Science 9 Q2 Module 1 Week 1 PDF

9 Science Quarter 2-Module 1 Week 1 (Electronic Structure of Matter) 1 Science - Grade 9 Quarter 2 - Module 1 (Electronic Structure of Matter) Second Edition, Revised 2021 Republic Act 8293, section 176 states that: No copyright shall subsist in any w...

9 Science Quarter 2-Module 1 Week 1 (Electronic Structure of Matter) 1 Science - Grade 9 Quarter 2 - Module 1 (Electronic Structure of Matter) Second Edition, Revised 2021 Republic Act 8293, section 176 states that: No copyright shall subsist in any work of the Government of the Philippines. However, prior approval of the government agency or office wherein the work is created shall be necessary for exploitation of such work for profit. Such agency or office may, among other things, impose as a condition the payment of royalty. Borrowed materials (i.e., songs, stories, poems, pictures, photos, brand names, trademarks, etc.) included in this book are owned by their respective copyright holders. Every effort has been exerted to locate and seek permission to use these materials from their respective copyright owners. The publisher and authors do not represent nor claim ownership over them. Published by the Department of Education – Cebu City Division Schools Division Superintendent: Rhea Mar A. Angtud, EdD Development Team of the Module Writer/Compiler: Joan Jennel M. Cayme, T III, Inayawan National High School Content Editors: Neil Andrian A. Angtud, HT I, Assisting Principal CCDCAGMNHS Florenda G. Yap, MT II, Assistant Principal, Apas National High School Language Editor: Wilma Y. Villaflor, Principal III, Don Vicente Rama Memorial ES Management Team: Dr. Rhea Mar A. Angtud, Schools Division Superintendent Dr. Bernadette A. Susvilla, Assistant Schools Division Superintendent Mrs. Grecia F. Bataluna, CID Chief Dr. Raylene S. Manawatao, EPS-Science Mrs. Vanessa L. Harayo, EPS-LRMDS Printed in the Philippines by: Department of Education: Division of Cebu City Office Address: Imus Avenue, Cebu City Telephone Nos.: 032) 255-1516/ (032) 253-9095 E-mail Address: [email protected] 1  Lesson Bohr’s and Quantum Mechanical Model of an 1 Atom Hello there our young scientist! Welcome to the first module of this quarter. Before moving forward, please be guided with what’s waiting for you ahead! WHAT I NEED TO KNOW Learning Competency: Explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons. In Grade 8, you have learned that Rutherford’s atomic model which pictures the atom as mostly empty space and its mass is concentrated in the nucleus, where you find the protons and the neutrons. This model has worked well during his time, but it was only able to explain a few simple properties of atoms. However, it could not explain why metals or compounds give off characteristic colors when heated in a flame. A model different from Rutherford’s atomic model is necessary to describe the behavior of atoms. Niels Bohr refined Rutherford’s model of an atom. Based on his experiments, Bohr described the electron to be moving in definite orbits around the nucleus. Much later, scientists discovered that it is impossible to determine the exact location of electrons in an atom. Within this module, you will learn about the evidence that Bohr used to explain his model of the atom. You will do a task that will help you understand that there is a certain portion of space around the nucleus where the electron is most likely to be found. In addition, you will know more about the present model of the atom, which is called the quantum mechanical model. It is important for you to understand that the chemical properties of atoms, ions and molecules are related to how electrons are arranged in these particles of matter. You will find out the answers to the following questions as you perform the activities in this module. 2  How does the Bohr atomic model differ from Rutherford’s model? What is the basis for the quantum mechanical model of the atom? How are electrons arranged in the atom? Excited to discover the answers to the above cited questions? Before you start, answer the following pre-assessment first. WHAT I KNOW Direction: Choose the letter of the correct answer and write your answer on a separate sheet of paper. 