Science Chemistry Unit 04 Exemplar Lesson 01: Chemical Bonds and Molecular Geometry

Science Unit: 04 Lesson: 01 Suggested Duration: 13 days Science Unit 04 Exemplar Lesson 01: Chemical Bonds and Molecular Geometry This lesson is one approach to teaching the State Standards associated with this unit. Districts are encouraged to customize this lesson by supplementing with district-approved resources, materials, and activities to best meet the needs of learners. The duration for this lesson is only a recommendation, and districts may modify the time frame to meet students needs. To better understand how your district may be implementing CSCOPE lessons, please contact your child s teacher. (For your convenience, please find linked the TEA Commissioner s List of State Board of Education Approved Instructional Resources and Midcycle State Adopted Instructional Materials.) Lesson Synopsis This lesson is designed to introduce students to how atoms form ionic bonds, covalent bonds, and metallic bonds. Students will learn to write electron configurations and Lewis electron dot structures. Students will use valence electrons and electronegativities to predict the types of bonds that atoms form. Students will understand the delocalized nature of electrons in metallic bonding and its effect on the physical properties of metals. Lastly, students will learn VSEPR theory to predict molecular structures. TEKS The Texas Essential Knowledge and Skills (TEKS) listed below are the standards adopted by the State Board of Education, which are required by Texas law. Any standard that has a strike-through (e.g. sample phrase) indicates that portion of the standard is taught in a previous or subsequent unit. The TEKS are available on the Texas Education Agency website at http://www.tea.state.tx.us/index2.aspx?id=6148. C.5 Science concepts. The student understands the historical development of the Periodic Table and can apply its predictive power. The student is expected to: C.5C Use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity, and ionization energy. Readiness Standard C.6 Science concepts. The student knows and understands the historical development of atomic theory. The student is expected to: C.6E Express the arrangement of electrons in atoms through electron configurations and Lewis valence electron dot structures. Readiness Standard C.7 Science concepts. The student knows how atoms form ionic, metallic, and covalent bonds. The student is expected to: C.7C Construct electron dot formulas to illustrate ionic and covalent bonds. Readiness Standard C.7D Describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility. Supporting Standard C.7E Predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory. Scientific Process TEKS Supporting Standard C.1 Scientific processes. The student, for at least 40% of instructional time, conducts laboratory and field investigations using safe, environmentally appropriate, and ethical practices. The student is expected to: C.1A Demonstrate safe practices during laboratory and field investigations, including the appropriate use of safety showers, eyewash fountains, safety goggles, and fire extinguishers. C.1B Know specific hazards of chemical substances such as flammability, corrosiveness, and radioactivity as summarized on the Material Safety Data Sheets (MSDS). C.1C Demonstrate an understanding of the use and conservation of resources and the proper disposal or recycling of materials. C.2 Scientific processes. The student uses scientific methods to solve investigative questions. The student is expected to: C.2E Plan and implement investigative procedures, including asking questions, formulating testable hypotheses, and selecting equipment and technology, including,graphing calculators computers and probes, sufficient scientific glassware such as beakers, Erlenmeyer flasks, pipettes, graduated cylinders, volumetric flasks, safety goggles, and burettes, electronic balances, and an adequate supply of consumable chemicals. Last Updated 06/05/2013 page 1 of 29

C.2F Collect data and make measurements with accuracy and precision. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days C.2H Organize, analyze, evaluate, make inferences, and predict trends from data. C.2I Communicate valid conclusions supported by the data through methods such as lab reports, labeled drawings, graphs, journals, summaries, oral reports, and technology-based reports. C.3 Scientific processes. The student uses critical thinking, scientific reasoning, and problem solving to make informed decisions within and outside the classroom. The student is expected to: C.3F Research and describe the history of chemistry and contributions of scientists. GETTING READY FOR INSTRUCTION Performance Indicators High School Science Unit 04 PI 01 Given three compounds (one ionic, two covalent), use the concepts of VSEPR theory, intermolecular forces, electronegativity, and chemical bonds to explain the properties of each of the substances. Represent each compound on a visual display, such as a poster, including the electron configuration, Lewis valence electron dot structure, and shape and polarity of each. In addition, summarize the nature of metallic bonding for a given metal and how the bonding relates to the properties of the metal. Standard(s): C.2I, C.6E, C.7C, C.7D, C.7E ELPS ELPS.c.1C, ELPS.c.4G, ELPS.c.5G Key Understandings Arrangement of the electrons in atoms can be shown using electron configurations and Lewis dot structures. How can the electron configuration and Lewis dot structure for an atom be represented? The type of bond that can be produced between two elements can be predicted by the element s position on the Periodic Table. Which properties of elements and periodic trends can be used to predict the type of bond that can be made between two elements? The VSEPR theory can be used to identify a molecule s geometric shape. How can the VSEPR theory be applied to predict shapes of molecules? The geometric shape of a molecule can be used to determine polarity. How is polarity related to the geometric shape of a molecule? Metallic elements have unique properties as a result of metallic bonding. How does metallic bonding create the unique properties of metals? Vocabulary of Instruction valence electrons oxidation number ion electronegativity ionization energy electron affinity energy levels electron orbitals electron configuration orbital diagrams Lewis dot structures covalent bond ionic bond metallic bond polar non-polar VSEPR theory molecular geometry Materials aluminum foil (small piece per student, see Advance Preparation, 1 per teacher) aluminum shot (1 per teacher) aluminum wire (1 per student, 1 per teacher) balances (triple beam or electronic, 1 per group) balloon (9-in. latex,1 per teacher) beakers (250 ml, 5 per group) burettes (2 per teacher) card sets from Teacher Resource: It s Elementary Cards (from previous activity) ceramic (or plastic) plate (1 per teacher) colored pencils (per group) conductivity probes or testers (1 per group) conductivity tester (1 per teacher) containers (labeled, see Advance Preparation, 1 each per group) deionized water, 1 L Last Updated 06/05/2013 page 2 of 29

