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1 Instructor Lesson Plan – (PIB Chemistry) Unit 4: Chemical Names and Formulas and Bonding Instructor: D. Ditkowsky (Thanks J. Galinski) Introductory Resources: Addison-Wesley v.5 - Chapter 6, 14, 15, 16 Addison-Wesley v.4 - Chapter 5, 12, 14, 15 Advanced Resources: Brown v.4 - Chapter 5, 7, 8, 9 Zumdahl v.3 - Chapter 2, 7, 8, 9 Zumdahl v.5 - Chapter 2, 7, 8, 9 Main Idea Summary: Every substance is either an element or a compound. A compound consists of more than one kind of atom. A compound is either molecular or ionic in nature. Molecular compounds are composed of two or more nonmetals. A molecular formula shows the number and kinds of atoms present in a molecule of a compound. Ionic compounds are composed of oppositely charged ions combined in electrically neutral groupings. A formula unit gives the lowest whole-number ratio of ions in the compound. The charges of the ions of the representative elements can be determined by the position of these elements in the periodic table. Most transition metals have more than one common ionic charge. A polyatomic ion is a group of atoms that behaves as a unit and has a charge. When a cation can have more than one ionic charge, a Roman numeral is used in the name. There are simple rules for naming ionic and molecular compounds. Elements that have similar properties also have similar electron configurations and are members of the same group. Regular changes in the electron configuration of the elements cause gradual changes in both the physical and chemical properties within a group and within a period. Atomic radii decrease as you move from left to right in a given period. Ionization energy increases as you move from left to right in a given period. Atomic radii increase within a given group because the outer electrons are farther from the nucleus as you go down the group. Ionization energy decreases as you move down through a group. 2 Atoms in compounds are held together by chemical bonds. Chemical bonds result from the sharing or transfer of valence electrons between pairs of atoms. Bonded atoms attain the stable electron configuration of a noble gas. The noble gases themselves exist as isolated atoms because that is their most stable condition. For the representative elements, the number of valence electrons is equal to the element’s group number in the periodic table. The transfer of one or more valence electrons between atoms produces positively and negatively charged ions, or cations and anions respectively. The attraction between an anion and a cation is an ionic bond. A substance with ionic bonds is an ionic compound. Nearly all ionic compounds are crystalline solids at room temperature. They have high melting points. The total positive charge cancels out the total negative charge, yielding a neutral compound. When melted or in solution, ionic compounds can conduct electricity, because the ions can move freely when a voltage is applied. Metals bond by packing positive metal ions together, surrounded by a sea of electrons. The sea of electrons in metallic bonding causes most of the properties of metals (malleability, ductility, electrical conductivity). Atoms form covalent bonds when they share electrons to form an octet. A shared pair of valence electrons constitutes a single covalent bond. Additional sharing of electron pairs constitute the formation of higher order bonds. Resonance structures help to visualize the bonding in molecules when more than one electron dot formula can be written. Molecular orbital theory is a logical extension of the quantum mechanical description of the electron structure of atoms. The Valence Shell Electron Pair Repulsion Theory explains expected molecular geometry of a molecule. 