chem 101 chapter09 lec

8/13/2019 Chem 101 chapter09 LEC

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Chapter 9

Chemical

Bonding I:

Lewis Theory

2008, Prentice Hall

Chemistry: A Molecular Approach, 1stEd.

Nivaldo Tro

Roy Kennedy

Massachusetts Bay Community College

Wellesley Hills, MA


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Tro, Chemistry: A Molecular Approach 2

Bonding Theories explain how and why atoms attach together

explain why some combinations of atoms are stableand 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 modelscalledLewis structuresthat allow us to predict many

properties of molecules

aka Electron Dot Structures

such as molecular shape, size, polarity


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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 composeatoms

the potential energy between charged particles isdirectly proportional to the product of the charges

the potential energy between charged particles isinversely proportional to the distance between thecharges


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Potential Energy Between

Charged Particles

0is a constant= 8.85 x 10-12C2/Jm

for charges with the same sign, Epotentialis + and themagnitude gets less positive as the particles get fartherapart

for charges with the opposite signs, Epotentialis and

the magnitude gets more negative as the particles getcloser together

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

r

qq 21

0potential

4

1E


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Potential Energy Between

Charged Particles

The repulsion between

like-charged particles

increases as theparticles get closer

together. To bring

them closer requires

the addition of moreenergy.

The attraction between

opposite-charged

particles increases asthe particles get closer

together. Bringing

them closer lowers the

potential energy of thesystem.


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Bonding

a chemical bond forms when the potentialenergy 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 repulsionnucleus-to-electron attraction


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Types of Bonds

Types of Atoms Type of BondBond

Characteristic

metals to

nonmetals Ionicelectrons

transferred

nonmetals to

nonmetalsCovalent

electrons

shared

metal to

metalMetallic

electrons

pooled


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8

Types of Bonding


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Ionic Bonds

when metals bond to nonmetals, some electronsfrom the metal atoms are transferred to the

nonmetal atomsmetals have low ionization energy, relatively easy to

remove an electron from

nonmetals have high electron affinities, relativelygood to add electrons to


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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 ofpotential energy for the atoms to share valenceelectrons

potential energy lowest when the electrons are between thenuclei

shared electrons hold the atoms together by attractingnuclei of both atoms


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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 valenceelectrons

pair first two electrons for thesorbital

put one electron on each open side forpelectrons

then pair rest of thepelectrons

LiBe

B

C

N

O

F

Ne


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Lewis Symbols of Ions

Cations have Lewis symbols withoutvalence electrons

Lost in the cation formation

Anions have Lewis symbols with 8 valenceelectrons

Electrons gained in the formation of the anion

Li Li+1

F

1

F


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Stable Electron Arrangements

And Ion Charge Metals form cations by losing

enough electrons to get thesame electron configurationas the previous noble gas

Nonmetals form anions bygaining enough electrons toget the same electronconfiguration as the next

noble gas The noble gas electronconfiguration must be verystable

Atom Atoms

Electron

Config

Ion Ions

Electron

ConfigNa [Ne]3s

1Na

+1[Ne]

Mg [Ne]3s2 Mg

+2 [Ne]

Al [Ne]3s23p

1Al

+3 [Ne]

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

F [He]2s22p

5 F

-1[Ne]


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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 electronH 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 6valence electrons

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


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Lewis Theory

the basis of Lewis Theory is that there arecertain electron arrangements in the atom thatare more stable

octet rule

bonding occurs so atoms attain a more stableelectron configuration

more stable = lower potential energyno attempt to quantify the energy as the calculation

is extremely complex


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Properties of Ionic Compounds

hard and brittle crystalline solidsall are solids at room temperature

melting points generally > 300C the liquid state conducts electricitythe solid state does not conduct electricity

many are soluble in waterthe solution conducts electricity well

Melting an Ionic Solid


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Conductivity of NaCl

in NaCl(s), the

ions are stuck in

position and not

allowed to moveto the charged

rods

in NaCl(aq), the

ions are

separated andallowed to move

to the charged

rods


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Lewis Theory and Ionic Bonding

Lewis symbols can be used to represent thetransfer of electrons from metal atom to

nonmetal atom, resulting in ions that areattracted to each other and therefore bond

