Science 9 Past Paper PDF

**Day/Time** **Learning Competency** **Learning Tasks** **Mode of Delivery** **4 Hours every week** **Explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons. (S9MT-IIa22)** **SLM9 Quarter 2, Week 1 page 1- 8** **- Recall lesson about Atoms (...

**Day/Time** **Learning Competency** **Learning Tasks** **Mode of Delivery** **4 Hours every week** **Explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons. (S9MT-IIa22)** **SLM9 Quarter 2, Week 1 page 1- 8** **- Recall lesson about Atoms (Get Refreshed)** **- Identify different colors of light emitted by firecrackers (Get** **Acquainted)** **- Perform Activity 1 about the probability of finding an electron** **(Get Involved)** **- Identify the different physicists that developed Planetary Model** **and Quantum Mechanical Model of an Atom (Get Informed)** **- Classify Main Energy Level, sublevel, and their maximum** **number of electrons** **- Define atomic orbital** **- Explain Quantum Numbers and Pauli Exclusion Principle** **- Recognize each set of quantum numbers** **- Write all answer on the attached Learner's Activity Sheet (LAS)** **Modular Distance Learning** **Personal submission of modules by parents/guardians to the class adviser/ co -- adviser** **every Wednesday at** **8:00AM -12:00NN** **1:00 -- 5:00PM.** **QUANTUM MECHANICAL MODEL OF AN ATOM**  **Get Started** Hello dear students! Happy New Year! Congratulations for successfully completed the first quarter modules and now Chemistry will lead the second quarter. In Grade 8, you have learned about the 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. However, it could not explain why metals or compounds of metals give off characteristics colors when heated in a flame, or why objects when heated to much higher temperatures first glow to dull red, then to yellow, and then to white. A model different from Rutherford's atomic model is necessary to describe the behavior of atoms. In this module, you are expected to **explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons**. Specifically, the Neils Bohr model will be discussed on how this model refined Rutherford's model of atom. Furthermore, you will gain an understanding about the quantum mechanical model of the atom. This is important so you will understand that the characteristics of matter are related to how electrons are distributed inside the atoms I know that you are excited to learn all of these. Prepare yourself as well as your pen. Turn to the next pages and move your way until the end of this module. Goodluck!  ***Get Refreshed (Elicit)***. Let's check first if you still remember your lesson about the 3 subatomic particles. Arrange the scrambled words to answer the questions. Complete the tables below. Try to answer then compare your answers with the answers written on the key to corrections. No need to copy and submit your answers on this part. 1. Name the three subatomic particles and describe them by their charges and location in the atom. -------------------- ------------------------ ------------ ------------ ----------------------------- ---------------- **Scrambled Word** **Subatomic Particle** **Symbol** **Charge** **Location in the Nucleus** **Discoverer** **1. TEUNNORS** **2. RECLEONTS** **3. NORPOTS** -------------------- ------------------------ ------------ ------------ ----------------------------- ---------------- Look at your periodic table. Fill up the table below by supplying the missing data. -------------- ---------------- ------------------- ----------------- ----------------------- ------------------------ ------------------------- **Elements** **Symbol** **Atomic number** **Atomic mass** **Number of protons** **Number of neutrons** **Number of electrons** Potassium **K** **19** **39** **\_\_\_\_\_** **\_\_\_\_\_** **\_\_\_\_\_** Mercury **\_\_\_\_\_** **\_\_\_\_\_** **\_\_\_\_\_** **80** **121** **\_\_\_\_\_** sodium **Na** **\_\_\_\_\_** **23** **\_\_\_\_\_** **\_\_\_\_\_** **11** -------------- ---------------- ------------------- ----------------- ----------------------- ------------------------ -------------------------  **Get Acquainted (Engage)** Have you seen striking display of fireworks last New Year's Eve? Have you observed the different colors of light emitted by these fireworks? Do you know what is responsible