tap water, 500 ml sucrose solid, ~100 g NaCl solid, ~100 g empty for rinse water collection copper shot (1 per teacher) copper strips (1 per student, 1 per teacher) copper wire (1 per student, 1 per teacher) de-ionized water (at least 50 ml per teacher) Erlenmeyer flasks (250 ml, 2 per group) glue or tape (per group) graduated cylinders or pipettes (100 ml, 1 per group) ice cubes (several per teacher) magnesium ribbon (1 per student, 1 per teacher) materials for 3 D modeling (per student) Optional metal cooking pan (1 per teacher) mini marshmallows (5 per pair) paper clips (1 per student, 1 per teacher) poster board (or other materials or software for student visuals, per student) research materials or access to student computers/internet ring stand with double burette clamp (1 per teacher) rinse bottle containing deionized water (1 per group) toothpicks (5 per pair) vegetable oil (at least 50 ml per teacher) weighing paper (several sheets per group) wool or fur (1 piece per teacher) Science Unit: 04 Lesson: 01 Suggested Duration: 13 days Attachments All attachments associated with this lesson are referenced in the body of the lesson. Due to considerations for grading or student assessment, attachments that are connected with Performance Indicators or serve as answer keys are available in the district site and are not accessible on the public website. Teacher Resource: It s Elementary Cards (see Adv Prep, 1 set per pair) Teacher Resource: Electron Configurations (see Adv Prep, 1 set per pair) Handout: Orbitals and Electrons Note-taking Guide (see Adv Prep, class set of page 1, 1 per stu Handout: Blank Periodic Table (1 per student) Teacher Resource: Lewis Dot Structures (see Adv Prep, 1 set per pair) Handout: Drawing Lewis Dot Structures for Compounds (1 per student) Handout: Molecular Geometry Determination Using VSEPR Theory (1 per student and 1 for projection) Teacher Resource: Performance Indicator Instructions KEY Handout: Visual Display PI (1 per student) Resources None Identified Advance Preparation 1. Prior to Day 1, create enough sets of cards from the Teacher Resource: It s ELEMENTary Cards to have a class set of cards with one set per pair of students. Use different colors of cardstock and laminate, and cut them out and bag. Make some blank cards as well to challenge some students. 2. Prior to Day 2, arrange for access to student computers for the Explore I and Explain I activities. 3. Prior to Day 2, locate and preview interactive websites for students to use in the web quest (Explore I) on electron configurations. Learner.org has interactive materials and charts relating to the Periodic Table, electron configurations, and Aufbau principle that may be helpful for this lesson. You will need to display a copy of a chart of the Aufbau principle during class. 4. Prior to Day 2, print enough sets of cards from the Teacher Resource: Electron Configurations to have a class set of cards with one set per pair of students. Use different colors of cardstock, laminate, cut out, and bag. 5. Prior to Day 2, print a class set of page 1 of the Handout: Orbitals and Electrons Note-taking Guide. Place this copy in a sleeve protector or laminate, and use with each class. Print enough copies of page 2 for each student to affix to their notebooks. 6. Prior to Day 3, locate and preview an online simulation or game to illustrate orbitals and orbital notation and to help students build orbitals. Again, learner.org has information and an interactive Periodic Table available that may be helpful for this purpose. 7. Prior to Day 4, print enough sets of the Teacher Resource: Lewis Dot Structures on different colors of cardstock to have a set per pair of students. Laminate, cut out, and bag the cards. 8. Prior to Day 5, obtain or a build conductivity tester(s) or set up conductivity probes. Prepare sets of labeled containers for testing for groups: sugar, NaCl, tap water, and de-ionized water. Put out sufficient 250 ml Erlenmeyer flasks, 250 ml beakers, 100 ml graduated cylinders, and 25 Last Updated 06/05/2013 page 3 of 29

ml pipettes. 9. Prior to Day 5, fill one burette with de-ionized water and a second with vegetable oil; clip both to a ring stand. Place a 250 ml beaker under each burette to collect liquid when the stopcocks are opened. 10. Prior to Day 5, locate and preview interactive website(s) to assist in guiding students through an overview of electronegativity, ionization energy, electron affinity, covalent and ionic bonding, and polarity. 11. Prior to Day 8, prepare materials for demonstration by cutting copper strips, wires, and magnesium ribbons for each student. Cut aluminum squares for each student. 12. Prior to Day 9, locate and preview an interactive website to introduce the basics of metallic bonds. Consider using the internet search terms metallic bonding and the properties of metals chemistry tutorial. Alternatively, consider using your locally adopted textbook and other resources. 13. Prior to Day 12, decide on the substances you want students to model (one ionic, one non-polar covalent, one polar covalent, and one metallic) in the Performance Indicator, and determine how you will assign them to students. Also, decide on the modes of presentation you will allow, and obtain needed materials for modeling, such as toothpicks and mini-marshmallows. 14. Prepare attachment(s) as necessary. Background Information Chemical compounds are formed when the atoms of two or more different elements join. A stable compound occurs when the energy of the combination is less than the total energy of the separated atoms. A net attractive force between the atoms is created and termed a chemical bond. Electronegativity differences between the atoms in a combination determine the type of chemical bond that will be formed. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days Two types of chemical bonds are explored in this unit: 1. Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. Sharing may be equal, such as between two of the same atoms, or unequal, sharing between two different atoms. Unequal sharing results in a polar bond, depending on the differences in the electronegativities of the two atoms. The polarity of the bond will be more negative toward the more electronegative atom. 2. Ionic bond: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other Covalent and ionic bonds are stable if their atoms share or give/receive electrons in such a way as to create a noble (inert) gas electron configuration for each atom. The study of electron configurations of atoms will give the student a better understanding of the interactions between atoms, ions, and molecules. The behavior of electrons, therefore, explains most aspects of how atoms interact with other in the living and nonliving world. Electron configurations represent the arrangement of electrons in an atom around its nucleus. Electrons fill s, p, d, and f orbitals according to the Aufbau principle, Hund s rule, and the Pauli Exclusion Principle. Since only the valence electrons are used in forming chemical bonds, Lewis electron dot structures are used to model these electrons in an atom. The attractive and repulsive forces in molecules influence the architecture or shape of molecules. The Valence-Shell Electron-Pair Repulsion Theory, or VSEPR theory, can be used to predict how these forces affect chemical bonding and the shapes of molecules. It is based on the premise that the unshared pairs of valence electrons of each adjacent atom will repel each other (electrostatic repulsion) strongly and therefore, cause the peripheral atoms to move as far from each other as possible. Predicting molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries and their resulting molecular structures are shown in the following table: The polarity of bonds in conjunction with the molecular structure determines whether a molecule is polar or not. Water, H 2 O, is a polar molecule, for example. It is not linear, but bent due to oxygen s two pairs of unshared electrons. Since each 0 H bond is polar (negative toward the O), the H 2 O molecule is polar as well. CO 2 is linear since there are no unshared pairs. Even though the C-O bond is polar, the molecule is not, since the bond polarities cancel. A third type of bond, the metallic bond, is one in which electrons can move freely in all directions. This movement results in attractions between positively charged metal ions and free-floating valence electrons in metals. The many physical properties of metallic substances, such as thermal and electrical conductivity, malleability, and ductility, are explained by the freely moving nature of electrons in metallic bonds. STAAR Notes: Using the Periodic Table to identify and explain periodic trends will be tested as Readiness Standards under Resporting Category 1: Matter and the Periodic Table. Electron configuration and Lewis valence electron dot structures will be tested as Readiness Standards under Resporting Category 2: Atomic Structure and Nuclear. Electron dot formulas illustrating ionic and covalent bonds will be tested as Readiness Standards under Reporting Category 3: Bonding and Chemical Reactions. Included in the STAAR Reference Materials is a Periodic Table of the elements that includes the lanthanide and actinide series. Shown within the table are the name, symbol, atomic number, and atomic mass of each element. Mass numbers in parentheses are those of the most stable or most common isotopes. INSTRUCTIONAL PROCEDURES Instructional Procedures Notes for Teacher Last Updated 06/05/2013 page 4 of 29

