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1 How to Use This Presentation
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2 Standardized Test Prep
Resources Chapter Presentation Bellringer Transparencies Sample Problems Visual Concepts Standardized Test Prep

3 Chapter 6 Table of Contents Section 1 Covalent Bonds
Covalent Compounds Table of Contents Section 1 Covalent Bonds Section 2 Drawing and Naming Molecules Section 3 Molecular Shapes

4 Chapter 6 Section 1 Covalent Bonds Bellringer Make a list of the elements that form ionic bonds. Note that most ionic bonds contain a metal and a nonmetal.

5 Chapter 6 Section 1 Covalent Bonds Objectives Explain the role and location of electrons in a covalent bond. Describe the change in energy and stability that takes place as a covalent bond forms. Distinguish between nonpolar and polar covalent bonds based on electronegativity differences.

6 Chapter 6 Objectives, continued
Section 1 Covalent Bonds Objectives, continued Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences. Ignore pgs.192,193

7 Chapter 6 Sharing Electrons
Section 1 Covalent Bonds Sharing Electrons When an ionic bond forms, electrons are rearranged and are transferred from one atom to another to form charged ions. In another kind of change involving electrons, the neutral atoms share electrons.

8 Sharing Electrons, continued
Chapter 6 Section 1 Covalent Bonds Sharing Electrons, continued Forming Molecular Orbitals A covalent bond is a bond formed when atoms share one or more pairs of electrons. The shared electrons move within a space called a molecular orbital. A molecular orbital is the region of high probability that is occupied by an individual electron as it travels with a wavelike motion in the three-dimensional space around one of two or more associated nuclei.

9 Chapter 6 Formation of a Covalent Bond

10 Chapter 6 Visual Concepts Chemical Bond

11 Chapter 6 Energy and Stability
Section 1 Covalent Bonds Energy and Stability Energy Is Released When Atoms Form a Covalent Bond

12 Energy and Stability, continued
Chapter 6 Section 1 Covalent Bonds Energy and Stability, continued Potential Energy Determines Bond Length When two bonded hydrogen atoms are at their lowest potential energy, the distance between them is 75 pm. The bond length is the distance between two bonded atoms at their minimum potential energy. However, the two nuclei in a covalent bond vibrate back and forth. The bond length is thus the average distance between the two nuclei.

13 Chapter 6 Visual Concepts Bond Length

14 Energy and Stability, continued
Chapter 6 Section 1 Covalent Bonds Energy and Stability, continued Bonded Atoms Vibrate, and Bonds Vary in Strength The bond length is the average distance between two nuclei in a covalent bond. At a bond length of 75 pm, the potential energy of H2 is –436 kJ/mol. Thus 436 kJ of energy must be supplied to break the bonds in 1 mol of H2 molecules. The energy required to break a bond between two atoms is the bond energy. Bonds that have the higher bond energies (stronger bonds) have the shorter bond lengths.

15 Chapter 6 Visual Concepts Bond Energy

16 Electronegativity and Covalent Bonding
Chapter 6 Section 1 Covalent Bonds Electronegativity and Covalent Bonding In covalent bonds between two different atoms, the atoms often have different attractions for shared electrons. Electronegativity values are a useful tool to predict what kind of bond will form.

17 Chapter 6 Visual Concepts Electronegativity

18 Electronegativity and Covalent Bonding, continued
Chapter 6 Section 1 Covalent Bonds Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally When the electronegativity values of two bonding atoms are similar, bonding electrons are shared equally. A covalent bond in which the bonding electrons in the molecular orbital are shared equally is a nonpolar covalent bond.

19 Electronegativity and Covalent Bonding, continued
Chapter 6 Section 1 Covalent Bonds Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally, continued When the electronegativity values of two bonding atoms are different, bonding electrons are shared unequally. A covalent bond in which the bonding electrons in the molecular orbital are shared unequally is a polar covalent bond.

