chemical bonding, any of the interactions that account for the association of atoms into molecules, ions, crystals, and other stable species that make up the familiar substances of the everyday world.
1. BONDING 1
2. Chemical Bond A Quick Review…. • A bond results from the attraction of nuclei for electrons – All atoms are trying to achieve a stable octet • IN OTHER WORDS – the protons (+) in one nucleus are attracted to the electrons (-) of another atom • This is Electronegativity !! 2
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4. Three Major Types of Bonding • Ionic Bonding – forms ionic compounds – transfer of valence e- • Metallic Bonding • Covalent Bonding – forms molecules – sharing of valence e- – This is our focus this chapter 4
5. Ionic Bonding • Always formed between metal cations and non-metals anions • The oppositely charged ions stick like magnets + - [METALS ] [NON-METALS ] Lost e- Gained e- 5
6. Metallic Bonding • Always formed between 2 metals (pure metals) – Solid gold, silver, lead, etc… 6
7. Covalent Bonding molecules • Pairs of e- are shared between 2 non- metal atoms to acquire the electron configuration of a noble gas. 7
8. Covalent Bonding • Occurs between nonmetal atoms which need to gain electrons to get a stable octet of electrons or a filled outer shell. no nm et a ls
9. Drawing molecules (covalent) using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner electrons) • dots represent valence electrons • The ones place of the group number indicates the number of valence electrons on an atom. • Draw a valence electron on each side (top, right, bottom, left) before pairing them. 9
10. Always remember atoms are trying to complete their valence shell! “2 will do but 8 is great!” The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four 10
11. Draw Lewis Dot Structures You may represent valence electrons from different atoms with the following symbols x, , x H or H or H 11
12. Covalent bonding • The atoms form a covalent bond by sharing their valence electrons to get a stable octet of electrons.(filled valence shell of 8 electrons) • Electron-Dot Diagrams of the atoms are combined to show the covalent bonds • Covalently bonded atoms form MOLECULES
13. Methane CH4 • This is the finished Lewis dot structure • Every atom has a filled valence shell How did we get here? OR 13
14. General Rules for Drawing Lewis Structures • All valence electrons of the atoms in Lewis structures must be shown. • Generally each atom needs eight electrons in its valence shell (except Hydrogen needs only two electrons and Boron needs only 6). • Multiple bonds (double and triple bonds) can be formed by C, N, O, P, and S. • Central atoms have the most unpaired electrons. • Terminal atoms have the fewest unpaired electrons. 14
15. • When carbon is one of you atoms, it will always be in the center • Sometimes you only have two atoms, so there is no central atom Cl2 HBr H2 O2 N2 HCl • We will use a method called ANS (Available, Needed, Shared) to help us draw our Lewis dot structures for molecules 15
16. EXAMPLE 1: Write the Lewis structure for H2O where oxygen is the central atom. Step 1: Determine the total number of electrons available for bonding. Because only valence electrons are involved in bonding we need to determine the total number of valence electrons. AVAILABLE valence electrons: Electrons available 2H Group 1 2(1) = 2 O Group 6 6 8 There are 8 electrons available for bonding. Step 2: Determine the number of electrons needed by each atom to fill its valence shell. NEEDED valence electrons Electrons needed 2H each H needs 2 2(2) = 4 O needs 8 8 12 There are 12 electrons needed. 16
