chapter 9 chemical bonding i: lewis theory 2008, prentice hall chemistry: a molecular approach, 1 st...

Chapter 9Chemical

Bonding I:Lewis Theory

2008, Prentice Hall

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

Tro, Chemistry: A Molecular Approach 2

Bonding Theories• explain how and why atoms attach together• explain why some combinations of atoms are stable

and others are notwhy is water H2O, not HO or H3O

• one of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory

• Lewis Theory emphasizes valence electrons to explain bonding

• using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many properties of moleculesaka Electron Dot Structuressuch as molecular shape, size, polarity


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


Potential Energy Between Charged 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

r

qq 21

0potential 4

1E


Potential Energy BetweenCharged Particles

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.


Bonding

• a 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 repulsionelectron-to-electron repulsionnucleus-to-electron attraction


Types of Bonds

Types of Atoms Type of BondBond

Characteristic

metals to

nonmetalsIonic

electrons

transferred

nonmetals to

nonmetalsCovalent

electrons

shared

metal to

metalMetallic

electrons

pooled

8

Types of Bonding


Ionic Bonds

• when metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atomsmetals have low ionization energy, relatively easy to

remove an electron fromnonmetals have high electron affinities, relatively

good to add electrons to


Covalent Bonds• nonmetals 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 electronspotential energy lowest when the electrons are between the

nuclei• shared electrons hold the atoms together by attracting

nuclei of both atoms


Determining the Number of Valence Electrons in an Atom

• the column number on the Periodic Table will tell you how many valence electrons a main group atom hasTransition Elements all have 2 valence electrons; Why?

1A 2A 3A 4A 5A 6A 7A 8A

Li Be B C N O F Ne

1 e-1 2 e-1 3 e-1 4 e-1 5 e-1 6 e-1 7 e-1 8 e-1


Lewis Symbols of Atoms• aka electron dot symbols

• use symbol of element to represent nucleus and inner electrons

• use dots around the symbol to represent valence electronspair first two electrons for the s orbitalput one electron on each open side for p electrons then pair rest of the p electrons

Li Be

B

C

N

O

F

Ne


Lewis Symbols of Ions• Cations have Lewis symbols without

valence electronsLost in the cation formation

• Anions have Lewis symbols with 8 valence electronsElectrons gained in the formation of the anion

Li• Li+1

F

1

F


Stable Electron ArrangementsAnd Ion Charge

• Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas

• Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas

• The noble gas electron configuration must be very stable

Atom Atom’s Electron Config

Ion Ion’s Electron Config

Na [Ne]3s1 Na+1 [Ne]

Mg [Ne]3s2 Mg+2 [Ne]

Al [Ne]3s23p1 Al+3 [Ne]

O [He]2s22p4 O-2 [Ne]

F [He]2s22p5 F-1 [Ne]


Octet Rule• when atoms bond, they tend to gain, lose, or share electrons to

result in 8 valence electrons• ns2np6

noble gas configuration

• many exceptions H, Li, Be, B attain an electron configuration like He

He = 2 valence electrons Li loses its one valence electron H shares or gains one electron

though it commonly loses its one electron to become H+ Be loses 2 electrons to become Be2+

though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons

B loses 3 electrons to become B3+

though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons

expanded octets for elements in Period 3 or below using empty valence d orbitals


Lewis Theory• the basis of Lewis Theory is that there are

certain electron arrangements in the atom that are more stableoctet rule

• bonding occurs so atoms attain a more stable electron configurationmore stable = lower potential energyno attempt to quantify the energy as the calculation

is extremely complex


Properties of Ionic Compounds

• hard and brittle crystalline solidsall are solids at room temperature

• melting points generally > 300C• the liquid state conducts electricity

the solid state does not conduct electricity

• many are soluble in waterthe solution conducts electricity well

Melting an Ionic Solid


Conductivity of NaCl

in NaCl(s), the ions are stuck in position and not allowed to move to the charged rods

in NaCl(aq), the ions are separated and allowed to move to the charged rods


Lewis Theory and Ionic Bonding

• Lewis 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

FLi +

1

F

Li +


Predicting Ionic FormulasUsing Lewis Symbols

• 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

O

Li

Li

2

O2 Li + Li2O


Energetics of Ionic Bond Formation• 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 exothermic½Cl2(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


Ionic Bonds• electrostatic attraction is nondirectional!!

