Chemical Bonds

1
Chemical Bonds

• Ionic Bonds
• Lewis Structures
• Nonpolar and Polar Covalent Bonds
• Electronegativity
• Bond Length, Energy and Order
• Formal Charges
• Resonance Structures
• Exceptions to the Octet Rule

2
Noble gas configuration

• The noble gases are noted for
• their chemical stability and
• existence as monatomic
• molecules.
• Except for helium,
• they share a common
• electron configuration
• that is very stable.
• This configuration has 8 valence-shell electrons.
  

valence e- He
2 Ne 8 Ar 8 Kr 8 Xe
8 Rn 8
3
The octet rule

• Atoms are most stable if they have a filled or
  empty outer layer of electrons.
• Except for H and He, a filled layer contains 8
  electrons - an octet.
• Atoms will
• gain or lose (ionic compounds)
• share (covalent compounds)
• electrons to make a filled or empty outer
  layer.

4
Ionic bonds

• Ionic bonds consist of the electrostatic
  attraction between positively charged and
  negatively charged ions.
• Ionic bonds are commonly formed between reactive
  metals and nonmetals.

Cl-
Na
5
Ions and the octet rule

• Simple ions are atoms that have gained or lost
  electrons to satisfy the octet rule.
• They will typically form based on what requires
  the smallest gain or loss of electrons to
  complete an octet.

Na Na e- group IA (1) e- Cl Cl-
group VIIA (17)
6
Ionic compounds

• Dont exist as individual molecules
• Tend to form crystals
• Ions touch many others
• Formula represents the average ion ratio

NaCl sodium chloride
7
Metals with multiple charges

• Group IA (1)
• All elements have a 1 charge. No numbers are
  required in the name.
• Group IIA (2)
• All have a single oxidation state of 2 No
  numbers are required in the name.
• Group IIIA (13)
• All have a 3 oxidation state. Tl also has a 1
  oxidation state.
• You must use numbers with thallium but not the
  other Group IIIA (13) metals.

8
Metals with multiple charges

• Group IV (14)
• All metals and semimetals have oxidation states
  of 2 and 4.
• Group V (15)
• All metals and semimetals have oxidation states
  of 3 and 5.
• Group VI (16)
• Metals and semimetals have oxidation states of
  4 and 6 except Po (2 only).
• Numbers must be used except for Po.

9
Why does this happen?

• There is an energy difference between p and s
  sublevels -
• p is greater.
• When a metal has electrons of both types, it is
  possible to lose the p electrons or all electrons
  in its outer shell.
• This is why there are two possible oxidation
  states for many representative elements.

10
Why does this happen?
11
Transition metals

• Remember, these elements have electrons fill
  inner levels so most have a ns2 electron
  configuration.
• Listing the ones that only have a single common
  ionic charge is easy.
• All Group III B - 3
• Ni, Zn, Cd - 2
• Ag - 1
• Lanthanides and actinides - 3
• The other elements are able to form two or more
  cations.

12
Why do most transition metals form two or more
cations?

• When a metal has d electrons, they can play a
  role in forming an ion.
• Examples.
• Fe loses two 4s electrons to form Fe2. It
  loses two 4s and one 3d electron to form Fe3.
• Cu loses on 4s electron to form Cu. It loses
  one 4s and one 3d to form Cu2.
• It would be difficult for you to predict all the
  possible charges and stay within the scope of
  this class.

13
Energy and ionic bond formation

• Born-Haber cycles
• Application of Hesss law that shows all steps
  involved in the formation of a compound.
• They are used to calculate lattice energies,
  which are difficult to measure experimentally.
• Lattice energy - The energy required to separate
  ions of an ionic solid to an infinite distance.

14
Energy and ionic bond formation

• Example - formation of sodium chloride.
• Steps DHo, kJ
• Vaporization of Na(s) Na(g) 92
• sodium
• Decomposition of 1/2 Cl2 (g) Cl(g) 121
• chlorine molecules
• Ionization of sodium Na(g)
  Na(g) 496
• Addition of electron Cl(g) e- Cl-(g) -349
• to chlorine
• Formation of NaCl Na(g)Cl-(g) NaCl -771

15
Energy and ionic bond formation
16
Lattice energy
Lattice energy Compound
kJ/mol LiCl 834 NaCl 769 KCl 701 NaBr 732
Na2O 2481 Na2S 2192 MgCl2
2326 MgO 3795

• The higher the lattice energy, the stronger the
  attraction between ions.

