1. Chemical Bonding
2014

2. What is a chemical bond? In order to bond atoms must interact.
Atoms interact to achieve a stable octet.
Atoms will lose, gain or share electrons to get an octet. This creates a bond.
The type of bond depends on what takes place with the valence electrons.

3. What is a chemical bond? Definition:
A force of attraction caused by a loss, gain or sharing of electrons that holds atoms together.

4. Energy of Bonds
Chemical energy (potential energy) is released when bonds are formed.
Chemical energy is added to break these bonds.
Forming a bond lowers potential energy.
weak bonds – need little energy to break
strong bonds – need a lot of energy to break.

5. Types of Bonds 1. Ionic – formed by metal with non-metal
or positive ion and negative ion
2. Covalent – between two non-metals
3. Metallic – force of attraction that hold
atoms of metals together

7. Ionic Bonds – between ions.
Transfer of electrons form ions (charged particles)
cations and anions attract each other by electrostatic forces.
For example:
metal forms + ion (oxidation)
Na  Na +1 (loses 1 e-)
non-metal forms – ion (reduction)
Cl  Cl-1 (gains 1 e-)
Lewis dot diagrams show how electrons are transferred.

8. Example: sodium reacts with chlorineCl Na
Has one valence electron, wants to get rid of it so that its valence shell is full
Has 7 valence electrons, needs one more to have a full valence shell
This reactions makes sodium chloride or table salt.
Na+1[ Cl ]-1

9. Ionic Bonds – Ionization Energy and electronegativity
In order for an ionic bond to be formed one atom must have a low ionization energy (low electronegativity) and the other must have a high ionization energy (high electronegativity)
p. 2

11. Reaction to form an ionic compound
In your notes, describe the reaction.
Are the reactive properties of these elements consistent with what we learned last chapter about the groups of the periodic table?

12. Ionic Bonds – electronegativity
p. 2
Electronegativity values are used to determine bond type.
Differences equal to or greater than 1.7 are considered ionic.
Example: Look up on Table S
Na = 0.9 Cl = 3.2
3.2 – 0.9 = since 2.3>1.7 ionic
The higher the END the more ionic the bond.

13. Try These:
Determine the electronegativity difference (END) for the following:
KCl
MgF2
Which is more ionic???

14. Try These: Answers
Determine the electronegativity difference (END) for the following:
KCl 3.2 – 0.8 = ionic
MgF – 1.3 = ionic
MgF2 is more ionic than KCl

15. Exceptions to END > 1.7 H = 2.2 END = 2.2 – 0.9 = 1.3
Metal Hydrides (metals bonded to hydrogen)
Example: sodium hydride
NaH Na = 0.9
H = 2.2
END = 2.2 – 0.9 = 1.3
Considered ionic because electrons are transferred

16. Properties of Ionic Solids
High melting points
Do not conduct electricity, in solid phase.
Crystalline structure.
As aqueous solutions or when heated ions are free to move so only then will they conduct electricity.
Examples: NaCl and MgO

17. Try These:
p. 3
Write Lewis dot structures for the following ionic compounds:
Barium oxide
Magnesium chloride
Aluminum oxide

18. Covalent Bonds – molecular compounds
A bond that exist between atoms caused by a sharing of electrons.
Usually occurs between two non-metals (neither is willing to lose electrons)
Electronegativity difference < 1.7
Example: H2
CO2

19. Two Hydrogen Atoms
The valence shells overlap and the electrons are shared making a more stable molecule.

20. REVIEW:
Diatomic molecules are atoms of the same element that covalently bond to meet the octet rule.
The following elements are diatomic:
H2, O2, F2 Br2, I2, N2, Cl2
Remember the trick to remember them: HOF BrINCl

21. Exceptions:
Hydrogen fluoride, HF, has an electronegativity difference of 4.0 – 2.2 = 1.8
This appears to be ionic but is really covalent.

22. Double Covalent Bonds
p. 3
Occurs when atoms share two pair of electrons (4 electrons).
Examples:
CO2
C2H4 O2

23. Triple Covalent Bond
p. 4
Occurs when atoms share three pair of electrons (6 electrons).
Examples: N2
HCN

24. Important Bonding Patterns
The following non-metals can bond in various ways:
H O N C Halogens
1 bond bonds bonds bonds 1 bond

25. Dash structures: (structural formulas)
Dashed structures only exist for covalently bonded molecules
Each shared pair of electrons is represented by a dashed line ( )
Single bond: one dash ( )
Double bond: two dashes ( )
Triple bond: three dashes ( )

26. Drawing Lewis Dot Diagrams for Moleculars
Add up all the valence electrons for an atom in the molecule.
The “central atom” is the one that is most electronegative or there will only be one of this atom.
Add the electrons until you reach the total remembering that all the atoms must obey the octet rule except for hydrogen which obeys the duet rule.
Also remember that one pair of electrons between two atoms is a single bond two is a double and three pairs is a triple bond. These are known as bonded or shared pairs of electrons.
Not every electron has to be bonded. We call these the lone or non-bonded pairs because they are “lonely”.

27. Lewis dot structures Symbol Valence electrons
Move electrons to satisfy octet/duet for all atoms involved
Try to write dashed structures for each….

