presentation for bonding in chemistry

1. Bonding and Molecular
Structure: Fundamental
Concepts
Chapter 9
Chapter 9 1

2. Valence Electrons
- The electrons involved in bonding are called valence
electrons.
- Valence electrons are found in the incomplete,
outermost orbital shell of an atom.
- We can represent the electrons as dots around the
symbol for the element.
- These pictorial representations are called Lewis
Structures or Lewis Dot Structures.
Chapter 9 2

3. Lewis Symbols and the Octet Rule
Chapter 9 3

4. Chemical Bond Formation
- There are three types of chemical bonds
Ionic Bond - electrostatic attraction between ions of
opposite charge (NaCl).
Covalent Bond - sharing of electrons between two
atoms (Cl2).
Metallic Bond - sharing of electrons between several
atoms (Ag).
Chapter 9 4

5. Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g) → NaCl(s)
Chapter 9 5

6. Ionic Bonding
Na(s) + ½Cl2(g) → NaCl(s) ∆H°f = -410.9 kJ
- This reaction is very exothermic
- Sodium loses an electron to become Na+
- Chlorine gains an electron to become Cl−
- Na+ has an Ne electron configuration and Cl− has an
Ar configuration
Chapter 9 6

7. Ionic Bonding
Energetics of Ionic Bond Formation
Lattice Energy (∆Hlattice) – The energy required to
completely separate one mole of a solid ionic compound
into its gaseous ions.
Lattice energy depends on
-the charge on the ions
-the size of the ions
Coulomb’s equation:
Q1Q2
E=k
d
Q1, Q2 = charge on ions
k = 8.99 x 109 J-m/c2
d = distance between ions
Chapter 9 7

8. Covalent Bonding
When similar atoms bond, they share pairs of electrons
to each obtain an octet.
Example
Cl + Cl Cl Cl
a pair of electrons connect the two nuclei.
Chapter 9 8

9. Covalent Bonding
Multiple Bonds
- It is possible for more than one pair of electrons to be
shared between two atoms (multiple bonds)
- One shared pair of electrons - single bond (H2)
- Two shared pairs of electrons - double bond (O2)
- Three shared pairs of electrons - triple bond (N2).
H H O O N N
- Generally, bond distances decrease as we move from
single through double to triple bonds.
Chapter 9 9

10. Lewis Symbols and the Octet Rule
Octet rule – Atoms tend to gain, lose or share electrons
until they are surrounded by eight valence electrons.
Chapter 9 10

11. Drawing Lewis Structures
1) Draw a skeleton structure of the molecule or ion
showing the arrangement of the atoms and the
connect each atom to another with a single bond.
2) Determine the total number of valence elections in
the molecule or ion.
3) Deduct 2 electrons for each single bond used in step
1.
4) Distribute the rest of the electrons so that each atom
(except H) has 8 electrons.
- If you are “short” electrons, form multiple bonds
- If you have “extra” electrons, one of the heavy atoms may
be able to hold more that eight electrons.
Chapter 9 11

12. Drawing Lewis Structures
PCl3
Cl P Cl
Cl
Element Number Electrons Total
P 1 5 5
Cl 3 7 21
Total Electrons 26
Chapter 9 12

13. Drawing Lewis Structures
PCl3
Cl P Cl
Cl
Element Number Electrons Total
P 1 5 5
Cl 3 7 21
Total Electrons 26
Electrons used 6
Electrons remaining 20
Cl P Cl
Cl
Chapter 9 13

14. Drawing Lewis Structures
Isoelectronic Species
Molecules or ions having the same number of
valence electrons and the same Lewis structure.
N O + N N C O
Chapter 9 14

15. Drawing Lewis Structures
Resonance Structures
- Some molecules are not well described by Lewis
Structures.
Example: Ozone
O O
1 O O3 1 O O3
Chapter 9 15

16. Drawing Lewis Structures
Resonance Structures
- Experimentally, ozone has two identical bonds
whereas the Lewis Structure requires one single and
one double bond.
O
O
O
Chapter 9 16