20 Chapter 6 Predicting Bond Character from Electronegativity Differences

21 Electronegativity and Covalent Bonding, continued
Chapter 6 Section 1 Covalent Bonds Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends In a polar covalent bond, the ends of the bond have opposite partial charges. A molecule in which one end has a partial positive charge and the other end has a partial negative charge is called a dipole. In a polar covalent bond, the shared pair of electrons is not transferred completely. Instead, it is more likely to be found near the more electronegative atom.

22 Electronegativity and Covalent Bonding, continued
Chapter 6 Section 1 Covalent Bonds Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends, continued The symbol  is used to mean partial. + is used to show a partial positive charge – is used to show a partial negative charge charge example: H+F– Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

23 Comparing Polar and Nonpolar Covalent Bonds
Chapter 6 Visual Concepts Comparing Polar and Nonpolar Covalent Bonds

24 Polarity Is Related to Bond Strength
Chapter 6 Section 1 Covalent Bonds Polarity Is Related to Bond Strength In general, the greater the electronegativity difference, the greater the polarity and the stronger the bond.

25 Electronegativity and Bond Types
Chapter 6 Section 1 Covalent Bonds Electronegativity and Bond Types Differences in electronegativity values provide one model that can tell you which type of bond two atoms will form. Another general rule states: A covalent bond forms between two nonmetals. An ionic bond forms between a nonmetal and a metal.

26 Properties of Substances Depend on Bond Type
Chapter 6 Section 1 Covalent Bonds Properties of Substances Depend on Bond Type The type of bond that forms (metallic, ionic, or covalent) determines the properties of the substance. The difference in the strength of attraction between the basic units of ionic and covalent substances causes these types of substances to have different properties.

27 Chapter 6 Properties of Substances with Metallic, Ionic, and Covalent Bonds

28 1. Describe the attractive forces and repulsive
Chapter 6 Section 6.1 review, pg 198 1. Describe the attractive forces and repulsive forces that exist between two atoms as the atoms move closer together. The positive nucleus of each atom attracts the electrons of the other atom. At the same time, the nuclei repel each other, as do the electron clouds. 3. In what two ways can two atoms share electrons when forming a covalent bond? The two atoms may share electrons equally, forming a nonpolar covalent bond, or unequally, forming a polar covalent bond.

29 5. How are the partial charges shown in a polar covalent molecule?
Chapter 6 5. How are the partial charges shown in a polar covalent molecule? The symbol is written as a superscript on the element with the partial positive charge. The symbol is written as a superscript on the element with the partial negative charge. 6. What information can be obtained by knowing the electronegativity differences between two elements? Knowing the electronegativity difference suggests what type of bond will form between the atoms of the two elements.

30 Chapter 6 7. Why do molecular compounds have low melting points and low boiling points relative to ionic substances? The attractive forces between individual molecules are weak, accounting for the low melting point of molecular compounds.

31 Section 2 Drawing and Naming Molecules
Chapter 6 Bellringer Classify the following compounds according to the type of bonds they contain: NO CO HF NaCl HBr NaI

32 Section 2 Drawing and Naming Molecules
Chapter 6 Objectives Draw Lewis structures to show the arrangement of valence electrons among atoms in molecules and polyatomic ions. Explain the differences between single, double, and triple covalent bonds. Draw resonance structures for simple molecules and polyatomic ions, and recognize when they are required.

33 Chapter 6 Objectives, continued
Section 2 Drawing and Naming Molecules Chapter 6 Objectives, continued Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.

34 Lewis Electron-Dot Structures
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures Valence electrons are the electrons in the outermost energy level of an atom. A Lewis structure is a structural formula in which valence electrons are represented by dots. In Lewis structures, dot pairs or dashes between two atomic symbols represent pairs in covalent bonds.

35 Chapter 6 Visual Concepts Valence Electrons

36 Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons As you go from element to element across a period, you add a dot to each side of the element’s symbol.

37 Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued You do not begin to pair dots until all four sides of the element’s symbol have a dot.

38 Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued An element with an octet of valence electrons has a stable configuration. The tendency of bonded atoms to have octets of valence electrons is called the octet rule.