1. On the basis of Rutherford’s model of an atom, which sub- atomic particle is present in the nucleus of an atom? A. proton only C. proton and neutron B. proton and electron D. neutron and electron 2. If the first and second energy levels of an atom are full, then what would be the total number of electrons in the atom? A. 6 B. 8 C. 10 D. 18 3. Which atomic model is proposed by Schrodinger? A. nuclear model C. raisin bread model B. planetary model D. quantum mechanical model 4. Which electron transition results in the emission of energy? A. 1s to 2s C. 3p to 3s B. 2s to 2p D. 3p to 4p 5. The symbol “n” in the Bohr theory of atomic structure refers to the _______ A. energy of electron B. total energy of the atom C. orbit in which an electron is found D. number of electrons in the energy level 6. Which of the following sublevels is correctly designated? A. 1p5 B. 2p6 C. 3f 9 D. 3d11 7. How many orbitals are in the third principal energy level? A. 3 B. 6 C. 9 D. 12 3  8. Which configuration is possible in an excited state of an electron? A. 1H : 1d1 C. 11Na : 1s2 2s2 2p6 3d1 B. 2He : 1s2 D. 10Ne : 1s2 2s2 2p5 3s1 9. What are the orbitals in the fifth principal energy level? A.s orbital C. s, p, d orbitals B. s, p orbitals D. s, p, d, f and g orbitals 10. For a neutral atom with the electron configuration of 1s 2 2s2 2p5 3s1, which statement is FALSE? A. The atomic number is ten. B. The atom is in ground state. C. The atom is in an excited state. D. The 1s and 2s orbitals are filled. 11. Rutherford’s model of the atom concentrated on the nucleus while Bohr’s model focused on the ______ A. electrons C. protons B. neutrons D. quarks 12. Which of the following statements about electrons refer to the Bohr Model of an atom? A. move at a very high velocity around the nucleus B. exist at different energy levels in orbit around the nucleus C. can move between energy levels when they gain or lose energy D. all of the above 13. The following are rules/principles used in arranging the electrons around the nucleus of an atom EXCEPT A. Hund’s Rule B. Aufbau Principle C. Pauli Exclusion Principle D. Heisenberg’s Uncertainty Principle 14. If a neutral atom has 18 electrons, what is the highest principal energy level that its outermost electrons will occupy? A. 1 B. 2 C. 3 D. 4 4 15. Which of the following is the correct electron configuration for neon ( 10Ne)? A. 1s2 2s2 2p6 C. 1s4 2s2 3s2 2p2 B. 1s2 2s2 3s2 2p4 D. 1s2 2s2 3s1 2p5 WHAT’S IN Let’s take a trip down the memory lane on the development of the atomic theory by completing the table below. Use a separate sheet of paper. Atomic Proponent Illustration Description Model Solid The atom is a solid sphere that Sphere could not be divided into smaller Model particles. The atom has negatively-charged J.J. Thomson electrons embedded within a positive sphere. Most of the atom’s mass is in the positively charged nucleus. Far Nuclear from the nucleus are the Model negatively charged electrons. Source: https://en.wikipedia.org/wiki/ Rutherford_model Guide Question: 1. Why were these atomic models being rejected? __________________________________________________________________ __________________________________________________________________ Since these models were not accepted by the scientific community, what could have transpired and led them to formulate a new atomic model? You will find out in the next activity. 5  WHAT’S NEW Activity 1: The Flame Test Objectives: ❖ Determine the characteristic colors that metal salts emit; and ❖ Relate the colors emitted by salt metals to the structure of the atom. Procedure: NOTE: Do not perform this activity at home. Just read the situation given below. 1. Assuming that you have observed your teacher performing the flame test experiment. Situation: The teacher will place metals salts on five watch glasses and then he will add 2 to 3 drops of 3 M hydrochloric acid. Using another watch glass, he will pour about 3-5 mL of ethanol, light with a match and observe the color of the flame (this will serve as the reference for comparison of the flame color). Next step, the teacher will now light the other metal salts in each watch glass. 2. Study the diagram and table below: Source: Science 9 LM Table 1: Color of flame of metal salts 6  Metal Salt Tested Element Producing Color Color of the Flame Boric Acid boron green Calcium chloride calcium orange Sodium chloride sodium yellow orange Potassium chloride potassium light violet Copper (II) sulfate copper blue green Guide Questions: Write your answers on a separate sheet of paper. 1. Why are there different colors emitted? 2. What particles in the heated compounds are responsible for the production of the colored light? 