Honors Chemistry: Chapter 7 - Bonding
OBJECTIVES
- Describe the nature of the chemical bond and its relationship to valence electrons.
- Use electronegativities to predict the type of bond that will form between given atoms.
- Define ionic bond, covalent bond and molecule.
- Differentiate between the properties of ionic and covalent bonds.
- Define the ionic radius, and use ionic radius to calculate the internuclear distance between ions in a crystal.
- Distinguish between non polar, polar and covalent bond.
- Use electronegativities to compare the polarities of bonds.
- Describe the coordinate covalent bond.
- Define polyatomic ions, and describe their bonding characteristics.
- Explain metallic properties using the metallic bond concept.
- Use Lewis dot structures to show molecular structures.
- Illustrate, using examples, the concept of resonance.
- 13. Be able to write Lewis dot structures for cases that do not obey the octet rule.
- Distinguish between shared and unshared electron pairs.
- 15. Explain how the shared and unshared pairs of electrons determine molecular structures.
- Predict the shapes and bond angles of simple molecules based on the VSEPR theory.
- Describe hybrid orbital.
- Use hybridization theory to explain the bond angles in compound with multiple bonds.
- Use hybridization theory to predict the shapes of molecules.
- Differentiate between sigma and pi bonding.
- Use the bonding theory to explain unsaturated bonds.
READING:
Textbook - Chapter 7
Section 7.1: Lewis Structures and the octet Rule, pages 180-187 (exclude Formal Charges), 188-191
Section 7.2: Molecular Geometry, pages 192-200
Section 7.3: Polarity of molecules, pages 200-201
Section 7.4: Atomic Orbitals; Hybridization, pages 203-209
WRITTEN ASSIGNMENTS:
Textbook, page 211:
Lewis Structures: Exercises 2- 20 even.
Resonance Structures: Exercises 22 - 26 even.
Molecular Geometry: Exercises 30 – 40 even.
Molecular Polarity: Exercises 42-46 even.
Hybridization: Exercises 48-60 even
Sigma and Pi bonds: Exercises 62, 64.
Unclassified: 68, 7.
Dittos in the packet as assigned.
Introduction to Bonding: General Concepts (Chemical Bond)
DIRECTIONS: Write the letter preceding the word or expression that best completes the statement.
1. In a chemical bond, the link between atoms results from the attraction between electrons and (a) Lewis structures; (b) nuclei; (c) van der Waals forces; (d) isotopes
2. A covalent bond consists of )a) shared electron; (b) a shared electron pair; (c) two electrovalent electrons; (d) an octet of electrons
3. If two covalently bonded atoms are identical, the bond is identified as (a) nonpolar covalent; (b) polar covalent; (c) nonionic; (d) coordinate covalent
4. A covalent bond in which there is unequal attraction for the shared electrons is (a) nonpolar; (b) polar; (c) ionic; (d) dipolar
5. Atoms with a strong attraction for electrons they share with another atom exhibit (a) zero electronegativity; (b) low electronegativity; (c) high electronegativity; (d) Lewis electronegativity.
6. Bonds with between 5% and 50% ionic character are considered to be (a) ionic; (b) pure covalent; (c) polar covalent; (d) nonpolar covalent.
7. A nonpolar covalent bond is likely to exist between (a) a metal and a nonmetal; (b) two ions; (c) two identical atoms; (d) an atom and an ion.
8. The greater the electronegativity difference between two bonded atoms, the greater the percentage of (a) ionic character; (b) metallic character; (c) covalent character; (d) electron sharing.
9. In which of these compounds is the bond between the atoms NOT a nonpolar bond? (a) Cl2; (b) H2; (c) HCl; (d) O2.
DIRECTIONS: Complete the following statements, forming accurate sentences.
10. The electrons involved in the formation of a chemical bond are called ______
11. A chemical bond resulting from electrostatic attraction between positive and negative ions is called a(n)
______
12. If the electrons involved in bonding spend most of the time close to one atom rather than the other, the bond is
______
DIRECTIONS: Consult your Periodic Table, specifically examine the graph of % ionic character, and answer the following questions.
13. The percentage of ionic character and the type of bond for the Li-Cl bond in LiCl (electronegativity for Li = 1.0; electronegativity for Cl = 3.0) is ______
14. The percentage of ionic character and the type of bond for the Br-Br bond in Br2 (electronegativity for Br = 2.8) is ______
Bonding – General Concepts (Ionic Bond)
DIRECTIONS: Write the letter preceding the word or expression that best completes the statement.