39 Chapter 6 Visual Concepts The Octet Rule

40 Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued When two chlorine atoms form a covalent bond, each atom contributes one electron to a shared pair. An unshared pair, or a lone pair, is a nonbonding pair of electrons in the valence shell of an atom.

41 Lewis Electron-Dot Structures, continued
Section 2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued A single bond is a covalent bond in which two atoms share one pair of electrons The electrons can pair in any order. However, any unpaired electrons are usually filled in to show how they will form a covalent bond.

42 Drawing Lewis Structures with Single Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Draw a Lewis structure for CH3I.

43 Drawing Lewis Structures with Single Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Solution Draw each atom’s Lewis structure, and count the total number of valence electrons. number of dots: 14 Arrange the Lewis structure so that carbon is the central atom.

44 Drawing Lewis Structures with Single Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Solution, continued Distribute one bonding pair of electrons between each of the bonded atoms. Then, distribute the remaining electrons, in pairs, around the remaining atoms to form an octet for each atom. Change each pair of dots that represents a shared pair of electrons to a long dash.

45 Chapter 6.2 Sample Problem A,practice pg.202 1)Draw the Lewis structures for H2S, CH2Cl2, NH3, and C2H6.

46 Chapter 6.2 Sample Problem A,practice pg.202 2)Draw the Lewis structures for methanol, CH3OH.

47 Drawing Lewis Structures for Polyatomic Ions
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Draw a Lewis structure for the sulfate ion,

48 Drawing Lewis Structures for Polyatomic Ions
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Solution Count electrons for all atoms. Add two additional electrons to account for the 2− charge on the ion. number of dots: = 32 Distribute the 32 dots so that there are 8 dots around each atom.

49 Drawing Lewis Structures for Polyatomic Ions
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Solution, continued Change each bonding pair to a long dash. Place brackets around the ion and a 2 charge outside the bracket to show that the charge is spread out over the entire ion.

50 Chapter 6.2 Sample Problem B, practice pg.203 1)Draw a Lewis structure for ClO3- 2) Draw the Lewis structure for the hydronium ion, H3O +

51 Chapter 6 Multiple Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Multiple Bonds For O2 to make an octet, each atom needs two more electrons. The two atoms share four electrons. A double bond is a covalent bond in which two atoms share two pairs of electrons.

52 Multiple Bonds, continued
Section 2 Drawing and Naming Molecules Chapter 6 Multiple Bonds, continued For N2 to make an octet, each atom needs three more electrons. The two atoms share six electrons. A triple bond is a covalent bond in which two atoms share three pairs of electrons.

53 Comparing Single, Double, and Triple Bonds
Chapter 6 Visual Concepts Comparing Single, Double, and Triple Bonds

54 Drawing Lewis Structures with Multiple Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Multiple Bonds Sample Problem C Draw a Lewis structure for formaldehyde, CH2O.

55 Drawing Lewis Structures with Multiple Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Multiple Bonds Sample Problem C Solution Draw each atom’s Lewis structure, and count the total dots. number of dots: 12 Arrange the atoms so that carbon is the central atom. O H C H

56 Drawing Lewis Structures with Multiple Bonds
Section 2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Multiple Bonds Sample Problem C Solution, continued Distribute one pair of dots between each of the atoms and the rest, in pairs, around the atoms. C does not have an octet. To get an octet, move an unshared pair from the O to between the O and the C. Change each bonding pair to a long dash. Two pairs of dots represent a double bond.

57 Chapter 6.2 Sample Problem C, practice pg. 205 Draw the Lewis structures for carbon dioxide, CO2, and carbon monoxide, CO. 2) Draw the Lewis structures for ethyne, C2H2, and Hydrogen cyanide, HCN.

58 Chapter 6 Resonance Structures
Section 2 Drawing and Naming Molecules Chapter 6 Resonance Structures Some molecules, such as ozone, O3, cannot be represented by a single Lewis structure. When a molecule has two or more possible Lewis structures, the two structures are called resonance structures.