3. How did the scientists explain the relationship between the colors observed and the structure of the atom? You have observed that each of the substances that has been tested showed a specific color of the flame. Why do certain elements give off light of specific color when heat is applied? These colors given off by the vapors of the elements can be analyzed with an instrument called the spectroscope. Figure 1. An Atomic Spectroscope Source: Science 9 LM A glass prism separates the light given off into its component wavelength. The spectrum produced appears as a series of sharp bright lines with characteristic colors and wavelengths on a dark background instead of being continuous like the rainbow. We call this series of lines the atomic spectrum of the element. The color, the number 7 and position of lines produced is called the “fingerprint” of an element. These are all constant for a given element. See figure 2. Figure 2. Atomic Spectra of H, Na and Ne Source: Science 9 LM How did Niels Bohr explain what you have observed in Activity 1 and the findings about the elements in a spectroscope? Individual lines in the atomic spectra of elements indicate definite energy transformations within the atom. Bohr considered the electrons as particles moving around the nucleus in fixed circular orbits. These orbits are found at definite distances from the nucleus. The orbits are known as the energy levels, where n is a whole number 1, 2, 3…and so forth. Electrons in each orbit have a definite energy, which increases as the distance of the orbit from the nucleus increases. As long as the electron stays in its orbit, there is no absorption or emission of energy. As shown in figure 3, when an electron of an element absorbs extra energy (from the flame or electric arc), this electron moves to a higher energy level. At this point the electron is at its excited state. Once excited, the atom is unstable. The same electron can return to any of the lower energy levels releasing energy in the form of light with a particular color and a definite energy or wavelength. Bohr’s model explained the appearance of the bright line spectrum of the hydrogen atom but could not explain for atoms that have more than one electron. Figure 3. Excited state of an electron Source: Science 9 LM 4. Explain how the observation in Activity 1 relates to Bohr’s model of the atom. 8  5. Which illustration below represents the energy of the electron as described by Bohr? Explain. The energy levels of electrons are like the steps of a ladder. The lowest step of the ladder corresponds to the lowest energy level. A person can climb up and down by going from step to step. Similarly, the electrons can move from one energy level to another by absorbing or releasing energy. Energy levels in an atom are not equally spaced which means that the amount of energy is not the same. The higher energy levels are closer together. If an electron occupies a higher energy level, it will take less energy for it to move to the next energy level. As a result of the Bohr model, electrons are described as occupying fixed energy levels at a certain distance from the nucleus of an atom. However, Bohr’s model of the atom was not sufficient to describe atoms with more than one electron. The way around the problem with Bohr’s model is to know the arrangement of electrons in atoms in terms of the probability of finding an electron in certain locations within the atom. In the next activity, you will use an analogy to understand the probability of finding an electron in an atom. Activity 2: Predicting the Probable Location of an Electron Objective: ❖ Describe how it is likely to find an electron in an atom by probability. Materials: one sheet of short bond paper, pencil or colored marker with small tip, compass, graphing paper, one-foot ruler Procedure: 1. Draw a dot on the center of the sheet of paper. 9  2. Draw 5 concentric circles around the dot so that the radius of each circle is 1.0 cm, 3.0 cm, 5.0 cm, 7.0 cm and 9.0 cm from the dot. 3. Tape the paper on the floor so that it will not move. 4. Stand on the opposite side of the target (the center which represents the nucleus of an atom). Hold a pencil or marker at chest level above the center of the circles you have drawn. 5. Drop the pencil or marker so that it will leave 100 dots on the circles drawn on paper. 6. Count the number of dots in each circle and record the number on the data table. 7. Calculate the number of dots per square centimeter (cm2). 8. Using a graphing paper, plot the average distance from the center on the x-axis and number of dots per square centimeter on the y-axis. Data Table: Circle Average Area of Difference Number Number Percent Number Distance Circle of Areas of Dots of Dots Probability 2 2 from the (cm ) Two in Circle per cm of Finding Center Consecutive (E/D) Dots (%) Circles (A) (B) (C) (D) (E) (F) (G) 1 1.0 3.14 25.13 (example (example (example only) only) only) 5 0.1920 19.20 2 3.0 28.27 50.27 3 5.0 78.54 75.40 4 7.0 153.94 100.53 5 9.0 254.47 125.66 10 Guide Questions: Write your answers on a separate sheet of paper. 1. What happens to the number of dots per unit area as the distance of the dots go farther away from the center? 