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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 H H
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Hydrogen They Pair!! H H
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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
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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 Energy in Energy out 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 Energy out Energy in Bond
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Section Two Naming Molecules
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Binary Molecular Compounds
Example: N2O 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
The second element in the formula is named using its root and adding the suffix –ide. What is the second element? Oxygen What will the name be? Oxide
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Binary Molecular Compounds
Prefixes are used to indicate the number of atoms of each element are present in the compound. How many nitrogens do we have? Two What will the prefix be? Di- What is the prefix plus the element? Dinitrogen
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Binary Molecular Compounds
How many oxygens do we have? One What will the prefix be? Mono What is the prefix plus the element? Monoxide
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Binary Molecular Compounds
What is the final answer? Dinitrogen monoxide
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How do we know what we are naming?
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··Hint·· ClO3 is chlorate
Pop Quiz Match the following correctly, also note if the acid is binary or an oxyacid: HCl HClO3 H2S H2SO4 H2ClO2 Chlorous acid Sulfuric acid Hydrosulfuric acid Chloric acid Hydrochloric acid ··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 H P
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Space-filling Molecular Model
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Ball-and-stick Molecular Model
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Structural Formula H P
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Molecular Formula CH4
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Lewis Structure H C
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Space-filling Molecular Model
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Ball-and-stick Molecular Model
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Structural Formula H C
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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 Page 274 BONUS: 5 pts #137 #53
#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 Total Pairs Shared Pairs Lone Pairs
Example BeCl2 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 2 sp 180
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Molecular Shapes Trigonal Planar Total Pairs Shared Pairs Lone Pairs
Example AlCl3 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 3 sp2 120
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Molecular Shapes Tetrahedral Total Pairs Shared Pairs Lone Pairs
Example CH4 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 4 sp3 109.5
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Molecular Shapes Trigonal Pyramidal Total Pairs Shared Pairs
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 Total Pairs Shared Pairs Lone Pairs
Example H2O Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 4 2 Sp3 104.5
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Molecular Shapes Trigonal Bipyramidal Total Pairs Shared Pairs
Example NbCl5 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 5 sp3d 90; 120
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Molecular Shapes Octahedral Total Pairs Shared Pairs Lone Pairs
Example SF6 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 6 sp3d2 90; 90
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Practice Problems Page 264 #56 through60
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Electronegativity & Polarity
Section Five Electronegativity & Polarity
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Electron Affinity, Electronegativity, and Bond Character
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
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
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
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 Nonpolar covalent
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Polar Covalent Bonds Polar Covalent Bonds Partial Charge
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 Nonpolar 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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Properties of Covalent Compounds
Forces and Properties Hardness Depends on strength of intermolecular forces Many covalent compounds are soft Example: Paraffin, found in candles
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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
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
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