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Transcript of 1 10 11 12 15 17 20 24 31 37 39 49 53 55 61 65 67 71 88 103.
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AP Chapter 8Chemical Bonding
HW: 1 10 11 12 15 17 20 24 31 37 39 49 53 55 61 65 67 71 88 103
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8.1 - Chemical Bonds• Three basic types of bonds:
– Ionic• Electrostatic attraction between ions
– Covalent• Sharing of electrons
– Metallic• Metal atoms bonded to several
other atoms
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8.1 – Lewis Dot Symbols
• Atoms combine in order to form more stable electron configurations. The maximum stability is when the atom is isoelectronic with a noble gas = octet rule. Bonds involve valence electrons.
• The Symbols use one dot per valence electron. It is possible to accurately draw for s and p block elements.
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8.2 - Ionic Bonding• Ionic Bond = The electrostatic (+/-) force
holding together the cations and anions of an ionic compound. Caused by a TRANSFER of electron(s)
• Ionic compounds are composed of a crystal lattice structure
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8.2 - Ionic Bonding
• Can imagine as separate steps:– Cation loses electron(s) – Ionization Energy– Anion takes in electron(s) – Electron Affinity– Electrostatic attraction occurs(see next slides for a description of each of these for
NaCl)
Bonding Video – 4:12 –7:47: Formation of NaCl
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Energetics of Ionic Bonding
It takes 495 kJ/mol to remove an electron from sodium.
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Energetics of Ionic Bonding
We get 349 kJ/mol back by giving an electron to chlorine.
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Energetics of Ionic Bonding
• But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!
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Energetics of Ionic Bonding
• The third part is the electrostatic attraction between the newly formed sodium cationand chloride anion.
The LATTICE ENERGY of NaCl is +788 kJ/mol
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8.2 – Lattice Energy of Ionic Compounds
• Lattice Energy:The energy required to completely separate a mole of a solid
ionic compound into its gaseous ions.It is always a + valueThe higher the Lattice Energy, the more stable the compoundCannot be measured directly, so use “Born-Haber Cycle” to
calculate• The energy associated with electrostatic interactions is
governed by Coulomb’s law:
E = Q1Q2
r
The energy between two ions is directly proportional to the product of their charges and inversely proportional to the distance separating them. K = 8.99 x 109 Jm/C2
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Some Lattice Energies:
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Lattice Energy and the Formula of Ionic Compounds
-Lattice Energy explains why multi-positive ions can form
-Stability of “noble gas configuration” is not enough – Ionization energy is always + and increases with each subsequent IE-The lattice energy of the IONIC COMPOUND FORMED is high enough to compensate for the energy needed to remove the electrons from the cation
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• Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy to remove beyond that would be expended that cannot be overcome by lattice energies.
• Mg+2 is “worth” being formed when a compound like MgCl2 is formed
• Na+2 is not “worth” being formed because the lattice energy of NaCl2 is not high enough to compensate
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Transition Metal Ions
• Lose valence s electrons first
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Born – Haber Cycle • Born-Haber Cycle – The Born-Haber cycle
involves the formation of an ionic compound from the reaction of a metal with a non-metal (DHf). Born-Haber cycles are used primarily as a means of calculating lattice enthalpies, which cannot otherwise be measured directly.
• Relates Ionization Energy, Electron Affinity, and other properties in a Hess’ Law-type calculation
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Example of Born-Haber Cycle
Calculate the Lattice Energy of LiF:Goal: LiF(s) -> Li+1
(g) + F-(g)
Use:
Enthalpy of sublimation Lithium = +155.2 kJ/mol
Bond Dissociation Energy Fluorine Gas = +150.6 kJ/mol
Ionization Energy of Lithium = +520 kJ/mol
Electron Affinity of Fluorine = -328 kJ/mol
Enthalpy of Formation of Lithium fluoride = -594.1
Reaction:
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Example of Born-Haber CycleWork:
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The Born-Haber cycle can also be represented schematically
DHf = IE + DHsub + 1/2DHdiss + EA + - LE
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Another Example the of Born-Haber Cycle (not +1/-1)
Calculate the Lattice Energy of CaCl2:
Goal: CaCl2(s) -> Ca+2(g) + 2Cl-
(g)
Use:
Enthalpy of sublimation Calcium = +121 kJ/mol
Bond Dissociation Energy Chlorine Gas = +242.7 kJ/mol
First Ionization Energy of Calcium= +589.5 kJ/mol
Second Ionization Energy of Calcium = +1145 kJ/mol
Electron Affinity of Chlorine = -349 kJ/mol
Enthalpy of Formation of Calcium chloride= -795 kJ/mol
Reaction:
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Work:
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8.3 – Covalent Bonding• Electrons are shared• Molecular Compounds have covalent bonds• Lone Pairs = Pairs of electrons not in a covalent
bond• Lewis Structure – Representation of Covalent
Bonding• Octet Rule = Atoms bond to form a noble gas
configuration (often 8 valence e-)• Single Bond = 1 pair electrons• Double Bond = 2 pairs electrons• Triple Bond = 3 pairs electrons
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8.3Bond Length = Distance between nuclei of
two covalently bonded atoms in a molecule.