for this array of colors? Would you believe that this is due to the arrangement of electrons within the atoms? Excited to discover more about the atom? Firework effects are produced by the combustion of explosive materials present in fireworks. These explosive materials are also called metal salts. These metal salts are responsible for emitting characteristic colors of light as shown in Table 1. -------------------------------- -------------------------- -------------------- **Metal salt** **Element giving color** **Color of flame** Sodium Nitrate (NaNO~3~) Sodium (Na) Yellow Barium Chloride (BaCl~2~) Barium (Ba) Yellow - green Boric acid (H~3~BO~3~) Boron (B) green Calcium Chloride (CaCl~2~) Calcium (Ca) Orange Copper (II) sulfate (CuSO~4~) Copper (Cu) Blue Lithium Carbonate (Li~2~CO~3~) Lithium (Li) red Sodium chloride (NaCl) Sodium (Na) Orange -yellow -------------------------------- -------------------------- -------------------- Table 1. Color emitted of some metal salts and its element responsible for its color Analyze the information given in the table 1 and answer the questions. Write your answers on you're the attached LAS. 1\. What can you conclude about the elements and the color of flame produced? 2\. If sodium nitrate is change to another salt of potassium, ex. Potassium chloride (Sylvite/ Sylvine), do you think the color of the flame would be the same or different? Why?  ***Get Informed (Explain)*** When you were in Grade 8, you have studied about the colors of light as shown in the different colors of the rainbow that each color of light has a specific frequency, wavelength, and energy. We concluded that among the visible light, red light has the longest wavelength and has the lowest energy and frequency, and violet light has the shortest wavelength and has the highest energy and frequency.. When compounds of different elements as shown in Table 1 are heated over a flame, it comes to a point where the hot gaseous atom begins to emit light of a definite color. Analysis of light given off by the vapors of elements can be done more precisely with an instrument called **spectroscope**. With the use of spectroscope, one can detect a series of narrow lines or line spectrum on the light given off by an element. The spectral lines suggest different energy levels in an atom. The color, number, and position of lines produced is called the "fingerprint "of an element. These are all constant for a given element as shown in the atomic spectra of Hydrogen (H), Sodium (Na) and Neon (Ne). Individual lines in the atomic spectra of elements indicate definite energy transformations within the atom. Neils Bohr, a Danish physicist proposed a theory about the atomic structure which explained the line spectra of elements. He considered the electrons as particles moving around the nucleus in fixed circular obits. These orbits are found at definite distances from the nucleus. The orbits are known as the energy levels, n, where n is a whole number 1,2,3... and so forth. He proposed that the movement of the electrons is analogous to the planets orbiting around the sun. 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 the Figure 3, when an electron of an element absorbed extra energy (from a 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 has more than one electron. 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 are not the same. The higher energy electrons are described as occupying fixed energy levels at a certain distance from the levels are closer together. If an electron occupies a higher energy level, it will take less energy for it to move to the next higher 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 (1) electron. Activity 1 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 scientists led the development of a better model of the atom. **Physicists who developed Quantum Mechanical Model of an Atom** 1\. **Louie de Broglie** -- a French physicist who proposed a hypothesis to explain the theory of the atomic structure. \- He proposed that matter and slow-moving particles (electron s) has dual nature (particle and wave nature). The electron (which is thought of as a particle) could also be thought of as a wave.  2\. **Erwin Schrodinger** -- a Nobel Prize -- winning Austrian -- Irish physicist who contributed to the wave theory