3 4 Common Ion List CATIONS: Aluminum Ammonium Barium Beryllium Calcium Cesium Copper (I) Copper (II) Francium Hydrogen Iron (II) Iron (III) Lead (II) Lead (III) Lead (IV) Lithium Magnesium Potassium Radium Rubidium Silver Sodium Strontium Zinc ANIONS: Al+3 NH4+1 Ba+2 Be+2 Ca+2 Cs+1 Cu+1 Cu+2 Fr+1 H+1 Fe+2 Fe+3 Pb+2 Pb+3 Pb+4 Li+1 Mg+2 K+1 Ra+2 Rb+1 Ag+1 Na+1 Sr+2 Zn+2 Acetate Bromide Carbonate Chlorate Chloride Chlorite Chromate Cyanide Dichromate Fluoride Hydride Hydroxide Hypochlorite Iodate Iodide Nitrate Nitrite Oxalate Oxide Perchlorate Permanganate Phosphate Silicate Sulfate Sulfide Sulfite Thiocyanate Thiosulfate C2H3O2-1 Br-1 CO3-2 ClO3-1 Cl-1 ClO2-1 CrO4-2 CN-1 Cr2O7-2 F-1 H-1 OH-1 ClO-1 IO3-1 I-1 NO3-1 NO2-1 C2O4-2 O-2 ClO4-1 MnO4-1 PO43SiO3-2 SO4-2 S-2 SO3-2 SCN-1 S2O3-2 5 Laboratory Activity 5A – Density is a Periodic Property NAME: _____________________ The Periodic Table was developed by Dmitri Mendeleev in the mid-1800s. Mendeleev put all of the elements in order based on mass, and noticed a periodic reoccurance of chemical and physical properties. He arranged the elements in columns. Elements in each column have similar properties. Occasionally, he would find a hole in the table. He could use the surrounding information to predict the properties of a yet-to-be-discovered element. As new elements were discovered, they neatly filled holes in the periodic table. You will be given samples of Silicon, Tin, and Lead. All of these elements are in the same column of the Periodic Table, and have similar properties. Using a balance, you will find the mass of a sample of each element. Using a graduated cylinder and some water, you will find the volume of each sample. Then, you can determine the density of each element using calculations you have already mastered. Your goal for this lab is to determine the density of Germanium, element #32, based on what you learn about the other elements in the same column. Watch for patterns. Use the data below. Symbol Element Si Silicon Ge Germanium Sn Tin Pb Lead Qualitative Observations Mass (g) Volume (cm3) XXXXXXX XXXXXXX XXXXXXX XXXXXX XXXXXX XXXXXX XXXXXXX XXXXXXX XXXXXXX Density (g/cm3) 6 Periodic Trends NAME: ____________________ 1. Restate in one or two words: “The amount of energy required to remove one electron from the valence shell of a neutral atom.” 2. Restate in one or two words: “The tendency of an atom to hold on to its valence electrons while engaged in a chemical bond.” 3. Restate in one or two words: “The actions of the non-valence electrons, diluting the force of the attraction between nucleus and valence electrons.” 4. Which has greater shielding, Au or Cu? 5. Which is larger, Au or Cu? 6. Which has greater ionization energy, Cu or Ag? 7. Which has greater shielding, Xe or Ar? 8. Which is larger, Ca or Cs? 9. Which has greater shielding, Se or Ra? 10. Which has greater nuclear charge, Zn or Se? 11. Which is larger, Mg or P? 12. Which has greater ionization energy, Fe or K? 13. Restate in one or two words: “The actions of the non-valence electrons, diluting the force of the attraction between nucleus and valence electrons.” 14. Restate in one or two words: “Half the distance between the nuclei of two like atoms.” 15. Restate in one or two words: “The tendency of an atom to attract electrons when it is chemically combined with another element.” 16. Which has greater ionization energy, Cl or Hf? 17. Which has greater shielding, P or Ar? 18. Which is larger, Os or Ta? 19. Which is a larger ion, sulfur ion or phosphorus ion? 20. Which has greater shielding, Ge or Ra? 21. Which has greater nuclear charge, Sb or Se? 7 Laboratory Activity 5B – Periodic Trends Page One NAME: ____________________ Part I: Periodic Trend Graphing Activity Atomic Number 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 Element H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ionization Potential (eV) 13.60 24.59 5.39 9.32 8.30 11.26 14.53 13.62 17.42 21.56 5.14 7.65 5.99 8.15 10.49 10.36 12.97 15.76 4.34 6.11 6.54 6.82 6.74 6.77 7.44 7.87 7.86 Atomic Radius (Å) 0.30 0.93 1.52 0.89 0.88 0.77 0.70 0.66 0.64 1.12 1.86 1.60 1.43 1.17 1.10 1.04 0.99 1.54 