Unit 5 Bonding Chemistry 1
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1 Name: Date: Types of Chemical Bonds 1. Define a chemical bond. Unit 5 Bonding Chemistry 1 2. Why are the valence electrons of an atom the only electrons likely to be involved in bonding to other atoms? 3. For each of the following atomic numbers, use the periodic table to write the formula (including the charge) for the simple ion that the element is most likely to form. a. 53 b. 38 c. 55 d. 88 e. 9 Compounds That Contain Ions 4. List some properties of a substance that would lead you to believe it consists of ions. How do these properties differ from those of non-ionic compounds? 5. Why does an ionic compound conduct an electric current when the compound is dissolved in water? 6. For the following pairs of ions, use the concept that a chemical compound must have a net charge of zero to predict the formula of the simplest compound that the ions are most likely to form. a. Fe 3+ and P 3- b. Fe 3+ and S 2- c. Fe 3+ and Cl - d. Mg 2+ and Cl - e. Mg 2+ and O 2- Ionic Bonding and Structures of Ionic Compounds 7. What types of elements react to form ionic compounds? 8. Give an example of the formation of an ionic compound from its elements.
2 9. Is the formula we write for an ionic compound the same as the formula we write for a covalent compound like water (H 2 O)? Why or why not? 10. Describe in general terms the structure of ionic solids such as NaCl. 11. Why are cations always smaller than the atoms from which they are formed? Lewis Structures 12. Explain what the "duet" and "octet" rules are and how they are used to describe the arrangement of electrons in a molecule. 13. Give the total number of valence electrons in each of the following molecules. a. CBr 4 c. C 6 H 6 b. NO 2 d. H 2 O Write a Lewis structure for each of the following simple molecules. a. NH 3 c. PCl 3 b. CI 4 d. H 2 O Write a Lewis structure for each of the following simple molecules or polyatomic ions. a. H 2 S d. Cl 2 O b. SiF 4 e. SO 3 2- c. C 3 H 8 f. PO 4 3-2
3 16. How many electrons are involved when two atoms in a molecule are connected by a "double bond"? Write the Lewis structure of a molecule containing a double bond. 17. Draw the Lewis Dot Structure for each of the following molecules or ions containing multiple bonds. a. SO 2 c. CO 3 2- b. O 3 d. HCN e. CO 2 g. N 2 f. C 2 H 4 Molecular Structure 18. Draw H 2 S. a. How many pairs of valence electrons are there around the sulfur atom in the molecule? How many are bonding pairs? How many are lone pairs? b. What is the molecular shape and the approximate H-S-H bond angle? 19. Draw PH 3. a. How many pairs of valence electrons are there around the phosphorus atom in the molecule? How many are bonding pairs? How many are lone pairs? b. What is the molecular shape and the approximate H-P-H bond angle? 3
4 20. Draw BF 3. a. How many pairs of valence electrons are there around the boron atom in the molecule? How many are bonding pairs? How many are lone pairs? b. What is the molecular shape and the approximate F-B-F bond angle? 21. Draw GeH 4. a. How many pairs of valence electrons are there around the germanium atom in the molecule? How many are bonding pairs? How many are lone pairs? b. What is the molecular shape and the approximate H-Ge-H bond angle? Molecular Structure: The VSEPR Model 22. What general principles determine the molecular structure (shape) of a molecule? 23. Although the valence electron pairs in ammonia (NH 3 ) have a tetrahedral arrangement, the overall geometric structure of the ammonia molecule is not described as being tetrahedral. Explain. 24. Although both the BF 3 and NF 3 molecules contain the same number of atoms, the BF 3 molecule is flat, whereas the NF 3 molecule is trigonal pyramidal. Explain. 4
5 25. Using the VSEPR theory, predict the molecular structure of each of the following molecules. a. CCl 4 b. HCl c. GeI Using the VSEPR theory, predict the molecular structure of each of the following polyatomic ions. a. dihydrogen phosphate ion, H 2 PO 4 (The hydrogens are bonded to oxygens, all oxygens are bonded to the phosphorus atom) b. sulfite ion, SO 3 2- Electronegativity 27. What does the property of electronegativity tell us about an atom? 28. What does it mean to say that a bond is polar? 29. What are the conditions that give rise to a bond being polar? 30. For each of the following sets of elements, identify the element expected to be most electronegative and that expected to be least electronegative. a. K, Sc, Ca b. Br, F, At c. C, O, N 31. On the basis of the electronegativity values on the back of your Periodic Tables, indicate whether each of the following bonds would be expected to be ionic, covalent, or polar covalent. a. S S c. S H b. S O d. S K 5
6 32. Which of the following molecules contain polar covalent bonds? a. phosphorus, P 4 c. ozone, O 3 b. oxygen, O 2 d. hydrogen fluoride, HF 33. On the basis of the electronegativity values indicate which is the more polar bond in each of the following pairs. a. H-O or H-N c. H-O or H-F b. H-N or H-F d. H-O or H-CI 34. On the basis of the electronegativity values, indicate which bond of the following pairs has a more ionic character. a. Na-O or Na-N c. Na-Cl or K-Cl b. K-S or K-P d. Na-Cl or Mg-Cl Bond Polarity and Dipole Moments 35. What is a dipole moment? 36. Draw water showing its dipole moment. 37. Why is the presence of a dipole moment in the water molecule so important? What are some properties of water that are determined by its polarity? 38. In each of the following diatomic molecules, which end of the molecule is positive relative to the other end? a. hydrogen fluoride, HF c. iodine monochloride, ICl b. chlorine monofluoride, ClF 39. For each of the following bonds, draw a figure indicating the direction of the bond dipole, including which end of the bond is positive and which is negative. a. P-F c. P-C b. P-O d. P-H 6
7 Unit5Hand-StudHand13.doc 7
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