"IONIC/COVALENT BONDING AND MOLECULAR GEOMETRY"
IONIC/COVALENT BONDING AND MOLECULAR GEOMETRY BONDING AND MOLECULAR GEOMETRY An Educational Resource Kit Developed by: Kris Kleeman Triad High School April 8, 2008 Libraries Share PERKS (Providing Educational Resource Kits) LSTA LINCC Project 08-5232 “Funding for this grant was awarded by the Illinois State Library (ISL), a division of the Office of Secretary of State, using funds provided by the Institute of Museum and Library Services (IMLS), under the federal Library Services and Technology Act (LSTA).” Table of Contents BONDING and MOLECULAR GEOMETRY Intended Audience……………………………………………………………………1 Objectives……………………………………………………………………………..1 Illinois Learning Standards……………………………………………………………2 Time Frame……………………………………………………………………………3 List of Activities ………………………………………………………………………3 Additional Materials Needed…………………………………………………………..4 Evaluation Tools……………………………………………………………………….5 Source Citations……………………………………………………………………….11 Appendices: A. Contents Checklist B. Lecture Notes C. Demonstrations D. Electronegativity Worksheet E. Comparing Ionic and Covalent Bonds Lab F. Molecular Model Lab Intended Audience The following unit can be used in physical science or chemistry classes in grades 9-12. Objectives 1. Define chemical bond. 2. Relate chemical bond formation to electron configuration. 3. Describe the formation of positive and negative ions. 4. Describe the formation of chemical bonds. 5. Account for the physical and chemical properties of ionic compounds. 6. Describe metallic bonds. 7. Explain the physical and chemical properties of metallic bonds. 8. Apply the octet rule to atoms that bond covalently. 9. Describe the formation of single, double, and triple covalent bonds. 10. List 5 basic steps used in drawing Lewis structures. 11. Explain why resonance occurs, and identify resonance structures. 12. Identify exceptions to the octet rule. 13. Discuss the VSEPR bonding theory. 14. Predict the shape of a molecule. 15. Describe how electronegativity is used to determine bond type. 16. Compare ad contrast polar, non-polar bonds and molecules. 17. Describe the characteristics of compounds that are covalently bonded. 18. Describe the structure and shape of a water molecule. 19. Identify properties of water due to its shape. 20. List the three types and strengths of intermolecular forces. 21. Describe how the formation of a hydrogen bond. Illinois Learning Standards and Benchmarks: State Goal 11: Understand the processes of scientific inquiry and technological design to investigate questions, conduct experiments and solve problems. LEARNING STANDARD A: Know and apply the concepts, principles, and processes of scientific inquiry. SCIENCE PERFORMANCE DESCRIPTORS • 11.A.3 – Stage I -Conduct inquiry investigation, using technologies for observing and measuring directly, indirectly, or remotely. State Goal 12: Understand the fundamental concepts, principles and interconnections of the life, physical and earth/space sciences. LEARNING STANDARD C: Know and apply concepts that describe properties of matter and energy and the interactions between them. • 12.C.5b – Analyze the properties of materials (eg. mass, boiling point, melting point, hardness) in relation to their physical and/or chemical structures. SCIENCE PERFOMRANCE DESCRIPTORS • 12.C.3 – Stage I – Investigate the atomic and nuclear structure of matter, predicting bonding and molecular structure. • 12.C.4 – Stage I – Explain how physical and chemical structure of matter affect its properties relating to bonding types and shapes of molecules to organic and inorganic compounds. • 12.C.1 – Stage J – Explain chemical bonding reactions referencing bonding potential and strengths within and between atoms and molecules. State Goal 13: Understand the relationships among science, technology, and society in historical and contemporary contexts. LEARNING STANDARD A: Know and apply accepted practices of science. SCIENCE PERFORMANCE DESCRIPTORS • 13.A.1 – Stage H& I – Apply appropriate principles of safety within and beyond the science classroom. Time Line Day 1-2: Bonding Notes • Ionic Bonds • Ionic Structures • Covalent Bonds • Metallic bonds • Electronegativity Differences DEMONSTRATIONS: Balloon & electrons, Bond with a Classmate Electronegativity Work sheet Day 2-3: Comparing Types of Bonds Day 3: LAB – Comparing Ionic and Covalent Bonds Day 4: Lewis Structures • VSEPR Theory • Formal Charge • Polarity • Bond Geometry DEMONSTRATIONS – Teacher Built Models Day 5: Models of Covalent Bonds Lab Day 6: Quiz – Lewis Structures Intermolecular Forces • Water Structure DEMONSTRATIONS: Surface Tension, Drops of Water on a Penny, Ice Model, Salt Model Day 7: Review – Finish UP Day 8: Test List of Activities Labs- Models of Covalent Compounds, Comparing Ionic and Covalent Bonds Demonstrations – WorkSheets Test/Quiz ADDITIONAL ITEMS NEEDED OUTSIDE MATERIALS NEEDED: LABORATORY 100 ml Beakers Watch glasses Lab Burners Scoopulas Potassium Iodide Potassium Chloride Sugar Benzoic Acid Camphor Sodium Chloride DEMONSTRATIONS Petri Dishes Pennies Needles Buttermilk Food Coloring Ethyl Alcohol Acetone Water Pippettes Index Cards Source Citations The following websites contain worksheets and exercises to enhance the unit. www.JCE.DivCHED.org A. “Chemical Bonding: The Ties That Bind” Helser, T.L. Journal of Chemical Education. 2003, 80, 414-416. B. “Water Word Search” Journal of Chemical Education. 2005, 82, 551. www.Flinnscientific.com A. Splatter Test – Investigative Strength of Intermolecular Forces Appendix A Contents Check List 12 – Student Model Kits 12 – Deflagrating Spoons 12 – Conductivity Testers w/batteries 1 – Ice structure model 1 – Salt structure model 1 – Reactions and Elements CD ROM 1- PERK CD Bonding • isolated & neutral atoms rarely exist in nature • only noble gases exist independently • all other elements form compounds OCTET RULE: elements form compounds to complete an octet in their outer energy level Chemical Bond: link between atoms that results in the attraction of e- by the nuclei of the atoms. A. IONIC BONDS – bonds that transfer e- from one atom to another – must have a donator & receiver. ** is always exothermic NaCl IONIC STRUCTURES - NaCl B. COVALENT BOND- bond that results from the sharing of electrons. 1. NON-POLAR – equal sharing a. Are symmetrical 2. POLAR – unequal sharing a. Are asymmetrical 3. Produce “molecules” 4. Unshared e- are called “lone pairs” **Usually exothermic, sometimes endothermic H2 HCl • The most electronegative element will attract the e- more, creating a partial (-) and partial (+) charge differential. - sharing 1 pair= single bond - 2 pair = double bond - sharing 3 pair = triple bond C. METALLIC BONDING: bond between metal atoms, d sublevels blend together in a hybridization, creating an e- sea. • Metals – Metals = metallic bond • Nonmetals –Nonmetals = covalent bond • Metals – Nonmetals = ionic bond PROPERTIES OF BONDS: Ionic covalent metallic Bond strong very strong variable Strength Hardness high to very hard low to Moderate brittle moderate Conductivity only in insulators very good Liquid state Melting Pt moderate to low generally High high Solubility polar solvents low insoluble Solubilities except in acid COVALENT COMPOUND STRUCTURE LEWIS STRUCTURES:representation of valence e- and bonding atoms. 1. Count the number of valence e- 2. Determine the total e- needed and subtract 3. Draw a skeleton and distribute the shared e- the central atom is usually written 1st 4. Distribute the rest of the electrons to other their Other positions. VSEPR THEORY: representative geometry of compounds • unshared e- have more energy than share pairs • represented by species types A –represents central atom X- represents number of atoms bonded to central atom E- represents unshared pairs of e- around the central atom FORMAL CHARGE: =(# of valence e- in free atom) -(# of valence e- in bonded atom) = (# of valence e- ) – (½ # of bonding e-) – (#of non-) in free atom bonding e- NH4+ SO2 INTERMOLECULAR FORCES: Forces of attraction between molecules – weaker than bonds that hold atoms together. 1. DIPOLE-DIPOLE FORCES – polar molecules • (-) region attracts (+) regions • polar molecules can induce a dipole on a nonpolar molecule 2. LONDON DISPERSON FORCES/VAN DER WAALS FORCES – nonpolar molecules • result from constant motion of e- and instantaneous dipoles. • Only forces between noble gases & polar molecules • Forces increase as mass increases 3. HYDROGEN BONDS: strongest type of dipole bond • attraction of H and unshared e- on another electronegative element • accounts for high boiling pt of H containing cmpds • strongest type of bond in nature- capable of absorbing a lot of energy. * COHESION * ADHESION WATER STRUCTURE: Bent, Covalent, Polar, Hydrogen bonded - form molecular grouping of 6-9 molecules - rigid solid structure - hexagonal PROPERTIES: due to H-bonding • high specific heat • high melting point • high boiling point • capillary action • high surface tension • high cohesion Lewis (Electron) Dot Diagrams example: NH3 1. Calculate valence electrons for each atom, then total those needed. 