28. Try These:
H2O
NH3
CCl4

29. Drawing Lewis structures for polyatomic ions.
Figure out the total number of electrons considering charge (add up all valence electrons from each atom and adjust according to the overall charge)
Example:
SO4 2-
OH-
PO4 3-

30. Drawing Lewis structures for polyatomic ions.
Figure out the total number of electrons considering charge (add up all valence electrons from each atom and adjust according to the overall charge)
Example:
SO electrons total
OH- 8 electrons total
PO electrons total

31. Lewis Dot Structures for polyatomic ions
Symbol
Valence electrons
Add or subtract electrons based on overall charge
Move electrons to satisfy octet/duet for all atoms involved
Brackets and charge

32. SO4 2- OH- PO4 3-

33. Drawing Lewis structures for polyatomic ions.-2
SO O
O S O O
OH-1
PO4 3- -1
O H -3 O
O P O

34. Coordinate Covalent Bonds
p. 6
When both electrons in a covalent bond are contributed by one of the two atoms or when an electron is donated by the environment.
Example: ammonium ion NH4 +1
sulfate SO4 2-

35. Resonance
p. 5
Two alternate Lewis structures are equivalent except for placement of electrons are called resonance forms.
Example:
To describe ozone O3 properly, the real molecule is an average of resonance structures.
Another example: NO3

36. Metallic Bonding
p. 5
Metals have a crystalline lattice structure in which positive ions are set in fixed positions, valence electrons are free to move belonging to the crystal as a whole not to individual atoms.
“sea of electrons”
Since electrons can move freely, metals are good conductors of heat and electricity.

37. Metallic bond

38. Polarity of Bonds (covalent only)
Ionic bonds by nature are always polar: END >1.7
A polar covalent bond occurs when two atoms do not share the electrons in a bond equally, resulting in a partial charge on each end of molecule.
This unequal sharing is a result of a difference in electronegativity.
END of : polar covalent bond
p. 6

39. Examples: HF , H2S , PI3
END:
* HF is an exception to ionic/covalent
bonding END rule

40. Non-Polar Covalent Bonds
Occurs when two atoms share electrons in a bond equally.
This is a result of little or no difference in electronegativity.
END of 0.0 –0.3
Examples: diatomics
F2 H2 N2
END’s are always zero!!!!!!!!!!!

41. Practice:
Indicate whether the following have polar covalent, non-polar covalent or ionic bonds.
CaO
NH3
CO2
PCl3
Br2 HF

42. Practice: Answers
Indicate whether the following have polar covalent, non-polar covalent or ionic bonds.
CaO ionic
NH3 polar covalent
CO2 polar covalent
PCl3 polar covalent
Br2 non-polar covalent
HF polar covalent (exception to 1.7)

43. Polarity of Molecules
p. 6
Polar molecule (dipole) : a molecule that has a slightly + end and a slightly – end.
In order for a molecule to be polar it must have polar covalent bonds and asymmetric shape .
asymmetric shapes: bent, pyramidal, linear without a central atom
Examples:
HF H2O NBr3

44. Linear without a central atom: asymmetrical
H Cl

45. Bent: asymmetrical O H H Must remember non-bonding electron pairs:
they still have a repulsion
force O
H H

47. pyramidal: asymmetrical
Must remember
non-bonding electron pairs:
they still have a repulsion
force N
H H H

49. Polar molecules
Examples:
HF H2O NBr3

50. Non-polar Molecules
p. 7
A molecule that does not have a slightly + and – end.
**Non-polar molecule with non-polar bonds
Examples: F2 , H2 , Br2
**Non-polar molecule with polar bonds can be non-polar if it has symmetry: the same on both sides or ends
symmetric shapes= tetrahedral, linear with a central atom, trigonal planar
examples: CH4 , CO2

51. Diatomics
All non-polar bonds in linear symmetrical : non-polar molecules.
Examples:
H H
F F

52. Tetrahedral: symmetrical
H C H

54. Linear with central atom: symmetrical
O = C = O

56. trigonal planar: symmetricalCl
Al Cl

59. Intermolecular forces-
p. 7
Definition:
A force of attraction between two molecules
Related to many physical properties

60. Types of Intermolecular Forces
Dipole-dipole Attraction
Hydrogen Bonding
Van Der Waals Dispersion Forces

61. Hydrogen Bonding
Attraction between the H of one molecule and the F, O, N of another molecule.
Strongest intermolecular force
Examples:
HF H2O NH3

62. Dipole-dipole Attraction
Attraction between polar molecules
Examples:
HCl
H2S

63. Van Der Waals Dispersion Forces
Attraction between non-polar molecules or between monoatomic atoms
Only attraction for non-polar substances
Examples:
F2 Ne

64. Van Der Waals Dispersion Forces
Van Der Waal Forces increase with:
Decreasing distance
Increasing mass
p. 8

65. Ionic Crystals
p. 8
Repeating particles in crystals are three-dimensional and formed by linking ion pairs.
Examples:
sodium chloride NaCl
Properties:
High melting/boiling points
Hard
Do not conduct electricity as solid
Very soluble (dissolve easily)
Does conduct as aqueous

66. A Collection of Crystals

67. Molecular Crystals p. 8 Repeating particles are electrically neutral.
Ex: sugar C6H12O6
Properties:
Low melting/boiling points
Usually below 300 C
High to low solubility
Poor conductor
Made of nonmetals

68. Covalent Network Solids
Repeating patterns are atoms covalently bonded
Ex: diamond, graphite, SiC
Properties:
Very high melting points
Very hard, abrasive
Insoluble

69. Allotropes
Different molecular forms of an element in the same physical state
Examples: diamond and graphite
Oxygen (O2) and ozone (O3)

70. Metallic Solid
Repeating particles in crystal are positive with “mobile valence electrons” moving about the crystal
Good conductor of heat and electricity
Examples: copper, gold, silver

72. P
Cl Cl Cl