Honors Chemistry - Unit 9 Chapter 6: Bonding & Molecular Structures. Unit 9 Packet Page 1 of 14

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1 Honors Chemistry - Unit 9 Chapter 6: Bonding & Molecular Structures Unit 9 Packet Page 1 of 14 Vocab Quiz: UT Due: Test Date: Quiz Date(s): FORMULAS/CONSTANTS Memorize VSEPR Chart First 6 Shapes!! OBJECTIVES: Chapter 6 Be able to describe the difference between polar and nonpolar molecular bonds. Be able to describe single, double and triple bonds. Be able to describe the three intermolecular forces (Van der Waal s forces): dipole-dipole, London dispersion and hydrogen bonding. Be able to draw Lewis structures for various compounds and polyatomic ions. Be able to apply the VSEPR theory in determining molecular geometry. Be able to describe hybridization and give an example. VOCABULARY: 1. Ionic bonds 2. Covalent bonds 3. Molecule 4. Metallic bond 5. Octet rule 6. VSEPR 7. Single bond 8. Double bond 9. Triple bond 10. Polar 11. Nonpolar 12. Hydrogen bonding 13. Dipole-dipole force 14. (London) dispersion force 15. Hybridization 16. Intermolecular force (IMF)

2 Unit 9 Packet Page 2 of 14 VSEPR & Molecular Geometry (You must memorize the first six (6) shapes in this table!!) Molecular Shape Type of Molecule AByEz Atoms Bonded to Central Atom Lone Pairs of e-s on Central Atom Linear AB2 2 0 Bent Trigonal Planar Tetrahedral Trigonal Pyramidal AB2E AB3 AB4 AB3E Bent Trigonal bipyramidal AB2E2 AB Octahedral AB6 6 0 Bond Type Bond Types Table Type(s) of Atoms involved Force Properties Examples IONIC transfer of electrons COVALENT Sharing of electrons METALLIC Free flow of Electrons Metal & Nonmetal Two Nonmetals Polar = unequal sharing = partial charge Nonpolar = equal sharing = no charge Two Metals Attraction between ions, opposite charges attract; transfer of electrons Sharing of electrons Sharing of electrons between all atoms High melting point Water soluble Crystalline Aquesous solutions conduct a current Low melting point Brittle Nonconductors Good conductors Malleable Ductile NaCl MgO CaS Water CO2 NH3 Copper wire Iron bar

3 Unit 9 Packet Page 3 of 14 CHAPTER 6 - CHEMICAL BONDING Pages Write the FULL electron configurations for the following elements: Calcium, Phosphorous, Iron, and Sodium 2. Write the noble gas notation electron configuration for the following elements: Tin, Arsenic, Barium, Lead 3. What kind of bonding is found in polyatomic ions? Ionic Bonds The forces that hold matter together are called chemical bonds. There are four major types of bonds. We need to learn in detail about these bonds and how they influence the properties of matter. The four major types of bonds are: I. Ionic Bonds III. Metallic Bonds II. Covalent Bonds IV. Intermolecular (van der Waals) forces Ionic Bonds The ionic bond is formed by the attraction between oppositely charged ions. Ionic bonds are formed between metals and nonmetals. Remember that metal atoms lose one or more valence electrons in order to achieve a stable electron arrangement. When a metal atom loses electrons it forms a positive ion or cation. When nonmetals react they gain one or more electrons to reach a stable electron arrangement. When a nonmetal atom gains one or more electrons it forms a negative ion or anion. The metal cations donate electrons to the nonmetal anions so they stick together in an ionic compound. This means that ionic bonds are formed by the complete transfer of one or more electrons. A structure with its particles arranged in a regular repeating pattern is called a crystal. Because opposite charges attract and like charges repel, the ions in an ionic compound stack up in a regular repeating pattern called a crystal lattice. The positive ions are pushed away from other positive ions and attracted to negative ions so this produces a regular arrangement of particles where each ion is surrounded by ions of the opposite charge. Each ion in the crystal has a strong