3 H need 2 electrons each 3*2 = 6 1 N needs 8 electrons 1*8 = 8 Total 14 electrons needed 2. Calculate valence electrons available. If molecule is charged, add an electron for each negative charge, subtract an electron for each positive charge. 3 H have 1 valence electron each 3*1 = 3 1 N has 5 valence electrons 1*5 = 5 Total 8 electrons available. 3. Find the difference between electrons needed and electrons available: Number of bonding electrons. Bonds = bonding electrons/2 14 electrons needed −8 electrons available 6 electrons shared, for 3 bonds 4. Sketch the molecule. 5. Fill in electron pairs to fill the octet (2 for H). 6. If the number of bonding pairs exceeds the minimum needed to form single bonds, double or triple bonds are used. ex: SO3 Valence Shell Electron Pair Repulsion (VSEPR) Model Rules: 1. The preferred arrangement of electrons in the valence shell is to maximize their distance apart. (This minimizes their repulsion.) 2. A nonbonded pair of electrons takes up more space than a bonded pair. 3. Electron pairs in multiple bonds are considered as one pair of electrons. 4. Electron pairs in a multiple bond occupy more space than a single bond pair. Predicted Arrangement of Electron Pairs: Number of Pairs Arrangement 2 linear 3 trigonal planar 4 tetrahedral 5 trigonal bipyramidal 6 octahedral 7 pentagonal bipyramidal 8 square antiprism Some Notation: A = Central Atom B = Bonding Pair of Electrons E = Lone Pair of Electrons Consider 2 electron pairs: AB2 linear B A B 180° Consider 3 electron pairs: B AB3 trigonal planar A B 120° B B AB2E bent A B Consider 4 electron pairs: B 109.5° AB4 tetrahedral A B B B A AB3E trigonal pyramidal B B B B A AB2E2 bent B Consider 5 electron pairs: B 90° axial to equitorial B AB5 trigonal bipyramidal A B B 120° equitorial B B to equitorial B AB4E see-saw A B B B AB3E2 T-shaped A B B B AB2E3 linear A B Consider 6 electron pairs: B B B AB6 octahedral A 90° B B B B B B AB5E square pyramidal A B B B B AB4E2 square planar A B B Procedure for Working Problems: “VSEPR formula” has the form ABnEm 1. Determine the number of bonding valence electron pairs. number of B’s = number of bonding electron pairs = n 2. Calculate the total number of valence electrons (TVE). (# of valence electrons on A) + (# of B’s)(# of valence electrons on B) + (anion/cation correction) = total number of valence electrons anion/cation correction: anions – add an electron for each negative charge cations – subtract an electron for each positive charge 3. Calculate the number of bonding valence electrons (BVE). (# of B’s)(8) = number of bonding valence electrons 4. Calculate the number of nonbonding valence electron pairs. (TVE – BVE)/2 = number of nonbonding pairs = m 5. The total number of electron pairs is n + m. An alternative way to predict shapes is to draw the correct Lewis (Electron) Dot Diagram and base the VSEPR formula on the observed number of bonding and nonbonding electron pairs. Examples: 1. BF3 B is in group 13, F is in group 17 Lewis Structure: F n = 3 3 bonding pairs F B F 3 + 3(7) = 24 TVE 3(8) = 24 BVE F 0/2 = 0 nonbonding pairs = m F B B F Therefore BF3 → AB3 → trigonal planar F F F 2. XeF4 Xe is in group 18, F is in group 17 Lewis Structure: F n = 4 4 bonding pairs F Xe F 8 + 4(7) = 36 TVE F 4(8) = 32 BVE F F 4/2 = 2 nonbonding pairs = m Xe F F Therefore XeF4 → AB4E2 → square planar 2− 2− O 3. SO3 S is in group 16, O is in group 16 Lewis Structure: n = 3 3 bonding pairs O S O 6 + 3(6) + 2 = 26 TVE 3(8) = 24 BVE 2− 2/2 = 1 nonbonding pair = m S O 2− O Therefore SO3 → AB3E → trigonal pyramidal O DEMONSTRATIONS: ELECTRICITY - MOVEMENT OF ELECTRONS Van Der Graff Generator Balloon – Rub a balloon on a student’s hair, preferably one without much hair spray or hair gel. Then, stick it on a wall or hold it next to a small stream of water. The charged balloon will stick to the wall or cause the water stream to bend. BOND WITH A CLASSMATE ACTIVITY A great way to demonstrate the concept of bonding is through the following: Give each student a different ion (ex. H+1, S-2, Ca+2) on an index card and have them tape it to the front of their shirt. They must them go around the room and find other