2.31 1.97 1.60 1.46 1.31 1.25 1.29 1.26 1.25 Atomic Number 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Element Ni Cu Zn Ca Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Ionization Potential (eV) 7.64 7.73 9.39 6.00 7.90 9.81 9.75 11.81 14.00 4.18 5.70 6.38 6.84 6.88 7.10 7.28 7.37 7.46 8.34 7.58 8.99 5.79 7.34 8.64 9.01 10.45 12.13 Atomic Radius (Å) 1.24 1.28 1.33 1.22 1.22 1.21 1.17 1.14 1.69 2.44 2.15 1.80 1.57 1.41 1.36 1.30 1.33 1.34 1.38 1.44 1.49 1.62 1.40 1.41 1.37 1.33 1.90 1. On a sheet of graph paper, graph the ionization potential (y-coordinate) versus atomic number (x-coordinate). Make sure to properly label the graph. 2. On a sheet of graph paper, graph the atomic radius versus the atomic number. Label the graph. 3. What do the units “eV” and “Å” stand for? What do these measurements mean, in your own words. 8 Part III: Periodic Table Identification Activity 1. On the blank Periodic Chart on Page Three, clearly locate the following, using the following color code: a. b. c. d. e. f. g. h. i. Representative elements : Outline in HEAVY BLACK Transition elements : (Sky Blue) Metallic elements: Color these in shades of (Blue) Nonmetallic elements: Color these in the shades of (Red) Metalloids: Purple (Red and Blue mixed) Alkali metals (Dark Blue) Alkaline-earth metals (Greenish Blue) Halogens (Pink) Noble gases (Yellow) 2. On the same periodic chart, locate these elements and write in their atomic symbols: a. b. c. d. e. sodium potassium chlorine nickel bromine f. phosphorus g. carbon h. magnesium i. sulfur j. calcium k. barium l. aluminum m. silicon n. zinc o. lead 9 Laboratory Activity 5B – Periodic Trends Page Three NAME: ____________________ 10 Laboratory Activity 5B – Periodic Trends Page Four NAME: ____________________ Part IV: Interpretation Questions 1. Notice that the graph of first ionization potential versus atomic number consists of generally rising values followed by sharp drops. List the elements on the five major peaks in this graph. What name is given to this group of elements? 2. List four elements located at the bottom of the sharp drops. What name is given to this group of elements? 3. Assuming that the periodic trends indicated on the graph continue, what value do you predict for the first ionization potential of cesium, Cs, atomic number 55? 4. What generalization can be made about the change in first ionization potential as the atomic number increases in a period (such as Na to Ar)? 5. What generalization can be made about the change in first ionization potential as the atomic number increases in a group (family)? 6. Looking at the atomic radius versus atomic number, what would you predict for the atomic radius of Cs, atomic number 55? (Use Cl-Ar-K and Br-Kr-Rb as examples.) 11 Electronegativity and types of bonds Carefully read text pages 405 – 406, 419, 422 for details about ionic bonds and electronegativity. The difference between the electronegativity of the anion minus the cation determines if a bond is ionic. If the difference in electronegativity is greater than 1.7, the bond is considered ionic because one atom possesses the electron more than 50% of the time. (Anything less than 0.3 is considered nonpolar covalent) H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na Mg 0.9 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr Mn Fe 1.6 1.5 1.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb Mo 1.6 1.8 Tc 1.9 Cs 0.7 Ba 0.9 La 1.0 Hf 1.3 Ta 1.5 Re 1.9 W 1.7 Co 1.9 Ni 1.9 Cu 1.9 Zn 1.6 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Ru 2.2 Rh 2.2 Pd 2.2 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Hg 1.9 Tl 1.8 Pb 1.9 Bi 1.9 Po 2.0 At 2.2 Percent Ionic Character 0% 5% non(p.c.) polar covalent 0 0.3 50% ionic 1.7 100% 3.3 1. What must always be true if a covalent bond is to be polar? 