4 Unit 9 Packet Page 4 of 14 electrical attraction to its oppositely charged neighbors so the whole crystal holds together as one giant unit. We have no individual molecules in ionic compounds, just the regular stacking of positive and negative ions. 1. Define the following terms: a) ionic bond b) cation c) anion d) crystal 2. What are the smallest units of an ionic bond? At room temperature ionic compounds are high melting point solids. They are usually white except for compounds of the transition metals that may be colored. They are brittle (break easily). They do not conduct electricity as solids, but do conduct electricity when melted or dissolved in water. 3. List several properties of ionic compounds: 4. When can electricity to be conducted in an ionic bond? Reviewing Lewis Dot Diagrams Write the Lewis Dot Diagrams for the following: helium atom: beryllium atom: beryllium ion: neon atom: aluminum atom: aluminum ion: magnesium atom: magnesium ion:

5 sodium atom: sodium ion: Unit 9 Packet Page 5 of 14 Write the Lewis Dot Diagrams for: oxygen atom: oxide ion: chlorine atom: chloride ion: phosphorus atom: phosphide ion: How would you describe (in general) the Lewis Dot Diagram for: a) a cation? b) an anion? What type of bonding would you expect in a compound that contains a metal and a nonmetal? Worksheet #4: Introduction to Covalent Bonds A covalent bond is formed between nonmetal atoms. The nonmetals are connected by a shared pair of valence electrons. Remember, nonmetals want to gain valence electrons to reach a stable arrangement. If there are no metal atoms around to give them electrons, nonmetal atoms share their valence electrons with other nonmetal atoms. Since the two atoms are using the same electrons they are stuck to each other in a neutral particle called a molecule. A molecule is a neutral particle of two or more atoms bonded to each other. Molecules may contain atoms of the same element such as N 2, O 2, and Cl 2 or they may contain atoms of different elements like H 2 O, NH 3, or C 6 H 12 O 6. Therefore, covalent bonding is found in nonmetallic elements and in nonmetallic compounds. Covalent bonds are intramolecular forces; that is, they are inside the molecule and hold the atoms together to make the molecule. Covalent bonds are strong bonds and it is difficult and requires a lot of energy to break a molecule apart into its atoms. However, since molecules are neutral one molecule does not have a strong electrical attraction for another molecule. The attractions between molecules are called intermolecular forces and these are weak forces. Covalent substances have low melting points and boiling points compared to ionic compounds or metals. At room temperature, covalent substances are gases, liquids or low melting point solids. They do not conduct electricity as solids or when molten and usually do not conduct when dissolved in water. 1. Define the following terms: a) covalent bond

6 Unit 9 Packet Page 6 of 14 b) molecule c) intramolecular force d) intermolecular force 2. List several properties of covalent compounds. There are many types of covalent bonds. A single covalent bond is when two atoms share one pair of valence electrons (see figure). A double covalent bond is when two atoms share two pairs of valence electrons. A triple covalent bond is when two atoms share three pairs of valence electrons. 3. Define polar and nonpolar in terms of BOTH bonds and molecules There is one last type of covalent bonding the bonding in network solids (macromolecules). In this type of bonding, atoms share valence electrons, but the atoms are arranged in a regular crystalline pattern in which each atom is covalently bonded to its neighbors in all directions. Therefore, you do not have a collection of small molecules that are easy to separate from each other; the whole system is one giant molecule or a macromolecule held together by this network of strong covalent bonds. Network solids are extremely hard, brittle, solids that do not conduct electricity. Diamonds (a form of pure carbon (see figure)), carborundum (silicon carbide) and quartz (silicon dioxide) are examples of macromolecules. 1. What is a network solid? 2. What type of bonding exists in network solids? 3. What are some properties of network solids? 4. What are some examples of network solids?

