chapter 8 covalent bonding. let’s review what do we already know? –what is a chemical bond?...
TRANSCRIPT
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Chapter 8
Covalent Bonding
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Let’s Review
• What do we already know?– What is a chemical bond?– What is an ionic bond?
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Section 1
The Covalent Bond
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Stability
• Lower energy is more stable
• Noble-Gas electron configuration
• Octet rule
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Covalent Bond
• Atoms in nonionic compounds share electrons
• Covalent bond is the bond that results from sharing valence electrons
• Molecule is formed when two or more atoms bond covalently
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Diatomic Molecules
• Two atom molecules are more stable than one atom
• H2, N2, O2, F2, Cl2, Br2, I2
HH
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Hydrogen
H H
They Pair!!
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Hydrogen
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Oxygen
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Fluorine
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Fluorine
F F
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Single Covalent Bonds
• One pair of valence electrons is shared– Pair may be referred to as “bonding” pair
• Also called sigma bonds– σ– Occurs when the shared pair is centered
between the two atoms
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Bonding Orbital
• Localized region where bonding electrons are most likely found
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Groups and Single Bonds
• Group 17
• Group 16
• Group 15
• Group 14
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Homework (due Tuesday)
• Draw the Lewis structures for the following molecules– PH3
– H2S
– HCl
– CCl4– SiH4
• Challenge– Draw a generic Lewis Structure for a molecule formed
between atoms of group 1 and group 16
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Homework continued
• Draw LDS for – CH4
– Br2
– C6H14 also written as CH3(CH2)4CH3
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Multiple Covalent Bonds
• Bond Order – Refers to the type of bond
• Single Bond– Shares ONE pair of electrons
• Double Bonds– Two pairs of electrons are shared
• Triple Bonds– Three pairs of electrons are shared
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The Pi Bond
• Multiple covalent bonds– Consist of at least one sigma and one pi bond
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Strength of Covalent Bonds
• CB involve attractive and repulsive forces
• Balance of the force is upset the bond can break
• Several factors influence strength of cb
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Bond Length
• Length depends on distance between bonded nuclei
• Bond length is the distance two nuclei at the position of maximum attraction– Determined by:
• Sizes of two bonding atoms• Number of electrons shared
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Bonds and Energy
• Energy changes occur– When bonds are broken
• Energy is released• Need energy put in to break it
– Bond-dissociation energy » is the energy required to break a specific bond» Indicates strength of the bond
– When bonds are formed
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Length and Energy
• Shorter the length the greater the energy
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Energies of Chemical Reactions
• Total energy is determined from energy of bonds broken and formed
• Two types– Endothermic– Exothermic
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Energies of Chemical Reactions
• Endothermic Reaction occurs when a greater amount of energy is required to break existing bonds in the reactants than is released when the new bonds formed.
• Endothermic Reaction– More energy to break a bond than energy
when bond is broken
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Energies of Chemical Reactions
• Exothermic
Energyin
Energyout
Bond
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Energies of Chemical Reactions
• Exothermic reaction occurs when more energy is released during product bond formation than is required to break bonds in reactants.
• Exothermic reaction– More energy is released than required to
break the bonds
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Energies of Chemical Reactions
• Endothermic
Energyout
Energyin
Bond
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Section Two
Naming Molecules
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Binary Molecular Compounds
Example: N2O
1. First element in the formula is always named first, using the entire element name.
• What is the first element?• Nitrogen
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Binary Molecular Compounds
2. The second element in the formula is named using its root and adding the suffix –ide.
1. What is the second element?• Oxygen
2. What will the name be?• Oxide
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Binary Molecular Compounds
3. Prefixes are used to indicate the number of atoms of each element are present in the compound.
1. How many nitrogens do we have?• Two
2. What will the prefix be?• Di-
3. What is the prefix plus the element?– Dinitrogen
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Binary Molecular Compounds
1. How many oxygens do we have?• One
2. What will the prefix be?• Mono
3. What is the prefix plus the element?• Monoxide
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Binary Molecular Compounds
• What is the final answer?
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How do we know what we are naming?
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Pop Quiz
1. HCl
2. HClO3
3. H2S
4. H2SO4
5. H2ClO2
A. Chlorous acid
B. Sulfuric acid
C. Hydrosulfuric acid
D. Chloric acid
E. Hydrochloric acid
Match the following correctly, also note if the acid is binary or an oxyacid:
··Hint·· ClO3 is chlorate
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Section Three
Molecular Structure
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Molecular Formula
• Shows the elements symbols and subscripts
• PH3
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Lewis Structure
HP
H
H
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Space-filling Molecular Model
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Ball-and-stick Molecular Model
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Structural Formula
HP
H
H
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Molecular Formula
• CH4
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Lewis Structure
HC
H
H
H
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Space-filling Molecular Model
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Ball-and-stick Molecular Model
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Structural Formula
HC
H
H
H
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Lewis Structures
• BH3
• Nitrogen trifluoride
• C2H4
• Carbon Disulfide
• NH4+
• ClO4-
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Announcement
• Print out chapter 8 review from teacher page.
• Complete by Friday (will have time in class tomorrow to work on it)
• Test Monday on sections 1,2,3
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Resonance Structures
• Resonance– A condition that occurs when more than one
valid Lewis structure can be written for a molecule or ion
– Molecules and ions that undergo resonance behave as if there is only one structure
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Classwork
• Page 260– #53
• Page 274– #84, 101, 102, 103, 104
• BONUS: 5 pts #137
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Exceptions to the Octet Rule
• Odd number of valence electrons• Suboctets and coordinate covalent bonds
– Stable configuration with fewer than eight electrons present
– BH3
– Coordinate Covalent bond• One atom donates both of the electrons to be
shared with an atom or ion that needs two electrons to form a stable electron arrangement with lower potential energy.
