ionic bonding covalent bonding

1. Ionic and Covalent Bonding
Atoms rarely exist in nature, in their pure unbonded state. Why?
• Atoms by themselves have higher potential
energy being in an unbonded state than a
bonded state; therefore, they bond with other
atoms to attain a lower potential energy state.
• Atoms tend to seek a natural lower energy state.
• It is the valence electrons in any atom that
determines whether or not a particular atom will
bond with other atoms and in what ratio.

2. The Bond Mechanism
• A chemical bond is formed when the nuclei of
two atoms mutually attract valence electrons of
the other atom. This occurs by the attraction of
the positive nuclei and the negative valence
electrons.
• Remember for this to occur one atom has to
have a higher electronegativity than the other
atom. This higher electronegative atom will
attract the other atoms valence electrons, while
the other atom will have a natural tendency to
lose electrons. Some atoms can bond together
with do difference in electronegativity. Ex
Hydrogen-Hydrogen
• In all cases, atoms will attain a lower potential
energy and become stable like a noble gas.

3. Two types of chemical bonds
Ionic bonds occur . . .
Covalent bonds occur . . .
when anions and
cations attract by
virtue of their
opposite charges
when atoms share
valence electrons.
between metal and
between two non metals
non metals

4. How do Cations & Anions Form?
• Metals lose electrons and form positive cations.
• Non metals gain electrons and form negative
anions.
• Remember the value of the charge is equal to
the number of electrons lost or gained; the
magnitude of the charge will be positive if the
atoms loses electrons and negative if the atom
gains electrons.

5. Ionic Bonds
• The mechanism in an ionic bond is the
attraction of the metal cation (+) to the
nonmetal anion (-).
• The ratio of cations to anions is determined
by the cancellation of positive and negative
charges so that the overall charge of the
compound is neutral or zero.
• The name of the cation-anion complex in its
lowest whole number ratio is a formula unit
and not a molecule.

6. Lewis Structures
• Also known as electron dot notation.
• Lewis structures are representations of the
atom and its valence electrons. In a Lewis
structure the nucleus is represented by the
symbol of the atom and the valence
electrons are represented by dots placed
around the atomic symbol

7. Review of how to draw Lewis Dot
Diagrams for the atoms.
Let’s say we want to draw the Lewis structure for
Sodium (Na)
First we must find Na’s column number on the periodic
Table. We find it and all alkali metals are in group 1.
This means we need to show 1 dot which stands for the
1 valence electron of the alkali metals and thus for Na.
So we write down the symbol for sodium and then put
the dot anywhere by the symbol, top, side or bottom, it
does not matter.
Na

8. How about an Alkaline Earth Metal like
Calcium (Ca)?
Looking at the periodic table of elements, we find
that Ca is in group 2 so it will need 2 valence
electrons.
We will there for begin placing the first electron dot at
any position around the atom, (top, bottom or either side,
it does not matter). The 2nd electron is placed either
clockwise or counterclockwise in another position from
where we placed the first electron dot
Ca

9. Try Aluminum
It’s a group 3 metal so it needs 3 dots for 3
valence electrons.
Don’t forget to place a dot in any position around the
atomic symbol (Al) and then follow with the next two
dots, in 2 other positions around the atoms in a clockwise
or counterclockwise rotation.
Al

10. Try carbon on for size
Carbon is a group 4 element so we will
place 4 dots around carbon, 1 dot at the 1 st
position and the next three at the three other
positions around the symbol C in a
clockwise or counter clockwise rotation.
C
Notice that carbon has 4 unpaired electrons

11. Do the Lewis Structure for Nitrogen
Nitrogen is a group 5 element. It has 5
valence electrons. We will need to place 5
dots around the symbol N. Each dot
represents a valence electron.
N
Notice that nitrogen has 1 pair
and three single electrons. It is
the single electrons that will
become involved in a covalent
bond. Paired electrons can be
part of a covalent bond but that is
beyond the scope of this course.

12. Can you do oxygen?
You should have determined that since
oxygen is a group 6 element, you will need
to place 6 dots around the symbol O. Each
dot stands for a valence electron and all
group 6 elements have 6 valence electrons.
O
Did you notice that there are 2 pairs and
2 singles? This means oxygen will form
2 bonds in most cases. Why?
Because paired
electrons usually
aren’t involved in
bonding.

13. Let’s do a Halogen
Since the Halogens are in group 7, they all
have 7 valence electrons. This means
they will all have 7 dots around their
atomic symbol, 1 dot representing each
valence electron.
Cl
Notice that there are 3 pairs of
electrons and only 1 single. This
means a halogen will have only 1
bond because the pairs rarely are
involved in bonding.

14. Examples
. Na. K. Rb. Fr.
•
Alkaline Earth .Be. .Mg. .Ca. .Sr.
Metals
.Ra.
•
Alkali Metals
Li
Cations they form:
Alkali Metals
Alkaline Earth
Metals
Li+ Na+ K+ Rb+ Fr+
Be2+ Mg 2+ Ca 2+ Sr 2+
Ra 2+

15. Halogens Atoms/Ions
Notice how each halogen above has 7 valence electrons
and just needs 1 more valence electron to make an octet
as shown in the anions that are formed from halogens below.
In each case above there is 1 more electron in all
combined energy levels than there is protons in the nucleus.
This makes the new ions anions.

