chapter 6 chemical bonding. types of chemical bonding chemical bond – mutual electrical attraction...
TRANSCRIPT
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Chapter 6
Chemical Bonding
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Types of Chemical Bonding
Chemical bond– Mutual electrical attraction between nuclei
and valence electrons of different atoms that binds the atoms together
– Creates more stable compoundsIonic bonds
– Electrical attraction between large numbers of cations and anions
Covalent bonds– Sharing electron pairs between two atoms
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Types of Chemical Bonds
Determine if bond is ionic or covalent by difference in electronegativities– Electronegativity difference 1.7 or less is
covalentBonding between diatomic molecule is
covalentNon-polar covalent - e- shared equally (≤0.3)Polar covalent - unequal attraction for e-
(0.3 - 1.7)– Electronegativity of 1.7 to 3.3 is ionic
Polar - uneven distribution of charge
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Predicting Bonds
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Problems
Use electronegativity differences to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?
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Problems
Use electronegativity differences to classify bonding between chlorine, Cl, and the following elements: calcium, Ca; oxygen, O; and bromine, Br. Indicate the more negative atom in each pair.
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Covalent Bonding
Molecule– Neutral group of atoms that are covalently
bonded together– H2O, sugar,O2
Chemical formula– Shows relative number of atoms in a
compoundDiatomic Molecule
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Formation of a Covalent Bond
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Energy and Stability
Energy released when covalent bond formed
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Covalent Bonding
Bond length– Distance between two bonded atoms at
minimum potential energyAtoms vibrate back and forthDepends on atoms that have
combinedBond energy
– Energy required to break a chemical bond
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Polarity and Bond Strength
The greater the electronegativity difference, the greater the polarity and the stronger the bond
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Problem
Arrange the following bonds in order of increasing bond length, from shortest bond to longest
Bond Bond Energy (kJ/mol)
H-F 569H-I 299
H-Cl 432H-Br 366
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Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level
Atoms with 8 electrons in outer shell are stable– Except:
Hydrogen and Helium - stable with 2 electronsBoron - forms bonds surrounded by 6 electronsBeryllium - forms bonds surrounded by 4 e-
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Electron-Dot Strucutres
Notation in which only valence electrons of element are shown by dots
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Electron-Dot Notation
Na Cl
Kr B
Ne1
2
34
5
67
8
H N
S Ba
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Lewis Structures
Combine 2 electon-dot structures to show shared electrons– Unshared or Lone pairs
– Indicates kind, number, arrangement, and bonds
– Shared electrons represented by dash– F-F
F F
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Lewis Structures
Least electronegative atom is central
– Except hydrogen– Carbon is usually central
Multiple bond represented– Single bond– Double bond– Triple bond
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Lewis Structures1. Determine the total number of valence
electrons in the compound.2. Arrange the atoms’ symbols to show how
they are bonded and show valence electrons as dots
3. Compare the number of valence electrons used in the structure to the number available from step 1.
4. Change to a single dash each pair of dots that represents two shared electrons.
5. Be sure that all atoms, with the exception of hydrogen, follow the octet rule.
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Problems
Draw the Lewis structure of iodomethane, CH3I.
Draw the Lewis structure of ammonia, NH3
Draw the Lewis structure for hydrogen sulfide, H2S
Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde.
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Problems
Draw the Lewis structure for:– Carbon dioxide, CO2
– Hydrogen cyanide, HCN– IBr– CH3Br– C2HCl– SiCl4– F2O
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Polyatomic Ions
A charged group of covalently bonded atoms
Combine with other ions to form ionic compounds
Add/subtract appropriate number of electrons
Place brackets around structureShow the charge of ion outside of
brackets
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Polyatomic Ions
Draw Lewis structures for:– NH4
+
– SO4-
– PO43-
– CO32-
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Ionic Bonding
Composed of cations and anions to make neutral compound– NaCl– Cannot be isolated and examined like
moleculesForm crystal lattice to stabilize
– Forces between like-charged ions and opposite-charged ions
– Na+ surrounded by 6 Cl-– Cl- surrounded by 6 Na+
Lattice Energy
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Ionic vs. Molecular Compounds
Ionic– Stronger bonds– Higher melting and
boiling points– Hard but brittle– Electrical
conductors when dissolved
– May separate when dissolved
Molecular– Weaker bonds– Lower melting and
boiling points
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Metallic Bonding
Chemical bonding the results from attraction between metal atoms and surrounding sea of electrons
Highest energy levels have few electrons
Many vacant orbitals
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Metallic Bonding
Metallic Properties– High electrical and thermal conductivity– Absorb many light frequencies– Shiny– Malleable– Ductile– Heat of vaporization → Bond strength
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VSEPR Theory
Valence-Shell Electron-Pair RepulsionRepulsion between sets of valence-
level electrons surrounding an atom causes sets to be oriented as far apart as possible– Electrons of bonded atoms want to be as
far away from each other as possible
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VSEPR Theory
CO2
– Shared pairs or oriented as far away from each other as possible
– 180 apart– Linear – AB2
BF3– 120 apart– Trigonal-Planar– AB3
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VSEPR Theory
CH4
– 109.5– Tetrahedral– AB4
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VSEPR & Unshared Electrons
Lone pairs occupy space and influence shape of molecule– H2O → AB2E2 → bent– NH3 → AB3E → trigonal-pyramidal
Unshared electrons repel electrons more strongly than shared electrons
Double and triple bonds treated like single bonds
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Problem
Use VSEPR theory to predict the molecular geometry of:– AlCl3– CBr4
– AlBr3
– SF6
– CH2Cl2
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Problem
Use VSEPR theory to predict the molecular geometry of:– CO2
– ClO3-
– SF2
– PCl3
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Hybridization
Mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies
Example → CH4
– 2s and 2p orbitals hybridize to form 4 identical orbitals called sp3 orbitals
Group 15 and 16– Nitrogen - sp3
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Hybrid Orbitals
sp sp2
sp3
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Intermolecular Forces
Forces of attraction between molecules
Boiling point is used to measure force of attraction between particles– Weaker than bonds– Boiling points of ionic compounds higher
than covalent moleculesDipole-Dipole ForcesHydrogen BondingLondon Dispersion Forces
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Dipole-Dipole Forces
Strongest force is between polar molecules
Dipole– Equal but opposite charges that are
separated by short distance– Direction is from positive pole to negative– Represented by arrow pointing toward
negative pole with a crossed tail
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Dipole-Dipole Forces
Forces of attraction between polar molecules are short-range forces
Dipole-Dipole forces cause higher boiling points
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Hydrogen Bonding
Many bonds with hydrogen are highly polar because of large electronegativity difference
Hydrogen therefore has positive charge in many compounds
Hydrogen atoms bonded to highly electronegative atom is attracted to an unshared pair of electrons– Example → H2O
Represented by dotted lines
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Hydrogen Bonding
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Hydrogen Bonding
IceIce Liquid WaterLiquid Water
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London Dispersion Force
Electrons are in continuous motionSlight uneven electron distributionTemporary dipoleIntermolecular attractions resulting
from constant motion of electrons and creation of instantaneous dipole
Only intermolecular force acting on noble gases
Strength increases with number of e-
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Chapter Review
Pg. 209– 3, 6acfg, 9, 16, 19abcfg, 21bde, 22, 24c,
29ab, 33, 45bc, 46bdf, 47cd, 48acdf, 49a