FLi +

1

F

Li +


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Predicting Ionic Formulas

Using 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 transfercomes out even

O

Li

Li

2

O2 Li + Li2O


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Energetics of Ionic Bond Formation

the ionization energy of the metal is endothermicNa(s) Na+(g) + 1 e DH= +603 kJ/mol

the electron affinity of the nonmetal is exothermicCl2(g) + 1 e

Cl(g) DH= 227 kJ/mol

generally, the ionization energy of the metal is largerthan the electron affinity of the nonmetal, therefore theformation of the ionic compound should beendothermic

but the heat of formation of most ionic compounds isexothermic and generally large; Why?

Na(s) + Cl2(g) NaCl(s) DHf = -410 kJ/mol


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Ionic Bonds

electrostatic attraction is nondirectional!!no direct anion-cation pair

no ionic molecule

chemical formula is an empirical formula, simplygiving the ratio of ions based on charge balance

ions arranged in a pattern called a crystal lattice

every cation surrounded by anions; and every anionsurrounded by cations

maximizes attractions between + and - ions


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Lattice Energy the lattice energyis 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 fromknowledge of other processes

lattice energy depends directly on size of charges andinversely on distance between ions


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Born-Haber Cycle

method for determining the lattice energy of anionic substance by using other reactions

use Hesss Law to add up heats of other processes

DHf

(salt) = DHf

(metal atoms, g) + DHf

(nonmetal atoms, g)

+ DHf(cations, g) + DHf(anions, g) + DHf(crystal lattice)

DHf(crystal lattice) = Lattice Energy

metal atoms (g) cations (g), DHf= ionization energy

dont forget to add together all the ionization energies to get to thedesired cation

M2+= 1stIE + 2ndIE

nonmetal atoms (g) anions (g), DHf= electron affinity


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Born-Haber Cycle for NaCl


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Practice - Given the Information Below,

Determine the Lattice Energy of MgCl2

Mg(s) Mg(g) DH1f= +147.1 kJ/mol Cl2(g) Cl(g) DH2f= +121.3 kJ/mol

Mg(g) Mg+1(g) DH3f= +738 kJ/mol

Mg+1(g) Mg+2(g) DH4f= +1450 kJ/mol

Cl(g) Cl-1(g) DH5f= -349 kJ/mol

Mg(s) + Cl2(g) MgCl2(s) DH6f= -641.3 kJ/mol

kJ2521H

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

H2HHH2HHH

HH2HHH2HH

energylatticef

energylatticef

f5f4f3f2f1f6energylatticef

energylatticeff5f4f3f2f1f6

DD

DDDDDDD

DDDDDDD


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Trends in Lattice Energy

Ion Size

the force of attraction between chargedparticles 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 Size

Metal ChlorideLattice Energy

(kJ/mol)

LiCl -834

NaCl -787

KCl -701

CsCl -657


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Trends in Lattice Energy

Ion Charge

the force of attraction betweenoppositely charged particles is

directly proportional to the product

of the charges larger charge means the ions aremore strongly attracted

larger charge = stronger attraction =

larger lattice energy

of the two factors, ion chargegenerally more important

Lattice Energy =

-910 kJ/mol

Lattice Energy =

-3414 kJ/mol

Example 9 2 Order the following ionic


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Example 9.2Order 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 chargesCa2+& 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


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Ionic Bonding

Model vs. Reality ionic compounds have high melting points and boiling

points

MP generally > 300C

all 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), thehigher the melting point


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Ionic Bonding

Model vs. Reality

ionic solids are brittle and hard the position of the ion in the crystal is critical to

establishing maximum attractive forcesdisplacing the ions from their positions resultsin like charges close to each other and therepulsive forces take over

+

-+ + + +

+ + + +- --

--

--

-

+ - + + + +

+ + + +- -

-

-

-

-

-

-

+ - + + + +

+ + + +- -

-

-

-

-

-

-


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Ionic Bonding

Model vs. Reality 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 chargedparticles that are able to flow through the material

