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CHAPTER 12 – CHEMICAL BONDING

CHAPTER 12 – CHEMICAL BONDING. CHEMICAL BOND – A force that holds two or more atoms together as a unit Individual atoms will naturally bond together to achieve a lower energy state (to be more stable). 3B-1 (of 42). TYPES OF BONDS. 1) METAL ATOMS AND NONMETAL ATOMS.

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CHAPTER 12 – CHEMICAL BONDING

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  1. CHAPTER 12 – CHEMICAL BONDING CHEMICAL BOND – A force that holds two or more atoms together as a unit Individual atoms will naturally bond together to achieve a lower energy state (to be more stable) 3B-1 (of 42)

  2. TYPES OF BONDS 1) METAL ATOMS AND NONMETAL ATOMS Metal atoms easily lose electrons forming positive ions, and nonmetal atoms easily gain electrons forming negative ions IONIC BOND – The electrostatic attraction between positive and negative ions Ionic bonding forms giant crystalline networks containing billions of positive and negative ions that are strongly attracted together Ionic bonding exists between metal and nonmetal ions 3B-2

  3. Elemental Iron Elemental Oxygen Rust + Fe atoms O atoms (molecules) Fe ions and O ions 3B-3

  4. 2) NONMETAL ATOMS Nonmetal atoms attract each other’s valence electrons, and share the valence electrons between pairs of atoms Covalent Bond – The electrostatic attraction of shared electrons to the nuclei of bonding nonmetal atoms Covalent bonding forms individual units called molecules, and while the atoms that covalently bond together strongly attract each other, the molecules that are created weakly attracted each other Covalent bonding exists between nonmetal atoms 3B-4

  5. Elemental Chlorine Elemental Carbon Carbon Tetrachloride + C atoms Cl atoms (molecules) CCl4 molecules 3B-5

  6. Picture NONPOLAR COVALENT BOND – A bond in which 2 atoms are sharing electrons equally POLAR COVALENT BOND – A bond in which 2 atoms are sharing electrons unequally IONIC BOND – A bond in which two atoms have transferred electrons 3B-6

  7. ELECTRONEGATIVITY – The attraction of an atom for shared electrons Atom with the highest EN? Atom with the lowest EN? The difference in the EN’s of 2 atoms tells the type of bond they make EN Difference Bond 0.0 0.1 to 1.6 1.7 to 3.3 Nonpolar Covalent Polar Covalent Ionic 3B-7

  8. N-N Bond EN for N = 3.0 3.0 – 3.0 = 0.0  Nonpolar Covalent Bond 3B-8

  9. C-O Bond EN for C = 2.5, O = 3.5 3.5 – 2.5 = 1.0  Polar Covalent Bond H-S Bond EN for H = 2.1, S = 2.5 2.5 – 2.1 = 0.4  Polar Covalent Bond Polar covalent bonds have partially positive and a partially negative ends C –– O d+ d- C –– O H –– S d+ d- H –– S DIPOLE MOMENT – The amount of separation of the positive and negative charges in a bond DIPOLE MOMENT ARROW – Shows the direction of the dipole moment, pointing toward the negative end of the bond 3B-9

  10. Na-Cl Bond EN for Na = 0.9, Cl = 3.0 3.0 – 0.9 = 2.1  Ionic Bond Na+ Cl- 3B-10

  11. BONDING IN IONIC COMPOUNDS Atoms form ions to obtain a stable, octet electron arrangement Sodium chloride . . . Cl : . . Na - + A sodium chloride crystal is a symmetrical array of sodium and chloride ions in a 1:1 ratio EMPIRICAL FORMULA – The simplest whole number ratio of atoms of different elements in a compound Empirical Formula: NaCl 3B-11

  12. Magnesium fluoride . . .F : . . . . .F : . . - 2 + - Mg Empirical Formula: MgF2 3B-12

  13. Potassium nitride Empirical Formula: 3B-13

  14. SIZES OF ATOMS AND IONS Positive ions are smaller than their neutral atoms and negative ions are bigger than their neutral atoms Na atom Cl atom Na+ ion Cl- ion 3B-14

  15. Sizes of atoms or ions are determined by 1) The more energy levels an atom or ion has the larger it will be 2) With the same number of energy levels, the more protons an atom or ion has the smaller it will be Li F Na Cl Li+ F- Na+ Cl- Electrons Energy Levels Protons Big to Small 3 2 3 3rd 9 2 9 4th 11 3 11 1st 17 3 17 2nd 2 1 4th 10 2 9 2nd 10 2 11 3rd 18 3 1st ISOELECTRONIC – Ions or atoms with the same number of electrons 3B-15

  16. BONDING IN COVALENT MOLECULES Atoms share electrons to obtain a stable, octet (or duet) arrangements Water (H2O) . . . O : . H H . . H – O : H ← LONE PAIR ← BONDING PAIR LEWIS STRUCTURE – A diagram using electron dot notation to show how the valence electrons are arranged among bonded atoms 3B-16

  17. To draw a proper Lewis Structure: 1 – Add up the valence e-s for all of the atoms in the molecule or ion 2 – Draw a skeletal structure by using pairs of electrons to make bonds 3 – Complete octets (or duets for H) for all atoms, outer atoms first, using the remaining valence e-s 4 – If octets are not produced, make the atoms that have octets share more e- pairs with atoms that do not have octets 3B-17

