SAS 4 Module 4 Quantum Numbers and Electronic Configuration CHE 026 PDF

Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _____...

Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ Lesson title: Quantum Numbers and Electronic Configuration Materials: Lesson Objectives: ⚫ Pen, SAS, highlighter At the end of the module, you should be able to: 1. Differentiate the different quantum numbers. References: 2. Write the correct electronic configuration of elements. ⚫ General, Organic and 3. Apply the principles involved in quantum numbers Biological Chemistry 6th Ed 4. Determine the number of valence electrons based on electronic Stoker, 2013) configuration. Productivity Tip: Try the Pomodoro Technique! Set your timer to 25 minutes and work during this time period. After this period, take a 5-minute break. A. LESSON PREVIEW/REVIEW 1) Introduction (2 mins) Recently, we have dealt with how the periodic table was organized and arranged according to their properties. Many proponents have contributed in the arrangement of periodic table. One difference among its arrangement is their electronic configuration. These configurations is based on how electrons move within atoms. In this module, you will be challenged in correlating different arrangements of electrons based on shells, subshells and orbitals. Quantum numbers represent this arrangements. Also electronic configurations are based on how these numbers affect the placement of electrons. 2) Activity 1: What I Know Chart, part 1 (5 mins) Instructions: In this chart, reflect on what you know. Do not worry if you are not sure of your answers. This activity simply serves to get you started on thinking about our topic. Answer only the first column “What I Know”. Leave the “What I learned” blank at this time. What I Know Questions: What I Learned (Activity 4) 1. What is an electronic configuration:? 2. What’s the difference between shells, subshells and orbitals? 3. What are the different quantum numbers that represent the placement of electrons? This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ B. MAIN LESSON 1) Activity 2: Content Notes (40 mins) Instructions: Please take note, underline, highlight or outline some important and striking information. You can use yellow highlighter for definitions and pink highlighter for enumerated examples or items. As electrons move about an atom’s nucleus, they are restricted to specific regions within the extranuclear portion of the atom. Such restrictions are determined by the amount of energy the electrons possess. Furthermore, electron energies are limited to certain values, and a specific “behavior” is associated with each allowed energy value. The space in which electrons move rapidly about a nucleus is divided into subspaces called shells, subshells, and orbitals. Electrons that Electron Shells occupy the first Electrons within an atom are grouped into main energy levels called electron shell are electron shells. An electron shell is a region of space about a nucleus that closer to the contains electrons that have approximately the same energy and that spend most nucleus and have of their time approximately the same distance from the nucleus. Electron shells are numbered 1, 2, 3, and so on, outward from the nucleus. Electron energy increases a lower energy as the distance of the electron shell from the nucleus increases. An electron in shell than electrons in 1 has the minimum amount of energy that an electron can have. The maximum the second number of electrons that an electron shell can accommodate varies; the higher the electron shell. shell number (n), the more electrons that can be present. In higher-energy shells, the electrons are farther from the nucleus, and a greater volume of space is available for them; hence more electrons can be accommodated. (Conceptually, electron shells may be considered to be nested one inside another, somewhat like the layers of flavors inside a jawbreaker or similar type of candy.) The lowest- energy shell (n= 1) accommodates a maximum of 2 electrons. In the second, third, and fourth shells, 8, 18, and 32 electrons, respectively, are allowed. The relationship among these numbers is given by the formula 2n2, where n is the shell number. For example, when n=4, the quantity 2n2 = 2(4)2 = 32. Electron Subshells Within each electron shell, electrons are further grouped into energy sublevels called electron subshells. An electron subshell is a region of space within an electron shell that contains electrons that have the same energy. We can draw an analogy between the relationship of shells and subshells and the physical layout of a high-rise apartment complex. The shells are analogous to the floors of the apartment complex, and the subshells are the counterparts of the various Figure 1. apartments on each floor. The number of subshells within a shell is the same as Electron Shell the shell number. Shell 1 contains one subshell, shell 2 contains two subshells, Diagram shell 3 contains three subshells, and so on. Subshells within a shell differ in size (that is, the maximum number of electrons they can accommodate) and energy. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ The higher the energy of the contained electrons, the larger the subshell. Subshell size (type) is designated using the letters s, p, d, and f. Listed in this order, these letters denote subshells of increasing energy and size. The lowest-energy subshell within a shell is always the s subshell, the next highest is the p subshell, then the d subshell, and finally the f subshell. An s subshell can accommodate 2 electrons, a p subshell 6 electrons, a d subshell 10 electrons, and an f subshell 14 electrons. Both a number and a letter are used in identifying subshells. The number gives the shell within which the subshell is located, and the letter gives the type of subshell. Shell 1 has only one subshell — the 1s. Shell 2 has two subshells — the 2s and 2p. Shell 3 has three subshells — the 3s, 3p, and 3d, and so on. Electron subshells have within them a certain, definite number of locations (regions of space), called electron orbitals, where electrons may be found. In our apartment complex analogy, if shells are the counterparts of floor levels and subshells are the apartments, then electron orbitals are the rooms of the apartments. An electron orbital is a region of space within an electron subshell where an electron with a specific energy is most likely to be found. An electron orbital, independent of all other considerations, can accommodate a maximum of 2 electrons. Thus, an s subshell (2 electrons) contains one orbital, a p subshell (6 electrons) contains three orbitals, a d subshell (10 electrons) contains five orbitals, and an f subshell (14 electrons) contains seven orbitals. Orbitals have distinct shapes that are related to the type of subshell in which they are found. Note that it is not the shape of an electron, but rather the shape of the region in which the electron is found that is being considered. An orbital in an s subshell, which is called an s orbital, has a spherical shape. Orbitals found in p subshells — p orbitals — have shapes similar to the “figure 8” of an ice skater More complex shapes involving four and eight lobes, respectively, are associated with d and f orbitals. Some d and f orbitals have shapes related to, but not identical to, those shown in Figure 3. Figure 2. The number of subshells within a shell is equal to the shell number, as shown here for the first four shells. Each individual subshell is denoted with both a number (its shell) and a letter (the type of subshell it is in). This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ Figure 3. An s orbital has a spherical shape, a p orbital has two lobes, a d orbital has four lobes, and an f orbital has eight lobes. The f orbital is shown within a cube to illustrate that its lobes are directed toward the corners of a cube. Some d and f orbitals have shapes related to, but not identical to, those shown. Figure 7. Common observable changes in matter Electron Spin Experimental studies indicate that as an electron “moves about” within an orbital, it Figure spins on its own axis 8. in Classification of Matter either a clockwise according to composition or a counterclockwise direction. Furthermore, when two electrons are present in an orbital, they always have opposite spins; that is, one is spinning clockwise and the other counterclockwise. This situation of opposite spins is energetically the most favorable state for two electrons in the same orbital. Quantum Numbers A total of four quantum numbers are used to describe completely the movement and trajectories of each electron within an atom. The combination of all quantum numbers of all electrons in an atom is described by a wave function that complies with the Schrödinger equation. Each electron in an atom has a unique set of quantum numbers; according to the Pauli Exclusion Principle, no two electrons can share the same combination of four quantum numbers. Quantum numbers are important because they can be used to determine the electron configuration of an atom and the probable location of the atom's electrons. Quantum numbers are also used to understand other characteristics of atoms, such as ionization energy and the atomic radius. A summary of how quantum number of an electron, you can refer to the summary table provided in the next page. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ Quantum number (symbol) Values Function/Description Principal Quantum Number (n) 1, 2, 3, 4 It is the main energy level 1=K shell (closest to the nucleus) The principal energy level is 2=L shell related to the average distance 3=M shell from the nucleus. 