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Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Patterns in Electron Configuration One of the many patterns contained in the periodic table is that of electron configuration. In this activity, you will identify these patterns. Later, you will use these patterns to determine the order in which electrons fill the orbitals of an atom. As you complete the activity, keep the following in mind: • Period = row, Group = column • Use the table on your book cover, which shows only valence electrons. • There are two number systems for the Groups. We will focus on the A/B system. 1. Which Groups have an s-orbital as the last orbital? 2. Which Groups have a p-orbital as the last orbital? 3. Which Groups have a d-orbital as the last orbital? 4. Which section of the table is left? This section corresponds to the f-orbitals. 5. Look at Group 1A. What is the relationship between the Period number and the energy level of the valence electrons? 6. Look at Group 3A. What is the relationship between the Period number and the energy level of the valence electrons? 7. Look at Group 3B. What is the relationship between the Period number and the energy level of the d-orbitals? 8. Look at the Inner Transition Metals (bottom section). The Lanthanide series (58-71) is part of Period 6. The Actinide series (90-103) is part of Period 7. What is the relationship between the Period number and the energy level of the f-orbitals? 9. Look at all of the A Groups. What is the relationship between the Group number (1A, 2A, etc.) and the total number of valence electrons for each element? (Add up the exponents to find the total number of valence electrons.) Page 1 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Quantum Mechanics and Split Peas I. Introduction Bohr was oh-so-close to explaining the true nature of the electron. However, Bohr’s theories were too simplistic for multi-electron systems. It was evident that a new type of thinking would be needed to describe the “what’s” and “where’s” of the electron. Luckily, some pretty famous scientists (including one named Einstein) were already on the task. Quantum mechanics basically threw out Newtonian physics and applied completely different rules to describe things. The important thing to remember is that “quantum” describes things that are really small and really fast. II. Purpose: To demonstrate ideas of quantum mechanics, orbitals, and probability models. III. Materials: * 40 ml of split peas * funnel * target IV. Procedure 1. Separate the peas into equal volumes in the two graduated cylinders (20 ml in each). 2. Position the target so that the funnel is directly over the center of the target. 3. Have one member of the group place a finger over the end of the funnel and then add 20 ml of the peas into the funnel 4. Release the peas so that they fall onto the target. If they clump too much, re-do the drop from a higher distance. Record the distance you dropped from. 5. Use the rules below to count how many peas are in each area of the target. Record your data on the accompanying data sheet. 6. After completing the counting of the first run, collect the peas and repeat steps 1 – 5 from a different distance, recording this as the 2nd trial. 7. Repeat steps 1-5 with all 40 ml of peas for the 3rd trial. 8. When done, place all of the peas back in a container and clean up. **Rules for counting: o If a pea is completely within an area, it belongs to that area. o If the pea is on a line, it belongs to the area that the greater portion of its volume occupies. o If the pea is on a line, and seems to be equally in two areas, it belongs to the area nearest the center. o Any part of a pea counts as a pea. 6 5 4 3 2 1 The areas are numbered as this picture. Page 2 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Quantum Mechanics and Split Peas Data Table: Trial 1 – Distance ___________ Area # of peas % of total # Trial 2 – Distance ___________ Area # of peas % of total # Trial 3 – Distance ___________ Area # of peas % of total # 1 1 1 2 2 2 3 3 3 4 4 4 5 5 5 6 6 6 7 7 7 Total Total Total Answer the following questions to draw conclusions from this lab: 1. In the first run, which area did the most peas land in? 2. Did this change for the second run? The third? 3. If you were to do this again, only you had to predict where 90% of the peas would fall, what would your prediction be? 4. How do these results mimic the orbitals described in the quantum mechanical model of the atom? Page 3 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: ELECTRON ORBITAL ACTIVITY Background Info: The arrangement of electrons within the orbitals of an atom is known as the electron configuration. The most stable arrangement is called the ground-state electron configuration. This is the configuration where all of the electrons in an atom reside in the lowest energy orbitals possible. Keeping in mind several “rules” that the electrons must obey, we are able to predict the electron configurations of elements using the electron orbital diagrams and periodic table. Electron Orbital Diagram The distributions of orbitals can be laid out in the electron orbital diagram pictured to the left. You read it from the bottom (or closest to the nucleus) to the top in order of increasing energy. 