Q2-M1-LESSON-1-ELECTRONIC-STRUCTURE-OF-MATTER-1 PDF

Notebook Act. 1: Atomic Models GUIDE QUESTIONS: 1 Who was the ancient Greek philosopher that first proposed the idea of atoms? 2 How did Democritus describe these tiny particles that make up everything in the world? 3 What theory did Aristotle propose regarding what matter is made of? 4 What signifi...

Notebook Act. 1: Atomic Models GUIDE QUESTIONS: 1 Who was the ancient Greek philosopher that first proposed the idea of atoms? 2 How did Democritus describe these tiny particles that make up everything in the world? 3 What theory did Aristotle propose regarding what matter is made of? 4 What significant discovery did John Dalton make in 1808 related to atoms? 5 How did J.J. Thompson visualize the structure of an atom in his model? 6 What experiment did Ernest Rutherford conduct to further understand the structure of an atom? 7 How did Niels Bohr expand on Rutherford's nuclear model in 1913? 8 What complications did Bohr's planetary model encounter? 9 How did Werner Heisenberg's uncertainty principle impact our understanding of electrons within an atom? 10 What theory replaced Bohr's planetary model and introduced a new set of complexities? 9 SCIENCE #MATATAG Pangasinan National High School Electronic structure of matter MELC: Explain how the Quantum Mechanical Model of the atom describes the energies and positions of the electrons. Quarter2: Module 1Week 1 (S9MT-IIa-22) MR. OHMARK V. VELORIA, LPT S c i e n c e Te a c h e r  Learning Objective/s: a) Distinguish the different atomic models proposed by the scientist. b) Describe how Bohr Model of atom improved Rutherford’s Atomic Model. Atomic Models Atomic Models Atomic Models Atomic Models Atomic Models Atomic Models  a) Describe how Bohr Model of atom improved Rutherford’s Atomic Model. ✓ Atoms as empty space ✓ Its mass concentrated in the center nucleus ❖ Did not explain why electrons orbits the nucleus Group Activity No. 1: The Flame Test Objectives: Determine the characteristic colors that metal salts emit; and Relate the colors emitted by metal salts to the structure of the atom as describe by Niels Bohr.  Activity No. 1: The Flame Test Procedure 1. Place each metal salt on a watch glass and add 2 to 3 drops of 3M hydrochloric acid. 2. Pour about 3 - 5 mL or enough ethyl alcohol to cover the size of a 1 peso-coin in the first watch glass. Light with a match and observe the color of the flame. (This will serve as reference for comparison of the flame color). Wait for the flame to be extinguished or put out on its own. 3. Repeat procedure No. 2 for each salt. Observe the color of the flame. Copper II Potassium Sodium Calcium Boric Sulfate Chloride Chloride Ethanol Chloride Acid  Activity No.1: The Flame Test 4. Write your observation in a table similar to the one below. Table 1. Color of flame of metal salts Element Color of the Metal salt tested producing color flame Zinc Chloride Zinc Manganese Dioxide Manganese Sodium Chloride Sodium Potassium Chloride Potassium Copper Sulfate Copper Copper II Potassium Sodium Calcium Boric Ethanol Sulfate Chloride Chloride Chloride Acid  Activity No.1: The Flame Test 5. Answer the following questions then you will present your observations and answers later. Q1. Why do you think are there different colors emitted? Q2. What sub-atomic particles in the heated compounds are responsible for the production of the colored light? Q3. How did the scientists explain the relationship between the colors observed and the structure of the atom? Q4. Explain how your observation in Activity 1 relates to Bohr’s model of the atom. Q5. Which illustration below represents the energy of the electron as described by Bohr? Explain your answer. Copper II Potassium Sodium Calcium Boric Ethanol Sulfate Chloride Chloride Chloride Acid Presentation of the Activity Result Group Activity 1: Flame Test Rubric for Presentation of Output Activity Recall How was the Flame Test Activity related to the Bohr Model of Atom? Discuss How was the Flame Test Activity related to the Bohr Model of Atom? Discuss How was the Flame Test Activity related to the Bohr Model of Atom? Electron Transitions Discuss How was the Flame Test Activity related to the Bohr Model of Atom? ABSORPTION Discuss How was the Flame Test Activity related to the Bohr Model of Atom? EMISSION Discuss How was the Flame Test Activity related to the Bohr Model of Atom? Discuss How was the Flame Test Activity related to the Bohr Model of Atom? These colors that you have observed are given off by the vapors of the metallic elements that can be analyzed using an instrument called SPECTROSCOPE. Discuss How was the Flame Test Activity related to the Bohr Model of Atom? This visible light under a spectroscope produced characteristic color and wavelength which is called atomic spectrum of element. The color, number and position of lines produced is called the “fingerprint” of an element.  The problem with Bohr’s Model of Atom is to know the arrangement of electrons in atoms in terms of probability of finding the electron in certain location within the atom.  Activity No. 1: Predicting the Probable Location of an Electron Objective: Describe how it is likely to find the electron in an atom by probability. Procedure: 1. Draw a dot on the center of the sheet of paper or folder. 2. Draw 5 concentric circles around the dot so that the radius of each circle is 1.0 cm, 3 cm, 5 cm, 7 cm, and 9 cm from the dot. 3. Tape the paper on the floor so that it will not move. 4. Stand and target the center which represent 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 or folder. 