Basic Concepts of Chemical Bonding

1
Chapter 7

• Basic Concepts of Chemical Bonding

2
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
3
Ionic Electrostatic attraction between ions
Mg2O-2
K2(Cr2O7)-2
Ni2O-2
4
Covalent Shared pairs of electrons between
adjacent atoms
Br2 (l g)
C12H 22O 11 (s)
S8 (s)
5
Metallic Shared conduction band electrons
throughout crystal lattice
Mg0 (s)
Au20 (s)
Cu0 (s)
6
(No Transcript)
7
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 Ne3s23p4. Therefore it has the
  following Lewis symbol

which clearly depicts the six valence electrons.
8
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.
9
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.

10
Energetics of Ionic Bonding

• These numbers dont explain why the reaction of
  sodium metal and chlorine gas to form sodium
  chloride is so exothermic!

Figure 7.2
11
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.

12
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 Coulombs law

13
Lattice Energy

• Lattice energy increases with the charge on the
  ions.
• It also increases with the decreasing size of
  ions.

Table 7.2
14
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!

15
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.

16
Electron Configuration of Ions of the
Representative Elements

• Na 1s22s22p63s1 Ne3s1
• Na 1s22s22p63s1 Ne
• Cl 1s22s22p63s23p5 Ne 3s23p5
• Cl_ 1s22s22p63s23p6 Ne 3s23p6 Ar

17
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

18
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
  Ar3d64s2. To form the Fe2 ion, the two 4s
  electrons are lost giving the Ar3d6 electron
  configuration.
• To form the Fe3 ion, the two 4s electrons and
  one d electron are lost giving the Ar3d5
  electron configuration.

19
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
20
Lewis Structures

• Lewis structures are representations of molecules
  showing all electrons - bonding and nonbonding.
• --- denotes 1 pair of electrons.

21
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
22
(No Transcript)
23
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

24
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 column

Figure 7.5
25
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
26
Polar Covalent Bonds

• Although atoms often form compounds by sharing
  electrons, the electrons are not always shared
  equally.

Figure 7.6

• 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.

27
Polar Covalent Bonds

• The greater the difference in electronegativity,
  the more polar is the bond.

Table 7.3
Figure 7.8
28
Explaining Polar Molecules HCl
O.N. Cl 7-8 -1 H 1-0 1
29
Drawing Lewis Structures

• 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.

30
Drawing Lewis Structures

1. 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).

31
Drawing Lewis Structures

1. Complete the octets around all the atoms bonded
   to the central atom.

32
Drawing Lewis Structures

1. Place any leftover electrons on the central atom.

Remember we have 26 electrons
33
Drawing Lewis Structures

1. If there are not enough electrons to give the
   central atom an octet, try multiple bonds.

34
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.

35
Drawing Lewis Structures

• The best Lewis structure is the one
• with the fewest charges
• that puts a negative charge on the most
  electronegative atom.

36
Resonance Structures

• This is the Lewis structure we would draw for
  ozone, O3.

37
Resonance Structures

• This is at odds with the true, observed structure
  of ozone, in which
• both OO bonds are the same length, and
• both outer oxygens have a charge of ?½.

Figure 7.10
38
Resonance Structures

• One Lewis structure cannot accurately depict a
  molecule such as ozone.
• We use multiple structures called resonance
  structures, to describe the molecule.

39
Resonance Structures

• Just as green is a synthesis of blue and yellow,
  ozone is a synthesis of these two resonance
  structures.

Figure 7.11
40
Resonance Structures

• In truth, the electrons that form the second CO
  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.

41
Exceptions to the Octet Rule

• There are three types of ions or molecules that
  do not follow the octet rule
• Molecules and polyatomic ions containing an odd
  number of electrons.
• Molecules and polyatomic ions in which an atom
  has fewer than an octet of valence electrons.
• Molecules and polyatomic ions in which an atom
  has more than an octet of valence electrons.

42
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.

43
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.

44
Exceptions to the Octet RuleLess than an Octet
of Valence Electrons

• 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.

45
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, dont fill the octet of the central atom.

46
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.

47
Exceptions to the Octet RuleMore than an Octet
of Valence Electrons

• Even 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.

48
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.

49
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 ClCl bond,
• D(ClCl), is measured to be 242 kJ/mol.

50
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.

Table 7.4
NOTE These are average bond enthalpies, not
absolute bond enthalpies the CH bonds in
methane, CH4, will be a bit different than the
CH bond in chloroform, CHCl3.
51
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)
52
Bond Enthalpies and Enthalpies of Reaction
CH4(g) Cl2(g) ??? CH3Cl(g) HCl(g)

• In this example, one
• CH bond and one
• ClCl bond are broken
• and one CCl and one
• HCl bond are formed.

Figure 7.12
53
Bond Enthalpies and Enthalpies of Reaction

• ?(bond enthalpies of bonds broken) - ?(bond
  enthalpies of bonds formed)

?Hrxn D(CH) D(ClCl) ? D(CCl)
D(HCl) (413 kJ) (242 kJ)
? (328 kJ) (431 kJ) (655 kJ)
? (759 kJ) ?104 kJ
54
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
55
Nitroglycerin (1,2,3-Trinitroxypropane) Diatomaceous Earth Sodium Carbonate Dynamite Nitroglycerin (1,2,3-Trinitroxypropane) Diatomaceous Earth Sodium Carbonate Dynamite