2. The bonds between the following pairs of elements are covalent. Arrange them according to polarity, naming the most polar bond first. H—Cl, H—C, H—F, H—O, H—H, S—Cl 3. Arrange the following bonds in order of increasing ionic character. Cl—F N—N K—O C—H S—O Li—F 12 More Bonding Concepts NAME: ______________________ 1. Write electron configurations for the following and comment on the result: a. N3− ________________________ b. O2− ________________________ c. F1− ________________________ d. Ne ________________________ 2. Explain why molten MgCl2 does conduct an electric current although crystalline MgCl2 does not. 3. Write the correct chemical formula (the formula unit) for the compounds formed from each pair of ions. a. K+, S2− ____________ b. Ca2+, O2− ____________ c. Na+, SO42− ____________ d. Al3+, PO43− ____________ 4. Write formulas for each compound: a. potassium nitrate ___________________________ b. barium chloride ___________________________ c. magnesium sulfate ___________________________ d. lithium oxide ___________________________ e. ammonium carbonate ___________________________ 13 Worksheet 5.4 – Lewis Structures of Atoms NAME: ______________________ Draw the appropriate number of dots around the following atoms to create correct Lewis Dot Structures: H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te I Xe Cs Ba Tl Pb Bi Po At Rn Write lewis structures for the following: CH4 CCl4 H H C H CF4 PCl3 OF2 Cl H Cl C Cl CI4 Cl Cl P Cl F O Cl PBr3 SF2 F 14 Lewis Structures of Ions NAME: ______________________ Draw the appropriate number of dots around the following ions to create correct Lewis Dot Structures: [H]+ Heo [Li]+ [Be]++ [B]3+ [C]4+ [N]3- [O]2- [F]1- Neo [Na]+ [Mg]++ [Al]3+ [Si]4+ [P]3- [S]2- [Cl]1- Aro [K]+ [Ca]++ [Ga]3+ [Ge]4+ [As]5+ [Se]4+ [Br]1- Kro [Rb]+ [Sr]++ [In]3+ [Sn]4+ [Sb]3+ [Te]4+ [I]1- Xeo [Cs]+ [Ba]++ [Tl]3+ [Pb]2+ [Bi]4+ [Po]4+ [At]4+ Rno Write the Dot Diagrams for the following: LiCl Li3N Ca3P2 Use the example on page 439 to show what happens to the electrons in these bonds. Show the electrons of the cations in one color and the electrons of the anion a different color. Li __ __ Li __ __ Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 1s 2s 1s 2s 2p 3s 3p 4s Cl __ ___ __ __ __ Li __ __ Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 1s 2s 1s 2s 2p 3s 3p 4s Li __ __ Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 1s 2s 2p 3s 3p 4s P __ __ __ __ __ __ __ __ __ N __ ___ __ __ __ 1s 2s 2p 3s 3p 1s 2s 2p P __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 15 Group 1: How do I Identify general categories of elements? What are they? Group 2: How are electron configurations related to the position of an element on the periodic table? Group 3: How is the atomic size related to the position of an element on the periodic table? Group 4: How is the ionic size related to the position of an element on the periodic table? How is it related to the electron configuration. Group 5: How is ionization energy related to the position of an element on the periodic table, and its electron configuration? Group 6: How is electronegativity related to the position in the periodic table? Group 7: Make a powerpoint with animation to summarize all of the previous presentations. 16 Cations Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + Potassium K+1 + Potassium K+1 + Potassium K+1 + Potassium K+1 + Potassium K+1 + Potassium K+1 + Potassium K+1 + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Sodium Na+1 + Silver Ag+1 + + Calcium Ca+2 + + Magnesium Mg+2 + + Iron (II) Fe+2 + + Aluminum Al+3 + Aluminum Al+3 + Aluminum Al+3 + Aluminum Al+3 + Aluminum Al+3 + Aluminum Al+3 + Aluminum Al+3 + + + + + + + + Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + NH4+1 + Potassium Potassium Potassium Potassium Potassium Potassium Potassium +1 +1 +1 +1 +1 +1 K + K + K + K + K + K + K+1 + Cut out the cations, store these in a plastic bag. 17 Anions Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 Nitrate NO3-1 Acetate C2H3O2-1 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Dichromate Cr2O7-2 - Sulfate SO4-2 