2. Determine the percent probability of finding a dot in each of the circles drawn on the target by multiplying the number of dots /cm 2 (column D) by the total number of dots (100). For example, in circle 1 (A): Percent Probability = no. of dots/cm2 × 100 = (0.1920/100) ×100 = 19.20% 3. Based on your graph, what is the distance with the highest probability of finding a dot? Show this in your graph. 4. How many dots are found in the area where there is the highest probability of finding dots? 5. How are your results similar to the distribution of electrons in an atom? This activity demonstrates what scientists found out that it is not possible to know the exact position of the electron. So Bohr’s idea that electrons are found in definite orbits around the nucleus was rejected. Three physicists led the development of a better model of the atom. Physicists Contribution Louie de Broglie Proposed that the electron (which is also thought as a particle) could also be a wave. Erwin Schrodinger Used Broglie’s idea to develop a mathematical equation to describe the hydrogen atom. Werner Karl Heisenberg Discovered that for a very small particle like the electron, its location cannot be exactly known and how it is moving. This is called the uncertainty principle. These scientists believed that there is only a probability that the electron can be found in a certain volume in space around the nucleus. This volume or region of space around the nucleus where the electron is most likely to be found is called an 11 atomic orbital. Thus, we could only guess the most probable location of the electron at a certain time to be within a certain volume of space surrounding the nucleus. The quantum mechanical model of the atom comes from the mathematical solution to the Schrodinger equation. The quantum mechanical model views an electron as a cloud of negative charge having a certain geometrical shape. This model shows how likely an electron could be found in various locations around the nucleus. However, the model does not give any information about how the electron moves from one position to another. Figure 4 shows that the Figure 4: Average Distance of electrons having darker an area, the greater is the high and low energies probability of finding the electron Source: Science 9 LM in that area. The quantum mechanical model also gives information about the energy of the electron. This model also describes the region of space around the nucleus as consisting of shells. These shells are also called principal or main energy levels. The principal energy levels or shells may have one or more sublevels. These sublevels are assigned with letters: s, p, d, f and g as shown in Table 2. Table 2: Principal Energy Levels and Sublevels of Electrons. Princi Numb Maximum Number Total Number of pal er of of Electrons Per Electrons Type of Sublevel Energ Suble Orbital and Number of y vels Orbitals Level (n) 1 1 1s (1 orbital) 1s = 2 2 2 2 2s (1 orbital), 2p (3 2s = 2 8 orbitals) 2p = 6 3 3 3s (1 orbital), 3p (3 3s = 2 18 orbitals), 3d (5 3p = 6 orbitals) 12  3d = 10 4 4 4s (1 orbital), 4p (3 4s = 2 32 orbitals), 4d (5 4p = 6 orbitals), 4d = 10 4f (7 orbitals) 4f = 14 5 5 5s (1 orbital), 5p (3 5s = 2 50 orbitals), 5d (5 5p = 6 orbitals), 5d = 10 5f (7 orbitals), 5g (9 orbitals) 5f = 14 5g = 18 As shown in Table 2, the principal quantum number always equals the number of sublevels within that principal energy level. The maximum number of electrons that can occupy a principal energy level is given by the formula 2n2, where n is the principal quantum number. Guide Questions: 6. Based on Table 2, how many types of orbitals are in principal energy level three (3)? 7. How many atomic orbitals are in the highest sublevel of principal energy level three (3)? Figure 5. Shapes of s Orbital and p orbitals Source: Science 9 LM 13  Orbitals have specific energy values. They have particular shapes and direction in space. The s orbitals are spherical, and p orbitals are dumbbell-shaped, as shown in Figure 5. Because of the spherical shape of the s orbital, the probability of finding an electron at a given distance from the nucleus in an s orbital does not depend on direction, unlike the three kinds of p orbitals which are oriented along the x, y and z axes. So their different orientations in space are px, py and pz. The shapes of other orbitals (d and f) were derived from complex calculation and will not be discussed in this module. In an atom, the electrons and the nucleus interact to make the most stable arrangement possible. The way in which electrons are distributed in the different orbitals around the nucleus of an atom is called the