PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.
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Transcript of PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.
![Page 1: PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.](https://reader035.fdocuments.us/reader035/viewer/2022070409/56649e995503460f94b9c933/html5/thumbnails/1.jpg)
PowerPoint to accompany
Chapter 7
Basic Concepts of Chemical Bonding
![Page 2: PowerPoint to accompany Chapter 7 Basic Concepts of Chemical Bonding.](https://reader035.fdocuments.us/reader035/viewer/2022070409/56649e995503460f94b9c933/html5/thumbnails/2.jpg)
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical Bonds
Three basic types of bonds: Ionic
Electrostatic attraction between ions
Covalent Sharing of electrons
Metallic Metal atoms bonded to several other
atoms
Figure 7.1
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Ionic: Electrostatic attraction between ions
Mg+2O-2
K+2(Cr2O7)-2
Ni+2O-2
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Covalent: Shared pairs of electrons between adjacent atoms
Br2 (l & g)C12H 22O 11 (s)
S8 (s)
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Metallic: Shared conduction band electrons throughout crystal lattice
Mg0 (s)
Cu0 (s)
Au20 (s)
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Lewis Symbols
The Lewis symbol of an element consists of the chemical symbol for the element plus a dot for each valence electron.
Sulfur, for example, has the electron configuration [Ne]3s23p4. Therefore it has the following Lewis symbol:
S• •
• •• •
which clearly depicts the six valence electrons.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Octet Rule
Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.
An octet of electrons consists of full s and p subshells in an atom.
Note: There are many exceptions to the octet rule, but it remains useful for introducing many important concepts of bonding.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Energetics of Ionic Bonding
As we have seen, it takes 495 kJ/mol to
remove the outer 3s1 electron from sodium.
We get 349 kJ/mol back by giving the 3p6
electron to chlorine.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Energetics of Ionic Bonding
These numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!
Figure 7.2
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Energetics of Ionic Bonding
There must be a third piece to the puzzle.
What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Lattice Energy
This third piece of the puzzle is the lattice energy:
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 Coulomb’s law:
Eel = Q1Q2d
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Lattice Energy
Lattice energy increases with the charge on the ions.
It also increases with the decreasing size of ions.
Table 7.2
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Energetics of Ionic Bonding
By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.
In practice it is easier to measure such things empirically (heat of melting, heat of dissolution etc.) than to figure out why they are that way!
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Energetics of Ionic Bonding
These phenomena also help explain the “octet rule”.
Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.
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Electron Configuration of Ions of the Representative Elements
Na 1s22s22p63s1 = [Ne]3s1
Na+ 1s22s22p63s1 = [Ne]
Cl 1s22s22p63s23p5 = [Ne] 3s23p5
Cl_
1s22s22p63s23p6 = [Ne] 3s23p6 = [Ar]
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Electron Configuration of Ions of the Representative Elements
In polyatomic ions, two or more atoms are bound together by predominantly covalent bonds and the whole acts as a charged species when the ion forms an ionic compound with an ion of opposite charge.
E.g. NH4+ and NH4Cl
CO32- and Na2CO3
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Electron Configuration of Ions of the Representative Elements In forming ions, transition metals lose the valence-shell
s electrons first, then as many d electrons as are required to reach the charge of the ion.
For example, Fe has the electron configuration [Ar]3d64s2. To form the Fe2+ ion, the two 4s electrons are lost giving the [Ar]3d6 electron configuration.
To form the Fe3+ ion, the two 4s electrons and one d electron are lost giving the [Ar]3d5 electron configuration.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Covalent Bonding
In these bonds, atoms share electrons.
There are several electrostatic interactions in these bonds:
Attractions between negative electrons and positive nuclei
Repulsions between -electrons Repulsions between +nuclei
Figure 7.4
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Lewis Structures
Lewis structures are representations of molecules showing all electrons - bonding and nonbonding.
--- denotes 1 pair of electrons.
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Lewis Structures: Multiple Bonds
When two electron pairs are shared, two lines are drawn, i.e. double bond.
O :: C :: O or O = C = O
When three electron pairs are shared, three lines are drawn, i.e. triple bond.
