chem ch 8 new

8/8/2019 Chem Ch 8 new

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Concepts of ChemicalBonding

Chemistry, The Central Science , 10th editionTheodore L. Brown; H. Eugene LeMay , Jr.; and Bruce E. Bursten


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Chemical Bonds

Three basic types of bonds:

IonicElectrostatic attractionbetween ions

CovalentSharing of electrons

MetallicMetal atoms bonded toseveral other atoms


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Energetics of Ionic Bonding

As we saw in thelast chapter , it takes

495 kJ/mol toremove electronsfrom sodium.


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We get 349 kJ/molback by giving

electrons tochlorine.


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But these numbersdont explain why

the reaction of sodium metal andchlorine gas to formsodium chloride isso exothermic!


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There must be athird piece to thepuzzle.What is as yetunaccounted for isthe electrostatic

attraction betweenthe newly formedsodium cation andchloride anion.


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Lattice Energy

This third piece of the puzzle is the latticeenergy:The energy required to completely separate a mole of

a solid ionic compound into its gaseous ions.

The energy associated with electrostatic

interactions is governed by Coulombs law:

E el = O Q1Q2

d


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Lattice Energy

Lattice energy , then , increases with the charge onthe ions.

It also increaseswith decreasingsize of ions.


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By accounting for allthree energies

(ionization energy , electron affinity , andlattice energy) , wecan get a good idea

of the energeticsinvolved in such aprocess.


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These phenomenaalso helps explain theoctet rule.

Metals , for instance , tend to stop losing electronsonce they attain a noble gas configurationbecause energy would be expended that cannotbe overcome by lattice energies.


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Covalent Bonding

In these bonds atoms shareelectrons.There are severalelectrostatic interactions inthese bonds:

Attractions between electronsand nucleiRepulsions between electronsRepulsions between nuclei


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P olar Covalent Bonds

Although atoms oftenform compounds by

sharing electrons , theelectrons are notalways shared equally.

Fluorine pulls harder on the electrons itshares with hydrogen than hydrogen does.Therefore , the fluorine end of the moleculehas more electron density than the

hydrogen end.


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When two atoms shareelectrons unequally , a bond

dipole results.The dipole moment , Q, produced by two equal butopposite charges separated

by a distance , r , is calculated: Q = Qr

It is measured in debyes (D).


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The greater thedifference in

electronegativity , the more polar isthe bond.


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Lewis Structures

Lewis structures are representations of

molecules showing all electrons , bonding andnonbonding.


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Writing Lewis Structures

1. Find the sum of valence electrons of allatoms in thepolyatomic ion or molecule.

If it is an anion , add oneelectron for eachnegative charge.If it is a cation , subtractone electron for eachpositive charge.

PCl3

5 + 3(7) = 26


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2. The central atom isthe least

electronegativeelement that isnthydrogen. Connectthe outer atoms to it

by single bonds.Keep track of the electrons:

26 6 = 20


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3. Fill the octets of theouter atoms.

Keep track of the electrons:

26 6 = 20 18 = 2


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5. If you run out of electrons before the

central atom has anoctet

form multiple bonds

until it does.


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The best Lewis structureis the one with the fewest charges.puts a negative charge on the mostelectronegative atom.


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Resonance

This is the Lewisstructure wewould draw for ozone , O 3. -

+


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Resonance

But this is at oddswith the true ,

observed structureof ozone , in which

both OO bondsare the same length.

both outer oxygens have acharge of 1/2.


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Resonance

One Lewis structurecannot accurately

depict a moleculesuch as ozone.We use multiplestructures , resonancestructures , to describethe molecule.


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Resonance

Just as green is a synthesisof blue and yellow

ozone is a synthesis of these two resonancestructures.


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Resonance

In truth , the electrons that form the second CObond in the double bonds below do not always sit

between that C and that O , but rather can moveamong the two oxygens and the carbon.They are not localized , but rather are delocalized .


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Resonance

The organic compoundbenzene , C 6H6, has tworesonance structures.It is commonly depictedas a hexagon with acircle inside to signifythe delocalizedelectrons in the ring.


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Exceptions to the Octet Rule

There are three types of ions or molecules that do not follow the octetrule:

Ions or molecules with an odd number of electrons.

Ions or molecules with less than an octet.Ions or molecules with more than eightvalence electrons (an expanded octet).


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Odd Number of Electrons

Though relatively rare and usually quiteunstable and reactive , there are ionsand molecules with an odd number of electrons.


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Fewer Than Eight Electrons

Consider BF 3:G iving boron a filled octet places a negativecharge on the boron and a positive charge onfluorine.This would not be an accurate picture of thedistribution of electrons in BF 3.


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Therefore , structures that put a double bondbetween boron and fluorine are much lessimportant than the one that leaves boron withonly 6 valence electrons.


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The lesson is: If filling the octet of the centralatom results in a negative charge on thecentral atom and a positive charge on themore electronegative outer atom , dont fill theoctet of the central atom.


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More Than Eight Electrons

The only way P Cl5 canexist is if phosphorushas 10 electronsaround it.It is allowed to expandthe octet of atoms onthe 3rd row or below.

P resumably d orbitals inthese atoms participatein bonding.


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Even though we can draw a Lewis structure for thephosphate ion that has only 8 electrons around thecentral phosphorus , the better structure puts a

double bond between the phosphorus and one of the oxygens.


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This eliminates the charge on the phosphorusand the charge on one of the oxygens.The lesson is: When the central atom is on the3rd row or below and expanding its octeteliminates some formal charges , do so.


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Average Bond Enthalpies

This table lists theaverage bond

enthalpies for manydifferent types of bonds.Average bondenthalpies arepositive , becausebond breaking is anendothermic process.


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Average Bond Enthalpies

NOTE: These areaverage bond

enthalpies , notabsolute bondenthalpies; the CHbonds in methane ,

CH 4, will be a bitdifferent than theCH bond inchloroform , CHCl 3.


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Enthalpies of Reaction

Yet another way toestimate ( H for a

reaction is to comparethe bond enthalpies of bonds broken to thebond enthalpies of the

new bonds formed.

In other words , ( H rxn = 7 (bond enthalpies of bonds broken)

7 (bond enthalpies of bonds formed)


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CH 4(g ) + Cl 2(g ) pCH 3Cl(g ) + HCl (g )

In this example , oneCH bond and one

ClCl bond are broken;one CCl and one HClbond are formed.


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So ,

( H rxn = [D (CH) + D (ClCl) [D (CCl) + D (HCl)

= [(413 kJ) + (242 kJ)] [(328 kJ) + (431 kJ)]= (655 kJ) (759 kJ)= 104 kJ


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Bond Enthalpy and Bond Length

We can also measure an average bondlength for different bond types.As the number of bonds between two atomsincreases , the bond length decreases.

chem ch 8 new

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