ENGAGE It s ELEMENTary NOTE: 1 Day = 50 minutes Suggested Day 1 Science Unit: 04 Lesson: 01 Suggested Duration: 13 days 1. Provide pairs of students a deck of playing cards from the Teacher Resource: It s ELEMENTary (See the Instructional Notes and Advance Preparation.). 2. Instruct students to arrange the cards in some kind of logical order. Circulate around the classroom to facilitate the process as needed and probe student logic. 3. Allow students to compare their arrangements with a nearby group and discuss similarities/differences. Before proceeding, verify that all groups have been successful in sequencing their cards into four rows, representing the first 20 elements of the Periodic Table. Pose the following question: What was your reasoning for deciding how to group the cards? Answers will vary. 4. Ask students to copy the following terms at the top of a page in their science notebooks, leaving ample room below: protons, electrons, neutrons, valence electrons, and energy levels. 5. Instruct students to copy four cards, one from each row of their arrangement, into their science notebooks. They should copy no more than two per page, leaving ample room to add labels as they proceed through the lesson. 6. Instruct students to label any one of the illustrations using the terms provided in the previous step. 7. Faciltate a discussion using the following questions: What do you think the p on the card represents? Protons What do you think the n on the card represents? Neutrons How are electrons represented? Black dots What do the circles represent? Energy levels How do the energy levels relate to placement on the Periodic Table? The energy levels correspond to the rows (periods) of the Periodic Table. What do you notice about the valence electrons in each column? The elements in each column (families) have the same number of valence electrons. From the information provided, what element(s) do your 2-D models recorded in your science notebook represent? Answers will vary, depending on the element card they chose. Attachments: Teacher Resource: It s ELEMENTary Cards (see Advance Preparation, 1 set per pair) Instructional Notes: Today s activity extends the deck of cards analogy from the Las Vegas Periodic Table Engage activity from Unit 03 and has students prepare their science notebooks for this unit. To really challenge students, exchange blank cards for three or four of the It s ELEMENTary playing cards. Challenge students to predict the structures of the missing atoms by drawing them on the blank cards with dry erase markers. These cards should NOT show what atoms look like; they are 2-D models of atoms, indicating their subatomic particles and energy levels. STAAR Notes: Students learn about elements and symbols in Grade 6 (6.5A). Students describe the structure of atoms (mass, charge, protons, neutrons, and electrons) (8.5A). They are also introduced to valence electrons and reactivity in Grade 8 (8.5B) Students learn that protons determine an element s identity and its valence electrons determine its properties in Grade 8 (8.5B). It may be helpful to review these concepts as part of Question 7. Science Notebooks: Students record atomic diagrams of four elements in their science notebooks. These drawings will be referenced multiple times as this lesson progresses. 8. Ask students to record the atomic symbol for each of the four cards they recorded in their notebooks. Remind students to use their Periodic Table as a reference tool to help with identification. EXPLORE I Electron Configurations Suggested Days 2 and 3 1. Introduce your choice of interactive websites on electron configuration to students. Inform them they will be conducting a web quest in order to learn about electron configurations (see Advance Preparation). 2. Inform students that they will be exploring the interactive website(s), and they should take notes of their explorations in their science notebooks. 3. Divide the class into groups, based on the number of computers you have available. 4. Display the following web quest questions on the board, and instruct students to work together to find the answers. Then, students should individually answer the questions in their notebooks. What is the Aufbau Principle? Who was Wolfgang Pauli? What is his contribution to our understanding of the atom? Who was Friedrich Hund and what did he contribute? What is the maximum number of electrons each energy level can hold and the number of orbitals for distribution at each level? 5. Allow a reasonable time for completion of the web quest, and then facilitate a class discussion in which students reflect upon the information they have researched. Instruct students to make further entries in their notebooks during the discussion. 6. Display a chart of the Aufbau Principle (see Advance Preparation), and instruct students record the principle in their notebooks: Materials: card sets from Teacher Resource: It s Elementary Cards (from previous activity) colored pencils (per group) Attachments: Teacher Resource: Electron Configurations (see Advance Preparation, 1 set per pair) Instructional Notes: If computer access is limited, use a projector to introduce the interactive website(s), and then have students explore the sites in teams. Alternatively, go through the site as a modified web quest. Students may feel uncomfortable making the matches and trusting their problem-solving ability. DO NOT make the matches for them. Allow students to develop their own logic for the match-making activity. Allow students to develop their own understanding of electron-configurations and electron notation during the matching activity. Check student understanding and use guided questions to Last Updated 06/05/2013 page 5 of 29