4=N shell (farthest) It is responsible for the SIZE of the electron cloud Should be an integer, it can’t be zero The number of electrons per energy level can be computed using the formula: 2n2 Azimuthal/Angular Momentum l= 0 → n-1 Sublevel/subshell is a set of Quantum Number (l) n=1, l=0 equal-energy orbitals within a n=2, l=0,1 principal energy level. n=3, l=0,1,2 It is responsible for the SHAPE n=4, l=0,1,2,3 of the electron cloud. 0= s subshell (Spherical) 1= p subshell (dumbbell-shaped) 2= d subshell (cloverleaf/x- shaped) 3= f subshell (complex) Magnetic Quantum number (ml) m l= - l → l It is responsible for the n=1, l=0, ml = 0 (1 orbital) orientation of ORIENTATION n=2, l=0, ml = 0 (1 orbital) of electron cloud. l=1, ml = -1, 0, 1 (3 orbitals) This identifies the ORBITAL in n=3, l=0, ml = 0 an atom. l=1, ml = -1, 0, 1 An atomic orbital is a specific l=2, ml = -2, -1, 0, 1, 2 region of a sublevel containing n=4, l=0, ml = 0 a maximum of two electrons. l=1, ml = -1, 0, 1 Higher-energy orbitals (f, g, l=2, ml = -2, -1, 0, 1, 2 and so forth) also exist, but l=3, ml = -3, -2, -1, 0, 1, 2 they are important only in the description of the electron s = 1 orbital = 2e- arrangements of the heaviest p = 3 orbitals = 6e- elements. d = 5 orbitals = 10e- If n=2, then there is a total of 4 orbitals f = 7 orbitals = 14e- in the shell which also means it can carry 8 electrons in the given shell. Spin Quantum Number (s/ms) + ½ (spin up), -½ (spin down) It is responsible for the DIRECTION of the electron cloud. Only 2 electrons per orbital. It must start first with + ½ spin up. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________. Principles to take note in quantum numbers: ▪ PAULI’S EXCLUSION PRINCIPLE - No two electrons can have the same set of quantum numbers. ▪ AUFBAU’S PRINCIPLE - “Building up principle” - Lower electron orbitals are filled first ▪ HUND’S RULE OF PAIRING/MULTIPLICITY - Orbitals are filled up singly before pairing Pairing terms: Paramagnetism – attracted to a magnet, unpaired electrons Diamagnetism – repelled by a magnet, paired electrons ▪ HEISENBERG’S UNCERTAINTY PRINCIPLE - It is impossible to determine simultaneously the momentum and position of an electron. Writing Electron Configurations An electron configuration is a statement of how many electrons an atom has in each of its electron subshells. Because subshells group electrons according to energy, electron configurations indicate how many electrons of various energies an atom has. Electron configurations are not written out in words; rather, a shorthand system with symbols is used. Subshells containing electrons, listed in order of increasing energy, are designated by using number–letter combinations (1s, 2s, and 2p). A superscript following each subshell designation indicates the number of electrons in that subshell. Guidelines for Writing Electron Configurations of Atoms Obtain the total number of electrons in the atom from the atomic number found on the periodic table. The number of electrons equals the number of protons for an atom. Electrons in atoms occupy the lowest energy orbitals that are available, beginning with 1s. Fill subshells according to the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,. Remember: The s sublevel has one orbital and can hold two electrons. The p sublevel has three orbitals. The electrons will half-fill before completely filling the orbitals for a maximum of six electrons. ° The d sublevel has five orbitals. Again, the electrons will half-fill before completely filling the orbitals for a maximum of ten electrons VALENCE ELECTRONS A valence electron is an electron that is the most likely to be involved in a chemical reaction. They are typically the electrons with the highest value of the principal quantum number, n. It's worth noting the IUPAC definition of valence is for the single highest valence value that is displayed by an atom of an element. However, in practical use, main group elements of the periodic table may display any valence from 1 to 7 (since This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Student Activity Sheet #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ 8 is a complete octet). Most elements have preferred values of valence electrons.Valence electrons can be found by determining the electronic configurations of elements. Thereafter the number of electrons in the outermost shell gives the total number of valence electrons in that element. Ex. Si 1s 2 2s 2 2p 6 3s 2 3p 2 The outer most shell is 3, which all in all has 4 electrons. This means that Silicon has 4 valence electron. Write the electron configurations for the following elements. A. Strontium (atomic number = 38) B. Lead (atomic number = 82) a. The number of electrons in a strontium atom is 38. Remember that the atomic number gives the number of electrons. Subshells are filled, in order of increasing energy, until 38 electrons have been accommodated. The 1s, 2s, and 2p subshells fill first, accommodating a total of 10 electrons among them. 1s22s22p6 Next, the 3s subshell fills and then the 3p subshell. 