Chapter 12 Chemical Bonding

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1 Chapter 12 Chemical Bonding

2 Chapter 12 Review p Key Terms bond (12.1) bond energy (12.1) ionic bonding (12.1) ionic compound (12.1) covalent bonding (12.1) polar covalent bond (12.1) electronegativity (12.2) dipole moment (12.3) Lewis structure (12.6) duet rule (12.6) octet rule (12.6) bonding pair (12.6) lone (unshared) pair (12.6) single bond (12.7) double bond (12.7) triple bond (12.7) resonance (12.7) resonance structure (12.7) molecular (geometric) structure (12.8) bond angle (12.8) linear structure (12.8) trigonal planar structure (12.8) Tetrahedral structure (12.8) valence shell electron pair repulsion (VSEPR) model (12.9) tetrahedral arrangement (12.9) trigonal pyramid (12.9)

3 Summary (p.374) 1. Chemical bonds hold groups of atoms together. They can be classified into several types. An ionic bond is formed when a transfer of electrons occurs to form ions; in a purely covalent bond, electrons are shared equally between identical atoms. Between these extremes lies the polar covalent bond, in which electrons are shared unequally between atoms with different electronegativities. 2. Electronegativity is defined as the relative ability of an atom in a molecule to attract the electrons shared in a bond. The difference in electronegativity values between the atoms involved in a bond determines the polarity of that bond.

4 Ch.12 Chemical Bonding Summary (cont d) 3. In stable chemical compounds, the atoms tend to achieve a noble gas electron configuration. In the formation of a binary ionic compound involving representative elements, the valence-electron configuration of the nonmetal is completed: it achieves the configuration of the next noble gas. The valence orbitals of the metal are emptied to give the electron configuration of the previous noble gas. Two nonmetals share the valence electrons so that both atoms have completed valence-electron configurations (noble gas configurations). 4. Lewis structures are drawn to represent the arrangement of the valence electrons in a molecule. The rules for drawing Lewis structures are based on the observation that nonmetal atoms tend to achieve noble gas electron configurations by sharing electrons. This leads to a duet rule for hydrogen and to an octet rule for many other atoms.

5 Summary (cont d) 5. Some molecules have more than one valid Lewis structure, a property called resonance. Although Lewis structures in which the atoms have noble gas electron configurations correctly describe most molecules, there are some notable exceptions, including O2, NO, NO2, and the molecules that contain Be and B. 6. The molecular structure of a molecule describes how the atoms are arranged in space. 7. The molecular structure of a molecule can be predicted by using the valence shell electron pair repulsion (VSEPR) model. This model bases its prediction on minimizing the repulsions among the electron pairs around an atom, which means arranging the electron pairs as far apart as possible.

6 Active Learning Questions 1. Using only the periodic table, predict the most stable ion for Na, Mg, Al, S, Cl, K, Ca, and Ga. Arrange these elements from largest to smallest radius and explain why the radius varies as it does. 2. Write the proper charges so that an alkali metal, a noble gas, and a halogen have the same electron configurations. What is the number of protons in each? The number of electrons in each? Arrange them from smallest to largest radii and explain your ordering rationale. 3. What is meant by a chemical bond? 4. Why do atoms form bonds with one another? What can make a molecule favored compared with the lone atoms? 5. How does a bond between Na and Cl differ from a bond between C and O? What about a bond between N and N?

7 Active Learning Questions (cont d) 6. In your own words, what is meant by the term electronegativity? What are the trends across and down the periodic table for electronegativity? Explain them, and describe how they are consistent with trends of ionization energy and atomic radii. 7. Why are some bonds ionic and some covalent? 8. True or false? In general, a larger atom has a smaller electronegativity. Explain. 9. Why is there an octet rule (and what does octet mean) in writing Lewis structures? 10. Does a Lewis structure tell which electrons came from which atoms? Explain.

8 12.1 Types of Chemical Bonds - QUESTIONS 1. A chemical represents the force that holds together groups of two or more atoms and allows them to function as a unit. 2. The represents the quantity of energy required to break a chemical bond. 3. A(n) compound results when a metallic element reacts with a nonmetallic element. 4. When electrons in a molecule are shared between atoms, either evenly or unevenly, a(n) bond is said to exist.