17. Step 3: More electrons are needed then there are available. Atoms therefore make bonds by sharing electrons. Two electrons are shared per bond. SHARED (two electrons per bond) # of bonds = (# of electrons needed – # of electrons available) = (N-A) = (12 – 8) = 2 bonds. 2 2 2 Draw Oxygen as the central atom. Draw the Hydrogen atoms on either side of the oxygen atom. Draw the 2 bonds that can be formed to connect the atoms. OR Step 4: Use remaining available electrons to fill valence shells for each atom. All atoms need 8 electrons to fill their valence shell (except hydrogen needs only 2 electrons to fill its valence shell, and boron only needs 6). For H2O there are 2 bonds, and 2 electrons per bond. # available electrons remaining = # electrons available – # electrons shared = A-S = 8 – 2(2) = 4 extra e-s 17
18. Sometimes multiple bonds must be formed to get the numbers of electrons to work out • DOUBLE bond – atoms that share two e- pairs (4 e-) O O • TRIPLE bond – atoms that share three e- pairs (6 e-) N N 18
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20. Step 3: SHARED (two electrons per bond) # of bonds = (N – A) = (20 – 12) = 4 bonds. 2 2 Draw carbon as the central atom (Hint: carbon is always the center when it is present!). Draw the Hydrogen atoms and oxygen atom around the carbon atom. Draw 2 bonds of the 4 bonds that can be formed to connect the H atoms. Draw the remaining 2 bonds to connect the O atom (oxygen can form double bonds) Step 4: Use remaining available electrons to fill valence shell for each atom. # electrons remaining = Available – Shared = A – S = 12 – 4(2) = 4 extra e-s 20
21. Let’s Practice A=1x2=2 N=2x2=4 S = 4 - 2= 2 ÷ 2 = 1 bond Remaining = A – S = 2 – 2 = 0 21
22. Let’s Practice A = C 4x1 = 4 H 1x4 = 4 4 + 4 = 8 N = C 8x1 = 8 H 2x4 = 8 8 + 8 = 16 S = (A-N)16 – 8 = 8 ÷2 = 4 bonds Remaining = A-S = 8 – 8 = 0 22
23. Let’s Practice A = N 5x1 = 5 H 1x3 = 3 = 8 N = N 8x1 = 8 H 2x3 = 6 = 14 S = 14-8 = 6 ÷2 = 3 bonds Remaining = (A-S) 8 – 6 = 2 23
24. Let’s Practice A = C 4x1 = 4 O 6x2 = 12 = 16 N = C 8x1 = 8 O 8x2 = 16 = 24 S = 24-16 = 8 ÷ 2 = 4 bonds Remaining = (A-S) 16 – 8 = 8 not bonding DRAW – carbon is the central atom 24
25. Let’s Practice BCl3 boron only needs 6 valence electrons, it is an exception. A = B 3 x 1 = 3 Cl 7 x 3 = 21 = 24 N = B(6) x 1 = 6 Cl 8 x 3 = 24 = 30 S = 30-24 = 6 ÷ 2 = 3 bonds Remaining = 24 – 6 = 18 e- not bonding 25
26. Naming Molecular Compounds (Covalent) Type III Nonmetal + nonmetal
27. The Covalent Bond Sharing of electrons
28. Properties of Molecular or Covalent Compounds • Made from 2 or more nonmetals • Consist of molecules not ions
29. Molecular Formulas Show the kinds and numbers of atoms present in a molecule of a compound. Molecular Formula = H2O
30. Structural formula H N H H Molecular formula NH3
32. Rules for Naming Molecular compounds • The most “metallic” nonmetal element is written first (the one that is furthest left) • The most nonmetallic of the two nonmetals is written last in the formula • NO2 not O2N • All binary molecular compounds end in -ide
33. Molecular compounds • Ionic compounds use charges to determine the chemical formula • The molecular compound‘s name tells you the number of each element in the chemical formula. • Uses prefixes to tell you the quantity of each element. • You need to memorize the prefixes !