of matter and to other fundamentals of Quantum Mechanics, the Schrodinger equation provides a way to calculate the wave function of a system and how it changes dynamically in time. 3\. **Werner Karl Heisenberg** -- a Nobel Prize in Physics awardee himself is a German theoretical physicist and one of the pioneers of Quantum Mechanics and Nuclear physics. \- He 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 Heisenberg 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 **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 (Electron Cloud Model). However, the model does not give any information about how the electron moves from one position to another. Figure 4 shows that the darker an area, the greater is the probability of finding the electron in that area. The quantum mechanical model also gives information about the energy of the electron. The 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. --------------------------------------- ------------------------- --------------------------------------------------------------------------------------------------------------------- --------------------------------- **Principal or Main Energy Level, n** **Number of Sublevels** **Type of Sublevel and number of orbitals** **Maximum number of electrons** 1 or K 1 1s (1 orbital) 2 e- 2 or L 2 2s (1 orbital), 2p (3 orbitals) 8 e- 3 or M 3 18 e- 4 or N 4 4s (1 orbital), 4p (3 orbitals) 4d (5 orbitals), 4f (7 orbitals) 32 e- 5 or O 5 5s (1 orbital), 5p (3 orbitals) 5d (5 orbitals), 5f (7 orbitals) 5g (9 orbitals) 50 e- 6 or P 6 6s (1 orbital), 6p (3 orbitals) 6d (5 orbitals), 6f (7 orbitals) 6g (9 orbitals) 6h (11 orbitals) 72 e- 7 or Q 7 7s (1 orbital), 7p (3 orbitals) 7d (5 orbitals), 7f (7 orbitals) 7g (9 orbitals) 7h (11 orbitals), 7i (13 orbitals) 98 e- --------------------------------------- ------------------------- --------------------------------------------------------------------------------------------------------------------- --------------------------------- **Shapes of Atomic Orbitals:** 1\. **s -- orbital or sharp lines** -- sphere shape; it means that the electrons  in the s-orbital move around the nucleus in a spherical pattern. 2\. **p -- orbital or principal lines** -- dumb-bell shape (bilobate) which lie perpendicular to one another. 3\. **d -- orbital or diffuse lines** -- clover shape, quadrilobate (four lobed) 4\. **f -- orbital or fundamental lines** -- more complex shape but same with d-orbital if full but cut in half (eight lobes instead of four). **Quantum Number -** is a value that is used when describing the energy levels available to atoms and molecules. An electron in an atom or ion has four quantum numbers to describe its state and yield solutions to the Schrödinger wave equation for the hydrogen atom. There are **Four (4) Quantum Numbers**: 1\. **Principal Quantum Number (n)** -- refers to the main energy level of the orbital. n= 1 to 7 First Energy Level or K -- shell, n= 1 Fifth Energy Level or O -- shell, n= 5 Second Energy level or L -- shell, n= 2 Sixth Energy Level or P -- shell, n= 6 Third Energy level or M -- shell, n=3 Seventh Energy Level or Q -- shell, n= 7 Fourth Energy level or N -- shell, n=4 2\. **Secondary Quantum Number or Azimuthal or Angular momentum quantum number (ℓ)** \- describes the subshell or sublevel. s -- sublevel, ℓ = 0 p -- sublevel, ℓ = 1 d -- sublevel, ℓ = 2 f -- sublevel, ℓ=3 3\. **Magnetic Quantum Number (mℓ or m)** - describes the orbital of the subshell s -- sublevel, mℓ = 0 d -- sublevel, mℓ = -2, -1, 0, 1, 2 p -- sublevel, mℓ = -1, 0, 1 f -- sublevel, mℓ = -3, -2, -1, 0, 1, 2, 3 4\. **Spin Quantum Number (ms or s**) -- is a half -- integer value which describes the spin of the electrons either spin up (+½) or spin down (-½ ) s -- sublevel, ms = s^1^+½, s^2^ -½ p -- sublevel, ms = p^1^ =+½, p^2^ =+½, p^3^= +½ p^4^ = -½, p^5^ = -½, p^6^ = -½ d -- sublevel, ms = d^1^ =+½, d^2^ =+½, d^3^= +½, d^4^= +½, d^5^= +½ d^6^ = -½, d^7^ = -½, d^8^ = -½, d^9^ = -½, d^10^ = -½ f -- sublevel, ms = f^1^ =+½, f^2^ =+½, f^3^= +½, f^4^= +½, f^5^= +½, f^6^= +½, f^7^= +½ f^8^ = -½, f^9^ = -½, f^10^ = -½, f^11^ = -½, f^12^ = -½, f^13^ = -½, f^14^ = -½ Examples: Supply the set of quantum number for the given outermost electron of an element.  **Wolfgang Pauli** or in full **Wolfgang Ernst Friedrich Pauli** is Austrian-born physicist and recipient of the 1945 Nobel Prize for Physics for his discovery in 1925 of the Pauli Exclusion Principle, which states that in an atom no two electrons can occupy the same quantum state simultaneously. Pauli made major contributions to quantum mechanics, quantum field theory, and solid-state physics, and he successfully hypothesized the existence of the neutrino. **Pauli Exclusion Principle -** states that in a single atom [no two electrons will have an identical set or the same quantum numbers (n, l, ml, and ms)]. To put it in simple terms, every electron should have or be in its own unique state (single state). There are two salient rules that the Pauli Exclusion Principle follows: 1\. Only two electrons can occupy the same orbital. 2\. The two electrons that are present in the same orbital must have opposite spins or it should be antiparallel. The box represents the orbitals and the arrow represents the electrons. **s** -- has 1 orbital, so it has only 1 box having a maximum of 2 electrons in opposite spins. **p** -- has 3 orbitals, so it has 3 boxes representing px, py, pz. Each box should have a maximum of 2 electrons in opposite spins. **d** - has 5 orbitals, so it has 5 boxes representing d1, d2, d3, d4, d5. Each boxes should have a maximum of 2 electrons in opposite spins. **f** - has 7 orbitals, so it has 7 boxes representing f1, f2, f3, f4, f5, f6, f7. Each boxes should have a maximum of 2 electrons in opposite spins. Examples: 1\. 4p^4^ 3. 3s^2^ 4px 4py 4pz s 2\. 5f^5^ 4. 1s^2^ 2s^1^  **Get Involved (Explore)** Which subatomic particle was involved in emitting the specific color among the metal salts? Yes, it is the electron. Electrons are found outside the nucleus. Can you find exactly the location of an electron within the atom? Let's have the activity to understand the probability of finding an electron in an atom. **NOTE:** Please refer to the attached LAS. Activity 1 **Predicting the Probable Location of an Electron** Objective: Describe how it is likely to find the electron in an atom by probability. Materials: 1 short coupon bond, pencil, compass or any round object, ruler, graphing paper Procedure: 1\. Get the center of the coupon bond and draw a dot on the center. 2\. From the dot at the center, measure 1cm, 3 cm, 5 cm, 7 cm, and 9 cm. 3\. Draw a concentric circle around the dot as shown in the figure on the right side. 4\. Tape the paper on the floor so that it will not move. 5\. Stand in front of the paper. Hold your ballpen at chest level above the center of the circles you have drawn. (Target is the center which represent the nucleus of an atom). 6\. Drop the ballpen at chest level so that it will leave a small dot/mark on the paper. Cross out (X) the dot that marked on your paper after each drop for easy counting afterwards. You need to complete 50 drops. Submit the paper with marks. This will be your **OUTPUT NO.1.** 7\. Count the number of dots (marked by X) in each circle and record that number on the data table below. 8\. Calculate the number of dots per square centimeter (cm^2^) by dividing the results of Column E with Column D. 9\. Using a graphing paper, plot the average distance from the center (Column B) on the x-axis and number of dots (Column E\) on the y-axis**.** **10. Complete the items in Column E, F and G.** +---------+---------+---------+---------+---------+---------+---------+ | Circle | Average | Area of | Differe | Number | Number | Percent | | | Distanc | | nce | of Dots | of Dots | | | Number | e | Circle, | of | in | per | Probabi | | | | | Areas | Circle | cm^2^ | lity | | (A) | from | cm2 | of the | | | of | | | Center | | Two | (E) | (E)/(D) | Finding | | | | (C) | | | | dots, | | | cm | | Consecu | | (F) | | | | | | tive | | | \% | | | (B) | | | | | | | | | | Circles | | | (G) | | | | | , | | | | | | | | cm^2^ | | | | | | | | | | | | | | | | (D) | | | | +---------+---------+---------+---------+---------+---------+---------+ | 1 | 1.0 | 3.14 | 25.13 | | | | +---------+---------+---------+---------+---------+---------+---------+ | 2 | 3.0 | 28.27 | 50.27 | | | | +---------+---------+---------+---------+---------+---------+---------+ | 3 | 5.0 | 78.54 | 75.40 | | | | +---------+---------+---------+---------+---------+---------+---------+ | 4 | 7.0 | | 100.53 | | | | +---------+---------+---------+---------+---------+---------+---------+ | 5 | 9.0 | | 125.66 | | | | +---------+---------+---------+---------+---------+---------+---------+ Column F = [Column E ( No. of Dots in Circle)] Percent Probability = Column F (No. of dots per cm^2^) Column D (Difference of Areas) X 100 **Guide Questions:** 1\. What happens to the number of dots per unit area as the distance of the dots go farther from the center? 