1. Chapter 6 Objectives Section 1 Introduction to Chemical Bonding
Define chemical bond.
Explain why most atoms form chemical bonds.
List the major categories of chemical bonds.

2. Main Principles of Chemical Bonding
Chapter 6
Main Principles of Chemical Bonding
1. Bonding is a byproduct of atoms making change in their neutral ground state valency to achieve maximum charge-cloud stability ( i.e. match charge cloud character of a noble gas atom)
2. All chemical bonds are electrostatic forces of attraction [ which is the attractive force between opposing electrical charges (+/-) ].
During chemical bond formation amongst atoms, the atoms involved will lose, gain, or share valence electrons so that each atom involved in the bonding aquires a stable electron configuration matching that of a noble gas atom. This is the driving force that caused chemical bonds to form.

3. Chapter 6
The Octet Rule
Noble gas atoms are unreactive because their electron configurations are especially stable.
This stability results from the fact that the noble-gas atoms’ outer s and p orbitals are completely filled by a total of eight electrons. (With the exception of Helium that only has 2 valence electrons.)
Most bond formation follows the octet rule: Chemical bonds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level.

4. Visual Concepts
Chapter 6
The Octet Rule

5. Main Principles of Chemical Bonding
Chapter 6
Main Principles of Chemical Bonding
All chemical bond formation results in the release of energy from the matter being bonded. (i.e. Bond formation is an exergonic process.)
Matter is able to reduce its potential (stored) energy when it forms chemical bonds which further increases its stability.

6. Visual Concepts
Chapter 6
Chemical Bond

7. Main Categories of Chemical Bonds
Chapter 6
Main Categories of Chemical Bonds
Categorization or distinctions amongst chemical bonds are primarily based upon:
1. Location of the bond.
2. How the electrostatic force was brought into existence.
I. Intrastructural Bonds- bonds found within a structure between individual atoms that link individual atoms to build independent structures such as salt crystals, molecules, or metal crystals.
A. Ionic Bonds (a.k.a. Electrovalent Bonds)
B. Covalent Bonds
C. Metallic Bonds

8. Main Categories of Chemical Bonds
Chapter 6
Main Categories of Chemical Bonds
II. Interstructural Bonds (a.k.a. Intermolecular Forces or Van der Waals Forces- bonds found between molecules that link individual molecules together.
A. Hydrogen Bonds
B. Dipole-Dipole Interactive Forces
C. London Dispersion Forces

9. Chapter 6 Objectives Describe ionic and covalent bonding.
Describe formula units for ionic compounds.
Define molecule and molecular formula
Predict/Classify bonding type according to electronegativity differences.

10. Chapter 6 Ionic Bonding Section 3 Ionic Bonding and Ionic Compounds
Ionic bonding primarily occurs between the atoms of metallic and nonmetallic elements.
Ionic bonding involves the loss and gain of valence electrons by neutral atoms to become net charged ions.
The chemical process directly responsible for the formation of ions from neutral atoms is valence electron transfer.
The atoms of metal elements transfer away valence electrons to become net positive ions called cations.

11. Chapter 6 Section 3 Ionic Bonding and Ionic Compounds
The atoms of nonmetal elements receive the transferred electrons into their valence shell and become net negative ions called anions.
3 atomic properties are used to help predict the bonding nature of elements
1. Electronegativity- a measure of how an atom’s nucleus “pulls” on the valence electrons of nearby atoms trying to draw them into its valence shell.
2. Ionization Energy- a measure of the required amount of energy that an atom must absorb to transfer away an electron
3. Electron Affinity- a measure of the strength of attraction an atom’s nucleus has for its own electrons

12. General Values for atomic properties of metals and nonmetals
Atomic Property
Metal
Nonmetal
Electronegativity
low
high
1st Ionization Energy
Electron Affinity
A quick comparison of these properties reveals why the “fixed” direction of valence electron transfer is always
from metal to nonmetal.

