Chem Ch 8 new
Transcript of Chem Ch 8 new
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ChemicalBonding
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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Energetics of Ionic Bonding
We get 349 kJ/molback by giving
electrons tochlorine.
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Energetics of Ionic Bonding
But these numbersdont explain why
the reaction of sodium metal andchlorine gas to formsodium chloride isso exothermic!
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Energetics of Ionic Bonding
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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Energetics of Ionic Bonding
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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Energetics of Ionic Bonding
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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P olar Covalent Bonds
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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P olar Covalent Bonds
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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Writing Lewis Structures
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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Writing Lewis Structures
3. Fill the octets of theouter atoms.
Keep track of the electrons:
26 6 = 20 18 = 2
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Writing Lewis Structures
5. If you run out of electrons before the
central atom has anoctet
form multiple bonds
until it does.
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Writing Lewis Structures
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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Fewer Than Eight Electrons
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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Fewer Than Eight Electrons
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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More Than Eight Electrons
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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More Than Eight Electrons
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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Enthalpies of Reaction
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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Enthalpies of Reaction
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.