Chapter 8 Covalent Bonding

Let’s Review What do we already know? What is a chemical bond? What is an ionic bond?

Section 1 The Covalent Bond

Stability Lower energy is more stable Noble-Gas electron configuration Octet rule

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

Diatomic Molecules Two atom molecules are more stable than one atom H2, N2, O2, F2, Cl2, Br2, I2 H H

Hydrogen They Pair!! H H

Hydrogen

Oxygen

Fluorine

Fluorine F F

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

Bonding Orbital Localized region where bonding electrons are most likely found

Groups and Single Bonds

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

Homework continued Draw LDS for CH4 Br2 C6H14 also written as CH3(CH2)4CH3

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

The Pi Bond Multiple covalent bonds Consist of at least one sigma and one pi bond

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

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

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

Length and Energy Shorter the length the greater the energy

Energies of Chemical Reactions Total energy is determined from energy of bonds broken and formed Two types Endothermic Exothermic

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

Energies of Chemical Reactions Exothermic Energy in Energy out Bond

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

Energies of Chemical Reactions Endothermic Energy out Energy in Bond

Section Two Naming Molecules

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

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

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

Binary Molecular Compounds How many oxygens do we have? One What will the prefix be? Mono What is the prefix plus the element? Monoxide

Binary Molecular Compounds What is the final answer? Dinitrogen monoxide

How do we know what we are naming?

··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

Section Three Molecular Structure

Molecular Formula Shows the elements symbols and subscripts PH3

Lewis Structure H P

Space-filling Molecular Model

Ball-and-stick Molecular Model

Structural Formula H P

Molecular Formula CH4

Lewis Structure H C

Space-filling Molecular Model

Ball-and-stick Molecular Model

Structural Formula H C

Lewis Structures BH3 Nitrogen trifluoride C2H4 Carbon Disulfide NH4+ ClO4-

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

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

Classwork Page 260 Page 274 BONUS: 5 pts #137 #53 #84, 101, 102, 103, 104 BONUS: 5 pts #137

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.

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

Section Four Molecular Shape

Importance of Shape The shape can determine Physical properties Chemical properties Electron densities created by overlap of orbitals of shared electrons determine molecular shape

VSEPR Model Valence Shell Electron Pair Repulsion

VSEPR Model Arrangement that minimizes the repulsion of shared and unshared electron pairs around the central atom Bond Angle Angle between bonds

Hybridization Hybridization A process in which atomic orbitals mix and form new, identical hybrid orbitals

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

Hybridization Count like this. . . . 1 = s 2 = sp 3 = sp2 4 = sp3 5 = sp3d 6 = sp3d2

Molecular Shapes Linear Total Pairs Shared Pairs Lone Pairs Example BeCl2 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 2 sp 180

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

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

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

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

Molecular Shapes Trigonal Bipyramidal Total Pairs Shared Pairs Example NbCl5 Total Pairs Shared Pairs Lone Pairs Hybrid Orbitals Bond Angle 5 sp3d 90; 120

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

Practice Problems Page 264 #56 through60

Electronegativity & Polarity Section Five Electronegativity & Polarity

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

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

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

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

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

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

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

H2O & CCl4

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.

Properties of Covalent Compounds Covalent compounds have strong bonds between atoms Attraction forces between molecules are relatively weak Intermolecular forces Many types

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

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

Properties of Covalent Compounds Forces and Properties Hardness Depends on strength of intermolecular forces Many covalent compounds are soft Example: Paraffin, found in candles

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

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

The End