7 HYBRIDIZATION: Mixing of two or more atomic orbitals of similar energy to make new orbitals of equal energies. Example: Carbon Unit 9 Packet Page 7 of 14 You try: Boron MOLECULAR STRUCTURES: Review: Covalent bond =? Octet rule =? H and He are exceptions to octet rule H is stable with 2 e-s, b/c of their 1s 2 shell. LEWIS STRUCTURE: Element Symbol = nuclei and inner-shell electrons Dashes = shared electron pairs in covalent bond Dots = unshared electrons Single bond = Steps for Lewis Structure Problem (and an example) Example: Draw the Lewis Structure of iodomethane - CH3I 1. Determine the number and type of atoms. 2. Determine the total number of valence electrons in the atoms to be combined. (use periodic table) 3. Arrange the atoms to form a skeletal structure. (First atom is central unless otherwise specified or H is first). Add up the electrons used to form the bonds (each bond counts as 2 e- s). Subtract the number of electrons used from the total in step 2. Use these electrons in step Add electron pairs so that each atom is surrounded by 8 electrons (octet rule) (except H it can only hold 2 e-s) starting with the outer atoms and saving the central atom until last. 5. a. Count the number of e-s in the structure to be sure the number of valence electrons used equals the number available and the octet (duet) rule is satisfied for all atoms. **See step 5b below if not satisfied.** Practice: together: ammonia and silicon tetrafluoride

8 You try: BrI, CH3Br Unit 9 Packet Page 8 of 14 Double Bond = Triple Bond= Bond Strength? Bond Length? Need to use multiple bonds when there are not enough valence electrons to complete octets by adding unshared electron pairs (step 5 = numbers do not match) 5. b. If the octet rule has not been satisfied for all atoms, form multiple bonds (double or triple) by moving electron pair until all atoms have been satisfied. Example: (together) CH2O Practice: carbon dioxide and HCN POLYATOMIC IONS: 6. Additional Step for Completing Lewis Structures: when counting valence electrons must include the charge: If positive subtract electrons If negative add electrons Also, brackets must be placed around the structure with the charge of the ion outside the brackets to indicate the gain or loss of e - s in the total. Example: nitrate Practice: sulfate

9 Lewis Structures Practice Worksheet Unit 9 Packet Page 9 of 14 Draw Lewis Structures for the following molecules or ions: 1. TeCl2 2. PCl3 3. Phosphate 4. I2 5. ICl 6. H2S 7. Nitrite 8. Carbonate 9. Water 10. P2 11. Boron trifluoride 12. SF2 13. Bromine Monochloride 14. Ammonium VSEPR Theory valence-shell, electron pair repulsion Way to predict molecular geometry (shape) There is a repulsion between valence e - pairs Steps to Work Problems: (Be sure to refer to and memorize VSEPR table on page 3 of the packet first six shapes) 1) Draw Lewis Structure of the Molecule 2) Put molecule in AByEz form, where: A represents the central atom B represents the atoms bonded to A (y is the # of B atoms; B may different elements.) E represents the lone pair e - s on A (z is the # of lone pairs.) 3) Predict the shape based on the AByEz form Example 1: Use VSEPR Theory to predict the shape of CBr4. AB4 shape - tetrahedral Example 2: sulfur difluoride (SF2) AB2E2 shape bent Example 3: SF6 AB6 octahedral You try: Predict the VSEPR shapes of the molecules on the Lewis Structure worksheet 1-3, 6-9,11,12,14 (Why not try problem # s 4,5,10,13?)

