Chapter 8: Chemical Bonding

1. Chapter 8: Chemical Bonding

2. Questions for Consideration • How can we classify the types of bonding of different compounds? • What is the nature of the bonding in ionic compounds? • What is the nature of the bonding in molecular compounds? • How do we predict and explain the shapes of molecular compounds and polyatomic ions?

3. Chapter 8 Topics: • Types of Bonds • Ionic Bonding • Covalent Bonding • Bonding in Carbon Compounds • Shapes of Molecules

4. Introduction • How and why do atoms come together (bond) to form compounds? • Why do different compounds have such different properties? • What do molecules look like in three dimensions?

5. 8.1 Types of Bonds • Chemical bond • A force that holds atoms together in a molecule or compound • Two types of chemical bonds • Ionic Bonds • Covalent Bonds Figure 8.2 Figure 8.2

6. Table 8.1 Properties of Two Carbon Compounds Figure 8.2

7. General Properties of Ionic and Covalent Substances

8. Ionic and Covalent • In ionic compounds, ions are held together by electrostatic forces – forces between oppositely charged ions. • In molecular compounds, atoms are held together by covalent bonds in which electrons are shared. Figure 8.2

9. Ionic Bonds • A bond created by electrostatic attraction between oppositely charged ions • Occurs between a metal and a nonmetal • Electrons transferred between the cation (positively charged ion) and the anion (negatively charged ion) • Extremely strong bonds Figure 8.7

10. Covalent Bonds • A bond created by the sharing of electrons between atoms • Occurs between two nonmetals (resulting in a neutral overall charge) • Electrons not transferred in this case • Electrons typically shared in pairs • Weaker bonds than ionic bonds Figure 8.2

11. Activity: Identifying Types of Bonding • Identify the type of bonding in each of the following substances: • NaF • ClO2 • FeSO4 • SO2 • Ca(ClO2)2

12. Activity Solutions: Identifying Types of Bonding • Identify the type of bonding in each of the following substances: • NaF – Ionic bonding (metal + nonmetal) • ClO2 – Covalent bonding (2 nonmetals) • FeSO4 – Ionic bonding between the metal and polyatomic ion; covalent bonding between the nonmetals in the polyatomic ion • SO2 – Covalent bonding • Ca(ClO2)2 - Ionic bonding between the metal and polyatomic ion; covalent bonding between the nonmetals in the polyatomic ion

13. Polar vs. Nonpolar • Two general types of covalent bonds: • Polar covalent • Unequal sharing (or a partial transfer) of electrons • Occurs when different elements are covalently bonded to one another • Why different elements? • Because different elements have different electronegativities • Nonpolar covalent • Equal sharing (no transfer) of electrons • Occurs only when all of the atoms in a molecule belong to the same element

14. Polar vs. Nonpolar • Polar covalent bonds are: • Typically shorter bonds • Stronger bonds due to their increased ionic character • Nonpolar covalent bonds are: • Typically longer bonds • Weaker bonds • Polarity • Occurs in polar covalent molecules • Polarity is the degree of transfer of electrons in a covalently bonded molecule composed of different element’s atoms.

15. Classification of Bonding Figure 8.3

16. Polar Covalent Bonds • In polar covalent bonds, electrons are not shared equally. This result in unequal sharing which can be described as partial electron transfer. • In a polar bond there are partial charges on the atoms sharing electrons. Figure 8.4

17. Electronegativity • Ability of an atom to attract bonding electrons • Proposed by Linus Pauling in the early 1930’s • A difference in electronegativity between the atoms in a covalent bond results in: • A polar covalent bond • Increased ionic character • The greater the difference in electronegativity, the greater the ionic character and the more polar the bond that joins the atoms. • Decreased bond length and increased bond strength • No difference in electronegativity between atoms in a covalent bond results in a nonpolar covalent bond.

18. Electronegativity • The electronegativity scale tells us which elements have a greater pull on electrons in a covalent bond. • The more electronegative element pulls bonding electrons more strongly and obtains a partial negative charge (). • The less electronegative element loses electron density in the bond and therefore obtains a partial positive charge (+). Figure 8.4

19. Electronegativity Values Figure8.5

20. Electronegativity Trends Figure 8.6

21. Electronegativity Trends • The difference in electronegativity between metals and nonmetals is so large, that the electrons are transferred, not shared. • The greater the electronegativity difference, the more polar the bond. • Si-F > N-F> O-F >F-F • What partial charges go on each atom? Figure 8.6

22. Activity: Polar Bonds • Which of the following molecules have polar bonds? If a bond is polar, which atom has a partial negative charge? • SO2 • N2 • H2S • CCl4 • O3

23. Activity Solutions: Polar Bonds • Which of the following molecules have polar bonds? If a bond is polar, which atom has partial negative charge? • SO2 – Polar covalent bonds; O is more electronegative and has a partial negative charge • N2 – Nonpolar covalent bonds • H2S – Polar covalent bonds; S is more electronegative and has a partial negative charge • CCl4 – Polar covalent bonds; Cl is more electronegative and has a partial negative charge • O3 – Nonpolar covalent bonds

24. 8.2 Ionic Bonding • Each element immediately following a noble gas is a metal. • Metals lose electrons, forming a positive charge, to become cations. • Each element immediately preceding a noble gas is a nonmetal. • Nonmetals gain electrons, forming a negative charge, to become anions. • Formation of ions and ionic bonds relates to an element’s electron configuration. • Many main-group elements either lose or gain electrons to become isoelectronic with a noble gas (i.e. have the same electron configuration). As ions, they are known as the common ions.