13. Energetics of Ionic Bonding
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Energetics of Ionic Bonding
Cation formation is an energy requiring process (ionization energy), i.e. it is endergonic.
Anion formation is an energy releasing process, i.e. it is exergonic.

14. Visual Concepts
Chapter 6
Ionic Bonding

15. Chapter 6 Ionic Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Compounds
Most of the rocks and minerals that make up Earth’s crust consist of positive and negative ions (cations /anions) held together by ionic bonding.
example: table salt, NaCl, consists of sodium and chloride ions combined in a one-to-one ratio—Na+Cl–—so that each positive charge is balanced by a negative charge.
An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.

16. Chapter 6 Ionic Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Compounds
Most ionic compounds exist as crystalline solids.
A crystal of any ionic compound is a three-dimensional network of positive and negative ions mutually attracted to each other.
Once an ionic crystal is formed, all available areas of opposing charge are utilized in the ionic bonds between cations and anions; therefore, ionic crystals have no ability to engage in any further level of bonding (i.e. do not have Van der Waals Forces available).

17. Ionic Compounds, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Compounds, continued
The chemical formula of an ionic compound represents not molecules, but the simplest whole number ratio of the compound’s ions that allow their opposing charges to sum to zero (i.e. cancel each other out).
This type of formula to represent ionic compounds is known as a formula unit.

18. Formation of Ionic Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Formation of Ionic Compounds
The sodium atom has one valence electron and the chlorine atom has seven valence electrons.
Atoms of sodium and other alkali metals easily lose one electron to form cations with a fixed ionic charge of 1+.
Atoms of chlorine and other halogens easily gain one electron to form anions with a fixed ionic charge of 1-.

19. Formation of Ionic Compounds, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Formation of Ionic Compounds, continued
In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice.
Attractive forces exist between oppositely charged ions within the lattice.
Repulsive forces exist between like-charged ions within the lattice.
The combined attractive and repulsive forces within a crystal lattice determine:
the distances between ions
the pattern of the ions’ arrangement in the crystal

20. Characteristics of Ion Bonding in a Crystal Lattice
Visual Concepts
Chapter 6
Characteristics of Ion Bonding in a Crystal Lattice

21. Electron-Dot Notation
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation
To keep track of valence electrons, it is helpful to use electron-dot notation.
Electron-dot notation is an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. The inner-shell electrons are not shown.

22. The arrangement of elements in the Periodic Table is related to the electron structure of the element.
Valence electrons are the electrons in the outermost energy level.
Electron dot diagrams use the symbol of the element to represent the kernel (nucleus & nonvalence electrons) of the atom and dots to represent valence electrons.
In general, all elements in the same group have the same number of valence electrons.
Elements of the same Period have their valence electrons in the same energy level.

23. Electron-Dot Notation, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation, continued
Sample Problem B
a. Write the electron-dot notation for hydrogen.
b. Write the electron-dot notation for nitrogen.

24. Electron-Dot Notation, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Electron-Dot Notation, continued
Sample Problem B Solution
a. A hydrogen atom has only one occupied energy level, the n = 1 level, which contains a single electron.
b. The group notation for nitrogen’s family of elements is ns2np3.
Nitrogen has five valence electrons.

25. Electron-Dot Notation
Visual Concepts
Chapter 6
Electron-Dot Notation

26. LEWIS DIAGRAMS for IONS/IONIC COMPOUNDS
Metal cations are displayed with no remaining
valence (remember that they transfer them
away!!) . Brackets MUST be placed around the
element symbol for the metal and the correct
ionic charge MUST be placed outside the brackets
to the upper right .
Ex: NaCl , a
binary salt
Nonmetal anions are displayed with a complete octet
of valence (remember that they gain them !!!)
Brackets MUST be placed around the
element symbol for the nonmetal and the correct
ionic charge MUST be placed outside the brackets
to the upper right .

27. LEWIS DIAGRAMS for IONS/IONIC COMPOUNDS
EX: MgCl2 , magnesium chloride . A binary salt
The brackets are necessary to show that the electrons are not
being shared and that the bonding is caused by the
net opposing charge on the surface of the ions.