1. In the formula unit of sodium chloride, NaCl stands for one (a) formula unit; (b) molecule; (c) crystal; (d) atom
2. The chemical formula for an ionic compound represents the (a) number of atoms in each molecules; (b) the number of ions in each molecule; (c) simplest ratio of the combined ions that gives neutrality; (d) total number of ions in crystal structure.
3. A formula that shows the types and numbers of atoms combined in a single molecule is called (a) molecular formula; (b) ionic formula; (c) Lewis structure; (d) covalent formula.
4. In a crystal of an ionic compound, each cation is surrounded by a number of (a) molecules; (b) positive ions; (c) dipoles; (d) anions.
5. In a crystal, the valence electrons of adjacent ions (a) repel each other; (b) attract each other; (c) neutralize each other; (d) have no effect on each other.
6. Compared the neutral atoms involved in its formation, the crystal lattice that results is (a) higher in potential energy; (b) lower in potential energy; (c) equal in potential energy; (d) unstable.
7. The lattice energy of a compound A is greater that that of a compound B. What can be deducted from this fact? (a) Compound A is not an ionic compound. (b) It will be more difficult to break the bonds in compound A than in compound B. (c) Compound B is probably a gas. (d) Compound A has larger crystals than compound B.
8. Which of the following is NOT a property of an ionic compound? (a) vaporizes readily at room temperature; (b) brittle; (c) hard; (d) electrical conductor in molten state.
9. Compared to ionic compounds, molecular compounds (a) have higher boiling points; (b) are brittle; (c) have lower melting points; (d) are harder.
10. The forces of attraction between molecules in a molecular compound are (a) stringer that the forces of ionic bonding; (b) weaker that the forces of ionic bonding; (c) approximately equal to the forces of ionic bonding; (d) zero.
11. At room temperature, most ionic compounds will be (a) solids; (b) liquids; (c) gases; (d) molten.
DIRECTIONS: Complete each of the following statement with the best meaning word or expression.
12. A(n) ______is a shorthand representation of the composition of a substance using atomic symbols and numerical subscripts.
13. In the NaCl crystal, the packing of Na+ ions and Cl- ions is such that each ion has clustered around it ______of the oppositely charged ions.
14. In an ionic compound, the orderly arrangement of ions in a crystal is the state of ______energy.
15. ______energy is the type of energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
Bonding: Ionic Bond
1. Which of the following elements form positive ions when they react?
a. Li in Group 1 b. Ni, a transition element c. S in Group 16 d. C in Group 14
2. Explain, in terms of metallic character, why aluminum form a +3 ion much more readily than boron in the same group.
3. Explain why nitrogen forms a -3 ion, N3- , while bismuth, Bi, in the same group, forms a +3 ion, Bi3+ .
4. Why, in ionic crystal, are "+" ions surrounded by "-" ions rather than ions of the same charge?
5. How do you explain the fact that solid ionic compounds do not conduct an electrical current, yet they become conductors when melted?
6. In general, which is larger:
a. a positive ion or the corresponding atom? ______
b. a negative ion or the corresponding atom? ______
c. a metal atom or a nonmetal atom in the same period of the Periodic table? ______
d. a metal ion or a nonmetal ion in the same period? ______
7. Give the electron configuration of:
a. Na1+ b. Ca2+ c. Al3+ d. Cl1- e. O2-
8. What charges would you expect for ions of the following metals?
a. K b. Sc c. Li d. Al e. Sr
9. What charges would you expect for the ion of the following nonmetals?
a. S b. F c. N d. I e. O
10. Which of the following ions have a noble-gas configurations?
a. Cs1+ b. Be2+ c. In3+ d. Mn2+ e. Ca2+
11. Name the following ions:
a. NH41+ b. CO32- c. NO31- d. ClO31-
12. Write the formula, include the charge, for the following polyatomic ions:
a. sulfate b. phosphate c, hydrogen carbonate d. chromate
13. Give the simplest formula of a compound containing:
a. A+ and X- b. C3+ and X 1- c. B2+ and Y2- d. A1+ and Y2-
14. Using the electronegativity table, arrange the following compounds in order of decreasing ionic character:
LiI, BaO, AlCl3, CsF, RbBr, K2S, CaO, ClF, P2S3, F2, NaI, Mg3N2 .