1>2>3Bond Energy = Energy needed to break a
bond3>2>1
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8.4 - Electronegativity
• Ability to pull electrons; Pauli Scale (max 4.0)• Polar Covalent = Unequal sharing of electrons –
0.3-1.7 difference• Nonpolar Covalent = Equal sharing - <0.3
difference• Ionic = Transfer of electrons - >1.7 difference(Some books give different ranges)
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Comparison of Ionic and Molecular Compounds
Ionic MolecularIonic Bonds Covalent Bonds (Polar or
Nonpolar)Very Strong Bonds Strong Bonds
High mp/bp Low mp/bp
Crystal Lattice Individual Molecule
Conducts electricity when in water or melted
Do not conduct electricity in water or when melted
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8.4 – Bond Polarity and Electronegativity
• Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.
• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.
• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.
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Polar Covalent Bonds
• When two atoms share electrons unequally, a bond dipole results.
• The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:
= Qr• It is measured in debyes (D).
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Polar Covalent BondsThe greater the difference in electronegativity, the more polar is the bond.
H - F H - F
d+
-d
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8.5 – Drawing Lewis Structures
Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.
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Writing Lewis Structures1. Find the sum of
valence electrons of all atoms in the polyatomic ion or molecule.– If it is an anion, add
one electron for each negative charge.
– If it is a cation, subtract one electron for each positive charge.
PCl35 + 3(7) = 26
Old method of connecting dots does not always work. Here is the other method:
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Writing Lewis Structures
2. The central atom is the least electronegative element that isn’t hydrogen. Connect the 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 the outer atoms.
Keep track of the electrons:26 6 = 20 18 = 2
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Writing Lewis Structures
4. Fill the octet of the central atom.
Keep track of the electrons:26 6 = 20 18 = 2 2 = 0
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Writing Lewis Structures
5. If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.
HCN
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Lewis Structures
• HNO3
• CO3-2
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8.5 – Formal Charges• Then assign formal charges.
– For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
– Subtract that from the number of valence electrons for that atom: The difference is its formal charge.
Total formal charge on a neutral molecule = 0Total formal charge on a polyatomic cation /anion= charge
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Writing Lewis Structures
• The best Lewis structure…1. Most times FOLLOW THE OCTET RULE2. When more than one structure is possible:…is the one with the fewest formal charges.…puts a negative charge on the most
electronegative atom.
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8.6 - Resonance
-Structure has more than one equal placement of electrons.
This is the Lewis structure we would draw for ozone, O3.
-
+
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Resonance• But this is at odds with
the true, observed structure of ozone, in which…– …both O—O bonds are
the same length.– …both outer oxygens
have a charge of 1/2.
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Resonance• One Lewis structure
cannot accurately depict a molecule such as ozone.
• We use multiple structures, resonance structures, to describe the molecule.
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Resonance
Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.
Ozone resonance structure =
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Resonance• Formate Ion:
The electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
• They are not localized, but rather are delocalized.
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Resonance
• The organic compound benzene, C6H6, has two resonance structures.
• It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.
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8.7 - Exceptions to the Octet Rule
• There are three types of ions or molecules that do not follow the octet rule:– Ions or molecules with an odd number of
electrons.– Ions or molecules with less than an octet.– Ions or molecules with more than eight valence
electrons (an expanded octet).
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Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.
Ex – NO; NO2; ClO2
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Incomplete Octet
• Consider BF3:– Giving boron a filled octet places a negative charge on
the boron and a positive charge on fluorine.– This would not be an accurate picture of the
distribution of electrons in BF3.– B stays at less than an octet!!
OR
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Incomplete Octet
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.
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Incomplete OctetThe lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom. (Group 13 is stable with incomplete octet)
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Expanded Octet
• The only way PCl5 can exist is if phosphorus has 10 electrons around it.
• It is allowed to expand the octet of atoms on the 3rd row or below.– Presumably d orbitals in
these atoms participate in bonding.
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Expanded Octet
Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.
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Expanded Octet
• This eliminates the charge on the phosphorus and the charge on one of the oxygens.
• The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.
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8.8 – Strengths of Covalent Bonds
• Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.
• This is the bond enthalpy = Bond Dissociation Energy
• The bond enthalpy for a Cl—Cl bond,D(Cl—Cl), is measured to be 242 kJ/mol.
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Average Bond Enthalpies
• This table lists the average bond enthalpies for many different types of bonds.
• Average bond enthalpies are positive, because bond breaking is an endothermic process.
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Average Bond Enthalpies
NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than theC—H bond in chloroform, CHCl3.
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Enthalpies of Reaction
• Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed.
• In other words, Hrxn = (bond enthalpies of bonds broken) +
-(bond enthalpies of bonds formed)
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Enthalpies of ReactionCH4(g) + Cl2(g)
CH3Cl(g) + HCl(g)
In this example:Broken:Four C—H bondsOne Cl—Cl bondMade:One C—Cl One H—Cl bondThree C-H bonds
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Enthalpies of Reaction
So,Hrxn = [4D(C—H) + D(Cl—Cl) + -[3D(C—H) +D(C—Cl) + D(H—Cl)]
= [4(413 kJ) + 242 kJ] + -[3(413 kJ) + 431 kJ + 328 kJ]
= (1894kJ) + (-1998 kJ)= 104 kJ
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Bond Enthalpy and Bond Length
• We can also measure an average bond length for different bond types.
• As the number of bonds between two atoms increases, the bond length decreases.