electron configuration. WHAT IS IT Electron Configuration Electronic configuration or also known as electronic structure, is the arrangement of electrons in energy levels around an atomic nucleus. In terms of a more refined, quantum mechanical model, the energy levels are subdivided into a set of orbitals, each of which can be occupied by no more than a pair of electrons. When assigning electrons to orbitals, we must follow a set of three rules: the Aufbau Principle, the Pauli-Exclusion Principle and Hund's Rule. Rules for Electron Configuration ❖ Aufbau Principle- states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. ❖ Pauli Exclusion Principle- states that no two electrons in the same atom can occupy the same orbital and the two electrons in the same orbital must have opposite spins. ❖ Hund’s Rule- no two electrons can be identified by the same set of quantum numbers. For this module, you will focus on applying the Aufbau Principle only for the electron configuration. These are the steps in writing the electron configuration. STEPS: 1. Determine the number of electrons that an atom has. 2. Fill the s orbital in the first energy level (the 1s orbital) with the first two electrons. 3. Fill the s orbital in the second energy level (the 2s orbital) with the second two electrons. 4. Always remember the total number of electrons that each orbital can carry. Refer to figure 6. 5. You can only start filling in the next orbital once the lower orbital is already completely filled. 14  6. To be guided, you may refer to this mnemonic device (figure 7). In filling in the orbitals with electrons, just follow the direction of the arrow for the sequence of the orbitals. Figure 6. Maximum Number of Electrons per Orbital Figure 7 Sources: https://socratic.org/questions/please-solve http://astan.lk/al_virtualclassroom/electron-stability/ EXAMPLE: NOTE: For a neutral atom, its atomic number is equivalent to its number of protons and electrons. Element Symbol No. of Electron Configuration and Atomic Electron Number (s) Hydrogen 1H 1 1s1 Sodium 11Na 11 1s2 2s2 2p6 3s1 Vanadium 23V 23 1s2 2s2 2p6 3s2 3p6 4s2 3d3 Bromine 35Br 35 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 The exponent 1 in the orbital 1s1 of hydrogen means that there is only 1 hydrogen electron that must be distributed. For sodium with 11 electrons, the sum of its atomic orbital exponents 2, 2, 6 and 1 is equivalent to 11. Therefore, the sum of its exponents must be equal to the number of electrons of that specific atom. 15  WHAT’S MORE Activity 3: Writing Electron Configurations Objective: ❖ Write the electron configuration of the first 15 elements. Materials: pen and paper Procedure: 1. After learning how to write electron configurations using the Aufbau Principle, complete the table below and answer the guide questions. 2. Copy and answer the table on a separate sheet of paper. Element Symbol No. of Electron Configuration and Electron Atomic s Number example: Hydrogen 1H 1 1s1 Helium 2He Lithium 3Li Beryllium 4Be Boron 5B Carbon 6C Nitrogen 7N Oxygen 8O Fluorine 9F Neon 10Ne Sodium 11Na Magnesium 12Mg Aluminum 13Al Silicon 14Si Phosphorus 15P 16 Guide Questions: 1. Do you see patterns in the distribution of their electrons? 2. What are these patterns that you have observed? WHAT I HAVE LEARNED Upon learning new things about the structure of the atom, determine whether the following statements are TRUE or FALSE. Use a separate sheet of paper. 1. Bohr’s atomic model describes the atom like a solar system. Where the electron is found only in specific circular paths, or orbits, around the nucleus. 2. The quantum mechanical model of the atom describes the atom as having a nucleus at the center around which the electrons move. This model describes a region in space where the electron is most likely to be found. 3. The way in which electrons are distributed in the different orbitals around the nucleus of an atom is called the electron configuration. Filling the electrons starts from the highest energy level to the lower energy level. WHAT I CAN DO Using indigenous materials found in your home, make a model that represents the Quantum Mechanical Model of the atom. Refer to the rubric below: CATEGORY 25 20 15 10 Over-all It is clear that Somewhat neat More time and Needs to be Appearance time was taken and attractive, effort could neatly improved to make the but needs some have been put and lacks color structure neat, work. into improving attraction. colorful and the structure’s attractive. appearance. On-Time Whole project Project was Project was Project was was complete turned in a day turned in two turned in three and turned in by late. days late days late. the assigned due date. Model Project shows Project shows Project shows Model has more the the the than 3 errors. characteristics characteristics characteristics 17  of the Quantum of the Quantum of the Quantum Mechanical Mechanical Mechanical Model of the Model of the Model of the Atom and is Atom but it has Atom but it has labeled 1 missing label. 