Start by filling electrons in at the lowest energy level orbital, and build up to the higher energy level orbitals only after the lower are filled. 7. Ensure students correctly noted the maximum number of electrons each energy level can hold and number of orbitals available for distribution. 8. Provide pairs of students a deck of playing cards from the Teacher Resource: It s ELEMENTary (from previous activity) and the Electron Configurations for the first 20 elements of the Periodic Table. 9. Instruct students to match each Electron Configurations card with the corresponding It s ELEMENTary card. 10. Monitor student progress, and ask questions about student thinking as they complete the task. You may want to point out when a match is correct or incorrect, but do not offer any other assistance at this point. 11. Once all groups have made their matches, facilitate a discussion with students: How were you able to determine which electron configuration corresponded with which playing card? Answers will vary. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days help students make meaning and come to conclusions about the concepts. The letters s, p, d, and f, that are used to indicate orbitals, were derived from the words sharp, principal, diffuse, and fundamental, referring to properties of spectral lines. If only a classroom computer with projector is available for the atom building activity, call students up to work through the process as a class. Science Notebooks: Students continue to build their concept understanding of the atom, adding detail to the 2-D representations they have constructed in their science notebooks. Additionally, students record notes related to orbitals, electrons, and electron configurations. 12. Instruct students to add electron configurations to the diagrams they made in their notebooks during the Engage activity. EXPLAIN I Electron Configurations 1. Facilitate a discussion about orbitals and electrons. Consider using an interactive website to present the orbitals information (see Advance Preparation). 2. Use the Handout: Orbitals and Electrons Note-taking Guide to provide a construct for students to record pertinent notes in their science notebooks (see Advance Preparation). 3. Additionally, distribute a copy of the Handout: Blank Periodic Table to each student to affix into their notebooks. This template will be used to label and shade each block of the Periodic Table, s, p, d, and f. 4. Briefly introduce students to the concept of orbital notation. Again, an interactive website may be helpful to illustrate concepts. 5. Review Hund s Rule with students. Formulated by German physicist Friedrich Hund around 1927, the rule states that each p, d, or f orbital must receive one electron before any p, d, or f orbital can receive a second filling electron. 6. Review the Pauli Exclusion Principle with students. An orbital can hold zero, one, or two electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins. This is indicated by either an up or down arrow on the orbital diagram. Suggested Day 3 (continued) Materials: Attachments: glue or tape (per group) Handout: Orbitals and Electrons Note-taking Guide (see Advance Preparation, class set of page 1, 1 per student of page 2) Handout: Blank Periodic Table (1 per student) Instructional Note: Though orbital notation is not mentioned specifically in the TEKS (C.6E), inclusion of the concept in this lesson supports student understanding of electron configuration and Lewis valence electron dot structures. 7. Allow students to play a game or simulation involving the building of orbitals/elements. There are several available online. 8. Instruct students to draw a diagram in their science notebooks for each atom they build in the simulation, complete with orbital notation and electron configuration. 9. Instruct students to affix page 2 of the Handout: Orbitals and Electrons Notetaking Guide to their notebooks. EXPLORE/EXPLAIN II Lewis Dot Structures Suggested Day 4 1. Provide pairs of students a deck of It s ELEMENTary playing cards and cards from the Handout: Lewis Dot Structures for the first 20 elements of the Periodic Table. 2. Instruct students to match each of the Lewis Dot Structures with the corresponding It s ELEMENTary playing card. 3. Monitor student progress, and ask questions about student thinking as they complete the task. You may want to point out when a match is correct or incorrect, but do not offer any other assistance at this time. 4. Once all groups have made the matches, facilitate a class discussion using the following question: How were you able to determine which Lewis dot structure corresponded with which playing card? Answers will vary. Materials: card sets from the Teacher Resource: It s ELEMENTary Cards (from previous activity) Attachments: Teacher Resource: Lewis Dot Structures (see Advance Preparation, 1 set per pair) Check For Understanding: Last Updated 06/05/2013 page 6 of 29