59 Chapter 6 Visual Concepts Atomic Resonance

60 Multiple Bonds, continued
Section 2 Drawing and Naming Molecules Chapter 6 Multiple Bonds, continued Naming Covalent Compounds The first element named is usually the first one written in the formula. It is usually the less-electronegative element. The second element named has the ending -ide. Unlike the names for ionic compounds, the names for covalent compounds must often distinguish between two different molecules made of the same elements.

61 Naming Covalent Compounds, continued
Section 2 Drawing and Naming Molecules Chapter 6 Naming Covalent Compounds, continued This system of prefixes is used to show the number of atoms of each element in the molecule.

62 Naming Covalent Compounds, continued
Section 2 Drawing and Naming Molecules Chapter 6 Naming Covalent Compounds, continued Prefixes can be used to show the numbers of each type of atom in diphosphorus pentasulfide.

63 Naming Compounds Using Numerical Prefixes
Chapter 6 Visual Concepts Naming Compounds Using Numerical Prefixes

64 Homework Section 6.2 Review, pg. 207 questions 1 to 10

65 Section 6.2 Review, pg. 207 Which electrons do a Lewis structure show? 2. In a polyatomic ion, where is the charge located? 3. How many electrons are shared by two atoms that form a triple bond? 4. What do resonance structures represent? 5. How do the names for SO2 and SO3 differ?

66 Section 6.2 Review, pg. 207 6. Draw a Lewis structure for an atom that has the electron configuration 1s22s22p63s23p3. 7. Draw Lewis structures for each compound: BrF c. Cl2O b. N(CH3) d. ClO2

67 Section 6.2 Review, pg. 207 8. Draw three resonance structures for SO3. 9. Name the following compounds. SnI c. PCl3 b. N2O d. CSe2

68 Section 6.2 Review, pg. 207 10. Write the formula for each compound: a. phosphorus pentabromide b. diphosphorus trioxide c. arsenic tribromide d. carbon tetrachloride

69 Chapter 6 Section 3 Molecular Shapes Bellringer Write a short paragraph telling what you think the “valence shell electron pair repulsion theory” might have to do with molecular shape.

70 Chapter 6 Section 3 Molecular Shapes Objectives Predict the shape of a molecule using VSEPR theory. Associate the polarity of molecules with the shapes of molecules, and relate the polarity and shape of molecules to the properties of a substance.

71 Determining Molecular Shapes
Chapter 6 Section 3 Molecular Shapes Determining Molecular Shapes The three-dimensional shape of a molecule is important in determining the molecule’s physical and chemical properties. A Lewis Structure Can Help Predict Molecular Shape You can predict the shape of a molecule by examining the Lewis structure of the molecule.

72 Determining Molecular Shapes, continued
Chapter 6 Section 3 Molecular Shapes Determining Molecular Shapes, continued A Lewis Structure Can Help Predict Molecular Shape, continued The valence shell electron pair repulsion (VSEPR) theory is a theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other.

73 Determining Molecular Shapes, continued
Chapter 6 Section 3 Molecular Shapes Determining Molecular Shapes, continued Electron Pairs Can Determine Molecular Shape According to the VSEPR theory, the shape of a molecule is determined by the valence electrons surrounding the central atom. Electron pairs are negative, so they repel each other. Therefore, the shared pairs that form different bonds repel each other and remain as far apart as possible.

74 Determining Molecular Shapes, continued
Chapter 6 Section 3 Molecular Shapes Determining Molecular Shapes, continued Electron Pairs Can Determine Molecular Shape, continued For CO2, the two double bonds around the central carbon atom repel each other and remain far apart. For BF3, the three single bonds around the central fluorine atom will be at a maximum distance apart.

75 Determining Molecular Shapes, continued
Chapter 6 Section 3 Molecular Shapes Determining Molecular Shapes, continued Electron Pairs Can Determine Molecular Shape, continued The four shared pairs of electrons in CH4 are farthest apart when each pair is positioned at the corners of a tetrahedron.