people that they would be able to bond with. (Charges must equal zero). They must record two or three options. SURFACE TENSION-INTERMOLECULAR FOCES Float a needle on water in a small Petri Dish. This can be done on the overhead. Students can see the needle independently moving around the dish. Add some talcum powder to a Petri dish of water. The powder will float on the top. Take a small drop of liquid soap and drip it in the center of the dish and watch the powder slowly sink to the bottom as the soap breaks the surface tension. Pour some buttermilk into a Petri dish. Put one drop of food coloring anywhere in the milk, except the center. Repeat this with three other colors, making sure not to get the colors too close to one another. Carefully drip one drop of dishwashing soap in the center of the milk and watch what happens. The more fat in the milk the more effective the results as the bonds of fat are broken by the soap. Drop small drops on the head of a penny. See how many drops can be placed. Experiment with different liquids to determine if they increase or decrease surface tension (intermolecular forces) Bonding Reactions Name:_____________ Hour:_____________ Determine the type of bonding to occur in each case considering the types of elements involved and electronegativities. Then, use electron dot diagrams to show how the bonding takes place. 1. sodium & chlorine Metal or Not? Electro- negativity Bond Type 2. potassium & fluorine Metal or Not? Electro- negativity Bond Type 3. lithium & sulfur Metal or Not? Electro- negativity Bond Type 4. carbon & hydrogen Metal or Not? Electro- negativity Bond Type 5. beryllium & sulfur Metal or Not? Electro- negativity Bond Type Bonding Reactions Name:_____________ Hour:_____________ 6. magnesium & nitrogen Metal or Not? Electro- negativity Bond Type 7. silicon & fluorine Metal or Not? Electro- negativity Bond Type 8. aluminum & nitrogen Metal or Not? Electro- negativity Bond Type 9. carbon & chlorine Metal or Not? Electro- negativity Bond Type 10. carbon & fluorine Metal or Not? Electro- negativity Bond Type COMPARING IONIC and COVALENT COMPOUNDS Name___________________ Hour:___________________ Partner:__________________ Purpose: • To compare physical properties of ionic and covalent compounds • To determine whether an unknown substance is ionic or covalent Pre-lab Questions: 1) Based on what you’ve learned about electronegativity differences and bond types, predict which group of elements (A-C) or (D-F) are ionic and which are covalent. 2) Outline the procedure. Materials: Lab burner sodium chloride (A) Conductivity apparatus potassium iodide (B) Scoopula potassium chloride (C) Deflagrating spoon benzoic acid (D) Watch glass sugar (E) Beaker camphor (F) Unknown solid Saftey precautions: Safety goggles must be worn. Use caution when heating compounds. Follow instructions for the conductivity meter. PROCEDURE: 1) Create a chart for your observations. 2) Collect a small pea size sample of each compound 3) Try to detect if any substances have an odor by using the wafting method. MELTING POINT 4) Light the burner. Place a small sample of substance A into a deflagrating spoon and a small sample of substance D into the other. Record the one that melts first as low and the other as high. You do not need to melt the higher one. Repeat with substances B and E, C and F. 5) Repeat step 4 with your unknown. HARDNESS 6) Determine whether each substance seems hard or soft. Do this by rubbing a small sample of each between your fingers. Try to crush a few crystals of each. Place each compound in the watch glass at a time and use a scoopula to crush them. CONDUCTIVITY COMPARING IONIC and COVALENT COMPOUNDS Name___________________ Hour:___________________ Partner:__________________ 7) Make a solution of each compound using distilled water in separate beakers. 8) Test the conductivity of each solutions, including your unknown. Post-Lab Questions: 1) Compare the physical properties of ionic and covalent substances. 2) What is the number of your unknown? Is it ionic or covalent? 