15. What tests might you perform to determine whether a substance is ionic?
16. Write a balanced equations for the reaction of oxygen with
a. Li b. Zn c. Sc d. Bi(III)
17. Give all the possible the charges of ions formed by the following transition elements:
Ag b. Zn c. Cu d. Fe
18. A certain ionic compound contains 60.7% O, 17.7% N, 15.2% C, 6.37% H. Find the simplest formula of the compound and identify the ions present.
Bond Classification
Classify the bonds between the following pairs of atoms as principally ionic or covalent.
1. Al -Si 4. C - H 7. Ca - Cl
2. Ba - O 5 . Li - S 8. F - S
3. Ca - P 6. B - N 9. Rb - Br
For each of the atom pair listed below, decide whether an ionic or a covalent bond would form between the elements.
10. hydrogen - fluorine 12. cobalt - fluorine 14. iron - fluorine
11. astatine - beryllium 13. bromine - cerium 15. calcium - fluorine
Using any available information, estimate the percent of ionic character in the following bonds.
16. Pb - S 18. Na - Br 20. Cu - I
17. Ag - Cl 19. C - N 21. H - O
Covalent Bond: Drawing Lewis Structures
Draw Lewis structures for the following molecules. The central atom is underlined. NOTE the number of
shared and unshared pairs of electrons.
1. PCl3 23. ClO31-
2. SO42- 24. GeF4
3. PO43- 25. HBr
4. CH2Cl2 26. PF3
5. N2O 27. NI3
6. CO2 28. SiH4
7. HCN 29. CH3Cl
8. CO 30. Br2
9. H2S 31. CHBr3
10. H2Te 32. CBr4
11. H2CO3 33. SO3
12. H2S
13. HBr
14. NO31-
15. PO2F21-
16. CO
17. CO2
18. SCl2
19. N31-
20. ClF3
21. CBr4
22. XeF2
Covalent Compounds: Geometry (1)
Determine the number of shared and unshared electron-pairs around the central atom for the following molecules.
1. H2Te 2. PF3 3. BeF2 4. CBr4
Determine the molecular geometry using the VSEPR theory.
5. AlCl3 13. PCl3
6. CBr4 14. H2O
7. BF3 15. H - Be - H
8. CO2 16. OF2
F F
9. HCN 17. Xe
F F
10. H2CO 18. ClO31-
11. SO3 19. NO31-
12. SF2 20. N = N = 0
Predict whether the bond angle of each of the following is greater than, less than, or equal to 109.5 o.
21. F - N - F in NF3 23. F - Te - F in TeF2 25. F- As - F in AsF5
22. F - Be - F in BeF2 24. F - O - F in OF2 26. O - Xe - O in XeO4
List all the bond angles in the following compounds:
H H
27. CH4 28. H - C = C - H 29. H - C = C - H
Covalent Bond: Polar Bonds
1. Arrange the following elements in order of increasing force of attraction for electrons in the bond.
Sb, F, In, O, N, Se, Cl
2. Which of the following bonds is most polar? In each bond, indicate the atom that carries the partial negative charge.
a. H - I b. P - I c. As - Br d. N - S
3. Predict whether the following molecules are polar or nonpolar.
a. SO2 b. AsCl3 c. H2Se d. SO3
e. H2S f. CO2 g. CO i. I2
H H
\ /
j. PH3 k. SiCl4 l. SCl2 m. C = C
/ \
Cl Cl
H H
\ /
n. Cl - C - C – Cl o. H2O
/ \
H H
OH H H H Cl
/ | | \ /
q. H - O - S p. H - C - C = O r. C = C
\\ / \
O Cl H
Covalent Bond: Hybridization
1. Mixing of atomic orbitals in an atom (usually the central atom) to generate a set of new atomic orbitals is called ______.
2. The ______orbitals will result in four equal-energy bonds.
3. When two equal-energy orbitals are formed as a result of mixing ______and _____ orbitals you obtain sp hybridization.
4. sp3d2 hybrid orbitals will form only in elements of which periods?
5. The most common hybridization for an atom in group 13 is the _____ hybridization; and in group 14 is the ______hybridization.
6. What is the usual hybridization in group 2? ______