1. Ionic and Covalent Bonding Electron and Lewis Dot Structures
James Hutchison

2. Valence Electrons are…?
The electrons responsible for the chemical properties of atoms, and are those in the outer energy level.
Valence electrons - The s and p electrons in the outer energy level
the highest occupied energy level
Core electrons – are those in the energy levels below.

3. Keeping Track of Electrons
Atoms in the same column...
Have the same outer electron configuration.
Have the same valence electrons.
The number of valence electrons are easily determined. It is the group number for a representative element
Group 2A: Be, Mg, Ca, etc.
have 2 valence electrons

4. Electron Dot diagrams are…
A way of showing & keeping track of valence electrons.
How to write them?
Write the symbol - it represents the nucleus and inner (core) electrons
Put one dot for each valence electron (8 maximum)
They don’t pair up until they have to (Hund’s rule) X

5. The Electron Dot diagram for Nitrogen
Nitrogen has 5 valence electrons to show.
First we write the symbol. N
Then add 1 electron at a time to each side.
Now they are forced to pair up.
We have now written the electron dot diagram for Nitrogen.

6. The Octet Rule
In Chapter 6, we learned that noble gases are unreactive in chemical reactions
In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules
The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable
Each noble gas (except He, which has 2) has 8 electrons in the outer level

7. Formation of Cations
Metals lose electrons to attain a noble gas configuration.
They make positive ions (cations)
If we look at the electron configuration, it makes sense to lose electrons:
Na 1s22s22p63s1 1 valence electron
Na1+ 1s22s22p6 This is a noble gas configuration with 8 electrons in the outer level.

8. Electron Dots For Cations
Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

9. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons Ca

10. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons
Forming positive ions
Ca2+
This is named the “calcium ion”.
NO DOTS are now shown for the cation.

11. Electron Dots For Cations
Let’s do Scandium, #21
The electron configuration is: 1s22s22p63s23p64s23d1
Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+)
Sc = Sc2+ Sc
= Sc3+
Scandium (II) ion
Scandium (III) ion

12. Electron Dots For Cations
Let’s do Silver, element #47
Predicted configuration is: 1s22s22p63s23p64s23d104p65s24d9
Actual configuration is: 1s22s22p63s23p64s23d104p65s14d10
Ag = Ag1+ (can’t lose any more, charges of 3+ or greater are uncommon)

13. Electron Dots For Cations
Silver did the best job it could, but it did not achieve a true Noble Gas configuration
Instead, it is called a “pseudo-noble gas configuration”

14. Electron Configurations: Anions
Nonmetals gain electrons to attain noble gas configuration.
They make negative ions (anions)
S = 1s22s22p63s23p4 = 6 valence electrons
S2- = 1s22s22p63s23p6 = noble gas configuration.
Halide ions are ions from chlorine or other halogens that gain electrons

15. Electron Dots For Anions
Nonmetals will have many valence electrons (usually 5 or more)
They will gain electrons to fill outer shell. 3- P
(This is called the “phosphide ion”, and should show dots)

16. Stable Electron Configurations
All atoms react to try and achieve a noble gas configuration.
Noble gases have 2 s and 6 p electrons.
8 valence electrons = already stable!
This is the octet rule (8 in the outer level is particularly stable). Ar

17. Ionic Bonding
Anions and cations are held together by opposite charges (+ and -)
Ionic compounds are called salts.
Simplest ratio of elements in an ionic compound is called the formula unit.
The bond is formed through the transfer of electrons (lose and gain)
Electrons are transferred to achieve noble gas configuration.

18. Ionic Compounds
Also called SALTS
Made from: a CATION with an ANION (or literally from a metal combining with a nonmetal)

19. Ionic Bonding Na Cl
The metal (sodium) tends to lose its one electron from the outer level.
The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

20. Ionic Bonding
Na+
Cl -
Note: Remember that NO DOTS are now shown for the cation!

21. Ionic Bonding
Lets do an example by combining calcium and phosphorus: Ca P
All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable).