25. Formation of Ionic Bonds • To understand ionic bonding, we can think about how ions and ionic bonds would be formed from neutral atoms of the elements. • Consider the formation of NaCl from its elements: • Electrons are transferred from the metal to the nonmetal to form ions, each with a noble-gas electron configuration. Figure 6.6

26. Lewis Dot Symbols • Lewis Dot symbol • Electron dot symbol • Dots placed around an element’s symbol represent valence electrons • Pair electrons as needed • Octet rule • Tendency of an atom to achieve an electron configuration having 8 valence electrons • Same as the electron configuration of a noble gas • The 8 electrons exist in 4 pairs • Ions achieve 8 electrons by losing or gaining electrons

27. Lewis Symbols • Lewis Symbols help us to focus on the valence electrons – those that can participate in bonding.

28. Activity: Lewis Symbols for Ions • Write the Lewis symbols for beryllium and nitrogen ions. Then write a formula for the compound that would form between them, using their Lewis symbols.

29. Activity Solutions: Lewis Symbols for Ions . . . . . . symbol • Write the Lewis symbols for the beryllium and nitrogen ions. Then write a formula for the compound that would form between them, using their Lewis symbols.

30. Ionic Compounds • Ions of like charge repel each other in ionic compounds, and opposite charged ions attract. This results in a 3-dimensional regular pattern called a crystal lattice. Figure 8.7

31. Structures of Ionic Compounds • Crystal lattice • The pattern obtained when an ion, represented as a charged sphere, exerts a force equally in all directions. • Thus, ions of equal and opposite charge surround it. • Cations and anions must come into contact for a crystal lattice to form. Figure 8.10 fluorite, CaF2

32. Structures of Ionic Compounds Sodium Chloride, NaCl, Crystal Figure 8.8

33. Cesium Chloride Crystal Figure 8.9

34. Calcium Fluoride Crystal Figure 8.10

35. Properties of Ionic Solids • Why are crystalline ionic solids hard and brittle? • Why are they poor conductors of electricity in the solid state? Figure 8.11

36. Structures of Ionic Crystals • Ionic crystal • Ions are arranged in a regular geometric pattern that maximizes the attractive forces and minimizes the repulsive forces. • Hard and brittle • Can shatter if struck forcefully • The charges and sizes of ions largely determine the characteristic patterns of ionic crystals Figure 8.11

37. 8.3 Covalent Bonding • When two nonmetals form a bond, the bond is covalent. They are both close to the noble-gas electron configuration, so sharing will allow both to obtain it. • In a covalent bond, each shared electron interacts simultaneously with two nuclei. Figure 8.12

38. Carbon Dioxide – Covalent bonds • The atoms of CO2 molecules are held together by strong covalent bonds. No bonds connect the molecules, so CO2 molecules separate from each other into the gas state at room temperature. Figure from p. 27 Figure 8.13

39. The Octet Rule • Just as in ionic bonding, covalent bonds are formed so that each atom can have the noble-gas electron configuration. Noble gases have 8 valence electrons, an octet. Figure 8.14

40. The Octet Rule • Octet rule • Tendency of an atom to achieve an electron configuration having 8 valence electrons • Same as the electron configuration of a noble gas • Covalently bonded atoms achieve 8 valence electrons by sharing electrons • The 8 electrons exist in 4 pairs • H atoms bond with other atoms to obtain a total of 2 electrons like He.

41. Lewis Formulas for the Diatomic Elements • How does hydrogen obtain a noble-gas electron configuration? Figure 8.15

42. The Halogens • Do the atoms in each of these molecules have an octet? • Why do the halogens exist as diatomic molecules? Figure 8.16

43. Multiple Bonds • How many valence electrons does an oxygen atom have? • How many does it need to obtain an octet? • Can a single covalent bond allow each oxygen atom to satisfy the octet rule? • O2 has a double bond, two pairs of shared electrons Figure from p. 315 Figure 8.17

44. Multiple Bonds • How many valence electrons does a nitrogen atom have? • How many does it need to obtain an octet? • N2 has a triple bond, three pairs of shared electrons Figure 8.17

45. P4 and S8 • How do phosphorus and sulfur obtain an octet in P4 and S8? Figure 8.18

46. Activity: Valence Electrons and Number of Bonds • How many bonds do each of the following atoms tend to form? • H • Cl • O • N • C 1 1 2 3 4

47. Activity: Lewis Structures • Draw the Lewis structures for each of the following based on the number of bonds that each tends to form. Remember to include the nonbonding electrons so that all have octets (except H). • C2H6 • C2H4 • C2H2 • HCN

48. Activity Solutions: Lewis Structures • Draw the Lewis structures for each of the following based on the number of bonds that each tends to form. • Carbon forms 4 bonds and hydrogen forms 1: • Carbon forms 4 bonds and hydrogen forms 1: • Carbon forms 4 bonds and hydrogen forms 1: • Hydrogen forms 1 bond, carbon forms 4 bonds, and nitrogen forms 3 bonds:

49. Extra Activity: Valence Electrons and Number of Bonds • Draw a Lewis structure for each of the following based on how many bonds each tends to form. Remember to include the nonbonding electrons so that all have octets (except H). • H2CO • NF3 • H2O

50. Steps for Writing Lewis Structures • Write an atomic skeleton. • Sum the valence electrons from each atom to get the total number of valence electrons. • Place two electrons, a single bond, between each pair of bonded atoms. • Add remaining electrons to complete the octet of each outer atom and then to the central atom if necessary and if there are electrons available. • If necessary to satisfy the octet rule, shift unshared electrons from non-bonded positions on atoms with completed octets to positions between atoms to make double or triple bonds.