2\. Determine the percent probability of finding a dot in each of the circle drawn on the target by multiplying No. of dots /cm^2^ (column F). For example: In circle 1(A) -- 5 dots/25.13 = 0.199 = \[0.199 / 100 \] X 100 = 19.9% 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 highest probability of finding dots? 5\. How are your results similar to the distribution of electrons in an atom? **Get Ahead (Elaborate)** Now that you are familiar with the different sets of quantum number, A. Identify the type of orbital in the following given. Write your answers in your LAS. The first one is done for you. 1\. n = 4, l = 3 \_\_\_\_\_ 3. n = 5, l = 1 \_\_\_\_\_ 5. n = 2, l = 2 \_\_\_\_\_\_ 2\. n = 1, l = 0 \_\_\_\_\_ 4. n = 3, l = 2 \_\_\_\_\_ 6. n = 6, l = 3 \_\_\_\_\_\_ B. From the following sets of four quantum numbers; identify those sets that are allowed by writing the type of orbital (i.e., 2s, 4d,...) and if the set is NOT allowed, state the reason why it is not allowed. The first two sets were done for you. Write your answers on the LAS.  1\. ANS: [1s^1^] 2. ANS: Principal quantum number is up to 7 only. 3\. ANS: 4\. ANS: 5. ANS: 6. ANS: 7\. ANS: C. Draw the arrow following Pauli Exclusion Principle. Write your answer in your LAS. The first one is done for you. 1\. 1s^2^ 2s^2^ 2p^4^ 3. 6d^8^ 1s 2s 2px 2py 2pz d1 d2 d3 d4 d5 2\. 1s^2^ 2s^2^ 2p^6^ 3s^1^ 1s 2s 2px 2py 2pz 3s ***Get Assessed (Evaluate)*** ***Multiple Choice: Select the letter of the best answer from among the given choices. Write it on a separate sheet of paper then compare your answers with the answers written on the key to corrections found at the end of this module. No need to submit your answers on this part.*** \_\_\_\_1. What does the flame test proved about the inner structure of atom? A. The atom has a nucleus. C. The electrons are found outside the nucleus. B. The nucleus is positively charged. D. The electrons carry discreet or fixed energy. \_\_\_\_2. What happens to the energy of an electron as it goes farther from the nucleus? A. Decreases B. Fixed C. Increases D. Does not change. \_\_\_\_3. Which of the following energy levels can accommodate a maximum of 18 electrons? A. 1^st^ B. 2^nd^ C. 3^rd^ D. 4^th^ \_\_\_\_4. An atom with 32 electrons has \_\_\_\_\_\_ energy levels. A. 2 B. 2 C. 3 D. 4 \_\_\_\_5. How many sublevels are there in L - energy level? A. 1 B. 2 C. 3 D. 4 \_\_\_\_6. What happens when an electron jumps from higher to lower energy level? A. Colored light is given off. C. The atom becomes excited. B. This process is not possible D. Another electron goes from low to high energy level \_\_\_\_7. Who among the scientist does not contribute to the development of the quantum mechanical model of the atom? A. Erwin Schrodinger B. Louie de Broglie C. Neils Bohr D. Werner Karl Heisenberg \_\_\_\_8. Which statement below supports the Bohr's model of the atom? A. The model was based on the wave properties of the electron. B. The model accounted for the absorption spectra of atoms but not for the emission spectra. C. The model accounted for the emission spectra of atoms, but not for the absorption spectra. D. The model was accounted for describing the electron to be moving in definite orbits around the nucleus. ***\_\_\_\_9. Which among the given set of quantum numbers is not allowed?*** ***A. n= 1 B. l = -1 C. ml = 1 D. ms = +1/2*** ***\_\_\_\_10. Which atomic orbital is characterized having a clover shape or a quadrilobate?*** ***A. d -- orbital B. f -- orbital C. p -- orbital D. s -- orbital*** ***Get Moving (Extend)*** **References:** **Grade 9 LM pages 98 - 107** **Chemistry Textbook by Amelia P. Mapa, et.al, pages 66 -- 74** **ADM Module Quarter 2 Week 1 Module 1 DepEd -- Division of Iligan City** +-----------------------------------+-----------------------------------+ | **Congrats! 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