1. Covalent Bonding Chapter 8

2. Chapter Main Idea Covalent bonds form when atoms share electrons!!!

3. Section 1:The Covalent Bond

4. Section 1: Essential Questions & Vocabulary ▪How does the octet rule apply to atoms that form covalent bonds? ▪Why do atoms form single, double, and triple covalent bonds? ▪What are sigma and pi bonds and how do they contrast? ▪How are the strength of a covalent bond, its bond length, and its bond dissociation energy related? ▪Chemical bond ▪Covalent bond ▪Molecule ▪Lewis structure ▪Sigma bond ▪Pi bond ▪Endothermic reaction ▪Exothermic reaction Vocabulary

5. Section 1: Main Idea Atoms gain stability when they share electrons and form covalent bonds.

6. Why do atoms bond? ▪The stability of an atom, ion or compound is related to its energy: – lower energy states are more stable. ▪Elements can gain stability in two ways: – Transferring electrons – elements gain or lose electrons to form ions with noble-gas configurations. (Ionic Bonds) – Sharing electrons – elements share valence electrons with other atoms resulting in noble-gas configurations. (Covalent Bonds)

7. What is a covalent bond? ▪Covalent bond - chemical bond that results from sharing electrons ▪Molecule is formed when two or more atoms bond covalently. – Most covalent bonds form between nonmetal - nonmetal elements. ▪Diatomic molecules (H 2, N 2, F 2, O 2, I 2, Cl 2, Br 2 ) exist because the two-atom molecules are more stable than the individual atoms.

8. Formation of Cl 2 Cl Each chlorine atom wants to gain one electron to have a full octet.

9. Formation of Cl 2 Chlorine atoms get close together until they share electrons. Cl Full Octet The octet is achieved by each atom sharing the electron pair in the middle.

10. Cl Single bonds are abbreviated with a dash Chlorine Molecule This is the bonding pair or a shared pair of electrons called a single bond. Cl

11. What is a covalent bond? ▪The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

12. Single Covalent Bonds ▪single covalent bond - one pair of electrons is shared between two atoms ▪Example: two hydrogen atoms forming a hydrogen molecule with a single covalent bond

13. Single Covalent Bonds by Group ▪Lewis structure – 2 dots or a line are used to symbolize a single covalent bond. Group 14Group 15Group 16Group 17 Each element has 4 valence electrons Will form 4 single bonds with other non-metals Each element has 5 valence electrons Will form 3 single bond with other non-metals Each element has 6 valence electrons Will form 2 single bond with other non-metals Each element has 7 valence electrons Will form 1 single bond with other non-metals

14. Sigma Bonds ▪Sigma bonds - single covalent bonds. ▪Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms.

15. Double Bonds & Triple Bonds ▪ Double Bond – two pairs of shared electrons – 1 sigma bond, 1 pi bond ▪ Triple Bond – three pairs of shared electrons – 1 sigma bond, 2 pi bonds

16. Formation of Oxygen Molecule OO Each atom has two unpaired electrons OO Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.

17. OO Both electron pairs are shared. Formation of Oxygen Molecule 6 valence electrons plus 2 shared electrons = full octet O O

18. two bonding pairs, or two pairs of shared electrons called a double bond. O O Formation of Oxygen Molecule O O = For convenience, the double bond can be shown as two dashes.

19. Nitrogen Molecule

20. Lewis Dot Structures: Practice ▪ PH 3 ▪ H 2 S ▪ HCl ▪ CCl 4 ▪ SiH 4

21. Strength of Covalent Bonds ▪ Bond Length – distance between the two bonded nuclei Single Bonds are the longest and weakest bonds. Double bonds are shorter but stronger than single bonds. Triple bonds are shorter than double but stronger than double bonds.

22. Bond Length & Bond Strength

23. Bonds and Energy ▪ Endothermic Reaction – greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the products. – Energy goes into the reaction. ▪ Exothermic Reaction – more energy is released during product bond formation than is required to break the bonds in the reactants. – Energy is released from the reaction. Exothermic

24. Section 2: Naming Molecules

25. Essential Questions and Vocabulary ▪ What rules do you follow to name a binary molecular compound from its molecular formula? ▪ How are acidic solutions named? ▪ Oxyanion ▪ Oxyacid Vocabulary

26. Section 2: Main Idea Specific rules are used when naming binary molecular compounds, binary acids, and oxyacids.

27. Naming Binary Molecular Compounds ▪ Binary Molecule – Compound composed of only two nonmetal atoms – Use of prefixes to indicate the number of each atom in the molecule.