10 MINI LAB THE SHAPE OF MOLECULES Applying the VSEPR theory Unit 9 Packet Page 10 of 14 OBJECTIVE: Through building models (using marshmallows!) we will enhance our understanding of determining molecular shape through use of the VSEPR theory. You will work in groups with your lab partner(s). MATERIALS: Large white marshmallows (central atom) Gum Drops/Gummy Bears (surrounding atoms) Toothpicks (bonds). Red-Hots (round red candy; unshared electron pairs) REVIEW - Pre-Lab Questions: (Before you begin your lab activity answer these questions as a group; one paper per group is fine.) Answer on your own paper. 1. What does the word VSEPR stand for? 2. What structure must you have first before applying the VSEPR theory? ACTIVITY: Before you begin, your teacher will demonstrate the procedure and one example on the board (example: water) PROCEDURE: For the following molecules: A. Draw the Lewis Structure (on your sheet with the pre-lab question answers.) B. Make models of their VSEPR geometry using your notes and supplies. C. Have your teacher initial your paper after she/he has seen your completed models! 1. Ammonia NH3 - (use large white marshmallow for nitrogen, gum drops for hydrogens, toothpicks for bonds and Red-Hots, that have been licked for electron pairs). 2. SiH4 (Si- marshmallow, H - gum drops, etc.) 3. NO3-1 CONCLUSION: Write a conclusion paragraph. Be sure that your teacher sees your finished models and initials your paper when you are finished. After your paper has been checked, you may eat your molecules; (you will probably want to eat the "electrons" you licked!). If you prefer, you may throw away your models. Return any extra unused supplies and turn in this sheet and your group lab sheet.

11 Bonding Pictures Review Sheet Unit 9 Packet Page 11 of 14 Draw Lewis dot diagrams for the following compounds. Remember that you must check the difference in electronegativity if the difference is less than 1.7 draw a covalent structure and if the difference is 1.7 or greater draw an ionic structure. a) water (H 2 O) b) potassium iodide (KI) c) nitrogen molecule (N 2 ) d) nitrate ion (NO 3 ) - e) calcium iodide (CaI 2 ) f) phosphorous trichloride (PI 3 ) g) aluminum fluoride (AlF 3 ) h) carbon dioxide (CO 2 ) i) oxygen molecule (O 2 ) j) magnesium nitride (Mg 3 N 2 )

12 Unit 9 Packet Page 12 of 14 Worksheet #12 Continued k) chlorine molecule (Cl 2 ) l) ethane (C 2 H 6 ) (carbons hook to each other with H s all around) m) hydroxide ion [OH] -1 n) sulfite ion [SO 2 ] -2 o) ammonium ion [NH 4 ] +1 p) sodium oxide (Na 2 O) q) sulfate ion [SO 4 ] -2 r) calcium bromide (CaBr 2 ) s) methane (CH 4 ) t) sulfur dioxide (SO 2 )

13 The Polarity of Molecules Unit 9 Packet Page 13 of 14 Many of the properties of a molecule are directly related to the shape and from there the polarity. For example, water is polar and is a liquid at room temperature. Larger molecules like carbon dioxide, CO 2 which is nonpolar, are gasses. To determine polarity, you must now bring several things together. You must be able to: 1. Find the Lewis structure and VSEPR shape of the molecule. 2. Determine if the molecule is symmetrical. If molecules are symmetrical: they are usually non-polar (if the non-central atoms are all the same, so they can cancel each other out). If a molecule is not symmetrical (the polarities do not cancel and) the overall molecule is polar. Symmetrical VSEPR shapes typically nonpolar (If all surrounding atoms are the same): linear, trigonal planar, tetrahedral. These 3 shapes are nonpolar as long as the surrounding atoms match if they don t match then the molecule is polar. Asymmetrical VSEPR shapes- always polar: bent (both), trigonal pyramidal. Examples together: draw the VSEPR shape first then determine polarity 1)water: 2)carbon dioxide 3)HCN You try: determine the molecule polarity for all of the molecules on your VSEPR worksheet

14 Bellwork Review Sheet Unit 9 Packet Page 14 of What is the total number of electrons needed to fill the fourth energy level? 2. Write the long form and noble gas notation of As 3. Draw the dot diagram for As 4. Given: C Mg S Ba place these elements in order of: A. increasing atomic radius C. increasing electron affinity B. decreasing electronegativity D. decreasing ionization energy 5. Fill in the following table: Type(s) of Atoms Bond Type involved Force Properties One Example 6. Explain why the oxidation number of Magnesium is 2+ in ionic compounds. (It is recommended you use an electron configuration diagram in your explanation.). 7. What type of bond(s) is found in lead (II) phosphate? 8. Draw the Lewis structures and predict the molecular geometries for: a. PCl3 b. Carbonate

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