3) Intermolecular forces exist between molecules or ions. Based on this lab, which do you think are stronger, intermolecular forces between covalent compounds or intermolecular forces between ions? 4) Covalent compounds can be found as single molecules, but ionic substances can’t be found as single ion pairs. What evidence do you have for this fact? 5) Which type of substance conducted electricity when dissolved in water? What is the difference between ionic and covalent substances that causes this? Models of Molecular Compounds InTRODUCTION: Geometries of molecules can be determined by examining the central atom and identifying the number of atoms bonded to it and the number of unshared electron pairs surrounding it. The shapes of molecules may be predicted using VSEPR theory. VSEPR theory states that electron pairs around the central atom will position themselves to allow for maximum amount of space between them. In addition, electrons unshared by the central atom will possess more energy that shared electron pairs, which will result in unshared electrons occupying more space than shared pairs. Covalent bonds can be classified by comparing electronegativity differences between the bonded atoms. If the electronegativity difference us less than or equal to 0.4, the bond is non-polar covalent. Differences between 0.5 and 1.7, a polar covalent bond exists. Differences greater than 1.7, result in an ionic bond. In a polar covalent bond, the electrons will be more attracted to the atom of greater electronegativity, resulting in a partial negative and partial positive charge on the atom. Molecules made up of covalently bonded atoms can be either polar or non-polar. One indication of polarity is symmetry. For example, if polar bonds are arranged symmetrically around a central atom, their charges cancel out and the molecule is non-polar. On the other hand, the arrangement of the polar bonds is asymmetrical, the molecule will be polar. Molecular geometry determines the molecule’s boiling point, freezing point, viscosity, and the amount and type of its reactions. In this experiment, you will construct models of covalent molecules and predict geometry and polarity of each molecule. PRE-LAB 1. Define covalent bond: __________________________________________________________________ 2. What is a dipole? _____________________________________________________________________ 3. What two factors determine whether a molecule is polar or non-polar? ___________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ 4. Calculate the electronegativity difference and predict the bond type for each of the following examples (use your textbook to find the electronegativities of the elements.) a. Na-Cl ________________________________________________________________ b. C-H ________________________________________________________________ c. S-O ________________________________________________________________ d. N-N ________________________________________________________________ OBJECTIVES Determine the molecular geometry of molecules. Determine the polarity of covalent molecules MATERIALS Safety Goggles ball-and-stick model kit PROCEDURE 1. Put on your goggles. Construct ball-and-stick models of the compounds listed in the data table. 2. Complete the Data Table and Observation section. As an example, the first line of the Data Table has been filled in for you. 3. When you have finished the investigation, take your models apart and return the set to your teacher. FORMULA ELECTRON STRUCTURAL SHAPE OF POLARITY DOT FORMULA MOLECULE STRUCTURE H2 H:H H-H LINEAR NON POLAR HCl H2O NH3 CH4 HClO N2 CH3NH2 CO2 H2CO C2H2 CH3Cl HCOOH HCN H 2O 2 POST LAB 1. List the advantages and disadvantages of using the models to construct molecules: ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ 2. Explain how you used the shapes to predict molecular polarity. Support your answer with examples from the results of this investigation: ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ 3. Based on your results, predict the geometry and polarity of the following molecules: a. HI ________________________________________________________ b. SCI2 ______________________________________________________ c. PH3 _______________________________________________________ d. SO2 _______________________________________________________ 4. The polarity of a substance can effect its reactivity and solubility. “Like dissolves Like” is a general rule for predicting solubility. What can you predict about the polarity of alcohol if you know that alcohol dissolves in water? Why do you thin that water is not used to dissolve greasy stains and dirt at dry cleaners? ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________ ______________________________________________________________________________________