17. Drawing Lewis Structures
Resonance Structures
- Resonance structures are attempts to represent a real
structure that is a mix between several extreme
possibilities.
- Each Lewis structure is call a Resonance Form
Resonance Form – Two or more Lewis structures
having the same arrangements of atoms but a
different arrangement of electrons
Chapter 9 17

18. Drawing Lewis Structures
Resonance Structures
- In ozone the resonance forms have one double and one
single bond.
O O
O O O O
- The actual structure of O3 is a combination (or
average) of the individual forms called a resonance
hybrid.
Chapter 9 18

19. Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule:
- Molecules with an odd number of electrons
- Molecules in which one atom has less than an octet
- Molecules in which one atom has more than an octet
Chapter 9 19

20. Exceptions to the Octet Rule
Odd Number of Electrons
- there are few molecules which fit this category
Examples. ClO2, NO, and NO2
O N O
Chapter 9 20

21. Exceptions to the Octet Rule
Less than an Octet
- This refers to the central molecule
- Typical for compounds of Groups 1A, 2A, and 3A.
Examples: LiH, BeH2, BF3
F
Li F H Be H B
F F
Chapter 9 21

22. Exceptions to the Octet Rule
More than an Octet
- This starts for atoms in the 3rd period onwards.
- This is due to vacant d orbitals which can hold the
“extra” electrons.
- Another factor is the size of the central atom, as they
get bigger, it gets easier to place additional atoms
around the central atom.
Chapter 9 22

23. Molecular Shapes
Lewis structures give atomic connectivity (which atoms
are connected to which).
Chapter 9 23

24. Molecular Shapes
Molecular Shapes are determined by:
Bond Distance – Distance between the nuclei of two
bonded atoms along a straight line.
Bond Angle – The angle between any two bonds
containing a common atom.
Chapter 9 24

25. Molecular Shapes
Chapter 9 25

26. Molecular Shapes
Valence Shell Electron Pair Repulsion Theory
(VSEPR)
- VSEPR theory is based on the idea that electrostatic
repulsion of the electrons are reduced to a minimum
when the various regions of high electron density
assume positions as far apart as possible.
Chapter 9 26

27. Molecular Shapes
Predicting Molecular Geometries
- draw the Lewis structure
- count the total number of bonding regions and lone
pairs around the central atom
- arrange the bonding regions and lone pairs in one
of the standard geometries to minimize e− -e−
repulsion
- multiple bonds count as one bonding region
Chapter 9 27

28. Molecular Shapes
Predicting Molecular Geometries
Common Configuration for saturated molecules.
Regions of Electron-Pair Bond Angle
Density Geometry
2 (AX2) Linear 180o
3 (AX3) Trigonal Planar 120o
4 (AX4) Tetrahedral 109.5o
5 (AX5) Trigonal 90o / 120o
Bipyramidal
6 (AX6) Octahedral 90o
Chapter 9 28

29. Molecular Shapes
- The “region of electron density” refers to:
- Lone pairs
- Covalent bonds (single, double, triple)
- Remember, you can’t “see” lone-pairs but they do
take-up space.
Chapter 9 29

30. Molecular Shapes
Predicting Molecular Geometries
Chapter 9 30

31. Molecular Shapes
Predicting Molecular Geometries
Chapter 9 31

32. Molecular Shapes
Molecules with Expanded Valence Shells
Chapter 9 32

33. Molecular Shapes
To minimize e− −e− repulsion,
lone pairs are always placed in
equatorial positions.
Chapter 9 33

34. Molecular Shapes
Molecules with Expanded Valence Shells
Chapter 9 34

35. Charge Distribution
Formal Charge
Used to predict the correct Lewis Structure.
1) Half of the electrons in a bond are assigned to each atom
in a bond.
2) Both electrons of an unshared pair of electrons are
assigned to the atoms to which the unshared pair belong.
3) The formal charge of an atom is equal to the valence
electrons minus the number of electrons assigned to each
atom.
Formal Charge = (group number) – (assigned electrons)
4) The sum of the formal charges will equal the charge on
the molecule or polyatomic ion.
Chapter 9 35