22. Ionic Bonding Ca P

23. Ionic Bonding
Ca2+ P

24. Ionic Bonding
Ca2+ P Ca

25. Ionic Bonding
Ca2+
P 3- Ca

26. Ionic Bonding
Ca2+
P 3- Ca P

27. Ionic Bonding
Ca2+
P 3-
Ca2+ P

28. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

29. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

30. Ionic Bonding
Ca2+
Ca2+
P 3-
Ca2+
P 3-

31. = Ca3P2 Ionic Bonding Formula Unit
This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance.
For an ionic compound, the smallest representative particle is called a: Formula Unit

32. Properties of Ionic Compounds
Crystalline solids - a regular repeating arrangement of ions in the solid: Fig. 7.9, page 197
Ions are strongly bonded together.
Structure is rigid.
High melting points
Coordination number- number of ions of opposite charge surrounding it

33. NaCl CsCl TiO2 - Page 198 Coordination Numbers:
Both the sodium and chlorine have 6
NaCl
Both the cesium and chlorine have 8
CsCl
Each titanium has 6, and each oxygen has 3
TiO2

34. Do they Conduct?
Conducting electricity means allowing charges to move.
In a solid, the ions are locked in place.
Ionic solids are insulators.
When melted, the ions can move around.
Melted ionic compounds conduct.
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to move in aqueous solutions)

35. - Page 198
The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

36. Metallic Bonds are… How metal atoms are held together in the solid.
Metals hold on to their valence electrons very weakly.
Think of them as positive ions (cations) floating in a sea of electrons: Fig. 7.12, p.201

37. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

38. Metals are Malleable
Hammered into shape (bend).
Also ductile - drawn into wires.
Both malleability and ductility explained in terms of the mobility of the valence electrons

39. Due to the mobility of the valence electrons, metals have:
- Page 201
Due to the mobility of the valence electrons, metals have:
Notice that the ionic crystal breaks due to ion repulsion!
1) Ductility
2) Malleability
and

40. Malleable +
Force

41. Malleable + + + + Force + + + + + + + +
Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. + + + +
Force + + + + + + + +

42. Ionic solids are brittle
Force + -

43. Ionic solids are brittle
Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. + -
Force + - + - + -

44. Covalent Bonds
The word covalent is a combination of the prefix co- (from Latin com, meaning “with” or “together”), and the verb valere, meaning “to be strong”.
Two electrons shared together have the strength to hold two atoms together in a bond.

45. Molecules Many elements found in nature are in the form of molecules:
a neutral group of atoms joined together by covalent bonds.
For example, air contains oxygen molecules, consisting of two oxygen atoms joined covalently
Called a “diatomic molecule” (O2)

46. The nuclei repel each other, since they both have a positive charge (like charges repel).
How does H2 form?
(diatomic hydrogen molecule) + + + +

47. How does H2 form? But, the nuclei are attracted to the electrons
They share the electrons, and this is called a “covalent bond”, and involves only NONMETALS! + +

48. Covalent bonds Nonmetals hold on to their valence electrons.
They can’t give away electrons to bond.
But still want noble gas configuration.
Get it by sharing valence electrons with each other = covalent bonding
By sharing, both atoms get to count the electrons toward a noble gas configuration.

49. Covalent bonding
Fluorine has seven valence electrons (but would like to have 8) F

50. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven F F

51. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

52. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

53. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

54. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

55. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

56. F F Covalent bonding …both end with full orbitals
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F

57. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

58. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

59. Molecular Compounds
Compounds that are bonded covalently (like in water, or carbon dioxide) are called molecular compounds
Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds – this is not as strong a bond as ionic

60. Molecular Compounds
Thus, molecular compounds tend to be gases or liquids at room temperature
Ionic compounds were solids
A molecular compound has a molecular formula:
Shows how many atoms of each element a molecule contains

61. Molecular Compounds The formula for water is written as H2O
The subscript “2” behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omitted
Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

62. A Single Covalent Bond is...
A sharing of two valence electrons.
Only nonmetals and hydrogen.
Different from an ionic bond because they actually form molecules.
Two specific atoms are joined.
In an ionic solid, you can’t tell which atom the electrons moved from or to