17
Lewis structures

• This is a simple system to help keep track of
  electrons around atoms, ions and molecules -
  invented by G.N. Lewis.
• If you know the number of electrons in the
  valence-shell of an atom, writing Lewis
  structures is easy.
• Lewis structures are used primarily for s- and
  p-block elements.

18
Lewis symbols
Basic rules Draw the atomic symbol. Treat each
side as a box that can hold up to two
electrons. Count the electrons in the valence
shell. Start filling box - dont make pairs
unless you need to.
X
19
Lewis symbols
O
Oxygen has 6 electrons in its valence -
VIA. Start putting them in the boxes.
20
Lewis symbols
21
Lewis Symbols
O
This is the Lewis symbol for oxygen.
22
Lewis symbols
Lewis symbols of second period elements
Li Be B C N O F Ne
23
Lewis dot formula andthe formation of NaCl
Na Cl Na Cl
The electron from Na moves over to the Cl. Now
both satisfy the octet rule. Na becomes Na -
a cation Cl becomes Cl- - an anion The and
- charges attract each other and form an ionic
bond.
24
Lewis dot structures of covalent compounds

• In covalent compounds atoms share electrons. We
  can use Lewis structures to help visualize the
  molecules.

• Lewis structures
• Multiple bonds must be considered.
• Will help determine molecular geometry.
• Will help explain polyatomic ions.

25
Types of electrons

• Bonding pairs
• Two electrons that are shared between two atoms.
  A covalent bond.
• Unshared pairs
• A pair of electrons that are not shared between
  two atoms. Lone pairs or nonbonding electrons.

Unshared pair
oo
H Cl
oo
oo
oo
Bonding pair
26
Single covalent bonds
H
H
H
C
H
H
H
Do atoms (except H) have octets?
27
Nonpolar and polar covalent bonds

• Nonpolar
• When two atoms share a pair of electrons
  equally.
• H H Cl Cl
• Polar
• A covalent bond in which the electron pair in
  not shared equally.
• H Cl
• Note A line can be used to represent a shared
  pair of electrons.

oo
oo
oo
oo
oo
oo
oo
oo
oo
d-
d
oo
oo
oo
28
Polar molecules

• Electrons in a covalent bond are rarely shared
  equally.
• Unequal sharing results in polar bonds.

oo
H F
oo
oo
oo

• Slight positive side
• Smaller electronegativity

• Slight negative
• Larger electronegativity

29
Electronegativity

• The ability of an atom that is bonded to another
  atom or atoms to attract electrons to itself.
• It is related to ionization energy and electron
  affinity.
• It cannot be directly measured.
• The values are unitless since they are relative
  to each other.
• The values vary slightly from compound to
  compound but still provide useful qualitative
  predictions.

30
Electronegativities
Electronegativity is a periodic property.
Electronegativity
Atomic number
31
Electronegativity

• Relative ability of atoms to attract electrons of
  bond.

32
Electronegativity

• The greater the difference in electronegativity
  between two bonded atoms, the more polar the
  bond.
• If the difference is great enough, electrons are
  transferred from the less electronegative atom to
  the more electronegative one.
• - Ionic bond.
• Only if the two atoms have exactly the same
  electronegativity will the bond be nonpolar.

33
Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl
• H2O
• CH4
• NO2

34
Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl ENK 0.9 ENCl 2.8 D 1.9
• H2O
• CH4
• NO2

35
Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl ENK 0.9 ENCl 2.8 D 1.9
• H2O ENH 2.2 ENo 3.5 D 1.3
• CH4
• NO2

36
Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl ENK 0.9 ENCl 2.8 D 1.9
• H2O ENH 2.2 ENo 3.5 D 1.3
• CH4 ENC 2.5 ENH 2.2 D 0.3
• NO2

37
Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl ENK 0.9 ENCl 2.8 D 1.9
• H2O ENH 2.2 ENo 3.5 D 1.3
• CH4 ENC 2.5 ENH 2.2 D 0.3
• NO2 ENN 3.1 ENO 3.5 D 0.4

38
Properties of ionic and covalent compounds

• Ionic compounds
• Held together by electrostatic attraction
• Exist as 3-D network of ions
• Empirical formula is used
• Covalent compounds
• Discrete molecular units
• Atoms held together by shared electron pairs
• Formula represents atoms in a molecule

39
Drawing Lewis structures

• Write the symbols for the elements in the correct
  structural order.
• Calculate the number of valence electrons for all
  atoms in the compound.
• Put a pair of electrons between each symbol, the
  bond between each.
• Beginning with the outer atoms, place pairs of
  electrons around atoms until each has eight
  (except for hydrogen).
• If an atom other than hydrogen has less than
  eight electrons, move unshared pairs to form
  multiple bonds.