35. More Molecular Compound Rules • If there is only one of the first element do not put (prefix) mono • Example: carbon monoxide (not monocarbon monoxide) • If the nonmetal starts with a vowel, drop the vowel ending from all prefixes except di and tri • monoxide not monooxide • tetroxide not tetraoxide
59. Bond Types 3 Possible Bond Types: • Ionic • Non-Polar Covalent • Polar Covalent 59
60. Use Electronegativity Values to Determine Bond Types • Ionic bonds – Electronegativity (EN) difference > 2.0 • Polar Covalent bonds – EN difference is between .21 and 1.99 • Non-Polar Covalent bonds – EN difference is < .20 – Electrons shared evenly in the bond 60
61. Ionic Character “Ionic Character” refers to a bond’s polarity –In a polar covalent bond, • the closer the EN difference is to 2.0, the more POLAR its character • The closer the EN difference is to .20, the more NON-POLAR its character 61
62. Place these molecules in order of increasing bond polarity using the electronegativity values on your periodic table • HCl 3 EN difference = 0.9 • CH4 2 EN difference = 0.4 • CO2 4 EN difference = 1.0 a.k.a. • NH3 3 EN difference = 0.9 “ionic character” • N2 1 EN difference = 0 • HF 5 EN difference = 1.9 62
63. Polar vs. Nonpolar MOLECULES • Sometimes the bonds within a molecule are polar and yet the molecule itself is non-polar 63
64. Nonpolar Molecules • Molecule is Equal on all sides – Symmetrical shape of molecule (atoms surrounding central atom are the same on all sides) H Draw Lewis dot first and see if equal on all sides H C H H 64
65. Polar Molecules • Molecule is Not Equal on all sides – Not a symmetrical shape of molecule (atoms surrounding central atom are not the same on all sides) Cl H C H H 65
66. Polar Molecule H Cl - Unequal Sharing of Electrons 66
67. Non-Polar Molecule Cl Cl Equal Sharing of Electrons 67
68. Polar Molecule H Cl B H Not symmetrical 68
69. Non-Polar Molecule H H B HSymmetrical 69
70. Water is a POLAR molecule ANY time there are unshared pairs of electrons on the central atom, the molecule is POLAR H H O 70
71. Making sense of the polar non-polar thing BONDS MOLECULES Non-polar Polar Non-polar Polar EN difference EN difference Symmetrical Asymmetrical 0 - .2 .21 – 1.99 OR Unshared e-s on Central Atom 71
72. 5 Shapes of Molecules you must know! (memorize) 72
73. Copy this slide • VSEPR – Valence Shell Electron Pair Repulsion Theory – Covalent molecules assume geometry that minimizes repulsion among electrons in valence shell of atom – Shape of a molecule can be predicted from its Lewis Structure 73
74. 1. Linear (straight line) Ball and stick model OR Molecule geometry X A X OR A X Shared Pairs = 2 Unshared Pairs = 0 74
75. 2. Trigonal Planar Ball and stick Molecule geometry X A X X Shared Pairs = 3 Unshared Pairs = 0 75
76. 3.Tetrahedral Ball and stick Molecule geometry model Shared Pairs = 4 Unshared Pairs = 0 76
77. 4. Bent Ball and stick .. Lewis Diagram A X X Shared Pairs = 2 Unshared Pairs = 1 or 2 77
79. • I can describe the 3 intermolecular forces of covalent compounds and explain the effects of each force. 79
80. Intramolecular attractions • Attractions within or inside molecules, also known as bonds. – Ionic – Covalent Roads within a state – metallic 80
81. Intermolecular attractions • Attractions between molecules – Hydrogen “bonding” • Strong attraction between special polar molecules (F, O, N, P) – Dipole-Dipole • Result of polar covalent Bonds – Induced Dipole (Dispersion Forces) • Result of non-polar covalent bonds 81
82. More on intermolecular forces Hydrogen “Bonding” • STRONG intermolecular force - - – Like magnets • Occurs ONLY + + + + between H of one - molecule and N, O, F of another molecule + + Hydrogen Hydrogen bonding “bond” 1 min 82
83. Why does Hydrogen “bonding” occur? • Nitrogen, Oxygen and Fluorine – are small atoms with strong nuclear charges • powerful atoms – Have very high electronegativities, these atoms hog the electrons in a bond – Create very POLAR molecules 83
84. Dipole-Dipole Interactions – WEAK intermolecular force – Bonds have high EN differences forming polar covalent molecules, but not as high as those that result in hydrogen bonding. .21– Partial negative and partial positive charges slightly attracted to each other. – Only occur between polar covalent molecules 84
85. Dipole-Dipole Interactions 85
86. Induced Dipole Attractions – VERY WEAK intermolecular force – Bonds have low EN differences EN < .20 – Temporary partial negative or positive charge results from a nearby polar covalent molecule. – Only occur between NON-POLAR & POLAR molecules Induced dipole video 86 30 sec
88. Intermolecular Forces affect chemical properties • For example, strong intermolecular forces cause high Boiling Point – Water has a high boiling point compared to many other liquids 88
89. Which substance has the highest boiling point? • HF • NH3 • CO2 • WHY? 89
90. Which substance has the highest boiling point? • HF The H-F bond has the highest • NH3 electronegativity difference SO • CO2 HF has the most polar bond • WHY? resulting in the strongest H bonding (and therefore needs the most energy to overcome the intermolecular forces and boil) 90
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