Science Unit: 04 Lesson: 01 Suggested Duration: 13 days 5. Reinforce that the Lewis dot structure is a simple model that includes the element symbol and valence electrons only. 6. Instruct students to add Lewis dot structures to the element diagrams they created in their notebooks during the Engage activity. 7. Instruct students to answer the following question in their notebooks. How can the electron configuration and Lewis dot structure for an atom be represented? Lines represent shared valence shell electrons and dots represent unshared valence shell electrons in a molecule. Use the closing question as an opportunity to informally assess student understanding. Science Notebooks: Students continue to build their concept of an atom, adding detail to the 2-D representations they have constructed in their science notebooks. 8. Give students a few minutes to answer the question, and then ask for volunteers to share their responses. Clarify any misconceptions at this time. Encourage students to add to or revise their answers. EXPLORE III Covalent and Ionic Compounds Suggested Day 5 1. Divide the class into groups based on how many conductivity testers or probes you have available. Inform students that they will be conducting an investigation to test different solutions for conductivity (see Advance Preparation). 2. Conduct a safety prelab discussion. Review the following with students before they begin: location of safety equipment reminder about wearing safety goggles hazards of chemicals to be used and location of msds safe disposal 3. Remind students you will be expecting them to demonstrate safe practices, conservation of resources, and proper disposal. 4. Demonstrate proper use of the 100 ml graduated cylinder (meniscus) and/or pipette with bulb. 5. Inform students their first task is to make 100 ml of 10% sugar solution and 100 ml of 10% NaCl solution using deionized water in the Erlenmeyer flasks at their stations. Provide calculation and other assistance, as needed, to the groups. 6. Next, demonstrate the proper technique for using a conductivity probe tester. Emphasize the importance of rinsing with deionized water thoroughly between solutions. 7. Instruct students to use the conductivity tester to test the deionized water, tap water, sucrose solution, and NaCl solution (in that order to minimize contamination) for conductivity and record the data in their science notebooks. 8. Ask students to create a data table in their notebooks to record the data that will be taken. In addition, instruct students to draw and label the glassware/equipment they will use in completing the investigation. 9. Additionally, instruct students to write the Lewis dot structures for C, H, O, Na, and Cl in their science notebooks. 10. Monitor and assist students as they complete the investigation. Materials: balances (triple beam or electronic, 1 per group) beakers (250 ml, 5 per group) weighing paper (several sheets per group) Erlenmeyer flasks (250 ml, 2 per group) graduated cylinders or pipettes (100 ml, 1 per group) containers (labeled, see Advance Preparation, 1 each per group) deionized water, 1 L tap water, 500 ml sucrose solid, ~100 g NaCl solid, ~100 g empty for rinse water collection conductivity probes or testers (1 per group) rinse bottle containing deionized water (1 per group) Safety Note: Review safety precautions with students prior to any investigation: location of safety equipment reminder about wearing safety goggles hazards of chemicals to be used and location of MSDS safe disposal Science Notebooks: Students draw equipment and record data collected in their notebooks. EXPLAIN III Covalent and Ionic Bonds Suggested Days 5 (continued), 6, and 7 1. Facilitate a post investigation discussion in which students reflect on the results of the previous day s investigation. Ask: How did the conductivity of the four solutions compare? Deionized water and sucrose solution are not conductive; tap water is slightly conductive and NaCl solution is highly conductive. What does it mean that the deionized water and sugar solutions don t conduct an electric current and the tap water and NaCl solutions do? Answers will vary, but you may have to prompt students by reminding them that an electric current is moving electrons. If students respond dissolved salt and dissolved sugar are different, affirm and ask What about the dissolved salt allows that solution to conduct a current while the dissolve sugar does not? Materials balloon (9-in. latex,1 per teacher) wool or fur (1 piece per teacher) ring stand with double burette clamp (1 per teacher) burettes (2 per teacher) vegetable oil (at least 50 ml per teacher) de-ionized water (at least 50 ml per teacher) beaker (250 ml, 2 per teacher) Attachments: Last Updated 06/05/2013 page 7 of 29

2. Guide students to an understanding of the importance of electrons in determining physical properties of substances, eventually reaching the understanding that electrical conductivity is a measurement of the relative mobility of electrons in the substance in water, which is different for covalent and ionic compounds. 3. Ask students how sugar and NaCl are different. Sugar is a covalent compound, and NaCl is ionic. 4. Use an interactive website to guide students through an overview of electronegativity, ionization energy, electron affinity, covalent and ionic bonding, and polarity (see Advance Preparation). 5. Instruct students to take notes in their notebooks. 6. Distribute the Handout: Drawing Lewis Dot Structures for Comounds to each student. 7. Facilitate a discussion of how to construct electron dot formulas for ionic and covalent compounds, guiding students as they work through it. 8. Conclude this segment of the lesson by demonstrating and providing students with evidence that polar and non-polar covalent molecules behave differently (see Instructional Notes). Begin by introducing the burette as a tool for chemists to accurately measure liquid volume and, in this case, to provide a fine stream of liquid. Ask students to make a diagram with labels of the setup in their notebooks and then to sketch their observations. Charge an inflated balloon by rubbing it on fur, wool, or your own (or a student s) hair. Open the stopcock of the water-filled burette until you get a fine, unbroken stream. Hold the charged balloon near it. Make sure that students can see what you are doing and the balloon is not blocking their view. Note that the stream is attracted to the balloon. Repeat the procedure using the vegetable oil burette. This time, the stream will be unaffected. Handout: Drawing Lewis Dot Structures for Compounds (1 per student) Instructional Notes: The octet rule states atoms in periods 1 3 with eight electrons in their valence shell will be stable, regardless of whether these electrons are bonding or nonbonding. The 18-electron rule is operative for atoms in period 4 and 5, which have to achieve 18 electrons in their orbitals to achieve a stable configuration - The same as a noble (inert) gas. Similarly for period 6, the atoms have to achieve 32 electrons to fill their orbitals. For high school level chemistry, it is appropriate to limit experiences to elements in periods through 1 3. Consider making the demonstration more entertaining by claiming that you have magical powers and you alone can bend water. Pretend to be a sorcerer. Make sure that the balloon comes close enough to the water to make the attraction obvious, but not so close that the balloon gets completely soaked. Check For Understanding: Use the closing question as an opportunity to informally assess student understanding. Science Notebooks: Students construct Lewis dot structures for covalent and ionic compounds. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days 9. Continue the discussion: Why do you think the charged balloon attracts the water but not the vegetable oil? The polar nature of the water allows the electrons to align with the charged surface of the balloon, whereas the non-polar nature of the oil does not. What other examples can you think of that might demonstrate the different behaviors of polar and non-polar substances? Oil slick on a puddle; oil and vinegar don t mix in salad dressing; easy to dissolve table salt in water; not easy to dissolve flour in water 10. Introduce intra and inter for students by saying: Intramolecular and intraionic bonds affect the intermolecular forces or attractions between particles of different substances. Discuss as needed. 11. Instruct students to answer the following question in their notebooks. Which properties of elements and periodic trends can be used to predict the type of bond that can be made between two elements? 12. Allow a few minutes for students to write their answers. Ask several volunteers to share their answer with the class, and clarify any misconceptions. Encourage students to add to or revise their answers. EXPLORE IV Metallic Bonds Suggested Days 8 and 9 1. Project the terms thermal conductivity, electrical conductivity, malleability, and ductility on the board, and instruct students to copy the terms into their science notebooks as headings, leaving ½ page available after each to record information. 2. To demonstrate thermal conductivity of metals, place an ice cube on both a ceramic plate and metal cooking pan. Ask students to predict what will happen. 3. Walk around the classroom showing students how the ice on the metal pan begins Materials ceramic (or plastic) plate (1 per teacher) metal cooking pan (1 per teacher) ice cubes (several per teacher) conductivity tester (1 per teacher) Last Updated 06/05/2013 page 8 of 29