63. H O Water Each hydrogen has 1 valence electron
- Each hydrogen wants 1 more
The oxygen has 6 valence electrons
- The oxygen wants 2 more
They share to make each other complete H O

64. H O Water Put the pieces together The first hydrogen is happy
The oxygen still needs one more H O

65. H O H Water So, a second hydrogen attaches
Every atom has full energy levels
Note the two “unshared” pairs of electrons H O H

66. Multiple Bonds
Sometimes atoms share more than one pair of valence electrons.
A double bond is when atoms share two pairs of electrons (4 total)
A triple bond is when atoms share three pairs of electrons (6 total)
Table 8.1, p Know these 7 elements as diatomic:
Br2 I2 N2 Cl2 H2 O2 F2
What’s the deal with the oxygen dot diagram?

67. Dot diagram for Carbon dioxide
CO2 - Carbon is central atom ( more metallic )
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 more C O

68. Carbon dioxide
Attaching 1 oxygen leaves the oxygen 1 short, and the carbon 3 short C O

69. Carbon dioxide
Attaching the second oxygen leaves both of the oxygen 1 short, and the carbon 2 short O C O

70. Carbon dioxide
The only solution is to share more O C O

71. Carbon dioxide
The only solution is to share more O C O

72. Carbon dioxide
The only solution is to share more O C O

73. Carbon dioxide
The only solution is to share more O C O

74. Carbon dioxide
The only solution is to share more O C O

75. Carbon dioxide
The only solution is to share more O C O

76. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond O C O

77. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

78. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

79. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

80. How to draw them? Add up all the valence electrons.
Count up the total number of electrons to make all atoms happy.
Subtract; then Divide by 2
Tells you how many bonds to draw
Fill in the rest of the valence electrons to fill atoms up.

81. N H Example NH3, which is ammonia
N – central atom; has 5 valence electrons, wants 8
H - has 1 (x3) valence electrons, wants 2 (x3)
NH3 has 5+3 = 8
NH3 wants 8+6 = 14
(14-8)/2= 3 bonds
4 atoms with 3 bonds N H

82. H H N H Examples Draw in the bonds; start with singles
All 8 electrons are accounted for
Everything is full – done with this one. H H N H

83. Example: HCN HCN: C is central atom
N - has 5 valence electrons, wants 8
C - has 4 valence electrons, wants 8
H - has 1 valence electron, wants 2
HCN has = 10
HCN wants = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds – this will require multiple bonds - not to H however

84. H C N HCN Put single bond between each atom Need to add 2 more bonds
Must go between C and N (Hydrogen is full) H C N

85. H C N HCN Put in single bonds Needs 2 more bonds
Must go between C and N, not the H
Uses 8 electrons – need 2 more to equal the 10 it has H C N

86. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add
Must go on the N to fill its octet H C N

87. Another way of indicating bonds
Often use a line to indicate a bond
Called a structural formula
Each line is 2 valence electrons H O H H O H =

88. Other Structural Examples
H C N H
C O H

89. A Coordinate Covalent Bond...
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO) is a good example:
Both the carbon and oxygen give another single electron to share O C

90. C O Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO) is a good example:
Oxygen gives both of these electrons, since it has no more singles to share.
This carbon electron moves to make a pair with the other single. C O

91. C O Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO)
The coordinate covalent bond is shown with an arrow as: C O
C O

92. Molecular Orbitals are...
The model for covalent bonding assumes the orbitals are those of the individual atoms = atomic orbital
Orbitals that apply to the overall molecule, due to atomic orbital overlap are the molecular orbitals
A bonding orbital is a molecular orbital that can be occupied by two electrons of a covalent bond

93. VSEPR: stands for... Valence Shell Electron Pair Repulsion
Predicts the three dimensional shape of molecules.
The name tells you the theory:
Valence shell = outside electrons.
Electron Pair repulsion = electron pairs try to get as far away as possible from each other.
Can determine the angles of bonds.

94. VSEPR
Based on the number of pairs of valence electrons, both bonded and unbonded.
Unbonded pair also called lone pair.
CH4 - draw the structural formula
Has 4 + 4(1) = 8
wants 8 + 4(2) = 16
(16-8)/2 = 4 bonds

95. VSEPR for methane (a gas):
Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
The furthest they can get away is 109.5º H H C H H
This 2-dimensional drawing does not show a true representation of the chemical arrangement.