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

in the liquid state, or when dissolved in water, the ionshave the ability to move through the structure and

therefore conduct electricity


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Single Covalent Bonds

two atoms share a pair of electrons2 electrons

one atom may have more than one single bond

F

F

F

F

HH O

HH O

F F


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Double Covalent Bond

two atoms sharing two pairs of electrons4 electrons

O O

O

O

O O


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Triple Covalent Bond

two atoms sharing 3 pairs of electrons6 electrons

N

N

N

N

N N


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Covalent Bonding

Predictions from Lewis Theory Lewis theory allows us to predict the formulas of

molecules

Lewis theory predicts that some combinations should bestable, while others should not

because 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 moleculesrather than an array


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Covalent Bonding

Model vs. Reality molecular compounds have low melting points and

boiling pointsMP generally < 300C

molecular compounds are found in all 3 states at roomtemperature

melting and boiling involve breaking the attractionsbetween the molecules, but not the bonds betweenthe atoms the covalent bonds are strong

the attractions between the molecules are generally weak

the polarity of the covalent bonds influences the strength ofthe intermolecular attractions


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Intermolecular Attractions vs. Bonding


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Ionic Bonding

Model vs. Reality

some molecular solids are brittle and hard, butmany are soft and waxy

the kind and strength of the intermolecular

attractions varies based on many factors

the covalent bonds are not broken, however, thepolarity of the bonds has influence on these

attractive forces


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Ionic Bonding

Model vs. Reality molecular compounds do not conduct electricity in the

liquid state

molecular acids conduct electricity when dissolved inwater, but not in the solid state

in molecular solids, there are no charged particlesaround 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 Polarity covalent bonding between unlike atoms results in

unequal sharing of the electronsone atom pulls the electrons in the bond closer to its side

one end of the bond has larger electron density than the

other the result is a polar covalent bondbond polarity

the end with the larger electron density gets a partial

negative charge

the end that is electron deficient gets a partial positivecharge


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HF

H F

d d

FH

EN 2.1 EN 4.0


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Electronegativity

measure of the pull an atom has on bondingelectrons

increases across period (left to right) and

decreases down group (top to bottom)fluorine is the most electronegative element

francium is the least electronegative element

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


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Electronegativity Scale


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49

Electronegativity and Bond Polarity If difference in electronegativity between bonded atoms

is 0, the bond is pure covalent equal sharing

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

If difference in electronegativity between bonded atoms0.5 to 1.9, the bond is polar covalent

If difference in electronegativity between bonded atomslarger 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


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Bond Polarity

ENCl= 3.0

3.0 - 3.0 = 0Pure Covalent

ENCl= 3.0

ENH= 2.13.02.1 = 0.9

Polar Covalent

ENCl= 3.0

ENNa= 1.03.00.9 = 2.1

Ionic


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

polarity of a bond a dipole is a material with positively and negatively charged ends

measured

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

directlyproportional to the distance between them

m= (q)(r)

not Coulombs Law

measured in Debyes, D

the percent ionic characteris the percentage of a bondsmeasured dipole moment to what it would be if full ions


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Dipole Moments


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Watera Polar Molecule

stream of

water

attractedto a

charged

glass rod

stream of

hexane

notattracted

to a

charged

glass rod


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Lewis Structures

of Molecules shows pattern of valence electron distribution in

the molecule

useful for understanding the bonding in manycompounds

allows us to predict shapes of molecules

allows us to predict properties of molecules andhow they will interact together


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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 patternsdifferent from common have formal charges

B C N O F


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Writing Lewis Structures of Molecules

HNO3

1) Write skeletal structure H always terminal

in oxyacid, H outside attached to Os

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 = 5

H = 1

O3= 36 = 18

Total = 24 e-


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HNO33) Attach central atom to the surrounding atoms with

pairs of electrons and subtract from the total

ONOH

O

Electrons

Start 24

Used 8

Left 16


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HNO34) Complete octets, outside-in H is already complete with 2

1 bond

and re-count electrons

:

::

ONOH

O

N = 5

H = 1

O3= 36 = 18

Total = 24 e-

Electrons

Start 24

Used 8

Left 16

Electrons

Start 16

Used 16

Left 0


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HNO35) If all octets complete, give extraelectrons to central atom.