2 missing correctly. labels. Creativity and Model shows Model shows Model is 2-D Project is poorly Explanation on that learner learner used and copied from made. Little how it was researched standard class notes. thought was put made examples and materials in 2-D Very little into model or it used creative or 3-D and imagination was was completed ideas and he/she can put into the at the last materials. explain how it model. minute. Learner Model is 3-D. was made. Explanation is is unable to Explanation on Some creativity basic. explain as how the model evident. taught in class. was made is clear and easy to understand. Total Points and Comments ASSESSMENT Write the letter of the correct answer on a separate piece of paper. 1. Who proposed the probability that electrons will be found in certain locations around the nucleus of an atom? A. Neils Bohr C. Erwin Schrodinger B. Ernest Rutherford D. J.J. Thomson 2. Which of the following statements is NOT TRUE of the atomic model of Bohr? A. An electron can absorb or emit a quantity of radiation. B. The energy of the electron in a given orbit is not fixed. C. The hydrogen is made up of a positively charged nucleus. D. The electron revolves around the nucleus in a circular orbit. 3. Which orbital designation has the highest energy? A. 2s B.2p C. 3s D. 3d 4. Which statement is INCORRECT? A. Orbital is a region in an atom where an electron can be found. B. An electron can emit energy when it jumps to a higher energy level. C. An electron can absorb energy when it jumps to a higher energy level. D. Filling of electrons in an atom starts from the low energy level to the highest energy level. 18  5. What occurs when an electron moves from a high energy level to a low one? A. The atom moves faster. B. Colored light is given off. C. Colored light is absorbed. D. Another electron goes from a low energy level to a high one 6. Which combination describes the flame color of the compound when heated? A. boric acid- red C. potassium chloride- blue B. copper (II) sulfate- violet D. sodium chloride- yellow orange 7. What happens to the energy of an electron if it moves from one energy level to another closer to the nucleus? A. stays constant C. gains energy B. becomes and ion D. loses energy 8. Which principle states that electrons occupy orbitals of lowest energy first before filling in the higher energy levels? A. Hund’s Rule C. Schrodinger’s Equation B. Aufbau Principle D. Pauli Exclusion Principle 9. What shape are p orbitals? A. cloverleaf shaped C. hybrid structure B. dumbbell shaped D. spherical shape 10. The exponent (superscript) in an electron configuration such as in 2p5 tells us the _____ A. number of positions in that orbital B. distance from the nucleus or level C. number of electrons of that orbital D. number of electrons in the nucleus 11. During a flame test, a copper (II) salt produces a bluish green color of flame. This color is produced when electrons in the excited copper II atoms _____ A. are lost by the atoms B. are gained by the atoms C. move to higher energy states within the atoms D. return to lower energy states within the atoms 12. Which of the following best describes the quantum mechanical model of the atom? A. It is the currently accepted atomic model. B. It was based on Schrodinger’s mathematical calculations. C. It describes an electron probability distribution through orbitals. D. all of the above 13. The f orbitals can have how many orientations? A. 1 B. 3 C. 5 D. 7 14. Which electron movement is accompanied by a release of energy? A. 1s to 2s C. 3p to 3s B. 2s to 2p D. 4px to 4pz 15. Which of the following statements about the orbital is FALSE? A. It can be unoccupied. B. It can be occupied by one electron. C. It can be occupied by two electrons. D. It can be occupied by more than two electrons. 19 References Books: □ Brown, Theodore et al (2009) Chemistry: The Central Science 11th Edition Pearson Education, South Asia PTE. LTD. Singapore □ Carmichaels, H. (1983). Laboratory Chemistry,. Columbos, Ohio: Merill Publishing Co. □ Department of Education, Culture and Sports (2004). Chemistry: Science and Technology Textbook for 3rd Year. (Revised Edition). Quezon City. □ Kotz, John C. et al (2010) Chemistry and Chemical Reactivity Enhanced Ed. Canada: Brooks/Cole Cengage Learning □ LeMay, E. et al (1996) Chemistry Connection to Our Changing World, Teacher Edition. New Jersey: Prentice Hall, Inc. □ Mendoza, E & Religioso, T. (2001). Chemistry. Quezon City: Phoenix-SIBS Publishing House, Inc. □ Silberberg, Martin S.