1s22s22p6 3s23p6. At this point, 18 electrons have been accommodated. To get the desired number of 38 electrons, 20 more electrons are still needed. The 4s subshell fi lls next, followed by the 3d subshell, giving a total of 30 electrons at this point. 1s22s22p63s23p6 4s23d10... Note that the maximum electron population for d subshells is 10 electrons. Eight more electrons are needed, which are added to the next two higher subshells, the 4p and the 5s. The 4p subshell can accommodate 6 electrons, and the 5s can accommodate 2 electrons. Answer is 1s22s22p63s23p64s23d104p65s2 To double-check that we have the correct number of electrons, 38, the superscripts in our final electron configuration are added together. 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 = 38 b. To write this configuration, the same procedures are followed as in part a, remembering that the maximum electron subshell populations are s =2, p = 6, d = 10, and f = 14. Lead, with an atomic number of 82, contains 82 electrons, which are added to subshells in the following order. Answer is 1s22s22p63s23p64s23d104p65s24d105p66s24f 145d106p2. Note in this electron configuration that the 6p subshell contains only 2 electrons, even though it can hold a maximum of 6. Only 2 electrons are added to this subshell because that is sufficient to give 82 total electrons. If the subshell had been completely filled, 86 total electrons would be present, which is too many. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ 2) Activity 3: Skill-building Activities (with answer key) (30 mins + 5 mins checking) A. Answer the given items in then space provide. 1. What are the four (4) quantum numbers? What are the abbreviations of each quantum number? 2. Distinguish between a principal energy level and a sublevel. Answer in 2 to 3 sentences only. 3. Distinguish between a sublevel and an orbital.Answer in 2 to 3 sentences only. 4. How many electrons can an orbital carry? 5. What are the shapes of the given sublevels? a. s b. p c. d d. f 6. What is the maximum number of electrons in each of the following energy levels? a. n = 1 b. n = 2 c. n = 3 7. For any given principal energy level, what is the maximum number of electrons that can exist in the following subshells? a. s b. p c. d 8. State the Pauli exclusion principle. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ 9. State Hund’s rule. Determine whether the following orbital diagrams violate Hund’s rule. You should write “Follow the Hund’s rule” or “Does not follow the Hund’s rule” after each item to determine if the example followed or violated the Hund’s rule. 10. Write the electronic configuration and valence electron of each element: a. Na (Z=11) b. S (Z=16) c. Ca (Z=20) 11. Which of the following electron configurations are not possible? Why? a. 1s2 1p2 c. 2s2 2s2 2p6 2d1 2 2 2 b. 1s 2s 2p d. 1s2 2s3 Note: Check your answers against the Key to Corrections found at the end of this SAS. Write your score on your paper. 3) Activity 4: What I Know Chart, part 2 (3 mins) Instruction: To review what was learned from this session, please go back to Activity 1 and answer the “What I Learned” column. Notice and reflect on any changes in your answers. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ 4) Activity 5: Check for Understanding (30 mins) Multiple choice. Encircle the letter of the best answer and correct answer in the following questions. 1. It describes the shape of the electron cloud. a. n c. ms b. l d. ml 2. If l is equals to 0, it describes what shape of the subshell? a. spherical c. dumbbell b. clover-leaf d. complex 3. What is the maximum number of electrons that shell 2 can carry? a. 2 c. 18 b. 8 d. 32 4. Which among the shells given below is farthest to the nucleus? a. M c. L b. K d. O 5. The d subshell has a maximum of ____ orbitals. a. 1 c. 7 b. 3 d. 5 6. Which of the following pairs is INCORRECT? a. s – spherical b. p – dumbbell c. d – dumbbell d. f – complex 7. It gives the orientation of the electron cloud. a. n c. ms b. l d. ml 8. It is a region of space within an electron shell that contains electrons that have the same energy. a. Orbital c. Shell c. Subshell d. Spin 9. If l is equals to 2, it describes what shape of the subshell? This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ a. spherical c. dumbbell b. clover-leaf d. complex 10. If l is 1, how many orbitals are present? a. 7 c. 1 b. 5 d. 3 11. What is the correct electronic configuration of Cl (Z=17) ? a. 1s2 2s2 2p6 3s2 3p4 b. 1s2 2s2 2p6 3s3 3p2 c. 1s2 2s2 2p3 3s2 3p5 d. 1s2 2s2 2p6 3s2 3p5 12. Which principle states that no two electrons can have the same set of quantum numbers? a. Aufbau´s principle b. Pauli´s exclusion