Chapter 9Chapter 9Chemical Chemical Bonding I:Bonding I:Lewis Lewis TheoryTheory

Why Do Atoms Bond?Why Do Atoms Bond?processes are spontaneous if they

result in a system with lower potential energy

chemical bonds form because they lower the potential energy between the charged particles that compose atoms

the potential energy between charged particles is directly proportional to the product of the charges

the potential energy between charged particles is inversely proportional to the distance between the charges

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Potential Energy Between Potential Energy Between Charged ParticlesCharged Particles

0 is a constant ◦= 8.85 x 10-12 C2/J∙m

for charges with the same sign, Epotential is + and the magnitude gets less positive as the particles get farther apart

for charges with the opposite signs, Epotential is and the magnitude gets more negative as the particles get closer together

remember: the more negative the potential energy, the more stable the system becomes

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0potential 4

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Potential Energy BetweenPotential Energy BetweenCharged ParticlesCharged Particles

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The repulsion between like-charged particles increases as the particles get closer together. To bring them closer requires the addition of more energy.The attraction between

opposite-charged particles increases as the particles get closer together. Bringing them closer lowers the potential energy of the system.

BondingBondinga chemical bond forms when the

potential energy of the bonded atoms is less than the potential energy of the separate atoms

have to consider following interactions: ◦nucleus-to-nucleus repulsion◦electron-to-electron repulsion◦nucleus-to-electron attraction

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Ionic BondsIonic Bondswhen metals bond to nonmetals,

some electrons from the metal atoms are transferred to the nonmetal atoms◦metals have low ionization energy,

relatively easy to remove an electron from

◦nonmetals have high electron affinities, relatively good to add electrons to

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Covalent BondsCovalent Bondsnonmetals have relatively high ionization

energies, so it is difficult to remove electrons from them

when nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons◦potential energy lowest when the electrons

are between the nucleishared electrons hold the atoms together by

attracting nuclei of both atoms

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Lewis Theory and Ionic Lewis Theory and Ionic BondingBondingLewis symbols can be used to

represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond◦ electrons are transferred until the metal

loses all its valence electrons and the nonmetal has an octet

◦numbers of atoms are adjusted so the electron transfer comes out even

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Energetics of Ionic Bond Energetics of Ionic Bond FormationFormation

the ionization energy of the metal is endothermic◦Na(s) → Na+(g) + 1 e ─ H° = +603 kJ/mol

the electron affinity of the nonmetal is exothermic◦½Cl2(g) + 1 e ─ → Cl─(g) H° = ─ 227 kJ/mol

generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal, therefore the formation of the ionic compound should be endothermic

but the heat of formation of most ionic compounds is exothermic and generally large; Why?◦Na(s) + ½Cl2(g) → NaCl(s) H°f = -410 kJ/mol

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Ionic BondsIonic Bondselectrostatic attraction is nondirectional!!◦no direct anion-cation pair

no ionic molecule◦ chemical formula is an empirical formula,

simply giving the ratio of ions based on charge balance

ions arranged in a pattern called a crystal lattice◦ every cation surrounded by anions; and

every anion surrounded by cations◦maximizes attractions between + and - ions

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Lattice EnergyLattice Energy the lattice energy is the energy released when

the solid crystal forms from separate ions in the gas state◦ always exothermic ◦hard to measure directly, but can be

calculated from knowledge of other processes lattice energy depends directly on size of

charges and inversely on distance between ions

Tro, Chemistry: A Molecular Approach 11

Born-Haber Cycle for NaClBorn-Haber Cycle for NaCl

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ΔHof = 410.9 kJ/mol

ΔHof =107.7 kJ/mol

ΔHof = 121.7 kJ/mol

ΔHof = 495.9 kJ/mol

Ea = -349 kJ/mol

Born-Haber CycleBorn-Haber Cycle method for determining the lattice energy of an ionic

substance by using other reactions ◦ use Hess’s Law to add up heats of other processes

H°f(salt) = H°f(metal atoms, g) + H°f(nonmetal atoms, g) + H°f(cations, g) + H°f(anions, g) + H°f(crystal lattice)

◦ H°f(crystal lattice) = Lattice Energy

◦ metal atoms (g) cations (g), H°f = ionization energy

don’t forget to add together all the ionization energies to get to the desired cation M2+ = 1st IE + 2nd IE

◦ nonmetal atoms (g) anions (g), H°f = electron affinity

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Trends in Lattice Energy Ion Trends in Lattice Energy Ion SizeSize the force of attraction between charged

particles is inversely proportional to the distance between them

larger ions mean the center of positive charge (nucleus of the cation) is farther away from negative charge (electrons of the anion)◦ larger ion = weaker attraction =

smaller lattice energy

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Lattice Energy vs. Ion Lattice Energy vs. Ion SizeSize

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Metal ChlorideLattice Energy

(kJ/mol)

LiCl -834

NaCl -787

KCl -701

CsCl -657

Trends in Lattice EnergyTrends in Lattice EnergyIon ChargeIon Charge

the force of attraction between oppositely charged particles is directly proportional to the product of the charges

larger charge means the ions are more strongly attracted◦ larger charge = stronger

attraction = larger lattice energy

of the two factors, ion charge generally more important

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Lattice Energy =-910 kJ/mol

Lattice Energy =-3414 kJ/mol

ExamplesExamplesArrange MgO, CaO, and SrO in order of

increasing lattice energy

Ionic Bonding Model vs. Ionic Bonding Model vs. RealityReality

ionic compounds have high melting points and boiling points◦MP generally > 300°C◦all ionic compounds are solids at room

temperaturebecause the attractions between ions

are strong, breaking down the crystal requires a lot of energy◦the stronger the attraction (larger the

lattice energy), the higher the melting point

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Ionic Bonding Model vs. Ionic Bonding Model vs. RealityReality ionic solids are brittle and hard the position of the ion in the crystal is critical

to establishing maximum attractive forces – displacing the ions from their positions results in like charges close to each other and the repulsive forces take over

