Ar] ___ 4s ___ ____ ___ ___ ____ 3d Cr: [Ar] ___ 4s ___ ____ ___ ___ ____ 3d 4s1 3d5 is more stable than 4s2 3d4, even though it follows our pattern… Let’s pick up the pattern again… Mn [Ar] 4s2 3d5 Fe Co Ni Watch out for copper!! It happens one more time in the 3rd period… Copper! Cu [Ar] 4s1 3d10 Cu: [Ar] ___ ___ ____ ___ ___ ____ There isn’t a simple explanation… Everybody else behaves and follows our neat pattern… Zn Ga Ge As Se Br Kr [Ar] 4s2 3d10 [Ar] 4s2 3d10 4p1 [Ar] 4s2 3d10 4p2 Two definitions for you… The highest occupied energy level of an electron is the energy level with the highest principal quantum number. For example H, n=1 and the highest energy level is the first main energy level. For Be, n=2 and the highest energy level is the second main energy level. The inner shell electrons are the electrons that are not in the highest occupied energy level. That totally makes sense! This is also a good time to introduce Valence electrons are the electrons in the outermost (highest) principal energy level of an atom. For example, N 1s2 2s2 2p3 has e- in n=1 and n=2. N=2 is the outermost so the valence e- are the 2s and 2p electrons. For sodium, 1s2, 2s2, 2p6 3s1 or [Ne] 3s1, the valence e- is the electron in the 3s orbital because n=3 is the outermost level. Valence electrons are vewwy, vewwy important! Valence electrons are the most important because they are the ones involved when atoms form bonds (attach to each other). And did you happen to notice a trend in the periodic table?? Indeed! The atoms of elements in the same group (vertical column) have the same number of electrons in a given type of orbital. The only difference is the principal energy level of that sublevel. Thank you, Mr. Mendeleev Remember that elements were originally organized into groups on the periodic table based on similar chemical properties. Now we can understand the reason behind these groupings! Elements with the same valence electron arrangement show very similar chemical behavior! Sample B pg. 115 Now it’s your turn! Chapter 4 Section 3 Electron Configurations Sample Problem B • a. Write both the complete electronconfiguration notation and the noble-gas notation for iron, Fe. • b. How many electron-containing orbitals are in an atom of iron? How many of these orbitals are completely filled? How many unpaired electrons are there in an atom of iron? In which sublevel are the unpaired electrons located? Chapter 4 Section 3 Electron Configurations Sample Problem B Solution • a. The complete electron-configuration notation of iron is 1s22s22p63s23p63d64s2. Iron’s noble-gas notation is [Ar]3d64s2. • • b. An iron atom has 15 orbitals that contain electrons. They consist of one 1s orbital, one 2s orbital, three 2p orbitals, one 3s orbital, three 3p orbitals, five 3d orbitals, and one 4s orbital. Eleven of these orbitals are filled, and there are four unpaired electrons. They are located in the 3d sublevel. The notation 3d6 represents 3d • • • Practice Problems 1a) Write the noble gas notation for I: 1b) How many unpaired e- are there? 2a) Write the noble gas notation for Sn: 2b) How many unpaired e- are there? Answers Practice B 1a) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5, [Kr]4d10 5s2 5p5 1 unpaired eb) 27, 26, 1 2 a) [Kr]4d10 5s2 5p2, 2 b) 10, germanium 3a) 1s2 2s2 2p6 3s2 3p6 3d5 4s2 b) manganese 4 a) 9, 1s2 2s2 2p6 3s2 3p6 b. argon Let’s practice together: 1a) How many inner shell electrons does an atom of silicon, Si, contain? b) How many electron containing orbitals are in an atom of silicon? How many of these orbitals are filled? How many unpaired electrons are there in an atom of silicon? c) How many electron containing p orbitals are there in an atom of silicon? Answers: a) 10 b) 8, 6, 2 c) 5 Let’s try this one from the book! Sample C, pg. 116 Chapter 4 Section 3 Electron Configurations Sample Problem C • a. Write both the complete electronconfiguration notation and the noblegas notation for a rubidium atom. • b. Identify the elements in the second, third, and fourth periods that have the same number of highest-energy-level electrons as rubidium. Chapter 4 Section 3 Electron Configurations Sample Problem C Solution • a. 1s22s22p63s23p63d104s24p65s1, [Kr]5s1 • b. Rubidium has one electron in its highest energy level (the fifth). The elements with the same outermost configuration are, in the second period, lithium, Li; in the third period, sodium, Na; and in the fourth period, potassium, K. • • • Practice problems 1 – 2 Let’s get it a go! Answers in Appendix E, in the back of the book… Next, we need to try Section 3, Formative Assessment… which can only lead to ONE THING: For practice: Textbook pg. 120 #26-31. Homework answers: Aufbau principle 26a) An electron occupies the lowest energy orbital that can receive it. b) The lowest energy orbital is filled first. Electrons are then added to the orbital with the next lowest energy and so on until all of the electrons in the atom have been placed in orbitals. Hund’s rule 27a) Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. b) By placing as many single electrons as possible in separate orbitals in the same energy level, electron-electron repulsion is minimized and favorable lower energy arrangements result. 28a) No two electrons in the same atom can have the same four quantum numbers. b) The 2 different values of the spin quantum number permit two electrons of opposite spin states to occupy the same orbital. 29a) the highest occupied energy level in an atom is the electron containing main energy level that has the highest principal quantum number. b) Inner shell electrons are electrons that are not in the highest occupied energy level. 30a) first main energy level n=1 b) 2nd main energy level n=2 c) 3rd main energy level n=3 d) 4th main energy level n=4 e) 5th main energy level n=5 Practice time! 1a) 1s2 2s2 2p6 3s2 3p6 3d2 4s2; [Ar]3d2 4s2 b) 12: one 1s orbital, one 2s orbital, three 2p orbitals, one 3s orbital, three 3p orbitals, one 4s orbital and two 3d orbitals; 10 2