melting quickly, forming a puddle, and the ice on the ceramic pan forms a puddle much more slowly. 4. Ask students to think silently about why they think the ice on the metal pan melts more quickly than on the ceramic plate. Allow a moment for students to reflect, and then ask them to record their thoughts in their notebooks under the term thermal conductivity. 5. Allow students to share and compare their ideas with those around them. Listen attentively for student thinking and the vocabulary terms they use to describe the situation. DO NOT provide them with an answer at this time. 6. Next, use a conductivity tester to determine the electrical conductivity of the variety of metal samples available. Ask students to predict what will happen when you touch the samples. 7. Repeat this for each sample. Students know that metals conduct electricity, but emphasize the point that every metal in every form conducts electricity. 8. Ask students to silently think about why they think the metals are such great electrical conductors. Allow a moment for students to reflect, and then ask them to record their thoughts in their notebooks under the heading electrical conductivity. 9. Allow students to share and compare their ideas with those around them. Listen attentively for student thinking and the vocabulary terms they use to describe the situation. DO NOT provide them the answer at this time. 10. Distribute the pieces of aluminum foil, copper strips/wire, magnesium ribbon, and paper clips around the room, asking students to bend their pieces. aluminum shot (1 per teacher) copper shot (1 per teacher) copper strips (1 per student, 1 per teacher) copper wire (1 per student, 1 per teacher) aluminum foil (small piece per student, see Advance Preparation, 1 per teacher) aluminum wire (1 per student, 1 per teacher) magnesium ribbon (1 per student, 1 per teacher) paper clips (1 per student, 1 per teacher) Instructional Note: This activity consists of a hands-on teacher demonstration of properties of metals including thermal conductivity, electrical conductivity, malleability, and ductility. STAAR Note: Students compare properties of metals, nonmetals, and metalloids in Grade 6 (6.6A). Science Notebooks: Science Unit: 04 Lesson: 01 Suggested Duration: 13 days Students should record thoughts and observations throughout the teacher-led demonstration and interaction with metal samples. 11. Instruct students to think about what it is about the atomic structure of metals that makes them so bendable. Allow a moment for students to reflect, and then ask them to record their thoughts in their notebooks under the heading malleability. 12. Allow students to share and compare their ideas with those around them. Listen attentively for student thinking and the vocabulary terms they use to describe the situation. DO NOT provide them the answer at this time. 13. Show students the pieces of wire, and discuss the property of ductility. Discuss the following: Ductility is a solid material's ability to deform under stress; this is often characterized by the material's ability to be drawn into a wire. Ductility is especially important in metalworking and jewelry-making, as metals that crack or break under stress cannot be manipulated using processes such as hammering, rolling, and drawing. Malleable materials can be formed using stamping or pressing, whereas brittle metals and plastics must be molded. It is important to note that while, generally speaking, metals are described as ductile, brittle metals do exist. There are indeed metals and alloys (mixtures of various metals) that are very brittle, under certain conditions, especially at low temperature. 14. Instruct students to record what they think the atomic structure of metals that makes them ductile. Allow a moment for students to reflect, and then ask them to record their thoughts in their notebooks under the term under the heading under the heading ductility. 15. Allow students to share and compare their ideas with those around them. Listen attentively for student thinking and the vocabulary terms they use to describe the situation. DO NOT provide them the answer at this time. 16. Ask for volunteers to share what they wrote about metals atomic structure that they think influences the properties of thermal and electrical conductivity, malleability, and ductility. Again, wait to confirm answers. EXPLAIN IV Metallic Bonding 1. Pose the following question: If you had to pick one subatomic particle responsible for the properties of metals, which one and how? Electrons. The electrons in metals are mobile. They act as electrical and heat energy conductors and allow metal atoms to move over each when a bending or other stress is applied. Day 9 (continued) Instructional Notes: The delocalized electrons in metallic bonds are free to move throughout the solid lattice. These mobile electrons can act as charge carriers in the conduction of electricity or as energy conductors in the conduction of heat. The delocalized electrons in the sea of electrons in the metallic bond also result in the metal Last Updated 06/05/2013 page 9 of 29