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Exceptions to the Octet Rule
• Expanded Octets– Central atoms contain more than eight
valence electrons– Considers the d orbital– Extra lone pairs are added to the central atom
for more bonds
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Section Four
Molecular Shape
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Importance of Shape
• The shape can determine– Physical properties– Chemical properties
• Electron densities created by overlap of orbitals of shared electrons determine molecular shape
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VSEPR Model
• Valence
• Shell
• Electron
• Pair
• Repulsion
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VSEPR Model
• Arrangement that minimizes the repulsion of shared and unshared electron pairs around the central atom
• Bond Angle– Angle between bonds
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Hybridization
• Hybridization– A process in which atomic orbitals mix and
form new, identical hybrid orbitals
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Hybridization
• With regards to molecules that have more than two atoms
• To determine the orbital hybrid– Determine the number of e- pairs shared, and
lone pairs
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Hybridization
• Count like this. . . . – 1 = s– 2 = sp– 3 = sp2
– 4 = sp3
– 5 = sp3d– 6 = sp3d2
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Molecular Shapes
• Linear– Example BeCl2
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
2 2 0 sp 180
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Molecular Shapes
• Trigonal Planar– Example AlCl3
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
3 3 0 sp2 120
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Molecular Shapes
• Tetrahedral– Example CH4
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
4 4 0 sp3 109.5
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Molecular Shapes
• Trigonal Pyramidal– Example PH3
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
4 3 1 Sp3 107.3
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Molecular Shapes
• Bent– Example H2O
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
4 2 2 Sp3 104.5
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Molecular Shapes
• Trigonal Bipyramidal– Example NbCl5
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
5 5 0 sp3d 90; 120
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Molecular Shapes
• Octahedral– Example SF6
Total Pairs
Shared Pairs
Lone Pairs
Hybrid Orbitals
Bond Angle
6 6 0 sp3d2 90; 90
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Practice Problems
• Page 264– #56 through60
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Section Five
Electronegativity & Polarity
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Electron Affinity, Electronegativity, and Bond Character
• Electron Affinity– The measure of the tendency of an atom to
accept electrons– How attractive an atom is to electrons– Increases with atomic number within a period– Decreases with atomic number within a group
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Electron Affinity, Electronegativity, and Bond Character
• Electronegativity– Derived by comparing an atom’s attraction for
shared electrons to that of a fluorine’s atom attraction for shared electrons
– Ability of an atom to attract electrons to itself within a covalent bond
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Electron Affinity, Electronegativity, and Bond Character
• Bond Character– Chemical bonds between atoms of different
elements is never completely ionic or covalent– Four Types
• Mostly ionic• Polar covalent• Mostly covalent• Nonpolar covalent
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Electron Affinity, Electronegativity, and Bond Character
• Bond Character– Can be predicted using the electronegativity
difference of the elements that bond
Electronegativity Difference
Bond Character
> 1.7 Mostly ionic
0.4 – 1.7 Polar covalent
< 0.4 Mostly covalent
0 Nonpolar covalent
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Polar Covalent Bonds
• Polar Covalent Bonds– An unequal sharing of valence electrons
• Partial Charge– Represented by δ (Greek letter delta) – Due to unequal sharing, partial charges result
• Partial positive—the atom with the lower electron affinity
• Partial negative—the atom with higher electron affinity
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Molecular Polarity
• Covalently bonded molecules– Either polar or nonpolar
• Depends on location and nature of bonds
• Nonpolar Molecules– Not attracted by electric field
• Polar Molecules– Dipoles, with charged ends– Uneven electron density = attracted by
electric field
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Polarity and Molecular Shape
• Let’s look at H2O and CCl4• What shape does water take?
– Bent
• What shape does carbon tetrachloride take?– Tetrahedral
• Draw them
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H2O & CCl4
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Polarity and Molecular Shape
• The symmetry in CCl4 allows for a nonpolar molecule.
• There is no symmetry in H2O, so it is polar.
• What about NH3?
– It is polar.
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Properties of Covalent Compounds
• Covalent compounds have strong bonds between atoms
• Attraction forces between molecules are relatively weak
• Intermolecular forces– Many types
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Properties of Covalent Compounds
• Intermolecular Forces– Between nonpolar molecules
• Force is weak• Called dispersion force or induced dipole
– Between opposite charged ends of two polar molecules
• dipole-dipole force• The more polar the molecule the stronger the force
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Properties of Covalent Compounds
• Intermolecular Forces– Between hydrogen end of one dipole and a F,
O, N atom on another dipole• Hydrogen bond
• Forces and Properties– Weak forces result in relatively low melting
points– Molecular substances as gases at room
temperature• O2, CO2, H2S
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• Forces and Properties– Hardness
• Depends on strength of intermolecular forces• Many covalent compounds are soft
– Example: Paraffin, found in candles
Properties of Covalent Compounds
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Properties of Covalent Compounds
• Forces and Properties – Solid Phase
• Molecules align to form a crystal lattice– Similar to ionic solid– Less attraction between particles– Shape affected by molecular shape– Most information has been determined by molecular
solids
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Covalent Network Solids
• Covalent Network Solids– Composed only of atoms interconnected by a
network of covalent bonds• Example: Quartz and diamonds
– Structure can explain properties• Diamond
– Tetrahedral– Strong bonds– High melting point, extremely hard
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The End