no direct anion-cation pair

• no ionic moleculechemical formula is an empirical formula, simply

giving the ratio of ions based on charge balance

• ions arranged in a pattern called a crystal latticeevery cation surrounded by anions; and every anion

surrounded by cationsmaximizes attractions between + and - ions


Lattice Energy• the lattice energy is the energy released when the

solid crystal forms from separate ions in the gas statealways 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


Born-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 energydon’t forget to add together all the ionization energies to get to the

desired cationM2+ = 1st IE + 2nd IE

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


Born-Haber Cycle for NaCl


Practice - Given the Information Below, Determine the Lattice Energy of MgCl2

Mg(s) Mg(g) H1°f = +147.1 kJ/mol½ Cl2(g) Cl(g) H2°f = +121.3 kJ/molMg(g) Mg+1(g) H3°f = +738 kJ/molMg+1(g) Mg+2(g) H4°f = +1450 kJ/molCl(g) Cl-1(g) H5°f = -349 kJ/molMg(s) + Cl2(g) MgCl2(s) H6°f = -641.3 kJ/mol


Practice - Given the Information Below, Determine the Lattice Energy of MgCl2

Mg(s) Mg(g) H1°f = +147.1 kJ/mol½ Cl2(g) Cl(g) H2°f = +121.3 kJ/molMg(g) Mg+1(g) H3°f = +738 kJ/molMg+1(g) Mg+2(g) H4°f = +1450 kJ/molCl(g) Cl-1(g) H5°f = -349 kJ/molMg(s) + Cl2(g) MgCl2(s) H6°f = -641.3 kJ/mol

kJ 2521H

kJ) 2(-349kJ) 1450(kJ) 738(kJ) 121.3(2kJ) 147.1(-kJ) 3.641(H

H2HHH2HHH

HH2HHH2HH

energy latticef

energy latticef

f5f4f3f2f1f6energy latticef

energy latticeff5f4f3f2f1f6


Trends in Lattice EnergyIon Size

• 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


Lattice Energy vs. Ion Size

Metal ChlorideLattice Energy

(kJ/mol)

LiCl -834

NaCl -787

KCl -701

CsCl -657


Trends in Lattice EnergyIon 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

Lattice Energy =-910 kJ/mol

Lattice Energy =-3414 kJ/mol


Example 9.2 – Order the following ionic compounds in order of increasing magnitude of

lattice energy.CaO, KBr, KCl, SrO

First examine the ion charges and order by product of the charges

Ca2+& O2-, K+ & Br─, K+ & Cl─, Sr2+ & O2─

(KBr, KCl) < (CaO, SrO)

Then examine the ion sizes of each group and order by radius; larger < smaller

(KBr, KCl) same cation, Br─ > Cl─ (same Group)

KBr < KCl < (CaO, SrO)

(CaO, SrO) same anion, Sr2+ > Ca2+ (same Group)

KBr < KCl < SrO < CaO


Ionic BondingModel vs. Reality

• ionic compounds have high melting points and boiling pointsMP generally > 300°Call ionic compounds are solids at room temperature

• because 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



• 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

+ - + + + +

+ + + +- --

--

--

-+ - + + + +

+ + + +- --

--

--

-

+ - + + + +

+ + + +- --

--

--

-



• 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


Single Covalent Bonds• two atoms share a pair of electrons

2 electrons

• one atom may have more than one single bond

F••

••

•• • F•••••••

F••

••

•• ••

••F•••• HH O

•• ••••

••

H•H• O••

• •

••

F F


Double Covalent Bond

• two atoms sharing two pairs of electrons4 electrons

O••••O••

••••••

O••

• •

••O••

• •

••

O O······ ··


Triple Covalent Bond

• two atoms sharing 3 pairs of electrons6 electrons

N••

• •

•N••

• •

•

N•••••••••• N

N N····


Covalent BondingPredictions from Lewis Theory

• Lewis theory allows us to predict the formulas of molecules

• Lewis theory predicts that some combinations should be stable, while others should notbecause the stable combinations result in “octets”

• Lewis theory predicts in covalent bonding that the attractions between atoms are directional the shared electrons are most stable between the bonding atoms resulting in molecules rather than an array