96. H C H H H 4 atoms bonded Basic shape is tetrahedral.
A pyramid with a triangular base.
Same shape for everything with 4 pairs. H
109.5º C H H H

97. - Page 232
Methane has an angle of 109.5o, called tetrahedral
Ammonia has an angle of 107o, called pyramidal
Note the unshared pair that is repulsion for other electrons.

98. Bond Polarity Covalent bonding means shared electrons
but, do they share equally?
Electrons are pulled, as in a tug-of-war, between the atoms nuclei
In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond

99. Bond Polarity
When two different atoms bond covalently, there is an unequal sharing
the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge
called a polar covalent bond, or simply polar bond.

100. Electronegativity?
The ability of an atom in a molecule to attract shared electrons to itself.
Linus Pauling

101. Table of Electronegativities
Higher electronegativity

102. Bond Polarity Refer to Handout Consider HCl
H = electronegativity of 2.1
Cl = electronegativity of 3.0
the bond is polar
the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

103. Bond Polarity d+ d- d+ and d-
Only partial charges, much less than a true 1+ or 1- as in ionic bond
Written as:
H Cl
the positive and minus signs (with the lower case delta: ) denote partial charges.
d+ d-
d+ and d-

104. Bond Polarity H Cl Can also be shown:
the arrow points to the more electronegative atom.
Table 8.3, p.238 shows how the electronegativity can also indicate the type of bond that tends to form
H Cl

105. Polar molecules Sample Problem 8.3, p.239
A polar bond tends to make the entire molecule “polar”
areas of “difference”
HCl has polar bonds, thus is a polar molecule.
A molecule that has two poles is called dipole, like HCl

106. Polar molecules
The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape
carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

107. Polar molecules
The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape
water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e- away from H = very polar!

108. Polar molecules
When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.

109. Attractions between molecules
They are what make solid and liquid molecular compounds possible.
The weakest are called van der Waal’s forces - there are two kinds:
#1. Dispersion forces
weakest of all, caused by motion of e-
increases as # e- increases
halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A

110. #2. Dipole interactions
Occurs when polar molecules are attracted to each other.
2. Dipole interaction happens in water
Figure 8.25, page 240
positive region of one molecule attracts the negative region of another molecule.

111. #2. Dipole interactions H F d+ d- H F d+ d-
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract, but not completely hooked like in ionic solids.
H F
d+ d-
H F
d+ d-

112. #2. Dipole Interactions
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-

113. #3. Hydrogen bonding
…is the attractive force caused by hydrogen bonded to N, O, F, or Cl
N, O, F, and Cl are very electronegative, so this is a very strong dipole.
And, the hydrogen shares with the lone pair in the molecule next to it.
This is the strongest of the intermolecular forces.

114. Order of Intermolecular attraction strengths
Dispersion forces are the weakest
A little stronger are the dipole interactions
The strongest is the hydrogen bonding
All of these are weaker than ionic bonds

115. #3. Hydrogen bonding defined:
When a hydrogen atom is: a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom.
The hydrogen is left very electron deficient (it only had 1 to start with!) thus it shares with something nearby
Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!

116. Hydrogen Bonding (Shown in water)
This hydrogen is bonded covalently to: 1) the highly negative oxygen, and 2) a nearby unshared pair.

117. Hydrogen bonding allows H2O to be a liquid at room conditions.

118. Attractions and properties
Why are some chemicals gases, some liquids, some solids?
Depends on the type of bonding!
Network solids – solids in which all the atoms are covalently bonded to each other

119. Attractions and properties
Network solids melt at very high temperatures, or not at all (decomposes)
Diamond does not really melt, but vaporizes to a gas at 3500 oC and beyond
SiC, used in grinding, has a melting point of about 2700 oC

120. Covalent Network Compounds
Some covalently bonded substances DO NOT form discrete molecules.
Diamond, a network of covalently bonded carbon atoms
Graphite, a network of covalently bonded carbon atoms