Sulfate SO4-2 Sulfate SO4-2 Sulfate SO4-2 Sulfate SO4-2 Sulfate SO4-2 Sulfate SO4-2 - - - - - - - - - - - - - - Carbonate CO3-2 - Carbonate CO3-2 - Carbonate CO3-2 - Carbonate CO3-2 - Carbonate CO3-2 - Carbonate CO3-2 - Carbonate CO3-2 - Phosphate PO4-3 Phosphate PO4-3 Phosphate PO4-3 Phosphate PO4-3 Phosphate PO4-3 Phosphate PO4-3 Phosphate PO4-3 Bromide Br-1 Chloride Cl-1 - Bromide Br-1 Chloride Cl-1 - Bromide Br-1 Chloride Cl-1 - Bromide Br-1 Chloride Cl-1 - Bromide Br-1 Chloride Cl-1 - Cut out the anions, store in a plastic bag. Bromide Br-1 Chloride Cl-1 - Bromide Br-1 Chloride Cl-1 - 18 Laboratory Activity 5C – Making Models of Compounds Chemical Formulas Match the appropriate cation to the appropriate anion so that the number of + charges = the number of anions. Chloride Bromide Acetate Sulfate Carbonate Nitrate Cl-1 Br-1 C2H3O2-1 SO4-2 CO3-2 NO3-1 Ammonium ex: NH4+1 NH4+ Cl↓ NH4Cl Potassium K+1 Sodium Na+1 Silver Ag+1 Calcium Ca+2 Magnesium Mg+2 Iron (II) Fe+2 Aluminum Al+3 Name ____________________ Phosphate PO4-3 Dichromate Cr2O7-2 19 Laboratory Activity 5C – Making Models of Compounds Ammonium NH4+1 Chloride Cl-1 example: ammonium chloride Potassium K+1 Sodium Na+1 Silver Ag+1 Calcium Ca+2 Magnesium Mg+2 Iron (II) Fe+2 Aluminum Al+3 Bromide Br-1 Acetate C2H3O2-1 Names Sulfate SO4-2 Carbonate CO3-2 Name ____________________ Nitrate NO3-1 Phosphate PO4-3 Dichromate Cr2O7-2 20 Electron Dot Structures and VSEPR Shapes NAME: ____________________ 1. Classify the following compounds as ionic or covalent. Underneath the covalent compounds, draw a plausible electron dot structure for the molecule. a. MgCl2 b. OCl2 c. Na2S d. SF2 2. Describe the difference between an ionic and a covalent bond. 3. Draw the electron dot structure of each of the following molecules. a. F2 b. HCl c. HCCH d. NI3 4. Draw resonance structures for the nitrite ion, NO2−1. The oxygen atoms are attached to the nitrogen. 5. Explain the difference between “paramagnetic” and “diamagnetic.” 6. Predict whether the following species are diamagnetic or paramagnetic. a. BF3 b. O2−1 c. NO2 d. F2 21 More Work With VSEPR, Sigma and Pi Bonds NAME: ____________________ 1. Use VSEPR theory to predict the shapes of the following species. Draw the electron dot structures at the right. b. CO2 c. SiCl4 d. SO3 e. SCl2 f. CO g. I3−1 2. Which of the following species would you predict to be diamagnetic? Paramagnetic? a. NO3−1 b. OH−1 c. H2O d. SO3 3. Suggest a reason why phosphorus and sulfur have more than an octet in many of their compounds. Explain why compounds of nitrogen and oxygen never do. 22 Laboratory Activity 5D – Candy Dot Molecule Building NAME: ______________________ Students will be required to use marshmallows (atoms) and toothpicks (bonds) to create molecule models. The table provided below should be expanded to include all of the molecules used in this activity. Students will complete a table for each of the following: CCl4 PH4+1 AsCl3 CO2 PCl5 SiO2 GeCl4 SCl2 3-Dimensional Drawing BeI2 BF3 VSEPR Shape Bonds Lewis Dot Structure H2S SCl6 SeCl6 Bonds Molecular Formula NF3 SO3 BBr3 23 24 Polarity of Molecules NAME: _____________________ 1. Not every molecule with polar bonds is polar. Explain this statement, using CCl4 as an example. 2. Draw the electron dot structure for each molecule. Identify polar covalent bonds by assigning slightly positive (δ+) and slightly negative (δ−) symbols to the appropriate atoms. a. HOOH c. HBr b. BrCl d. H2O 3. Based on the information you have amassed about molecular shapes, which of these molecules would you expect to be polar? a. SO2 c. CO2 b. H2S d. BF3 4. Use VSEPR theory to predict the geometry of the following. Draw electron dot structures of each compound at the right. a. SiCl4 b. CO32− c. CCl4 d. SCl2 25 Intermolecular Forces NAME: ___________________ 1. How does a network solid differ from most other covalent compounds? 