Covalent BondingModel vs. Reality

• molecular compounds have low melting points and boiling pointsMP generally < 300°Cmolecular compounds are found in all 3 states at room

temperature• melting 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


Intermolecular Attractions vs. Bonding



• 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



• molecular 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


Bond Polarity• covalent bonding between unlike atoms results in

unequal sharing of the electronsone atom pulls the electrons in the bond closer to its sideone end of the bond has larger electron density than the

other

• the result is a polar covalent bond bond polaritythe end with the larger electron density gets a partial

negative chargethe end that is electron deficient gets a partial positive

charge


HF

H F••

FH

EN 2.1 EN 4.0


Electronegativity• measure of the pull an atom has on bonding

electrons• increases across period (left to right) and• decreases down group (top to bottom)

fluorine is the most electronegative elementfrancium is the least electronegative element

• the larger the difference in electronegativity, the more polar the bondnegative end toward more electronegative atom


Electronegativity Scale

49

Electronegativity and Bond Polarity• If difference in electronegativity between bonded atoms is 0,

the bond is pure covalentequal sharing

• If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent

• If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent

• If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic

“100%”

0 0.4 2.0 4.0

4% 51%Percent Ionic Character

Electronegativity Difference


Bond Polarity

ENCl = 3.03.0 - 3.0 = 0

Pure Covalent

ENCl = 3.0ENH = 2.1

3.0 – 2.1 = 0.9Polar Covalent

ENCl = 3.0ENNa = 1.0

3.0 – 0.9 = 2.1Ionic


Bond Dipole Moments• the dipole moment is a quantitative way of describing the

polarity of a bonda dipole is a material with positively and negatively charged endsmeasured

• dipole moment, , is a measure of bond polarity it is directly proportional to the size of the partial charges and

directly proportional to the distance between them = (q)(r)not Coulomb’s Lawmeasured in Debyes, D

• the percent ionic character is the percentage of a bond’s measured dipole moment to what it would be if full ions


Dipole Moments


Water – a Polar Molecule

stream of water attracted to a charged glass rod

stream of hexane not attracted to a charged glass rod


Example 9.3(c) - Determine whether an N-O bond is ionic, covalent, or polar covalent.

• Determine the electronegativity of each elementN = 3.0; O = 3.5

• Subtract the electronegativities, large minus small(3.5) - (3.0) = 0.5

• If the difference is 2.0 or larger, then the bond is ionic; otherwise it’s covalent

difference (0.5) is less than 2.0, therefore covalent• If the difference is 0.5 to 1.9, then the bond is

polar covalent; otherwise it’s covalentdifference (0.5) is 0.5 to 1.9, therefore polar covalent


Lewis Structures of Molecules

• shows pattern of valence electron distribution in the molecule

• useful for understanding the bonding in many compounds

• allows us to predict shapes of molecules

• allows us to predict properties of molecules and how they will interact together


Lewis Structures• use common bonding patterns

C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs

often Lewis structures with line bonds have the lone pairs left off their presence is assumed from common bonding patterns

• structures which result in bonding patterns different from common have formal charges

B C N O F


Writing Lewis Structures of Molecules HNO3

1) Write skeletal structure H always terminal

in oxyacid, H outside attached to O’s

make least electronegative atom central N is central

2) Count valence electrons sum the valence electrons for each

atom add 1 electron for each − charge subtract 1 electron for each + charge

ONOH

O

N = 5H = 1O3 = 3∙6 = 18Total = 24 e-



3) Attach central atom to the surrounding atoms with pairs of electrons and subtract from the total

ONOH

O

———

ElectronsStart 24Used 8Left 16



4) Complete octets, outside-in H is already complete with 2

1 bond

and re-count electrons

:

::

——— ONOH

O

N = 5H = 1O3 = 3∙6 = 18Total = 24 e-




Writing Lewis Structures of Molecules HNO35) If all octets complete, give extra electrons to central

atom. elements with d orbitals can have more than 8 electrons

Period 3 and below

6) If central atom does not have octet, bring in electrons from outside atoms to share

follow common bonding patterns if possible

:

::

—— ONOH|

O


Practice - Lewis Structures

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4


Practice - Lewis Structures

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

:O::C::O:

::

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••

16 e-

26 e-

18 e-

26 e-

32 e-

14 e-H P P H

HH

•• ••


Formal Charge• during bonding, atoms may wind up with more

or less electrons in order to fulfill octets - this results in atoms having a formal charge

FC = valence e- - nonbonding e- - ½ bonding e-

left O FC = 6 - 4 - ½ (4) = 0

S FC = 6 - 2 - ½ (6) = +1

right O FC = 6 - 6 - ½ (2) = -1

• sum of all the formal charges in a molecule = 0 in an ion, total equals the charge