Note: At this point in the lesson, accept all answers then go on with the explain activity. This question is used as a focus question for the day s activity. 2. Use an interactive and/or tutorial website to explore the basics of metallic bonds with students (see Advance Preparation). 3. Continue the discussion: How does metallic bonding create the unique properties of metals? The many physical properties of metallic substances, such as thermal and electrical conductivity, malleability, and ductility, are explained by the freely moving nature of electrons in metallic bonds. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days atoms being able to roll over each other when a stress is applied, accounting for their ductility and malleability. Science Notebooks: Students take notes about metallic bonding and generate a statement summarizing the nature of electrons in metallic bonds. 4. Instruct students to write a summary statement in their notebooks, using 25 words or less, that describes the behavior of electrons in metallic bonds and resultant properties of metals. 5. Ask students to share with a partner, and then ask for volunteers to share out with the class. Address any student misconceptions at this time. Allow students to revise and/or add to their statements. ELABORATE Shape Matters Suggested Days 10 and 11 1. Divide the class into pairs. Distribute a copy of the Handout: Molecular Geometry Determination Using VSEPR Theory to each student. 2. Introduce the Valence Shell Electron Pair Repulsion (VSEPR) theory and how it can be applied to predict shapes of covalently bonded molecules. 3. Model the steps of determining the molecular geometry of carbon dioxide as shown on the Handout: Molecular Geometry Determination Using VSEPR Theory. 4. Ask students to review the Electron-Pair Geometry and Polarity Table in their handout. 5. Based on the information provided on this table, ask students to discuss what they think determines the polarity of a molecule with their partner. You might make the analogy that electrons not in a bond are repulsive, and non-central atoms will go as far away from them as they can. NOTE: Students may realize that the number of lone pairs of electrons affects the polarity of a molecule. If there are no lone pairs, the molecule is non-polar. One or more lone pairs may result in to a polar molecule - the more lone pairs, the more polar the molecule, depending on the shape. 6. Instruct students to work with their partner for the remainder of the Elaborate activity. 7. Give each pair five toothpicks and five mini marshmallows to use for construction of molecular models. 8. Project the formulas for BeH 2, BF 3, SO 2, CH 4, CCl 4, and NH 3 on the board. 9. Assign each partner group one of these six molecules and, using the process provided on the handout, instruct groups to determine the shape and polarity of their assigned molecule. Additionally, have groups construct a toothpick/marshmallow model. Students will draw labeled models in their notebooks. 10. Monitor and assist students as necessary. 11. Once the groups have completed the task, ask them to compare their outcomes with the other pairs assigned the same molecule. 12. Each of the molecule groups should come to a consensus as to its geometry, model, and polarity. Materials toothpicks (5 per pair) mini marshmallows (5 per pair) glue or tape (per group) Attachments: Handout: Molecular Geometry Determination Using VSEPR Theory (1 per student) STAAR Notes: Using the Periodic Table to identify and explain periodic trends will be tested as Readiness Standards under Resporting Category 1: Matter and the Periodic Table. Electron configuration and Lewis valence electron dot structures will be tested as Readiness Standards under Responding Category 2: Atomic Structure and Nuclear. Electron dot formulas illustrating ionic and covalent bonds will be tested as Readiness Standards under Reporting Category 3: Bonding and Chemical Reactions. Science Notebooks: Students should create data tables in their notebooks to record data and sketches. Additionally, handouts are also affixed in their notebooks. 13. Choose one member of each molecule group to present their results to the rest of the class. 14. Ask students to create a data table in their science notebooks to record the results for each molecule, including lone pair, bonding pair, total number of structural electron pairs, polarity, and a sketch for each molecule. 15. Clarify any misconceptions as groups present their results to the class. 16. Project and discuss the following questions: How can the VSEPR theory be applied to predict shapes of molecules? Unpaired electron repulsion determines the shape of a given molecule. Last Updated 06/05/2013 page 10 of 29

How is polarity related to the geometric shape of a molecule? If the geometrical shape is unbalanced when there are polar bonds, then the molecule will be polar. Science Unit: 04 Lesson: 01 Suggested Duration: 13 days NOTE: If the geometrical shape is balanced and there are polar bonds between different elements, then the molecule can also be polar. 17. Allow a moment for students to reflect, and then ask them to record their thoughts in their notebooks. 19. Ask students to share with their partner, and then ask for volunteers to share out with the class. Clarify misconceptions at this time. Allow students to revise and/or add to their statements. EVALUATE Performance Indicator Chemical Bonding Showcase Suggested Days 12 and 13 High School Science Unit 04 PI 01 Given three compounds (one ionic, two covalent), use the concepts of VSEPR theory, intermolecular forces, electronegativity, and chemical bonds to explain the properties of each of the substances. Represent each compound on a visual display, such as a poster, including the electron configuration, Lewis valence electron dot structure, and shape and polarity of each. In addition, summarize the nature of metallic bonding for a given metal and how the bonding relates to the properties of the metal. Standard(s): C.2I, C.6E, C.7C, C.7D, C.7E ELPS ELPS.c.1C, ELPS.c.4G, ELPS.c.5G 1. Refer to the Handout: Visual Display PI and Teacher Resource: Performance Indicator Instructions KEY for information on administering the assessment. Materials: research materials or access to student computers/internet poster board (or other materials or software for student visuals, per student) materials for 3 D modeling (per student) Optional Attachments: Teacher Resource: Performance Indicator Instructions KEY Handout: Visual Display PI (1 per student) Last Updated 06/05/2013 page 11 of 29

It s Elementary It s Elementary It s Elementary It s Elementary It s Elementary It s Elementary

HS/Science Unit: 04 Lesson: 01 Electron Configurations (First 20 Elements) 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 2 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 6 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2 2s 2 2p 6 3s 1 2012, TESCCC 08/16/12 page 1 of 1

Orbitals and Electrons Note-taking Guide HS/Science Unit: 04 Lesson: 01 Orbitals and Electrons: Record three key pieces of information: 1. 2. 3. Color and label the s-block elements on the blank periodic table provided. The S Orbital: Shape: Name the geometry of and draw the s orbital shape. # of s orbitals: Maximum # of electrons accommodated: The P Orbital: Shape: Name the geometry of and draw the p orbital shape. # of p orbitals: Maximum # of electrons accommodated in the p orbital: Color and label the p-block elements on the blank periodic table provided. The D Orbital: # of d orbitals: Maximum # of electrons accommodated in d orbital: Color and label the d-block elements on the blank periodic table provided. The F Orbital: # of F orbitals: Maximum # of electrons accommodated in f orbital: Color and label the f-block elements on the blank periodic table provided. 2012, TESCCC 04/19/13 page 1 of 2

Three important rules to remember when constructing orbital diagrams: HS/Science Unit: 04 Lesson: 01 1. Aufbau Principle: Start filling electrons in at the lowest energy level, and build up to the higher energy levels only after the lowest are filled. 2. Hund s Rule: States that each p, d, or f orbital must receive one electron before any p, d, or f orbital can receive a second filling electron Correct representation Incorrect representation 3. Pauli Exclusion Principle: An orbital can hold zero, one, or two electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins. Correct representation Incorrect representation 2012, TESCCC 04/19/13 page 2 of 2