elements with dorbitals can havemore than 8 electrons

Period 3 and below

6) If central atom does not haveoctet, bring in electrons fromoutside atoms to share

follow common bonding patternsif possible

:

::

ONOH

|

O


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Practice Lewis Structures


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Practice - Lewis Structures

CO2

SeOF2

NO2-1

H3

PO4

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


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Formal Charge during bonding, atoms may wind up with more

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

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

left O FC = 6 - 4 - (4) = 0

S FC = 6 - 2 - (6) = +1

right O FC = 6 - 6 - (2) = -1

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

O S O


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Writing Lewis Formulas of

Molecules (contd)7) Assign formal charges to the atoms

a) formal charge = valence e-- lone pair e-- bonding e-

b) follow the common bonding patterns

OSO

H

|

HOCCH

|||

OH

0 +1 -1

all 0


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Common Bonding Patterns

B C N O

C

+

N

+

O

+

C

-

N

-

O

-

B

-

F

F+

-F

Practice Assign Formal Charges


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Practice - Assign Formal Charges

CO2

SeOF2

NO2-1

H3

PO4

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


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Practice - Assign Formal Charges

CO2

SeOF2

NO2-1

H3

PO4

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

rest 0

S = +1Se = +1

-1

-1all 0

-1

-1-1


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Resonance


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Ozone Layer

Rules of Resonance Structures


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Rules of Resonance Structures

Resonance structures must have the same connectivity only electron positions can change

Resonance structures must have the same number ofelectrons

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

Drawing Resonance Structures


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O N

O

O

Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets

2. assign formal charges

3. move electron pairs from atoms

with (-) formal charge toward

atoms with (+) formal charge

4. if (+) fc atom 2ndrow, only movein 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


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Exceptions to the Octet Rule

expanded octetselements with empty dorbitals can have more

than 8 electrons

odd number electron species e.g., NOwill have 1 unpaired electron

free-radical

very reactive

incomplete octetsB, Al


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Practice -Identify Structures with Better or


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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


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Bond Energies

chemical reactions involve breaking bonds in reactantmolecules and making new bond to create the products

the DHreactioncan be calculated by comparing the cost

of breaking old bonds to the profit from making newbonds

the amount of energy it takes to break one mole of abond in a compound is called the bond energy

in the gas state

homolyticallyeach atom gets bonding electrons


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Trends in Bond Energies

the more electrons two atoms share, the strongerthe covalent bond

CC (837 kJ) > C=C (611 kJ) > CC (347 kJ)

CN (891 kJ) > C=N (615 kJ) > CN (305 kJ)

the shorter the covalent bond, the stronger thebond

BrF (237 kJ) > BrCl (218 kJ) > BrBr (193 kJ)

bonds get weaker down the column


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E ti t th E th l f th F ll i R ti


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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

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

(Appendix DHf= -136.3 kJ/mol)


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Bond Lengths

the distance between the nuclei ofbonded atoms is called the bondlength

because the actual bond lengthdepends on the other atoms aroundthe bond we often use the average

bond lengthaveraged for similar bonds from

many compounds


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Trends in Bond Lengths

the more electrons two atoms share, the shorter thecovalent bond

CC (120 pm) < C=C (134 pm) < CC (154 pm)

CN (116 pm) < C=N (128 pm) < CN (147 pm) decreases from left to right across periodCC (154 pm)> CN (147 pm)> CO (143 pm)

increases down the columnFF (144 pm)> ClCl (198 pm)> BrBr (228 pm)

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


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Bond Lengths


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Metallic Bonds

low ionization energy of metals allows them tolose 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

electrons

electrons delocalized throughout the metal structure

bonding results from attraction of cation for thedelocalized electrons

Metallic Bonding


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Metallic Bonding

Metallic Bonding


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g

Model 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 vibratefaster, making it harder for electrons tomake their way through the crystal

Metallic Bonding


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g

Model vs. Reality

metallic solids conduct heat

the movement of the small, light electronsthrough the solid can transfer kinetic energy

quicker than larger particles

metallic solids reflect light

the mobile electrons on the surface absorbthe outside light and then emit it at the samefrequency

Metallic Bonding


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g

Model vs. Reality

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 attractionsand the crystal structure


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chem 101 chapter09 lec

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