28. How to Identify a Compound as Ionic
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
How to Identify a Compound as Ionic

29. How to Identify a Compound as Ionic
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
How to Identify a Compound as Ionic

30. Covalent Bonds Chapter 6
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Covalent Bonds
Covalent bonding involves the sharing of valence electrons in pairs
between the nuclei of bonding atoms.
The shared pairs of valence electrons are established by the
overlapping of valence shell orbitals between bonding atoms.
Covalent bonding primarily occurs between atoms of
nonmetal elements. (nonmetal to nonmetal bonding)

31. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
A molecule is a neutral group of nonmetal atoms that are held together by covalent bonds.
A chemical compound whose simplest units are molecules is called a molecular compound.
The term MOLECULE should only be used to describe a chemical structure that is held together by COVALENT BONDS!!!!!!!!!!!!!!!

32. Visual Concepts
Chapter 6
Molecule

33. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond
Most atoms have lower potential energy when they are bonded to other atoms than they have as they are independent particles (the exceptions are the noble gas atoms.
The figure below shows potential energy changes during the formation of a hydrogen-hydrogen bond.

34. Characteristics of the Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Characteristics of the Covalent Bond
When two atoms form a covalent bond, their shared electrons form overlapping orbitals.
This achieves a noble-gas configuration.
The bonding of two hydrogen atoms allows each atom to have the stable electron configuration of helium, 1s2.

35. Chapter 6 Section 2 Covalent Bonding and Molecular Compounds
A single covalent bond, or single bond, is a covalent bond in which one pair of valence electrons is shared between two atoms.
The norm in single bond formation is each atom participating in the formation of a shared pair of valence electrons contributes one electron apiece into the creation of the shared pair (i.e. “equal donation into the creation”).
The Coordinate Covalent bond is a special case of forming a shared pair of valence electrons where only one of the atoms participating in the formation of a shared pair contributes both electrons ( i.e. “unequal donation into the creation”).

36. Coordinate bonds frequently appear in ternary molecules and polyatomic ions.

37. Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between the same two atoms.
Double bonds are often found in molecules containing carbon, nitrogen, and oxygen.
A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes.

38. Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between the same two atoms.
example 1—diatomic nitrogen:
example 2—ethyne, C2H2:

39. Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds.
In general, double bonds have greater bond energies and are shorter than single bonds.
Triple bonds are even stronger and shorter than double bonds.
When writing Lewis structures for molecules that contain carbon, nitrogen, or oxygen, remember that multiple bonds between pairs of these atoms are possible.

40. Comparing Single, Double, and Triple Bonds
Visual Concepts
Chapter 6
Comparing Single, Double, and Triple Bonds

41. Bond Prediction
The difference in electronegativity value between two atoms in a
chemical bond can be used to predict what type of bond will form
according to the following scale of calculated differences:
0.0  nonpolar covalent bond
.03  polar covalent bond
1.7  ionic bond
***Please note that the nature of the elements involved in the bond
supercedes these differences when characterizing a bond

42. Bond polarity is an indicator of the degree of sharing
of valence electrons in the bond.
In nonpolar covalent bonding the pair of electrons is shared equally
and therefore there is an even distribution of charge in the bonding.
In polar covalent bonding the pair of electrons is shared unequally and therefore there is an uneven distribution of charge in the bonding that creates bond polarity.

43. A dipole symbol is often used to show polar bonds and the partial positive and partial negative charge involved in the bonding.

44. Comparing Polar and Nonpolar Covalent Bonds
Visual Concepts
Chapter 6
Comparing Polar and Nonpolar Covalent Bonds

45. Using Electronegativity Difference to Classify Bonding
Visual Concepts
Chapter 6
Using Electronegativity Difference to Classify Bonding

46. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A
Use electronegativity values to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

47. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A Solution
The electronegativity of sulfur is 2.5. The electronegativities of hydrogen, cesium, and chlorine are 2.1, 0.7, and 3.0, respectively. In each pair, the atom with the larger electronegativity will be the more-negative atom.
Bonding between Electroneg. More-neg-
sulfur and difference Bond type ative atom
hydrogen 2.5 – 2.1 = 0.4 polar-covalent sulfur
cesium – 0.7 = 1.8 ionic sulfur
chlorine – 2.5 = 0.5 polar-covalent chlorine

48. Chapter 6 Objectives, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Objectives, continued
List the basic steps used in writing Lewis structures for molecules/polyatomic ions.
Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.

49. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
The composition of a compound is given by its chemical formula.
A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound.

50. Visual Concepts
Chapter 6
Chemical Formula

51. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
Electron-dot notation can also be used to represent molecules.
The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond.
For a molecule of fluorine, F2, the electron-dot notations of two fluorine atoms are combined.

52. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots between the two symbols represents the shared pair of a covalent bond.
In addition, each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds.
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

53. Visual Concepts
Chapter 6
Lewis Structures

54. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash.
example:
A structural formula indicates the kind, number, and arrangement, and bonds but not the unshared pairs of the atoms in a molecule.
example: F–F H–Cl

55. Visual Concepts
Chapter 6
Structural Formula

56. Drawing Lewis Structures with Many Atoms
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Drawing Lewis Structures with Many Atoms

57. Drawing Lewis Structures with Many Atoms
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Drawing Lewis Structures with Many Atoms

58. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C
Draw the Lewis structure of iodomethane, CH3I.

59. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution
1. Determine the type and number of atoms in the molecule.
The formula shows one carbon atom, one iodine atom, and three hydrogen atoms.
2. Write the electron-dot notation for each type of atom in the molecule.
Carbon is from Group 14 and has four valence electrons.
Iodine is from Group 17 and has seven valence electrons.
Hydrogen has one valence electron.

60. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution, continued
3. Determine the total number of valence electrons available in the atoms to be combined. C
1 × 4e– =
4e– I
1 × 7e–
7e– 3H
3 × 1e–
3e–
14e–

61. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Sample Problem C Solution, continued
4. If carbon is present, it is the central atom. Otherwise, the least-electronegative atom is central. Hydrogen, is never central.
5. Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons.

62. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D
Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde.

63. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution
1. Determine the number of atoms of each element present in the molecule.
The formula shows one carbon atom, two hydrogen atoms, and one oxygen atom.
2. Write the electron-dot notation for each type of atom.
Carbon is from Group 14 and has four valence electrons.
Oxygen, which is in Group 16, has six valence electrons.
Hydrogen has only one valence electron.

64. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
3. Determine the total number of valence electrons available in the atoms to be combined. C
1 × 4e– =
4e– O
1 × 6e–
6e– 2H
2 × 1e–
2e–
12e–

65. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
Arrange the atoms to form a skeleton structure for the molecule. Connect the atoms by electron-pair bonds.
Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons.

66. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
6a.Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the number available.
The structure has 14 electrons. The structure has two valence electrons too many.
6b.Subtract one or more lone pairs until the total number of valence electrons is correct.
Move one or more lone electron pairs to existing bonds until the outer shells of all atoms are completely filled.

67. Multiple Covalent Bonds, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds, continued
Sample Problem D Solution, continued
Subtract the lone pair of electrons from the carbon atom. Move one lone pair of electrons from the oxygen to the bond between carbon and oxygen to form a double bond.

68. Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Objectives
Compare a chemical formula for a molecular compounds with one for an ionic compound.
List and compare the distinctive properties of ionic and molecular compounds.
Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information.

69. Ionic Vs. Covalent Bonding
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Vs. Covalent Bonding

70. A Comparison of Ionic and Molecular Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
A Comparison of Ionic and Molecular Compounds
The force that holds ions together in an ionic compound is a very strong electrostatic attraction.
In contrast, the forces of attraction between molecules of a covalent compound are much weaker.
This difference in the strength of attraction between the basic units of molecular and ionic compounds gives rise to different properties between the two types of compounds.