HS/Science Unit: 04 Lesson: 01 2012 TESCCC 08/16/12 page 1 of 1

Lewis Dot Structures

Drawing Lewis Dot Structures for Compounds HS/Science Unit: 04 Lesson: 01 Background Information: In chemical compounds, atoms can either transfer or share their valence electrons to form bonds. When valence electrons are transferred, the bond is called ionic. When electrons are shared, the bond is called covalent. When one or more electrons from a metallic atom are transferred to a non-metallic atom, an ionic compound is formed. The metallic ion, the cation, is positively charged since it lost one or more electron. The non-metallic ion, the anion, is negatively charged since in gained one or more electrons. By transferring electrons, each ion achieves the full valence shell noble or inert gas electron configuration of its row in the periodic table. The positive and negative ions attract each other in ionic compounds, which are called salts. Note: all ionic compounds are called salts, not to be confused with table salt, NaCl. When one or more pairs of valence electrons are shared rather than transferred, a covalent bond is formed. By sharing electrons, each atom also achieves the noble or inert gas valence shell electron configuration of its row of the periodic table. Sharing of electrons in a covalent bond may be equal or unequal. When the electron sharing is equal, as in nitrogen gas (N 2 ), the bond is termed non-polar covalent. In hydrogen chloride gas (HCl), however, the sharing is unequal. The electrons are more strongly attracted to the Cl than the H, making the Cl slightly negative and the H slightly positive. The bond is termed polar covalent. Lewis dots represent the valence electrons in an atom, molecule, or ion. A Lewis Dot Structure then is made by arranging the atomic symbols and dots to show the sharing or transfer of valence electrons. Process for a Covalent Compound: 1. Draw the Lewis dot structure for each atom of the molecule (remember: valence electrons only). For example, the carbon atom in CO 2 has four valence electrons, and the oxygen atom has six valence electrons (1, 2) 2013, TESCCC 04/19/13 page 1 of 3

HS/Science Unit: 04 Lesson: 01 2. Determine the placement of the atoms in your molecule using electronegativity (EN) values. If there is more than one type of atom, place the least electronegative atom in the center of your diagram. NOTE: Electronegativity is a measure of the ability of an atom in a molecule to attract electrons. Electronegativity values can be found on the periodic table available at http://upload.wikimedia.org/wikipedia/commons/4/42/electronegative.jpg. Note how electronegativity decreases from right to left and top to bottom of the periodic table. Based on electronegativity values (O = 3.44, C = 2.55), the arrangement of atoms for CO 2 is: 3. Count the electrons shared around each atom; are the octets complete? If so, your Lewis dot structure is complete. 4. If not, arrange the Lewis dots so that each atom has eight dots in their shared valence shells. In the graphic (3) below, each carbon dot has been replaced by X. 5. Redraw the dots so that electrons, on any given atom, are in pairs wherever possible. When the pair is shared between the atoms, draw a line indicating a bond. One bond between adjacent atoms is termed a single bond. Two bonds between adjacent atoms are termed double bonds. Three bonds between adjacent atoms are termed triple bonds. (4) 6. Practice by drawing Lewis dot structures for the following molecules in your science notebook: H 2 O, BeH 2, BF 3, SO 2, CH 4, CCl 4, NH 3. 2013, TESCCC 04/19/13 page 2 of 3

Procedure for an Ionic Compound: HS/Science Unit: 04 Lesson: 01 1. Draw the Lewis dot structure for each atom of the compound to show how many valence electrons are present in each atom. For example, the sodium atom in sodium chloride, NaCl, has one valence electron and the chlorine atom has seven valence electrons (5, 6). 2. When sodium (Na) and chlorine (Cl 2 ) react, each sodium transfers its one valence electron to a chlorine atom. It becomes a sodium cation with a 1+ charge, and each chlorine atom becomes a chloride anion, with a 1- charge(7). The Lewis structures below show the sodiums with no dots (electrons), and the chloride ions with a complete octet. Notice the placement of the charge notation on the ions and the use of brackets for NaCl. 3. Practice by drawing Lewis dot structures for the following ionic compounds in your science notebook: MgCl 2, LiI, BeO. Image Sources Gancher, M. (Artist). (2011, February 12). Carbone lewis [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:carbone_lewis.svg Adrignola. (Artist). (2011, September 30). Lewis dot O [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:lewis_dot_o.svg Mills, B. (Artist). (2010, March 3o). Carbon-dioxide-octet-dot-cross-2D [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:carbon-dioxide-octet-dot-cross-2d.png Johnson, J. (Artist). (2006, April 27). Lewis carbon dioxide [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:lewis_carbon_dioxide.gif Adrignola. (Artist). (2011, September 30). Lewis dot Na [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:lewis_dot_na.svg Apostoloff. (Designer). (2009). Lewis dot Cl [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:lewis_dot_cl.svg Eloy. (Designer). (2006, February 26). NaCl-Obtención-2 [Print Graphic]. Retrieved from http://commons.wikimedia.org/wiki/file:nacl-obtención-2.svg 2013, TESCCC 04/19/13 page 3 of 3

Molecular Geometry Determination Using VSEPR Theory HS/Science Unit: 04 Lesson: 01 The geometry of a molecule (or ion) is determined by how many structurally significant electron pairs (lone pairs and bonding groups of electrons) there are in the central atom. Follow the procedure below to determine the molecular geometry of a molecule or ion. Carbon dioxide, CO 2, is provided as an example to step through the process. STEP 1 Write the Lewis dot structure for the molecule or ion. STEP 2 Determine the total number of lone pairs and bonding groups of electrons on the central atom. Bonding groups # of lone pairs on central atom # of bonding groups of e - on central atom Total # of groups of e - on central atom 0 2 2 STEP 3 Identify the electron pair geometry from the chart on the reverse of this page: 2012, TESCCC 06/05/13 page 1 of 2

HS/Science Unit: 04 Lesson: 01 Electron-pair Geometry Table and Polarity Structural e- pairs (total) 2 3 4 Electron-pair Geometry and Corresponding Polarity Linear (non-polar) 0 lone pairs of e- 2 bonding pairs of e- Trigonal Planar (non-polar) 0 lone pairs of e- 3 bonding pairs of e- Tetrahedral (non-polar) 0 lone pairs of e- 4 bonding pairs of e- Bent (polar) 1 lone pair of e- 2 bonding pairs of e- [no image available] Trigonal Pyramidal (polar) 1 lone pair of e- 3 bonding pairs of e- Bent (polar) 2 lone pairs of e- 2 bonding pairs of e- 2012, TESCCC 06/05/13 page 2 of 2