•• •• ••••••••

••O S O••••


Writing Lewis Formulas of Molecules (cont’d)

7) Assign formal charges to the atoms a) formal charge = valence e- - lone pair e- - ½ bonding e-

b) follow the common bonding patterns

OSO

H

|

HOCCH

|||

OH

0 +1 -1

all 0


Common Bonding Patterns

B C N O

C+

N+

O+

C-

N-

O-

B-

F

F+

-F


Practice - Assign Formal Charges

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••


Practice - Assign Formal Charges

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••

all 0

-1

P = +1rest 0

S = +1Se = +1

-1

-1all 0

-1

-1-1


Resonance• when there is more than one Lewis structure for a

molecule that differ only in the position of the electrons, they are called resonance structures

• the actual molecule is a combination of the resonance forms – a resonance hybridit does not resonate between the two forms,

though we often draw it that way

• look for multiple bonds or lone pairs

•••• •• ••••••••

•• ••O S O O S O•••••• ••••

••••

••••


Resonance


Ozone Layer


Rules of Resonance Structures• Resonance structures must have the same connectivity

only electron positions can change• Resonance structures must have the same number of

electrons• Second row elements have a maximum of 8 electrons

bonding and nonbonding third row can have expanded octet

• Formal charges must total same• Better structures have fewer formal charges• Better structures have smaller formal charges• Better structures have − formal charge on more

electronegative atom


O N

O

O·· ··

········

··

··

Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. move electron pairs from atoms

with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

-1

-1

+1

O N

O

O

·· ····

····

······

-1

-1 +1


Exceptions to the Octet Rule

• expanded octetselements with empty d orbitals can have more

than 8 electrons

• odd number electron species e.g., NOwill have 1 unpaired electronfree-radicalvery reactive

• incomplete octetsB, Al


Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. move electron pairs from atoms

with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

O S

O

O

O

HH

·· ··

········

··

······

-1

-1

+2

O S

O

O

O

HH

··

······

··

······

0

0

0


Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••

all 0

-1

P = +1

S = +1Se = +1

-1

-1all 0

-1

-1-1


Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

F Se

O

F

••

•• •

•••

••

••

••

••

••O S

O

O

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

• O S

O

O

••

••

•• •

•

O S

O

O

••

••

•• •

•

••

••

••

••

••

••

••••

••

••

••

••

••

••

••

O N O ••

••

••

••

••••O N O •

•••

••

••

••••

H P P H

HH

•• ••

none

-1

-1

-1

+1

all 0

+1

all 0

-1

none

S = 0in allres. forms


Bond Energies• chemical 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 statehomolytically – each atom gets ½ bonding electrons


Trends in Bond Energies• the more electrons two atoms share, the stronger

the covalent bondC≡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 bondBr−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ)bonds get weaker down the column


Using Bond Energies to Estimate H°rxn

• 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))

82

Estimate the Enthalpy of the Following Reaction

H H + O O H O O H


Estimate the Enthalpy of the Following Reaction

H2(g) + O2(g) H2O2(g)

reaction involves breaking 1mol H-H and 1 mol O=O and making 2 mol H-O and 1 mol O-O

bonds broken (energy cost)

(+436 kJ) + (+498 kJ) = +934 kJ

bonds made (energy release)

2(464 kJ) + (142 kJ) = -1070

Hrxn = (+934 kJ) + (-1070. kJ) = -136 kJ

(Appendix H°f = -136.3 kJ/mol)


Bond 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 lengthaveraged for similar bonds from

many compounds


Trends in Bond Lengths• the more electrons two atoms share, the shorter the

covalent bondC≡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 periodC−C (154 pm) > C−N (147 pm) > C−O (143 pm)

• increases down the columnF−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)

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


Bond Lengths


Metallic Bonds• low ionization energy of metals allows them to

lose electrons easily• the simplest theory of metallic bonding involves

the metals atoms releasing their valence electrons to be shared by all to atoms/ions in the metalan organization of metal cation islands in a sea of

electronselectrons delocalized throughout the metal structure

• bonding results from attraction of cation for the delocalized electrons


Metallic Bonding


Metallic BondingModel vs. Reality

• 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



• 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



• metals generally have high melting points and boiling pointsall 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 column• the cations get larger down the column, resulting in a

larger distance from the nucleus to the free electrons

chapter 9 chemical bonding i: lewis theory 2008, prentice hall chemistry: a molecular approach, 1 st...

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