28. Prefixes in Covalent Compounds ▪ Exception: The first element NEVER uses the prefix mono-

30. Naming Binary Molecules - RULES 1- The name of the first element does not get changed, only a prefix is added if it is more than one. Ex: Trisulfur tetrafluoride S 3 F 4 2- The second element gets a different ending, as well as a prefix. Ex : phosphorus trichloride PCl 3 (NOT: phosphorus trichlorine)

31. 3- If the first element in the compound is single you do not write mono: ex. CO 2 - carbon dioxide (NOT monocarbon dioxide) 4- If the element’s name starts with a vowel, and the prefix ends with a vowel, you can drop the last vowel on the prefix: ex. S 5 O 10 - pentasulfur decoxide (NOT pentasulfur decaoxide) Naming Binary Molecules – Rules cont.

32. 5- Covalent compounds do not get simplified. Name the compound exactly as you see it: Ex. H 2 F 2 = dihydrogen difluoride (NOT hydrogen monofluoride HF) Naming Binary Molecules – Rules cont.

33. Naming Binary Molecules: Examples 33 What is the name of SO 3 ? STEP 1 : The first nonmetal is S sulfur. STEP 2 : The second nonmetal is O, named oxide. STEP 3 : The subscript 3 of O is shown as the prefix tri. SO 3 → sulfur trioxide The subscript 1(for S) or mono is understood.

34. 34 Name P 4 S 3 STEP 1 : The first nonmetal P is phosphorus. STEP 2 : The second nonmetal S is sulfide. STEP 3 : The subscript 4 of P is shown as tetra. The subscript 3 of O is shown as tri. P 4 S 3 → tetraphosphorus trisulfide Naming Binary Molecules: Examples

35. 35 Names of Binary Molecules Remember Prefixes are used  in the names of covalent compounds  because two nonmetals can form two or more different compounds Examples of compounds of N and O: NO nitrogen oxide NO 2 nitrogen dioxide N 2 O dinitrogen oxide N 2 O 4 dinitrogen tetroxide N 2 O 5 dinitrogen pentoxide

36. Naming Binary Molecules: Practice 36 Select the correct name for each compound. A.SiCl 4 1) silicon chloride 2) tetrasilicon chloride 3) silicon tetrachloride B. P 2 O 5 1) phosphorus oxide 2) phosphorus pentoxide 3) diphosphorus pentoxide C.Cl 2 O 7 1) dichlorine heptoxide 2) dichlorine oxide 3) chlorine heptoxide

37. Naming Binary Molecules: Practice Solutions 37 A.SiCl 4 3) silicon tetrachloride B. P 2 O 5 3) diphosphorus pentoxide C.Cl 2 O 7 1) dichlorine heptoxide

38. 38 Write the correct formula for each of the following: A. phosphorus pentachloride B. dinitrogen trioxide C. sulfur hexafluoride Naming Binary Molecules: Practice

39. Naming Binary Molecules: Practice Solutions A. phosphorus pentachloride 1P penta = 5ClPCl 5 B. dinitrogen trioxide di = 2N tri = 3 O N 2 O 3 C. sulfur hexafluoride 1S hexa = 6FSF 6

40. 40 Write the name of each covalent compound. CO_____________________ CO 2 _____________________ PCl 3 _____________________ CCl 4 _____________________ N 2 O_____________________ Naming Binary Molecules: Practice

41. 41 CO carbon monoxide CO 2 carbon dioxide PCl 3 phosphorus trichloride CCl 4 carbon tetrachloride N 2 Odinitrogen oxide Naming Binary Molecules: Practice Solutions

42. TRY THESE: a)Diphosphorus pentoxide b)Nitrogen monoxide c)Xenon hexafluoride d)Tetrarsenic decoxide e) SiO 2 f) N 2 S 5 g) P 4 O 10 h) BrCl 6

43. Naming Acids ▪ Acids produce Hydrogen ions (H 1+ ) ions in solution. ▪ Two common types of acids: Binary Acids & Oxyacids – Binary Acids - An acid that contains hydrogen and one other element – Oxyacids – acid that contains hydrogen and an oxyanion. ▪ Oxyanions – polyatomic ion containing one or more oxygens.

44. Naming Binary Acids Rules 1.The first word a)Begins with the prefix Hydro- to name the hydrogen part of the acid compound. b)Rest of the word consists of the root of the second element and the suffix –ic. 2.The second word is acid. Example: HCl - Hydrochloric acid HF - Hydrofluoric acid

45. Naming Oxyacids Rules 1.Name of the oxyanion with the changes described below. a)Identify the oxyanion present. ▪If the oxyanion ends in –ate, change the ending to –ic. ▪If the oxyanion ends in – ite, change the ending to –ous. 2. The second word is acid. Examples: HNO 3  Oxyanion nitrate (NO 3 ) becomes nitric acid. H 2 SO 3  Oxyanion sulfite (SO 3 ) becomes sulfurous acid. NO 3 – the nitrate ion, becomes nitric. – The second word of the name is always acid, HNO 3 (hydrogen and nitrogen ion) becomes nitric acid.