  18. Sulfur dichloride, SCl2 6 + 7 + 7 = 20 valence e-s Cl S Cl 3B-18

  19. Phosphorus tribromide, PBr3 3B-19

  20. Ammonia, NH3 5 + 1 + 1 + 1 = 8 valence e-s H N H H 3B-20

  21. Methane, CH4 3B-21

  22. Fluorine, F2 7 + 7 = 14 valence e-s F F SINGLE BOND – One shared pair of e-s between two atoms 3B-22

  23. Oxygen, O2 6 + 6 = 12 valence e-s O O DOUBLE BOND – Two shared pairs of e-s between two atoms 3B-23

  24. Nitrogen, N2 5 + 5 = 10 valence e-s N N TRIPLE BOND – Three shared pairs of e-s between two atoms 3B-24

  25. Hydrogen cyanide, HCN Carbon disulfide, CS2 1 + 4 + 5 = 10 valence e-s H C N 3B-25

  26. Sulfate, SO42- Ammonium, NH4+ 6 + 4(6) + 2 = 32 valence e-s 5 + 4(1) - 1 = 8 valence e-s 2- O O S O O H H N H H + 3B-26

  27. Ozone, O3 6 + 6 + 6 = 18 valence e-s or O O O O O O O O O O O O ↔ RESONANCE – When more than one Lewis structure can be drawn for a molecule or ion RESONANCE STRUCTURES – The Lewis structures that can be drawn for the molecule or ion The real ozone molecule is a average of its resonance structures 3B-27

  28. O O O O O O ↔ O O O 2 “1½” bonds 3B-28

  29. MOLECULAR SHAPE VSEPR THEORY (Valence Shell Electron Pair Repulsion) – All atoms and lone pairs attached to a central atom will spread out as far as possible to minimize repulsion A Lewis structure must be drawn to use the VSEPR Theory 3B-29

  30. H C H H H H H C H H STERIC NUMBER (SN) – The sum of the bonded atoms and lone pairs on a central atom The steric number of carbon is 4 (SN = 4): 4 bonded atoms and no lone pairs Tetrahedral Bond angle is 109.5° 3B-30

  31. N H H H The steric number of nitrogen is 4 (SN = 4): 3 bonded atoms and 1 lone pairs H N H H Trigonal Pyramidal Bond angle is 108° 3B-31

  32. O H H . . H – O : H The steric number of oxygen is 4 (SN = 4): 2 bonded atoms and 2 lone pairs Bent Bond angle is 105° 3B-32

  33. O C H H Formaldehyde, H2CO 1 + 1 + 4 + 6 = 12 valence e-s H H C O The steric number of carbon is 3 (SN = 3): 2 bonded atoms and 1 lone pairs Trigonal Planar Bond angle is 120° 3B-33

  34. S Si S SiS2 4 + 6 + 6 = 16 valence e-s S Si S The steric number of silicon is 2 (SN = 2): 2 bonded atoms and 0 lone pairs Linear Bond angle is 180° 3B-34

  35. SN 4 4 4 3 3 2 Atoms 4 3 2 3 2 2 Lone Pairs 0 1 2 0 1 0 Shape Tetrahedral Trigonal Pyramidal Bent (109.5°) Trigonal Planar Bent (120°) Linear 3B-35

  36. MOLECULAR POLARITY A BOND is polar if it has a positive end and a negative end A MOLECULE is polar if it has a positive end and a negative end • To determine if a molecule is polar or nonpolar: • 1) Draw the correct Lewis structure • Draw its correct shape • Use EN’s to determine if the BONDS in the molecule are polar or nonpolar • For the polar bonds, label the positive and negative ends with δ+ and δ- • If a line can be drawn separating all δ+’s from all δ-’s, the molecule is polar, it not its nonpolar 3B-36

  37. O H H . . H – O : H EN’s: O = 3.5, H = 2.1 3.5 – 2.1 = 1.4  the O-H BONDS are polar All of the δ+’s can be separated from all of the δ-’s,  the H2O MOLECULE is polar δ- δ- δ+ δ+ 3B-37

  38. H N H H EN’s: N = 3.0, H = 2.1 3.0 – 2.1 = 0.9  the N-H BONDS are polar All of the δ+’s can be separated from all of the δ-’s,  the NH3MOLECULE is polar δ- δ- δ- N H H δ+ δ+ H δ+ 3B-38

  39. F C F F F Carbon tetrafluoride, CF4 4 + 4(7) = 32 valence e-s F F C F F δ- EN’s: C = 2.5, F = 4.0 4.0 – 2.5 = 1.5  the C-F BONDS are polar All of the δ+’s cannot be separated from all of the δ-’s,  the CF4MOLECULE is nonpolar δ+ δ+ δ+ δ- δ- δ+ δ- 3B-39

  40. REVIEW FOR TEST Electromagnetic Radiation, Photons Ground State, Excited State Orbital Energy Levels Sublevels Orbital Notation Electron Configuration Notation Electron Dot Notation Valence Electrons Octet Electron Pair 2B-40

  41. REVIEW FOR TEST Periodic Trends in Metal, Nonmetal Activity Atomic Radii Ionization Energy Electron Affinity Ionic Bonds, Covalent Bonds Electronegativity Bond Polarity from Electronegativities Ion Sizes 2B-41

  42. REVIEW FOR TEST Lewis Structures for Ionic Compounds Covalent Compounds Resonance Molecular Shapes Molecular Polarity 2B-42

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