40
Lewis structures

• Example CO2
• Step 1
• Draw any possible structures

C-O-O O-C-O
You may want to use lines for bonds. Each
line represents 2 electrons.
41
Lewis structures

• Step 2
• Determine the total number of valence electrons.
• CO2 1 carbon x 4 electrons 4
• 2 oxygen x 6 electrons 12
• Total electrons 16

42
Lewis structures

• Step 3
• Try to satisfy the octet rule for each atom
• - all electrons must be in pairs
• - make multiple bonds as required

Try the C-O-O structure
No matter what you try, there is no way satisfy
the octet for all of the atoms.
C O O
43
Lewis structures
This arrangement needs too many electrons.
O C O
How about making some double bonds?
That works!
is a double bond, the same as 4 electrons
44
Ammonia, NH3
45
Multiple bonds

• So how do we know that multiple bonds really
  exist?
• The bond energies and lengths differ!
• Bond Bond Length Bond energy
• type order pm kJ/mol
• C C 1 154 347
• C C 2 134 615
• C C 3 120 812

46
Formal Charges

• A bookkeeping system for electrons.
• They are used to show the approximate
  distribution of electron density in a molecule or
  polyatomic ion.
• Assign each atom half of the electrons in each
  pair it shares.
• Also give each atom all electrons from unshared
  pairs it has.
• Subtract the number of electrons assigned to each
  atom from the number of valence electrons for an
  uncombined atom of the element.

47
Formal charges

• Example. CO2
• For each oxygen
• 4 electrons from unshared electrons
• 2 from the bonds
• 6 total
• Formal charge 6 - 6 0
• For carbon
• 4 from the bonds
• 4 total
• Formal charge 4 - 4 0

48
Formal charges

• Another example - CO
• For oxygen
• 2 from unshared electron pairs
• 3 from bonded electron pairs
• 5 total
• Formal charge 6 - 5 1
• For carbon
• 2 from unshared electron pairs
• 3 from bonded electron pairs
• 5 total
• Formal charge 4 - 5 -1

49
Resonance structures

• Sometimes we can have two or more equivalent
  Lewis structures for a molecule.
• O - S O O S - O
• They both - satisfy the octet rule
• - have the same number of bonds
• - have the same types of bonds
• Which is right?

50
Resonance structures

• They both are!
• O - S O O S -
  O
• O S O
• This results in an average of 1.5 bonds between
  each S and O.

51
Resonance structures

• Benzene, C6H6, is another example of a compound
  for which resonance structure must be written.
• All of the bonds are the same length.

or
52
Exceptions to the octet rule

• Not all compounds obey the octet rule.
• Three types of exceptions
• Species with more than eight electrons around an
  atom.
• Species with fewer than eight electrons around an
  atom.
• Species with an odd total number of electrons.

53
Atoms with more than eight electrons

• Except for species that contain hydrogen, this is
  the most common type of exception.
• For elements in the third period and beyond, the
  d orbitals can become involved in bonding.
• Examples
• 5 electron pairs around P in PF5
• 5 electron pairs around S in SF4
• 6 electron pairs around S in SF6

54
An example SO42-

• 1. Write a possible
• arrangement.
• 2. Total the electrons.
• 6 from S, 4 x 6 from O
• add 2 for charge
• total 32
• 3. Spread the electrons
• around.

55
Atoms with fewer than eight electrons

• Beryllium and boron will both form compounds
  where they have less than 8 electrons around them.

FBF F




56
Atoms with fewer than eight electrons

• Electron deficient. Species other than hydrogen
  and helium that have fewer than 8 valence
  electrons.
• They are typically very reactive species.

57
Species with an odd total number of electrons

• A very few species exist where the total number
  of valence electrons is an odd number.
• This must mean that there is an unpaired
  electron which is usually very reactive.
• Radical - a species that has one or more
  unpaired electrons.
• They are believed to play significant roles in
  aging and cancer.

58
Species with an odd total number of electrons

• Example - NO
• Nitrogen monoxide is an example of a compound
  with an odd number of electrons.
• It is also known as nitric oxide.
• It has a total of 11 valence electrons six
  from oxygen and 5 from nitrogen.
• The best Lewis structure for NO is