46. Naming Oxyacids Examples

47. Writing Formulas from molecular names. ▪ The name of a molecular compound reveals its composition. ▪ The flowchart can help you determine the name of a molecular covalent compound.

48. Section 3: Molecular Structures

49. Essential Questions and Vocabulary ▪ What are the basic steps used to draw Lewis structures? ▪ Why does resonance occur, and what are some resonant structures? ▪ Which molecules are exceptions to the octet rule, and why do these exceptions occur? Vocabulary ▪ Ionic bond ▪ Structural formula ▪ Coordinate covalent bond ▪ Resonance

50. Section 3: Main Idea Structural formulas show the relative positions of atoms within a molecule.

51. Structural Formulas ▪ Structural Formulas – use of letter symbols and bonds to show relative positions of atoms.

52. Lewis Structures Gilbert Lewis ▪ Lewis developed the idea of the octet and coined the term Lewis dot structures ▪ He was nominated 35 times for the Nobel Prize in chemistry, but never won.

53. Drawing Lewis Structures - Guidelines 1.Predict the location of each atom a)The atom that has the least attraction for shared electrons will be the central atom b) All other atoms become terminal atoms. Note: Hydrogen is always a terminal atom. 2.Determine the number of electrons available for bonding. a)Total number of valence electrons for all atoms in molecule.

54. Drawing Lewis Structures - Guidelines 3.Determine the number of bonding pairs. – Divide the total number of electrons available for bonding by two. 4.Place the bonding pairs. – Place one bonding pair (single bond) between the central atom and each of the terminal atoms.

55. Drawing Lewis Structures - Guidelines 5.Determine the number of electron pairs remaining. – Subtract the number of pairs used in step 4 from total number of bonding pairs in step 3. 6.Determine whether the central atom satisfies the octet rule. – Central atom should be surrounded by 4 electron pairs. – If not, convert one or two lone pairs on the terminal atoms to form a double bond or triple bond.

56. Lewis Structure for Ammonia (NH 3 )

57. 5. Number of electron pairs remaining. 6. The remaining lone pair must go on the central atom because hydrogen can only have one bond. Lewis Structure for Ammonia (NH 3 ) 4 pairs total − 3 pairs used = 1 pair available

58. Lewis Structure for Carbon Dioxide (CO 2 )

59. 5. Number of electron pairs remaining. 6. Add three lone pairs to each terminal oxygen atom. ▪ Carbon atom does not meet the octet. Subtract a lone pair from each oxygen to form a double bond with oxygen. 8 pairs total − 2 pairs used = 6 pair available Lewis Structure for Carbon Dioxide (CO 2 )

60. Resonance Structures ▪Resonance – a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. ▪This figure shows three correct ways to draw the structure for nitrate ion (NO 3 ) -1.

61. Resonance Structures ▪The molecule behaves as though it has only one structure. ▪The bond lengths are identical to each other and intermediate between single and double covalent bonds.

62. Exceptions to the Octet Rule ▪Suboctet - A few compounds form stable configurations with less than 8 electrons around the atom ▪Coordinate covalent bond - forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

63. Exceptions to the Octet Rule ▪Expanded Octet – Occurs when central atom has more than eight valence electrons – Can occur in period 3 or higher because of the presence of a d- orbital which allows for more than four covalent bonds.

64. Lewis Structures Practice & Solutions ▪ NH 3 ▪ CH 4 ▪ H 2 O 2 ▪ PCl 3

65. Section 4: Molecular Shapes

66. Essential Questions & Vocabulary ▪ What is VSEPR bonding theory? ▪ How can you use the VSEPR model to predict the shape of, and the bond angles in a molecule? ▪ What is hybridization? ▪ Atomic Orbitals ▪ VSEPR model ▪ Hybridization Vocabulary

67. Section 4: Main Idea The VSEPR model is used to determine molecular shape.

68. VSEPR Theory ▪ Valence-Shell-Electron-Pair-Repulsion theory ▪ This theory helps us understand the 3D structure of molecules and their properties. ▪ Bonding and unshared pairs of valence electrons become very important to uswithin VSEPR theory! ▪ The shapes of molecules are determined because electron pairs want to be far apart from each other (repulsion).