36. Charge Distribution
Using Formal Charge
1) A Lewis structure in which all formal charges in a molecule
are equal to zero is preferable to one in which some formal
charges are not zero.
2) If a Lewis structure has non-zero formal charges, the one
with the fewest nonzero formal charges is preferred.
3) A Lewis structure with one large formal charge is preferable
to one with several small formal charges.
4) A Lewis structure with adjacent formal charges should have
opposite signs.
5) When choosing between several Lewis structures, the
structure with negative formal charges on the more
electronegative atom is preferable.
Chapter 9 36

37. Bond Polarity and Electronegativity
- Electrons in a covalent bond may not be shared
evenly.
Electronegativity – The ability of an atom in a molecule
to attract electrons to itself.
- The periodic trend for electronegativity is up and to
the right across the periodic table.
Chapter 9 37

38. Bond Polarity and Electronegativity
Electronegativity
Chapter 9 38

39. Bond Polarity and Electronegativity
Electronegativity and Bond Polarity
- A chemical bond between elements with large
differences in eletronegativity will shift the electrons
to the atom with the higher electronegativity.
- The positive end (or pole) in a polar bond is
represented δ+ and the negative pole δ-.
- This is called a polar covalent bond.
- If the electronegativity difference is small, the bond is
nonpolar; if it is large, it is a polar bond.
Chapter 9 39

40. Polarity of Molecules
- To determine if a molecule is polar, you need to know
two things:
- polarity of the bonds in a molecule
- how the bonds are arranged
- A molecule is considered polar if its center of negative
and positive charge do not coincide.
δ+ δ−
H F H F
- Polar molecules have a dipole (a vector quantity)
- If these dipoles act equally and in opposition to each
other, the dipoles cancel-out and the molecule is
considered nonpolar.
Chapter 9 40

41. Polarity of Molecules
Dipole Moments of Polyatomic Molecules
Example:
CO2, each C-O dipole is canceled because the
molecule is linear.
H2O, the H-O dipoles do not cancel because the
molecule is bent.
Chapter 9 41

42. Polarity of Molecules
Dipole Moments of Polyatomic Molecules
Two simple rules to help determine molecular polarity
(most of the time)
1. If there are lone pairs on the central atom – the
molecule is polar.
2. If there is more than one type of bond on the central
atom – the molecule is polar.
Chapter 9 42

43. Polarity of Molecules
Dipole Moments of Polyatomic Molecules
Chapter 9 43

44. Strengths of Covalent Bonds
Bond Enthalpy (Energy) - The energy required to break
a covalent bond of a gaseous substance.
Cl2(g) → 2Cl(g) ∆H = DCl-Cl
- When more than one bond is broken the bond
enthalpy is a fraction of ∆H for the atomization
reaction :
CH4(g) → C(g) + 4H(g) ∆H = 1660 kJ
DC-H = ¼∆H = ¼(1660 kJ) = 415 kJ.
- Bond enthalpies can either be positive or negative.
Chapter 9 44

45. Strengths of Covalent Bonds
Chapter 9 45

46. Strengths of Covalent Bonds
Bond Enthalpies and the Enthalpies of Reaction
- Bond enthalpies can be used to calculate ∆Hrxn.
∆Hrxn = ∑D(bonds broken) − ∑D(bonds formed).
Chapter 9 46

47. Bond Enthalpies and the Enthalpies of Reaction
H H
H H
C C + H O O H H O C C O H
H H H H
Bonds Broken Bonds Formed
C=C C-C
614 kJ/mol 348 kJ/mol
O-O C-O
146 kJ/mol 358 kJ/mol
∆Hrxn = [1mol(614kJ/mol)+1mol(146kJ/mol)]-[2mol(358kJ/mol)+1mol(348kJ/mol)]
= -304 kJ
Chapter 9 47

48. Bond Enthalpies and the Enthalpies of Reaction
Cl
N Cl
2 N
N
+ 3
Cl
Cl Cl
Bonds Broken Bonds Formed
N-Cl N=N
200 kJ/mol 941 kJ/mol
Cl-Cl
242 kJ/mol
∆Hrxn = [2(3mol(200kJ/mol))]-[1mol(941kJ/mol)+3mol(242kJ/mol)]
= -467 kJ
Chapter 9 48

49. Homework
2, 6, 18, 22, 26, 32, 46, 57
Chapter 9 49