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+ - + + + +

+ + + +- --

--

--

-+ - + + + +

+ + + +- --

--

--

-

+ - + + + +

+ + + +- --

--

--

-

Ionic BondingModel vs. Ionic BondingModel vs. RealityReality ionic compounds conduct electricity in the liquid

state or when dissolved in water, but not in the solid state

to conduct electricity, a material must have charged particles that are able to flow through the material

in the ionic solid, the charged particles are locked in position and cannot move around to conduct

in the liquid state, or when dissolved in water, the ions have the ability to move through the structure and therefore conduct electricity

some molecular solids are brittle and hard, but many are soft and waxy

the kind and strength of the intermolecular attractions varies based on many factors

the covalent bonds are not broken, however, the polarity of the bonds has influence on these attractive forces

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Covalent BondingModel vs. Covalent BondingModel vs. RealityRealitymolecular compounds have low melting

points and boiling points◦MP generally < 300°C◦molecular compounds are found in all

3 states at room temperaturemelting and boiling involve breaking the

attractions between the molecules, but not the bonds between the atoms◦the covalent bonds are strong◦the attractions between the molecules

are generally weak◦the polarity of the covalent bonds

influences the strength of the intermolecular attractions

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Ionic BondingModel vs. Ionic BondingModel vs. RealityRealitymolecular compounds do not conduct

electricity in the liquid state molecular acids conduct electricity when

dissolved in water, but not in the solid state

in molecular solids, there are no charged particles around to allow the material to conduct

when dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity

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Bond EnergiesBond Energieschemical reactions involve breaking

bonds in reactant molecules and making new bond to create the products

the H°reaction can be calculated by comparing the cost of breaking old bonds to the profit from making new bonds

the amount of energy it takes to break one mole of a bond in a compound is called the bond energy◦in the gas state◦homolytically – each atom gets ½

bonding electrons

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Trends in Bond EnergiesTrends in Bond Energies

the more electrons two atoms share, the stronger the covalent bond◦C≡C (837 kJ) > C=C (611 kJ) > C−C (347

kJ)◦C≡N (891 kJ) > C=N (615 kJ) > C−N (305

kJ) the shorter the covalent bond, the stronger

the bond◦Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br

(193 kJ)◦bonds get weaker down the column

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Using Bond Energies to Using Bond Energies to Estimate Estimate H°H°rxnrxn

the actual bond energy depends on the surrounding atoms and other factors

we often use average bond energies to estimate the Hrxn

◦works best when all reactants and products in gas state

bond breaking is endothermic, H(breaking) = +bond making is exothermic, H(making) = −

Hrxn = ∑ (H(bonds broken)) + ∑ (H(bonds formed))

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Estimate the Enthalpy of the Estimate the Enthalpy of the Following ReactionFollowing Reaction

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ExamplesExamplesUse bond energies to estimate the enthalpy of

reaction for the combustion of methane:CH4(g) + 2 O2(g) CO2(g) + 2H2O(l)

Bond LengthsBond Lengths the distance between the

nuclei of bonded atoms is called the bond length

because the actual bond length depends on the other atoms around the bond we often use the average bond length◦ averaged for similar bonds

from many compounds

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Trends in Bond LengthsTrends in Bond Lengthsthe more electrons two atoms share, the

shorter the covalent bond◦C≡C (120 pm) < C=C (134 pm) < C−C (154

pm)◦C≡N (116 pm) < C=N (128 pm) < C−N (147

pm)decreases from left to right across period◦C−C (154 pm) > C−N (147 pm) > C−O (143 pm)

increases down the column◦F−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)

in general, as bonds get longer, they also get weaker

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Bond LengthsBond Lengths

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Metallic BondsMetallic Bondslow ionization energy of metals allows

them to lose electrons easilythe simplest theory of metallic bonding

involves the metals atoms releasing their valence electrons to be shared by all to atoms/ions in the metal◦an organization of metal cation islands in

a sea of electrons◦electrons delocalized throughout the

metal structurebonding results from attraction of cation

for the delocalized electrons

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Metallic BondingModel vs. Metallic BondingModel vs. RealityReality

metallic solids conduct electricity because the free electrons are mobile, it allows the electrons

to move through the metallic crystal and conduct electricity as temperature increases, electrical conductivity decreases heating causes the metal ions to vibrate faster, making it

harder for electrons to make their way through the crystal metallic solids conduct heat the movement of the small, light electrons through the solid

can transfer kinetic energy quicker than larger particles metallic solids reflect light the mobile electrons on the surface absorb the outside light

and then emit it at the same frequency

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Metallic Bonding Model vs. Metallic Bonding Model vs. RealityReality

metallic solids are malleable and ductile

because the free electrons are mobile, the direction of the attractive force between the metal cation and free electrons is adjustable

this allows the position of the metal cation islands to move around in the sea of electrons without breaking the attractions and the crystal structure

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Metallic Bonding Model vs. Metallic Bonding Model vs. RealityRealitymetals generally have high melting points and

boiling points◦all but Hg are solids at room temperature

the attractions of the metal cations for the free electrons is strong and hard to overcome

melting points generally increase to right across period

the charge on the metal cation increases across the period, causing stronger attractions

melting points generally decrease down columnthe cations get larger down the column,

resulting in a larger distance from the nucleus to the free electrons

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