7d Increasing energy 6f 7p 6d 5f 7s 6p The bottom energy level is level 1 - it has the lowest energy. 5d 4f 6s 5p Each " " represents an orbital. 4d 5s 4p You can see that there is 1 orbital for an s subshell. There are 3 orbitals for a p subshell, 5 for a d, and 7 for an f subshell. Each orbital can hold up to 2 electrons. 3d 4s 3p 3s 2p 2s Therefore, the s subshell can hold 2 electrons; the p can hold 6; the d can hold 10; and the f can hold 14. See chart below for a summary: Nucleus 1s Subshell s p d f Number of Orbitals 1 3 5 7 Shape of orbital sphere dumbells Max # of electrons 1x2 = 2 3x2 = 6 5x2 = 10 7x2 = 14 Example of the Configuration of H: Atomic number 1 means at ground state, 1 electron and 1 proton. Therefore, 1 electron in the first energy level, sub level s, orbital 1 - 1s1 Recall, however, the “d-block” anomaly – the following diagram may also be helpful in understanding how electrons fill energy levels once the “(n – 1)” d-block comes into play. Page 4 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Procedure: 1. For each of the following atoms or ions, determine the number of protons and electrons contained and record them in the Data Table. 2. Determine the correct number of electrons (chips) according to the following three rules. A. Place markers (electrons) in the energy levels from low energy to higher energy (starting at level 1 which is closest to the nucleus) B. Place one marker (electron) in each box (orbital) before doubling up C. Each box has a max capacity of two. When doubling up within a box (orbital), be sure the one marker (electron) has its arrow facing up while the other is down. 3. Record how your completed card looks (draw in arrows) on the electron orbital diagrams. 4. Also record the information for the electrons’ energy level, subshell, and number of electrons in the data table. 5. Convert the information from the table into the electron configuration. Electron configurations give the address information for the electron’s location for the ground state of an atom. See below for a key to writing electron configurations. The first one, carbon, has been done for you as an example. 6. Remove the chips and repeat for the next atom. Atom # of protons Carbon 6 Energy level Subshell 1, 2, 3, … s, p, d or f 1 s 2 s 2 p Electron Configuration: Example: # of electrons 6 # of electrons 2 2 2 1s2 2s2 2p2 Page 5 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Data Tables: 1. B - Boron Period (row) # of protons # of electrons Energy level 1, 2, 3, … Subshell s, p, d or f # of electrons Name: Electron Orbital Diagrams: Electron Configuration: 2. Ne - Neon Period (row) # of protons # of electrons Energy level 1, 2, 3, … Subshell s, p, d or f # of electrons Electron Configuration: 3. Cl - Clorine Period (row) # of protons # of electrons Energy level 1, 2, 3, … Subshell s, p, d or f # of electrons Electron Configuration: Page 6 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure 4. V - Vanadium Period (row) # of protons # of electrons Energy level Subshell # of electrons 1, 2, 3, … s, p, d or f Name: Electron Configuration: 5. Cu - Copper Period (row) # of protons # of electrons Energy level Subshell 1, 2, 3, … s, p, d or f # of electrons Electron Configuration: 6. Br - Bromine Atom # of protons # of electrons Energy level Subshell 1, 2, 3, … s, p, d or f # of electrons Electron Configuration: Page 7 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Now let’s try some ions: Ions are formed when atoms lose or gain electrons in an effort to create more stable electron arrangements. Data Tables: 7. Atom # of protons # of electrons +2 Be Energy Subshell # of electrons level 1, 2, 3, … s, p, d or f Electron Orbital Diagrams: Electron Configuration: 8. Atom N-3 Energy level 1, 2, 3, … # of protons # of electrons Subshell # of electrons s, p, d or f Electron Configuration: Mystery Elements: Identify the following elements based on their electron configurations. Then re-write the configuration using Nobel Gas abbreviated notation. Page 8 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Concluding Questions 1. Electrons can be thought of “obeying” three basic rules when it comes to their location within an atom. Match the following rules with the three from procedure THREE that you just used to complete the lab. Rule Hund’s Rule Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. Pauli Exclusion Principle A maximum of two electrons may occupy a single orbital, but only if the electrons have opposite spins. The Aufbau Principle: Each electron occupies the lowest possible energy orbital Letter? 