:N ::: N: or :N ≡ N:
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Bond Polarity and Electronegativity
Nonpolar covalent bond the electrons are shared equally between
the two atoms, e.g. in Cl2 and N2
Polar covalent bond one of the atoms exerts a greater attraction
for the bonding electrons than the other, e.g. in H2O
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Bond Polarity and Electronegativity
The ability of an atom in a molecule to attract electrons to itself.
On the periodic chart, electronegativity increases as you go: from left to right across
a row from the bottom to the
top of a columnFigure 7.5
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Dipole Moments 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).
Figure 7.7
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Polar Covalent Bonds
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.
Figure 7.6
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Polar Covalent Bonds
The greater the difference in electronegativity, the more polar is the bond.
Figure 7.8
Table 7.3
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Explaining Polar Molecules: HCl
O.N.:Cl 7-8 = -1H 1-0 = +1
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Drawing Lewis Structures
1. 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.
PCl3
5 + (3 x 7) = 26
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Drawing Lewis Structures
2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond, a dash (representing two electrons).
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Drawing Lewis Structures
3. Complete the octets around all the atoms bonded to the central atom.
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Drawing Lewis Structures
4. Place any leftover electrons on the central atom.
Remember we have 26 electrons
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Drawing Lewis Structures
5. If there are not enough electrons to give the central atom an octet, try multiple bonds.
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Drawing Lewis Structures - Formal Charge 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.
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Drawing Lewis Structures
The best Lewis structure is the one: with the fewest charges that puts a negative charge on the most
electronegative atom.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Resonance Structures
This is the Lewis structure we would draw for ozone, O3.
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Resonance Structures
This is at odds with the true, observed structure of ozone, in which: both O—O bonds are
the same length, and both outer oxygens have
a charge of ½.
Figure 7.10
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Resonance Structures
One Lewis structure cannot accurately depict a molecule such as ozone.
We use multiple structures called resonance structures, to describe the molecule.
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Resonance Structures
Just as green is a synthesis of blue and yellow, ozone is a synthesis of these two resonance structures.
Figure 7.11
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Resonance Structures In truth, 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 localised, but rather are delocalised.
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Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:1. Molecules and polyatomic ions containing
an odd number of electrons.
2. Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons.
3. Molecules and polyatomic ions in which an atom has more than an octet of valence electrons.
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Exceptions to the Octet RuleOdd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons, e.g. NO contains 5 + 6 = 11 valence electrons.
N = O...
..
.. N = O....
...and
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Exceptions to the Octet RuleLess than an Octet of Valence Electrons 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.
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Exceptions to the Octet RuleLess than an Octet of Valence ElectronsTherefore, 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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Exceptions to the Octet RuleLess than an Octet of Valence Electrons
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.
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Exceptions to the Octet RuleMore than an Octet of Valence Electrons
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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Exceptions to the Octet RuleMore than an Octet of Valence ElectronsEven though we can draw a Lewis structure for the phosphate ion that has only eight electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.
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Exceptions to the Octet RuleMore than an Octet of Valence Electrons This eliminates the charge on the phosphorus
and the charge on one of the oxygens, i.e. 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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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.
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.NOTE: These are average bond enthalpies, not
absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than the C—H bond in chloroform, CHCl3.
Table 7.4
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Bond Enthalpies and 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, i.e.
Hrxn = (bond enthalpies of bonds broken) - (bond enthalpies of bonds
formed)
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Bond Enthalpies and Enthalpies of Reaction
In this example, one
C—H bond and one
Cl—Cl bond are broken
and one C—Cl and one
H—Cl bond are formed.
Figure 7.12
CH4(g) + Cl2(g) CH3Cl(g) + HCl(g)
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Bond Enthalpies and Enthalpies of Reaction
(bond enthalpies of bonds broken) - (bond enthalpies of bonds formed)
Hrxn = [D(C—H) + D(Cl—Cl)] [D(C—Cl) + D(H—Cl)]= [(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 bond length for different bond types.
As the number of bonds between two atoms increases, the bond length decreases.
Table 7.5
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Nitroglycerin(1,2,3-Trinitroxypropane)
+Diatomaceous Earth
+Sodium Carbonate=
Dynamite