76 VSEPR and Lone Electron Pairs
Chapter 6 Chapter 6 Visual Concepts VSEPR and Lone Electron Pairs

77 VSEPR and Basic Molecular Shapes
Chapter 6 Visual Concepts VSEPR and Basic Molecular Shapes

78 Predicting Molecular Shapes
Chapter 6 Section 3 Molecular Shapes Predicting Molecular Shapes Sample Problem D Determine the shape of H2O.

79 Predicting Molecular Shapes
Chapter 6 Section 3 Molecular Shapes Predicting Molecular Shapes Sample Problem D Solution Draw the Lewis structure for H2O. Count the number of shared and unshared pairs of electrons around the central atom. H2O has two shared pairs and two unshared pairs.

80 Predicting Molecular Shapes
Chapter 6 Section 3 Molecular Shapes Predicting Molecular Shapes Sample Problem D Solution, continued Find the shape that allows the shared and unshared pairs of electrons to be as far apart as possible. The water molecule will have a bent shape.

81 Molecular Shape Affects a Substance’s Properties
Chapter 6 Section 3 Molecular Shapes Molecular Shape Affects a Substance’s Properties Shape Affects Polarity One property that shape determines is the polarity of a molecule. The polarity of a molecule that has more than two atoms depends on the polarity of each bond and the way the bonds are arranged in space.

82 Molecular Shape Affects a Substance’s Properties, continued
Chapter 6 Section 3 Molecular Shapes Molecular Shape Affects a Substance’s Properties, continued Shape Affects Polarity, continued If two dipoles are arranged in opposite directions, they will cancel each other. If two dipoles are arranged at an angle, they will not cancel each other. Compare the molecules of nonpolar carbon dioxide, CO2, which has a linear shape, and polar water, H2O, which has a bent shape.

83 Chapter 6 Molecular Shape Affects Polarity

84 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 1. Which of these combinations is likely to have a polar covalent bond? A. two atoms of similar size B. two atoms of very different size C. two atoms with different electronegativities D. two atoms with the same number of electrons

85 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 1. Which of these combinations is likely to have a polar covalent bond? A. two atoms of similar size B. two atoms of very different size C. two atoms with different electronegativities D. two atoms with the same number of electrons

86 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 2. According to VSEPR theory, which of these is caused by repulsion between electron pairs surrounding an atom? F. breaking of a chemical bond G. formation of a sea of electrons H. formation of a covalent chemical bond I. separation of electron pairs as much as possible

87 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 2. According to VSEPR theory, which of these is caused by repulsion between electron pairs surrounding an atom? F. breaking of a chemical bond G. formation of a sea of electrons H. formation of a covalent chemical bond I. separation of electron pairs as much as possible

88 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 3. How many electrons are shared in a double covalent bond? A. 2 B. 4 C. 6 D. 8

89 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 3. How many electrons are shared in a double covalent bond? A. 2 B. 4 C. 6 D. 8

90 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 4. How can the difference in number of valence electrons between nitrogen and carbon account for the fact that the boiling point of ammonia, NH3, is 130°C higher than that of methane, CH4.

91 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 4. How can the difference in number of valence electrons between nitrogen and carbon account for the fact that the boiling point of ammonia, NH3, is 130°C higher than that of methane, CH4. Answer: Ammonia is a polar molecule because nitrogen has a pair of electrons that are not involved in a covalent bond, while methane is a nonpolar molecule. The attraction between polar ammonia molecules causes the higher boiling point.

92 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 5. Why don’t scientists need VESPR theory to predict the shape of HCl?

93 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 5. Why don’t scientists need VESPR theory to predict the shape of HCl? Answer: Because HCl has two atoms, the shape can be only linear.

94 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 6. What are the attractive and repulsive forces involved in a covalent bond and how do their total strengths compare?

95 Understanding Concepts
Chapter 6 Standardized Test Preparation Understanding Concepts 6. What are the attractive and repulsive forces involved in a covalent bond and how do their total strengths compare? Answer: Attractive forces exist between each electron and each nucleus. Repulsive forces exist between electrons and between nuclei. In a covalent bond, total attractive and repulsive forces are balanced.