2. What is the maximum number of electrons that can be in one orbital (or box)? ____ 3. How many electrons are needed to fill the first energy level?____ How many elements are there in the first period or row of the periodic table? ____ 4. How many electrons are needed to fill the second energy level (2s + 2p)? ____. How many elements are there in the second period of the periodic table? ____ 5. How many electrons are needed to fill the third energy level (3s + 3p + 3d)? ____. How many elements are there in the third period of the periodic table? ____ Wait a second…that seems strange! 6. What element(s) did you write the configuration for that ended with an electron configuration of 3d? ________________________________ What row or period is it in? ________ 7. Something is unusual about the 4s and 3d orbitals. What do you notice about the order in which they are filled? 8. Why do you think Groups 1 and 2 referred to as the s-block of the periodic table? (look at the ending of the electron configurations you would write for elements K and Be) 9. Why are Groups 13 through 18 referred to as the p-block of the periodic table? (look at the ending of the electron configurations you wrote for elements C, Ne, Al, and Cl) Page 9 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: 10. Why are Groups 3 through 12 referred to as the d-block of the periodic table? (look at the ending of the electron configurations you wrote for element V and Cu) 11. The most stable elements have full outer sub shells. Next most stable are those with half filled shells. Take a look at the electron configuration you wrote for Copper. Copper’s actual configuration is: 1s2 2s2 2p6 3s2 3p6 4s1 3d10. Give one possible reason for this. 12. Using what you now know about electron configurations explain the notion that elements in the same column in the periodic table have similar chemical and physical properties. 13. What element would have the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d5. (For example: 1s2 means that in the first energy level there are two electrons in an s orbital.) 14. Write out the electron configuration of O2-. Notice that this is an ion! 15. Write out electron configuration of Na+. Notice that this is a ion! 16. Look at the electron configurations for O2- and Na+ which you figured out above. Compare them to the other electron configurations you did. Do they have anything in common with one or two of them? Which one(s) and Why? 17. Look at the electron configuration for Neon, a noble gas. Why do you think they are more stable than the other elements? Page 10 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: 7d 6f 7p 6d 5f 7s 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Page 11 Apologia Chemistry – Module 7 In-Class Labs – Atomic Structure Name: Atomic Structure and Theory Magic Square Directions: Put the number of the definition from the list below into the square with the appropriate term. Check your answers by adding the numbers to see if all the sums of all rows, both across and down add up to the same number, the Magic #. Democritus Dalton Thomson Chadwick Total _____ _____ _____ ______ _____ Rutherford Proton Atom Bohr _____ _____ _____ _____ Wave Model Neutron Nucleus Alpha particle _____ _____ _____ _____ Electron Model Energy levels Electron cloud _____ ______ _____ _____ Total _____ _____ _____ _____ _____ _____ ______ Magic Number ______ 1. Represented by a symbol; all are found on the Periodic Table 2. Made a mental model of the atom; Greek philosopher 3. Used by Rutherford in his experiment; made of two protons and two neutrons 4. The paths in which electrons circle the nucleus according to the Bohr model 5. The positive particle in the nucleus of an atom 6. The tiny positive core of an atom; contains protons and neutrons 7. Formed the atomic theory model of the atom; English schoolteacher 8. Discovered the nucleus using his gold foil experiment 9. Current explanation of where electrons might be found in the atom 10. Used by scientists to explain something we cannot see or understand 11. The smallest particle of an element that has the properties of that element 12. Discovered the neutron 13. Current model of the atom; proposed by Schrodinger 14. Mass of protons and neutrons 15. Developed the model of the atom in which electrons orbit the nucleus in energy levels 16. The negative particle that circles the nucleus 17. The neutral particle in the nucleus of an atom 18. Proposed the “plum-pudding” model of the atom; discovered the electron Page 12