(2009) Chemistry: The Molecular Nature of Matter and Change 5th Edition, International Ed. 2010 McGraw-Hill Companies, Inc. New York □ Smot, R.C. et al (1995) Chemistry Wraparound Teacher's Edition, Glencoe/McGraw-Hill, Merill Publishing Co., Ohio □ University of the Philippines National Institute for Science and Mathematics Education Development. (2001). Practical Work in High School Chemistry: Sourcebook for Teachers. Quezon City □ Wilbraham, A.C. et al (1997). Chemistry Expanded (4th Edition) Teacher Edition. California: Addison-Wesley Publishing Co. Electronic Sources: □ The ekShiksha Team, Affordable Solutions Lab (ASL), Indian Institute of Technology, Bombay, India. Matter in our Surrounding. Retrieved: October 3, 2013. http://www.it.iitb.ac.in//ekshiksha/eContent-Show.do?document!d=88 Internet Sources: □ https://www.sparknotes.com/chemistry/fundamentals/atomicstructure/section1 / □ https://www.clutchprep.com/chemistry/practice-problems/106936/list-the-four- different-sublevels-and-given-that-only-a-maximum-of-two-electrons □ https://en.wikipedia.org/wiki/List_of_chemistry_mnemonics 20  21 What’s In Atomic Model Proponent Illustration Description Solid Sphere John Dalton The atom is a solid sphere that could Model not be divided into smaller particles. Plum- The atom has negatively-charged Pudding J.J. Thomson electrons embedded within a Model positive sphere. Most of the atom’s mass is in the Ernest Nuclear Model positively charged nucleus. Far from Rutherford the nucleus are the negatively charged electrons. Guide Question: 1. (Answers may vary) Possible answer: These atomic models were rejected because it wasn’t not able to support the new discoveries about the structure of the atom. Activity 1: The Flame Test 1. Metal salts emitted different colors because of the absorption of heat from the flame. 2. The outermost particles in the metallic element are responsible for the production of the colored light. 3. The colors observed is an indication that definite energy transformations occur inside the atom emitting light. It follows that electrons must occupy orbits of fixed energy. 4. The electrons are moving around the nucleus in circular orbits. When an electron absorbs extra energy from an outside source (flame), the electron moves to a higher orbit. The colored light is emitted when the electron falls back to a lower orbit. The light is the difference between the energies of the two orbits involved. 5. The energy levels of electrons are like the steps of a ladder. The lowest step of the ladder corresponds to the lowest energy level. The electrons can move from one energy level to another by absorbing or releasing energy. Energy levels in an atom are not equally spaced which means that the amount of energy is not the same. The higher energy levels are closer together. If an electron occupies a higher energy level, it will take less energy for it to move to the next energy level. Activity 2: Predicting the Probable Location of an Electron *Answers may vary in the table based on their performance of the activity. Guide Questions: 1. The number of dots increases abruptly and then decreases as the dots go farther from the center. 2. Answer will vary 3. Answer will vary 4. Answer will vary 5. The results of the activity are similar to the structure of the atom because the probability of finding an electron (dot) increases abruptly then decreases as it goes farther from the nucleus (target). 6. There are 3 types of orbitals (s, p and d) in the principal energy level 3. 7. There are 5 atomic orbitals in the highest sublevel of the principal energy level 3. Answer Key:  22 E-mail Address: [email protected] Telephone Nos.: (063)255-1516, (032)253-9095 Office Address: Imus Avenue, Cebu City Department of Education: Cebu City Division For inquiries or feedback, please write or call: Activity 3: Electron Configurations Element Symbol No. of Electron Configuration and Atomic Electrons Number example: Hydrogen 1H 1 1s1 Helium 2He 2 1s2 Lithium 3Li 3 1s2 2s1 Beryllium 4Be 4 1s2 2s2 Boron 5B 5 1s2 2s2 2p1 Carbon 6C 6 1s2 2s2 2p2 Nitrogen 7N 7 1s2 2s2 2p3 Oxygen 8O 8 1s2 2s2 2p4 Fluorine 9F 9 1s2 2s2 2p5 Neon 10Ne 10 1s2 2s2 2p6 Sodium 11Na 11 1s2 2s2 2p6 3s1 Magnesium 12Mg 12 1s2 2s2 2p6 3s2 Aluminum 13Al 13 1s2 2s2 2p6 3s2 3p1 Silicon 14Si 14 1s2 2s2 2p6 3s2 3p2 Phosphorus 15P 15 1s2 2s2 2p6 3s2 3p3 Guide Questions: 1. Yes 2. Filling the orbitals with electrons starts from the lowest energy level to the highest energy level. What I Have Learned 1. True 2. True 3. False