69. AXE Notation A represents the central atom X represents the bonding atoms E represents the lone pairs on the central atom AXE – Method to represent compounds

70. Linear Draw or build CO 2 (AX 2 ) Meaning 1 central atom, 2 bonded atoms It has a linear shape No lone pairs A bond angle of 180 o Bonding pairs are far apart from each other

71. Trigonal Planar ▪ Draw or build BF 3 ▪ This has a trigonal planar ▪ AX 3 – Meaning 1 central, 3 bonded ▪ Bond angle: 120 o ▪ Bonds pointing to the corners of a triangle

72. Tetrahedral ▪ Draw or build CH 4 ▪ This has a tetrahedral ▪ AX 4 – Meaning 1 central, 4 bonded ▪ Bond angle: 109.5 o

73. Use VSEPR theory to predict the shape of: 1. carbon tetrachloride, CCl 4 2. carbon tetrabromide, CBr 4 3. dichloromethane, CH 2 Cl 2 VSEPR Practice

74. Bent Draw or build H 2 O This has a bent shape AX 2 E 2 Meaning 1 central, 2 bonded, 2 lone pairs Bond angle: 105 o Similar angles to the tetrahedral bond angles

75. Trigonal pyramidal Draw or build NH 3 This has a trigonal pyramidal AX 3 E Meaning 1 central, 3 bonded, 1 lone pair Bond angle: 107 o Similar angles to the tetrahedral bond angles

76. VSEPR Theory – Practice & Solutions ▪ BeCl 2 ▪ OF 2 ▪ AlCl 3 ▪ PCl 3 ▪ CF 4 ▪ Linear ▪ Bent ▪ Trigonal Planar ▪ Trigonal pyramidal ▪ Tetrahedral

77. Hybridization ▪Hybridization – a process in which atomic orbitals mix and form new, identical hybrid orbitals. – Carbon often undergoes hybridization, which forms an sp 3 orbital formed from one s orbital and three p orbitals. – Lone pairs also occupy hybrid orbitals. – Single, double, and triple bonds occupy only one hybrid orbital (CO 2 with two double bonds forms an sp hybrid orbital).

78. Section 5: Electronegativity and Polarity

79. Essential Questions & Vocabulary ▪ How is electronegativity used to determine bond type? ▪ How do polar and nonpolar covalent bonds and polar and nonpolar molecules compare? ▪ What are the characteristics of covalently bonded compounds? ▪ Electronegativity ▪ Polar covalent bond Vocabulary

80. Section 5: Main Idea A chemical bond’s character is related to each atom’s attraction for the electrons in the bond.

81. Electronegativity ▪ Electronegativity – the ability of an atom to attract electrons in a chemical bond.

82. Electronegativity and Bond Character ▪ Bond character can be determined based on the difference in electronegativity.

84. Polar Covalent Bond ▪ Polar Covalent Bond – bond with unequal sharing of electrons ▪Bonding is often not clearly ionic or covalent.

85. Polar Covalent Bonds ▪Polar covalent bonds form when atoms pull on electrons in a molecule unequally. ▪Dipole – partial charges at the ends of the bond resulting from electrons spending more time around one atom than another

86. Polarity Practice & Solutions ▪ N-H ▪ F-F ▪ Ca- Cl ▪ Al- Cl ▪ 0.9 slightly polar ▪ 0 Nonpolar ▪ 2.0 ionic ▪ 1.5 very polar

87. Molecular Polarity ▪Covalently bonded molecules can be either polar or non- polar depending on the nature of the bonds and molecular geometry. – Non-polar molecules are not attracted by an electric field. – Polar molecules align with an electric field.

88. Polarity of Water Molecule Water has two polar covalent bonds. Water is polar molecule because its bent shape is asymmetrical.

89. Polar and Non-Polar Molecules ▪Molecules with nonpolar bonds are nonpolar. ▪Molecules with polar bonds can be both polar and non-polar depending on their geometry. ▪Note: If bonds are polar, asymmetrical molecules are polar and symmetrical molecules are nonpolar. – If the molecule is symmetrical, then the molecule is non-polar. – If the molecule is asymmetrical, then the molecule is polar. ▪Note: If bonds are polar, asymmetrical molecules are polar and symmetrical molecules are nonpolar.

90. Properties of molecular compounds ▪ Composed of nonmetals ▪ Can be a solid, liquid, or gas at room temperature ▪ Low melting point and boiling points (compared to ionic compounds) ▪ Poor to non-conductors of heat and electricity

92. Sodium chloride vs sugar NaCl ▪ Sodium is a metal ▪ Chloride is a nonmetal ▪ Melting point ~800 o C C 12 H 22 O 11 ▪ C, H, and O are nonmetals ▪ Melting point ~185 o C (Covalent Molecule) (Ionic Compound)

93. Properties of Covalent Bonds: Intermolecular Forces ▪Covalent bonds between atoms are strong, but attraction forces between molecules are weak. – Van der Waals Forces – the weak interaction forces between molecules ▪Dipole-Dipole Force – the force between two oppositely charged ends of two polar molecules – Hydrogen bond - an especially strong dipole-dipole force between a hydrogen end of one dipole and a fluorine, oxygen, or nitrogen atom on another dipole.

94. Hydrogen Bonding