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Ch 9 Molecular Bonding WARM-UP: 1. Why do atoms bond? 2. How do they bond? 3. What are the deciding factors in the type of bond? 4. How does the bond type affect structure and shape?
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Types of Bonds inside a compound Metallic Bond: a sea of electrons move freely between the __cations___ Ionic Bond: electrons are _transferred_ b/c of the large difference in electronegativities Covalent Bond: electrons are shared ; the closer the electronegativity values are to each other the more equal the sharing Polar Covalent Bond: electrons are shared unequally Non-Polar Covalent Bond: electrons are shared EQUALLY!
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Characteristics of each bond type Metallic crystalline solid (except Hg), high melt & boil points, good conductors of heat and electricity Ionic crystalline solids, high melt & boil points, conductors when molten or aqueous Molecular (Covalent and Polar Covalent) liquids, gases, waxy solids low mp & bp, do NOT conduct elec. forms molecules
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Determining Bond Type Predict based on Location in Periodic Table Metallic = M+M, Ionic = M+NM, Covalent = NM+NM Difference in Electronegativities Non Polar ≤ 0.4 > Polar Covalent > 1.7 ≤ Ionic Covalent Ex: prediction electronegativities Li I M+NM = I Li=.98 I = 2.66 2.66 -.98 = 1.71 CH 4 NH 3
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Polar – Inside? Outside? Both? Polar Bond – atom to atom inside a molecule You determine by difference in electronegativity Polar Molecule – the outside of the molecule has slightly + and – ends You determine by: 1.Has a polar bond inside O=3.44 2.Is asymmetrical in shape H=2.2 1.2 = polar covalent bond BENT shape allows for + & - poles
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Lewis Structures Ionic – show the transferred electron(s) and resulting charges Ex: magnesium bromide MgBr 2 Br Mg Br 2+
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Lewis Structures for Molecules Show the shared pairs w/ a dash & the lone pears w/ dots 1. least electroneg atom (furtherest left on P. Table) is the CENTRAL atom 2. H is always on the END as are terminal HALOGENS 3. Put DOTS around each atom to represent the VALENCE ELECTRONS 4. Use 1 pair of e- to connect each of the atoms any leftovers go to be LONE PAIRS 5. Make sure outer atoms have 8 ELECTRONS (except H can have 2 ) 6. if central atom still needs 4 pair (8 e-) add DOUBLE or TRIPLE bonds
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Now practice on your worksheet You can not do a formula for sulfur and oxygen because they are both negative. So just put SO in the blank Please STOP HERE TODAY with the power point
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Go back one slide
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STOP and go back Left arrow twice please
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The other arrow(left)
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Naming Covalent molecules Use your pre-fixes
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Lewis Structure Practice lithium fluoridecalcium chloride methane (CH 4 )ammonia (NH 3 ) ethane (C 2 H 6 )ethene (C 2 H 4 ) nitrogen trifluoridecarbon disulfide
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Polyatomic ions Phosphate ion Coordinate covalent bond = one of the atoms donates both electrons to the shared pair perchlorate ion ammonium ion nitrite ion
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Resonance Structures Nitrate ion Ozone
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Molecular Shapes VSEPR - Valence shell electron pair repulsion Look at the central atom 2 bonds and 0 lone pairs = LINEAR 180 o 3 bonds and 0 lone pairs = Trigonal PLANAR 120 o 4 bonds = TETRAHEDRAL 109.5 o
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More shapes with VSEPR 1. 2 lone pairs are the most repulsive 2. 1 lone pair and 1 shared pair (bond) next most repulsive 3. 2 shared pairs (bonds) least repulsive 3 bonds & 1 unshared (lone) pair PYRAMIDAL 107 o 2 bonds & 2 unshared (lone) pairs BENT 105 o
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Exceptions to the Octet Rule 1. Odd # of valence electrons (end up w/ x pairs and 1 leftover e-) the one leftover unpaired e- causes the cmpd to be paramagnetic 2.Boron cmpds are sometimes deficient in their octets and will be happy with 6 instead of 8 3. expanded octets: a few cmpds exceed the 8 e- P family sometimes 10 e- S can do 12 e- Xe can do 12 e-
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Polar – Inside? Outside? Both? Polar Bond – atom to atom inside a molecule You determine by difference in electronegativity Polar Molecule – the outside of the molecule has slightly + and – ends You determine by: 1.Has a polar bond inside O=3.44 2.Is asymmetrical in shape H=2.2 1.2 = polar covalent bond BENT shape allows for + & - poles
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Polarity Practice CH 4 NH 3 CaO H 2 O BF 3
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Intramolecular Bonds vs Intermolecular Forces Water molecules have a polar covalent bond between each H and O. The O is more electronegative so it keeps the e- more of the time, resulting in a slight – charge on O & a slight + for the H s. The + of one H 2 O is attracted to the – of another H 2 O
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Bond Strength Covalent Bonds have a balance between attractive and repulsive forces Bond Length = distance from center of 1 nucleus to center of the other depends on : 1. atom size 2. how many e- pairs are shared shorter bond length = stronger bond _ _ ++ ++
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Energy and Intramolecular Bonds Energy is given off when a bond forms ( a neg. number) Energy must be added to break a bond and is called bond dissociation energy ( a + number) Ex: 2H 2 + O 2 2H 2 O If nrg added = nrg released no nrg change overall If add is greater endothermic If release is greater exothermic
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Intermolecular Forces Attraction of one molecule for another molecule weaker than a bond 1. nonpolar – nonpolar (weak) dispersion forces 2. polar – polar (stronger) dipole - dipole 3. Hydrogen bond = H of one molecule attracted to F, O, or N of another molecule
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Molecular Orbitals Atomic orbitals describe e- location in 1 atom. orbital could be s,p,d,f, or some hybrid of these Molecular orbitals are formed when atomic orbitals overlap as atoms combine to make molecules. Ex: + overlapping s = σ H H H 2 ss
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Sigma bond (σ ) formed when 2 atomic orbitals combine to make a molecular orbital that is symmetrical along the axis connecting the 2 atomic nuclei (end to end) + overlapping s = σ overlapping p ends = σ ss
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Pi bond ( π ) Bonding e- are most likely found in sausage shaped regions above & below the nuclei They form between p – p orbitals that overlap side to side instead of end to end Ex: diatomic nitrogen N = 1s 2 2s 2 2p 3
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1 “s” & 1 “p” 2 “sp” HYBRID Atomic ORBITALS New atomic orbital made by combining s and p The hybrid looks more like a “p” but has both “s” and “p” character.
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HYBRID ORBITALS In a Carbon cmpd 1s 2s 2p
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Tetrahedral shape shows 4 equal bonds-- thus s and p orbitals must be hybridized into 4 equal orbitals.
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Shape determines the hybridization Pyramid, or bent
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