1. Bonding
Chapter 7

2. Unit Objectives 5.0 Define key terms and concepts
5.1 Draw Lewis Dot structures and line structures for simple chemical compounds.
5.2 Identify ionic and covalent compounds by their chemical formulas and their properties.
5.3 Determine the formal charge on an atom.
5.4 Explain the octet rule and exceptions to the rule.
5.5 Draw Resonance structures for chemical compounds.
5.6 Determine the lattice energy of a compound using the Born-Haber Cycle.
5.7 Calculate the bond enthalpy change in a reaction.

3. Stable Electron Configurations
Valence electrons interact between atoms of elements to form bonds when compounds are formed
This interaction involves either the sharing or transfer of electrons
Octet Rule
When electrons gain, lose, or share electrons to obtain 8 valence electrons (or a full outer shell)
Basically every element is trying to mimic a Noble Gas with their electron configuration

4. Lewis Electron Dot Structures
A way to represent the bonding and non-bonding electrons in a compound

5. Lewis Electron Dot Structures
Bonding Electrons
Involved in a bond
Non-Bonding Electrons (Lone Pairs)
Not involved in a bond
Line Structure
Replaces the bonding pair of electrons with a line

6. Lewis Electron Dot Structures
He S Ca Na Al O C N B

7. Stable Electron Configuration
Ionic Bonding
Covalent Bonding

8. Ionic Bonding
Bond formed when an electron is transferred from one atom to another resulting in the creation of two ions.
Typically between a metal and a non-metal
Ions
Element with a charge
Formed due to ionic bonding
Ionic Charge
An electrical charge on an element due to the lose/gain of an electron
This charge difference creates an electrostatic force, holding the two ions together.
The overall compound has a neutral charge.

9. Ions Cations Anions Ions with a positive charge Typically metals
Na+1, Al+3, C+4
Anions
Ions with a negative charge
Typically nonmetals
O-2, Cl-1, N-3

10. Lewis Dot Structures – Ionic Compounds

11. Draw Lewis Dot and line structures of the following compounds.
Be2O
HBr
MgF2
K2S

12. Covalent Bonds Formed when electrons are shared between two elements.
Created when two nonmetals bond together
Polar Covalent
Electrons are held slightly closer to one atom than another, creating an imbalance of electrons, resulting in a dipole moment
Dipole
Separation of positive and negative charges in a bond
Non-Polar Covalent
Equal sharing of electrons
Does not have a dipole moment

13. Covalent Bonds

14. Covalent Bonds

15. Covalent Bonds Diatomic elements
Formed when an 2 of the same elements bond covalently
Are gaseous
Non-polar

16. Covalent Bonds

17. Multiple Bonds Single Bond Double Bonds Triple Bonds
1 pair of electrons is involved in bonding
Double Bonds
Sharing of two electron pairs
Triple Bonds
Sharing of three electron pairs
Takes more energy to hold a triple bond that a single bond.

18. Lewis Dot Structures

19. Lewis Dot Structures

20. Draw Lewis Dot and line structures of the following compounds
Draw Lewis Dot and line structures of the following compounds. H2 SH2 CBr4 PCl3

21. Covalent Bonds Bond Length
The distance between the nuclei of two covalently bonded atoms
Single bonds have the longest bond length while triple bonds have the shortest
Bond length increases with atomic radius
Bond Type
Bond Length (pm)
C-C
154
C=C
133
C≡C
120

22. Electronegativity The tendency of an atom to attract electrons
Increases to the right and up on the periodic table
The greater the difference in charges between to atoms, the more the electrons are attracted to one atom over another

23. Electronegativity Nonpolar Covalent Bond Polar Covalent Bond
EN difference less than 0.5
Polar Covalent Bond
EN difference from 0.5 to 1.7
Ionic Bond
EN difference greater than 1.7

24. What Are Your Questions?

25. Formal Charge of an atom
The electrical charge difference between the valence electrons of an atom and the number of electrons assigned to that same atom in the Lewis structure.
To determine the formal charge of an atom
Assign all of the non-bonding electrons to the atom they are attached to
Assign half of the bonding electrons to each of the two atoms involved in a bond
Writing formal charges
The sum of the charges must equal zero for molecules
The sum of the charges must be a positive charge for cations
The sum of the charges must be a negative number for anions

26. Formal Charge of an atom
Determining the most plausible Lewis structure for a molecule
Having no charge is preferable to having a charge
Structures with small formal charges are preferable to those with large formal charges
Charges are usually limited to -1, 0, or +1 unless a metal is present
It is more plausible for the negative charges to be placed on the more electronegative atom, etc.