9 12.1 Types of Chemical Bonds QUESTIONS: 5. Describe the type of bonding that exists in the Cl2(g) molecule. How does this type of bonding differ from that found in the HCl(g) molecule? How is it similar? 6. Compare and contrast the bonding found in the H2(g) and HF(g) molecules with that found in NaF(s) Electronegativity QUESTIONS: 7. The relative ability of an atom in a molecule to attract electrons to itself is called the atom's

10 Most chemical bonds consist of electrostatic attractive forces and are called bonds, or of shared electrons and are called bonds. 1. electric; shared 2. ionic; covalent 3. ionic; molecular 4. electronic; coordinate

11 When electrons in a covalent bond are shared equally the bond is, but if the electrons are not shared equally the bond is, which means that it has a positive side and a negative side.

12 When considering a bond between two atoms, the greater the difference in, the more is the bond. 1. polarity; divided 2. atomic weight; nonpolar 3. electronegativity; polar 4. electronegativity; nonpolar

13 Using Fig in the text, we see that the difference in electronegativity between hydrogen and iodine ( ) = 1.5, which means that hydroiodic acid has a(n). 1. low boiling point 2. low acidity 3. unstable gas phase 4. dipole moment

14 If we consider the electron configuration of strontium, [Kr]5s 2, and that of oxygen, [He]2s 2 2p 4, both atoms will attain stable noble gas electron configurations by the transfer of electron(s). This will give Sr a charge of and O a charge of. Hence the ionic compound formed has the formula and is named strontium oxide ; 1+ ; 1- ; SrO 2. 2 ; 2+ ; 2- ; SrO 3. 2 ; 2+ ; 1- ; SrO 4. 2 ; 2- ; 2+ ; SrO

15 The structures of ionic compounds are usually described as the packing of with smaller fitting into the interstices. 1. anions; cations 2. anions; electrons 3. cations; anions 4. cations; electrons

16 ANSWER Choice #1 correctly describes the packing of anions (which tend to be larger) into specific patterns known as crystal lattices, while cations fit into the spaces or interstices between the packed anions. Section 12.5: Ionic Bonding and Structures of Ionic Compounds

17 Lewis structures show the arrangement of electrons in an atom or ion. 1. all 2. core 3. valence 4. missing

18 In the Lewis structure for H 2 S there are a total of electrons and pair(s) of nonbonding electrons ; ; ; ; 1

19 The Lewis structure for SO 2 contains total electrons and nonbonding pairs of electrons, as well as one bond and one bond between the central sulfur and the oxygen atoms ; 4 ; single ; double ; 6 ; single ; double ; 4 ; single ; triple ; 6 ; single ; double

20 In the Lewis structure of the SO 2 molecule, the central S is connected to one O by a single bond and to the other O by a double bond. Does this mean that the two bonds in this molecule are different from each other? Explain!

21 It is important to fully understand that Lewis structures are useful in determining the bonding relationships between atoms in a molecule, but that they do not directly provide a true picture of molecular shape. While the Lewis structure for methane, CH 4, an important greenhouse gas, suggests a flat structure with 4 hydrogens arranged around a central carbon, the methane molecule is actually. 1. square planar 2. trigonal 3. tetrahedral 4. octahedral

22 While the electron pair geometry of NH 3 is, VSEPR predicts the molecular shape as, due to the pair of nonbonding electrons on the central N. 1. tetrahedral; trigonal pyramidal 2. trigonal planar; tetrahedral 3. tetrahedral; trigonal planar 4. Both are trigonal planar.

23 In determining the shape of the SO 2 molecule we examine the Lewis structure and find the central S atom attached to O s via one single and one double bond. The electron pair geometry is and the molecular geometry is, with bond angles of degrees. 1. trigonal planar ; linear ; trigonal planar ; bent ; tetrahedral ; linear ; tetrahedral ; trigonal planar ; 120

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