71. A Comparison of Ionic and Molecular Compounds, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
A Comparison of Ionic and Molecular Compounds, continued
Molecular compounds have relatively weak forces between individual molecules.
They melt at low temperatures.
The strong attraction between ions in an ionic compound gives ionic compounds some characteristic properties, listed below.
very high melting points
hard but brittle
not electrical conductors in the solid state, because the ions cannot move

72. Melting and Boiling Points of Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Melting and Boiling Points of Compounds

73. Comparing Ionic and Molecular Compounds
Visual Concepts
Chapter 6
Comparing Ionic and Molecular Compounds

74. Comparing Ionic and Molecular Substances
Section 5 Molecular Geometry
Chapter 6
Comparing Ionic and Molecular Substances

75. Chapter 6 Polyatomic Ions
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Polyatomic Ions
Certain atoms bond covalently with each other to form a group of atoms that has both molecular and ionic characteristics.
A charged group of covalently bonded atoms is known as a polyatomic ion ( a net charged molecule).
Like other ions, polyatomic ions have a charge that results from either a shortage or excess of electrons.

76. Chapter 6 Polyatomic Ions
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Polyatomic Ions
An example of a polyatomic ion is the ammonium ion: It is sometimes written as to show that the group of atoms as a whole has a charge of 1+.
The charge of the ammonium ion is determined as follows:
The seven protons in the nitrogen atom plus the four protons in the four hydrogen atoms give the ammonium ion a total positive charge of 11+.

77. Polyatomic Ions, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Polyatomic Ions, continued
The charge of the ammonium ion is determined as follows, continued:
When nitrogen and hydrogen atoms combine to form an ammonium ion, one of their electrons is lost, giving the polyatomic ion a total negative charge of 10–.
The total charge is therefore (11+) + (10–) = 1+.

78. Polyatomic Ions, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Polyatomic Ions, continued
Some examples of Lewis structures of polyatomic ions are shown below.

79. Chapter 6 Multiple Choice
Standardized Test Preparation
Chapter 6
Multiple Choice
1. A chemical bond results from the mutual attraction of the nuclei for
A. electrons.
B. neutrons.
C. protons.
D. dipoles.

80. Chapter 6 Multiple Choice
Standardized Test Preparation
Chapter 6
Multiple Choice
1. A chemical bond results from the mutual attraction of the nuclei for
A. electrons.
B. neutrons.
C. protons.
D. dipoles.

81. Chapter 6 Multiple Choice
Standardized Test Preparation
Chapter 6
Multiple Choice
2. A polar covalent bond is likely to form between two atoms that
A. are similar in electronegativity.
B. are of similar size.
C. differ in electronegativity.
D. have the same number of electrons.

82. Chapter 6 Multiple Choice
Standardized Test Preparation
Chapter 6
Multiple Choice
2. A polar covalent bond is likely to form between two atoms that
A. are similar in electronegativity.
B. are of similar size.
C. differ in electronegativity.
D. have the same number of electrons.

83. Chapter 6 Multiple Choice 3. The Lewis structure of HCN contains
Standardized Test Preparation
Chapter 6
Multiple Choice
3. The Lewis structure of HCN contains
A. one double bond and one single bond.
B. one triple bond and one single bond.
C. two single bonds.
D. two double bonds.

84. Chapter 6 Multiple Choice 3. The Lewis structure of HCN contains
Standardized Test Preparation
Chapter 6
Multiple Choice
3. The Lewis structure of HCN contains
A. one double bond and one single bond.
B. one triple bond and one single bond.
C. two single bonds.
D. two double bonds.

85. Chapter 6 Multiple Choice 5. Which molecule contains a double bond?
Standardized Test Preparation
Chapter 6
Multiple Choice
5. Which molecule contains a double bond?
A. COCl2
B. C2H6
C. CF4
D. SF2

86. Chapter 6 Multiple Choice 5. Which molecule contains a double bond?
Standardized Test Preparation
Chapter 6
Multiple Choice
5. Which molecule contains a double bond?
A. COCl2
B. C2H6
C. CF4
D. SF2