27. Exceptions to the Octet rule
An Incomplete Octet
Applies to hydrogen to carbon
Aluminum
Some atoms with a formal charge

28. Exceptions to the Octet rule
Odd number of electrons
Radical
Has an unpaired electron
Expanded Octet
More than 8 valence electrons
Can occur in elements in the third period and beyond

29. Draw the Lewis dot structure and determine the formal charge of the elements in each of the following compounds:
H3O+
NO3-

30. Draw the Lewis dot structure and determine the formal charge of the elements in each of the following compounds:
CN-
H2SO4

31. Resonance Structures
When more than one Lewis structure can be drawn for a given compound.
Typically happens when a molecule contains 1 or more multiple bonds.
Common on ions

32. Draw the resonance structures for the following compounds.
NO3-
PO43-
HSO4-

33. What Are Your Questions?

34. ∆H° = ΣBE(reactants) – ΣBE(products)
Bond Enthalpy
The energy required to break a bond in 1 mole of gaseous molecules.
The bond enthalpies of liquids and solids are effected by the surrounding molecules.
Where BE stands for average bond energy
This applies only to the gas phase
Turn to Table 9.4, pg. 401.
∆H° = ΣBE(reactants) – ΣBE(products)

35. Estimate the enthalpy change for the following reaction: N2(g) + H2(g)  NH3(g)

36. Estimate the enthalpy change for the following reaction: CH4(g) + O2(g)  H2O(g) + CO2(g)

37. Estimate the enthalpy change for the following reaction: C2H4(g) + H2(g)  C2H6(g)

38. Ionic Compounds Lattice Energy Coulomb’s Law Born-Haber Cycle
The energy required to completely separate one mole of a solid ionic compound into gaseous ions.
Coulomb’s Law
The potential energy (E) between two ions is directly proportional to the product of their charges and inversely proportional to the distance between the two ions.
Born-Haber Cycle
Relates lattice energies of ionic compounds to other atomic and molecular properties, such as ionization energy and electron affinity.
Based on Hess’s Law

39. The Born-Haber Cycle

40. Calculate the enthalpy change for following reaction using the Born-Haber Cycle. Mg(s) + ½O2(g)  MgO(s)
∆Hatomization O = 249kJ/mole
∆Hatomization Mg = 148kJ/mole
∆H1st ionization energy Mg = 738kJ/mole
∆H2nd ionization energy Mg = 1451kJ/mole
∆H1st electron affinity O = -141kJ/mole
∆H2nd electron affinity O = 798kJ/mole
∆Hlattice energy MgO = -3791kJ/mole

41. Calculate the enthalpy change for following reaction using the Born-Haber Cycle. 2Li(s) + ½O2(g)  Li2O(s)
∆Hatomization O = 249kJ/mole
∆Hatomization Li = 162kJ/mole
∆H1st ionization energy Li = 540kJ/mole
∆H1st electron affinity O = -141kJ/mole
∆H2nd electron affinity O = 798kJ/mole
∆Hlattice energy Li2O = -2800kJ/mole

42. Calculate the enthalpy change for following reaction using the Born-Haber Cycle. Mg(s) + Cl2(g)  MgCl2(s)
∆Hatomization Cl = 122kJ/mole
∆Hatomization Mg = 148kJ/mole
∆H1st ionization energy Mg = 738kJ/mole
∆H2nd ionization energy Mg = 1451kJ/mole
∆Hformation MgCl2 = -641kJ/mole
∆Hlattice energy MgCl2 = kJ/mole

43. What Are Your Questions?