96 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills Read the passage below. Then answer the questions. Although water is a polar molecule, pure water does not carry an electric current. It is a good solvent for many ionic compounds, and solutions of ionic compounds in water do carry electric currents. The charged particles in solution move freely, carrying electric charges. Even a dilute solution of ions in water becomes a good conductor. Without ions in solution, there is very little electrical conductivity.

97 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills 7. Why is a solution of sugar in water not a good electrical conductor? F. Sugar does not form ions in solution. G. The ionic bonds of sugar molecules are too strong to carry a current. H. Not enough sugar dissolves for the solution to become a conductor. I. A solution of sugar in water is not very conductive because it is mostly water, which is not very conductive.

98 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills 7. Why is a solution of sugar in water not a good electrical conductor? F. Sugar does not form ions in solution. G. The ionic bonds of sugar molecules are too strong to carry a current. H. Not enough sugar dissolves for the solution to become a conductor. I. A solution of sugar in water is not very conductive because it is mostly water, which is not very conductive.

99 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills 8. Why do molten ionic compounds generally conduct electric current well, while molten covalent compounds generally do not? A. Ionic compounds are more soluble in water. B. Ionic compounds have more electrons than compounds. C. When they melt, ionic compounds separate into charged particles. D. Most ionic compounds contain a metal atom which carries the electric current.

100 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills 8. Why do molten ionic compounds generally conduct electric current well, while molten covalent compounds generally do not? A. Ionic compounds are more soluble in water. B. Ionic compounds have more electrons than compounds. C. When they melt, ionic compounds separate into charged particles. D. Most ionic compounds contain a metal atom which carries the electric current.

101 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills If water is not a good conductor of electric current, why is it dangerous to handle an electrical appliance when your hands are wet or when you are standing on wet ground?

102 Chapter 6 Reading Skills
Standardized Test Preparation Reading Skills If water is not a good conductor of electric current, why is it dangerous to handle an electrical appliance when your hands are wet or when you are standing on wet ground? Answer: Because even a small amount of ionic compounds dissolved in water makes it a good conductor. The salts in your body or on the ground are enough to cause the water to carry a current.

103 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics Use the diagram below to answer question 10.

104 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 10. The diagram above best represents which type of chemical bond? F. ionic G. metallic H. nonpolar covalent I. polar covalent

105 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 10. The diagram above best represents which type of chemical bond? F. ionic G. metallic H. nonpolar covalent I. polar covalent

106 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics The table below shows the connection between electronegativity and bond strength (kilojoules per mole). Use it to answer questions 11 through 13.

107 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 11.Which of these molecules has the smallest partial positive charge on the hydrogen end of the molecule? A. HF B. HCl C. HBr D. HI

108 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 11.Which of these molecules has the smallest partial positive charge on the hydrogen end of the molecule? A. HF B. HCl C. HBr D. HI

109 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 12. How does the polarity of the bond between a halogen and hydrogen relate to the number of electrons of the halogen atom? F. Polarity is not related to the number of electrons of the halogen atom. G. Polarity decreases as the number of unpaired halogen electrons increases. H. Polarity decreases as the total number of halogen atom electrons increases. I. Polarity decreases as the number of valence electrons of the halogen atom increases.

110 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 12. How does the polarity of the bond between a halogen and hydrogen relate to the number of electrons of the halogen atom? F. Polarity is not related to the number of electrons of the halogen atom. G. Polarity decreases as the number of unpaired halogen electrons increases. H. Polarity decreases as the total number of halogen atom electrons increases. I. Polarity decreases as the number of valence electrons of the halogen atom increases.

111 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 13. Based on the information in this table, how does the electronegativity difference in a covalent bond relate to the strength of the bond?

112 Interpreting Graphics
Chapter 6 Standardized Test Preparation Interpreting Graphics 13. Based on the information in this table, how does the electronegativity difference in a covalent bond relate to the strength of the bond? Answer: A stronger bond is indicated by greater bond energy, so the strength of the bond increases as electronegativity increases.