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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
2
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
3
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qq 21
0potential 4
1E
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
6
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
23
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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27
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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