2. Which of the following are characteristic of most covalent compounds? a. b. c. d. e. high melting point shared bonding electrons low water solubility existence as molecules composed of a metal and a nonmetal 3. Depict the hydrogen bonding between two ammonia molecules and between one ammonia molecule and one water molecule. 4. Which compound in each pair exhibits the stronger intermolecular hydrogen bonding? a. H2S, H2O c. HBr, HCl b. HCl, HF d. NH3, H2O 5. What is a hydrogen bond? 6. Why do compounds with strong intermolecular attractive forces have higher boiling points than compounds with weak intermolecular attractive forces? 26 Practice Test – Page One Naming, Formula Writing, Trends NAME ______________________ 1. Which is larger, an atom of Ni or an atom of Ag? 2. Which is smaller, an atom of Si or an atom of Ga? 3. Which has a higher electronegativity, Ti or Cr? 4. Which has more shielding, K or Zr? _____ _____ _____ _____ 5. Which is larger, an atom of Magnesium or a Magnesium ion? 6. Which is smaller, an ion of Sulfur or an ion of Sodium? _______________ 7. Which has greater nuclear charge, Sodium or Magnesium? 8. Which has a lower ionization energy, Ba or Rb? 9. Which holds electrons more tightly, N or B? _______________ _______________ _____ _____ 10. Which neutral atom has higher electronegativity, neutral atom “A,” which has an electron configuration of 1s22s22p63s23p64s23d104p3 or atom “X,” which has an electron configuration of 1s22s22p63s23p64s1? _____ 11. Define electronegativity. 12. Compound: Fe2(CO3)3 Type of bond (Ionic or Covalent): ________________ Type of structure (Molecule or Lattice): ________________ Name of compound: ________________ 13. Compound: N2O7 Type of bond (Ionic or Covalent): ________________ Type of structure (Lattice or Molecule): ________________ Name of compound: ________________ 27 Practice Test – Page Two Naming, Formula Writing, Trends NAME ______________________ 14. Compound: Manganese (II) Sulfate Type of bond (Covalent or Ionic): ________________ Type of structure (Molecule or Lattice): ________________ Formula of compound: ________________ 15. Compound: Oxygen Monochloride Type of bond (Ionic or Covalent): ________________ Type of structure (Molecule or Lattice): ________________ Formula of compound: ________________ 16. Compound: Al2S3 Type of bond (Ionic or Covalent): ________________ Action of electrons (Sharing or Transfer): ________________ Name of compound: ________________ 17. Compound: Tricarbon Heptachloride Type of bond (Ionic or Covalent): ________________ Action of electrons (Sharing or Transfer): ________________ Formula of compound: ________________ 18. Define shielding. 19. At the right, put the following atoms or ions in order, from smallest to largest: K, K+1, Ca, Mg+2, P, P-3, Cl, O ___ ___ ___ ___ ___ ___ ___ ___ 28 PRACTICE TEST: Page Three NAME: ____________________ I: DRAW the Lewis Dot Diagrams for the following. IDENTIFY the shape of the molecule just under each drawing. (Phosphorus CAN but DOESN’T HAVE TO always violate the octet rule). SCl2 OF2 AsCl3 CF4 SiO2 PI5 II: Which of the above compounds are polar molecules? III: HCN, has a few bonds. Draw the structure of HCN and identify the number of Sigma and Pi bonds it has. Sigma: Pi: IV: Draw the Lewis Dot structure for each of the resonance conformers of ozone, O3. 29 PRACTICE TEST Page Four NAME: ____________________ V: Draw Lewis structures for the following atoms and ions under their names. Sulfur Sulfide ion Sodium Sodium ion VI: Draw a Lewis structure of the compound formed when Sodium ionically bonds with Sulfur. VII: Explain the differences between iron and steel. VIII: What will form stronger bonds, NH3 or NF3? IX: What is stronger, Ionic Bond or Hydrogen Bond? X: What is stronger, Dipole-Dipole Interaction or Covalent Bonds? XI: What type of intermolecular bonds will hold molecules of PCl3 together? XII: What type of intermolecular bonds will hold molecules of CO2 together? XIII: Which covalent bond is most polar? H-H, H-N, H-F, or H-O? 30