6. Count the number of dots in each circle and record that number on the data table. 7. Calculate the number of dots per square centimeter (cm²). Answer the data table Answer the guide questions 5. The results of the activity are similar to the structure of the atom because the probability of finding an electron (dot) increases then decreases as it goes farther from the nucleus (target).  HISTORY OF THE QUANTUM MECHANICAL MODEL Lesson LOUIS DE ERWIN WERNER BROGLIE SCHRODINGER HEISENBERG Heisenberg's Refined the wave-particle λ= h / m⋅ν theory proposed by de Broglie. Uncertainty Principle Proposed that the electron Developed an equation that It is impossible to know (which is thought of as a treated an electron like a wave and both the position and predicted the probable location of particle) could also be velocity of an electron an electron around the nucleus thought of as a wave. simultaneously. called the atomic orbital. This volume or region of space around the nucleus where the electron is most likely to be found is called an atomic orbital. The QUANTUM MECHANICAL MODEL views an electron as a cloud of negative charge having a certain geometrical shape. There are different kinds of atomic orbitals that differ in the amount of energy and shapes (where the electron probably is). The atomic orbitals get filled by electrons Lesson in a certain order. s-orbitals are spherically shaped p-orbitals are “dumbbell” shaped d-orbitals are “cloverleaf” shaped  The quantum mechanical model of the atom treats an electron like a wave, with this, the quantum mechanical model describes the probable location of electrons in atoms by describing the Four Sets of Quantum Numbers: PRINCIPAL ENERGY LEVEL Principal Quantum Number (n) Lesson ENERGY SUBLEVEL Angular Momentum Quantum Number (l) ORBITAL (Each Sublevel) Magnetic Quantum Number (ml) SPIN Spin Quantum Number (ms)  PRINCIPAL ENERGY LEVEL Principal Quantum Number (n) -also called as "shells“. It indicates the relative size and energy of atomic orbitals. n=positve integers: n= 1, 2, 3, to 7. Lesson As value of n increases: orbital becomes larger electron spends more time farther away from nucleus atom's energy level increases  ENERGY SUBLEVEL Angular Momentum Quantum Number (l) Principal energy levels are broken down into sublevels. This can have any value from 0 to n-1. Sublevels define the orbital shape (s, p, d, f) n=1, l = 0 (s) orbital Lesson n=2, l = 0, 1 (s, p) orbitals n=3, l = 0,1, 2 (s, p, d) orbitals n=4, l = 0, 1, 2, 3 (s, p, d, f) orbitals  ORBITAL (Each Sublevel) Magnetic Quantum Number (ml) Each sublevel has a different number of orbitals. ml= - l to l Lesson s: 1 orbital p: 3 orbitals d: 5 orbitals f: 7 orbitals Sublevels of the Five Energy levels Type of Principal Maximum Number of Sublevel and Energy Number of Given the value of Sublevels (l) Number of Level (n) Electrons principal energy level, Orbitals (ml) you can use this n=1 l=0-----1 1s (1orbital) 2 formula to compute 2s (1orbital), for the maximum n=2 l=0,1 -----2 8 Lesson 2p (3 orbitals) number of electrons 3s (1orbital), 2n2 electrons. n=3 l=0, 1, 2 -----3 3p (3 orbitals) 18 3d (5 orbitals) 4s (1orbital), 4p (3 orbitals) n=4 l=0, 1, 2 3 -----4 32 4d (5 orbitals) 4f (7 orbitals) 5s (1orbital), 5p (3 orbitals) n=5 l=0, 1, 2, 3 4-----5 5d (5 orbitals) 50 5f (7 orbitals) 5g (9 orbitals)  SPIN Spin Quantum Number (ms) Lesson Electrons act like they are spinning on an axis could be 1/2 or -1/2 Generates a magnetic field No two electrons in the same orbital can have the same spin Electron Configuration Lesson  Electron Configuration It is the representation of the arrangement of electrons distributed among the orbital shells and subshells of an atom. It is used to describes the electron arrangement in atoms at ground state. With this, there are 3 rules for electron configuration at ground state: Aufbau Principle Lesson Pauli Exclusion Principle Hund’s Rule  AUFBAU PRINCIPLE It requires that the electrons occupy the Lesson lowest possible energy level before filling up the next.  PAULI EXCLUSION PRINCIPLE No two electrons in an atom can have precisely the same four quantum Lesson numbers; the spin quantum number limits the number of electrons in an orbital to a maximum of two.  HUND'S RULE A single electron with the same spin Lesson must occupy each orbital in a sublevel before they pair up with an electron with an opposite spin. ELECTRON CONFIGURATION The four different types of orbitals (s,p,d, and f) have different shapes, Lesson and one orbital can hold a maximum of two electrons. The p, d, and f orbitals have different sublevels, thus can hold more electrons. ELECTRON CONFIGURATION Pa’no=p S’ya=s Si = s Franky=f Susan= s Daddy= d Pumunta = p Lesson Sa= s Pa’no=p Party= p Siya=s Franky= f Si= s Daddy=d Daddy= d Pumunta=p Pa’no=p Sa=s Franky=f Disco=d Daddy=d Pa’no=p  4 Steps in Writing Electron Configurations 1. Determine the atomic number of the element from the Periodic Table of Elements. This gives the number of protons and electrons in Lesson the atom Ex. Mg (Z=12) , so Mg has 12 protons and 12 electrons 2. Draw boxes to represent the first 3 energy levels s and p orbitals Since there are only 12 electrons, 9 should be plenty. 1s 2s 2p 3s 3p 4 Steps in Writing Electron Configurations 3. Add one electron to each box in a set., then pair the electrons before Lesson going the next set until you use all the electrons. When pair, put in opposite arrows  4 Steps in Writing Electron Configurations 4. Use the diagram to write the electron configuration Write the number of electrons in Lesson each set as a superscript next the name of the orbital set Sample Electron Configurations Lesson Lesson