principle c. Heisenberg´s uncertainty principle d. Hund´s rule of multiplicity 13. The following statements are true, regarding an element with an electronic configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5 ,EXCEPT: a. The valence shell of this element is 5 b. There are 7 valence electrons c. The atomic number of this element is 53. d. None of the above 14. What is the valence electron of the atom with electronic configuration of: 1s22s22p63s23p2? a. 3 b. 2 c. 6 d. 4 15. Which element has a valence electron of 4? a. Carbon (Z=14) b. Aluminum (Z=13) c. Potassium (Z=19) d. Magnesium (Z=12) Note: Answers will be provided and discussed by the instructor. C. LESSON WRAP-UP 1) Activity 6: Thinking about Learning (5 mins) A. Work Tracker You are done with this session! Let’s track your progress. Shade the session number you just completed. This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ P1 P2 1 2 3 4 5 6 7 8 9 10 A. Think about your Learning Tell me about your thoughts in this module! What are your insights of this topic? Did you gain new information? FAQs 1. Do all elements follow the filling up of electrons in the electronic configuration? Not at all. An example for this is the element Copper (Cu). As observed in the periodic table, Copper has an electronic configuration 1s22s22p63s23p64s13d10 when we expect it to be 1s22s22p63s23p64s23d9. However, because the 3d orbital is so much larger then the 4s orbital and the 3d orbital only needs one more electron to be filled, the 3d orbital pulls an electron from the 4s orbital to fill this empty space.This is also according to Heisenberg’s Uncertainty principle in which movements of electrons are usually unpredictable that’s why it it has a tendency to occupy other shells. (https://terpconnect.umd.edu/) 2. How do valence electrons involve in chemical reactions. Valence electrons are the electrons of elements which are either donated or accepted by a specific element to make a stable compound. Combining elements to make a new compound is an example of chemical reaction. KEY TO CORRECTIONS Activity 3: Skill-building Activities (with answer key) (18 mins + 2 mins checking) A. Answer the given items in then space provide. 1. What are the four (4) quantum numbers? What are the abbreviations of each quantum number? This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ Prinicpal (n), Azimuthal/Angumar momentum (l), Magnetic(ml) and Spin (ms) 2. Distinguish between a principal energy level and a sublevel. The principal energy level is related to the average distance from the nucleus.It is responsible for the SIZE of the electron cloud. Sublevel/subshell is a set of equal-energy orbitals within a principal energy level. It is responsible for the SHAPE of the electron cloud. 3. Distinguish between a sublevel and an orbital. Sublevel/subshell is a set of equal-energy orbitals within a principal energy level. It is responsible for the SHAPE of the electron cloud. An atomic orbital is a specific region of a sublevel containing a maximum of two electrons. 4. How many electrons can an orbital carry? 2 electrons 5. What are the shapes of the given sublevels? A. s = speherical B. p = dumbbell C. d = cloverleaf/x-shaped D. f = complex 6. What is the maximum number of electrons in each of the following energy levels? a. n = 1 2(1)2=2 b. n = 2 2(2)2=8 c. n = 3 2(3)2=18 7. For any given principal energy level, what is the maximum number of electrons that can exist in the following subshells? b. s=2 b. p = 6 c. d = 10 8. State the Pauli exclusion principle. No two electrons can have the same set of quantum numbers. 12. State Hund’s rule. Determine whether the following orbital diagrams violate Hund’s rule. You should write “Follow the Hund’s rule” or “Does not follow the Hund’s rule” after each item to determine if the example followed or violated the Hund’s rule. Orbitals are filled up singly before pairing a. Follows the Hund’s rule b. Does not follow the Hund’s rule c. Does not follow the Hund’s rule This document is the property of PHINMA EDUCATION  Course Code: CHE 026 Teacher’s Guide Module #4 Name:_____________________________________________________________ Class number: _______ Section: ____________ Schedule: _____________________________________ Date: _______________ 9. Write the electronic configuration and valence electron of each element: A. Na (Z=11) 1s22s22p63s1 B. S (Z=16) 1s22s22p63s23p4 C. Ca (Z=20) 1s22s22p63s23p64s2 10. Which of the following electron configurations are not possible? Why? a. 1s2 1p2 c. 2s2 2s2 2p6 2d1 2 2 2 b. 1s 2s 2p d. 1s2 2s3 a. 1s2 1p2 is not possible c. 2s2 2s2 2p6 2d1 is not possible since it did not start with the lowest principal energy and the electronic configuration sequence is wrong. d. 1s2 2s3 is not possible since the s subshell can only carry 2 electrons. This document is the property of PHINMA EDUCATION