16. Ratios of Cations & Anions in Ionic
Compounds
• For Ca2+ and F1- the ratio will be:
CaF2 1(+2 Ca) + 2(-1 F) =0
• For Mg2+ and O2- the ratio will be:
MgO
1 (+2 Mg) + 1(-2 O) = 0
• For Al3+ and S2- the ratio will be:
• Al2S3 2(+3 Al) + 3(-2 S) = 0
• For determining these ratios, there is an
alternative method, called the crisscross
rule. Try looking this one up.

17. Differences in Electronegativities of
elements
• Its easy to see why metals and nonmetals form
ionic compounds; its because metals are
electropositive and non metals are
electronegative; they have large differences in
electronegativity.
• But bonding between 2 non metals results in
bonding between elements that don’t have a
large difference in electronegativity. This type
of bonding results in sharing of electrons and
creates covalent bonds which hold together
covalent compounds or molecular compounds.

18. Bond Character
If the differences in electronegativity of two
elements is greater than 1.7 then the bond is ionic.
You can look up electronegativity values.
If the difference is 1.7 or less then the bond is
said to be covalent; however,
Difference in
electronegativity
covalent bonds can be classified
as nonpolar covalent bonds
Ionic
3.3 to > 1.7
(with electronegative
Polar
differences of 0 to 0.3) or
Covalent 1.7 to > .3
polar covalent bonds (with
Covalant .3 to 0
electronegativity differences
above 0.3 to below 1.7)

19. Covalent Bonds
• The idea of sharing electrons between
two non-metals, in order to obtain an
octet is what covalent bonding is all
about.

20. Lewis structures showing Covalent
Bonds
In a Hydrogen-Hydrogen bond we begin by
showing each atom’s electron dot structure and
drawing them side by side.
Then we draw a circle around each hydrogen and include
in that circle two electrons, since that is what each
hydrogen needs for it to become stable like helium
.
Let’s see that in Click by click animation
H
H
Single Bond
Put the electron dots between the 2 atoms,
Since that is where the bond will be!
The number of pairs of electrons in the vin of
The vin diagram shows the type of bond

21. Single, Double & Triple Bonds
•
•
•
•
•
1 shared pair is a single bond
2 shared pairs is a double bond
3 shared pairs is a triple bond
See examples of these Lewis Structures
Single bonds have the longest bond length
have the weakest bond energy for
covalent bonds; triple bonds have the
shortest bond length and have the
strongest bond energy for covalent bonds.

22. Coordinate Covalent Bond
• A coordinate covalent bond occurs when
one nonmetal donates both electrons in
the bond to the shared pair.
• Look for a coordinate covalent bond
example on your own on the internet or in
your book

23. VSEPR
Valence Shell Electron Pair Repulsion Theory
This theory states that the 4 pairs of valence
electrons, around a bonded atom, all repel each
other such that they position themselves equal
distances, and as far as possible from each other; so
it is the positions of these electron pairs affects the
shapes of molecules.

24. Understand these basic shapes
Linear
example: any diatomic molecule
Triatomic Linear
example: carbon dioxide
Triatomic Bent
example: Water
H
H-H
O C O
O
H
Pyramidal
example: Ammonia
H
N
H
H
Tetrahedral
Cl
example: Carbon Tetrachloride
C
Cl
Cl
Cl

25. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Nitrogen
1)
Draw the Lewis structure for 2 Nitrogen atoms. Since nitrogen
is in group 5 of the periodic table show 5 valence electrons for
each atom.
Next place all the single electrons (there are 3 singles and one
Pair) between both atoms, because they need to share 3 each,
to make an octet (8) of valence electrons
N
N

26. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Nitrogen
1)
Draw the Lewis structure for 2 Nitrogen atoms. Since nitrogen
is in group 5 of the periodic table show 5 valence electrons for
each atom.
Next place all the single electrons (there are 3 singles and one
Pair) between both atoms, because they need to share 3 each,
to make an octet (8) of valence electrons
N
N

27. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Nitrogen
1) Finally circle 8 electrons for each nitrogen atom.Since nitrogen
Draw the Lewis structure for 2 Nitrogen atoms.
is in group 5 of the periodic table showelectrons in the vin for
Notice that you must have 3 pairs or 6 5 valence electrons
each atom.
diagram
Next place all the single electrons (there are 3 singles and one
Pair) between both atoms, because they need to share 3 each,
to make an octet (8) of valence electrons
N
N

28. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Fluorine
1)
Draw the Lewis structure for 2 Fluorine atoms. Since Fluorine
is in group 7 of the periodic table show 7 valence electrons for
each atom.
Next place the single electrons between both atoms,
(there is 1 single and three pairs). Place the single electrons
between the two atoms because that is where the bond will be
and because they need to share 1 each, to make a pair and to
make an octet (8) of valence electrons
F
F

29. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Fluorine
1)
Draw the Lewis structure for 2 Fluorine atoms. Since Fluorine
is in group 7 of the periodic table show 7 valence electrons for
each atom.
Next place the single electrons between both atoms,
(there is 1 single and three pairs). Place the single electrons
between the two atoms because that is where the bond will be
and because they need to share 1 each, to make a pair and to
make an octet (8) of valence electrons
F
F

30. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Fluorine
Finally circle 8 electrons for each fluorine atom.
Notice that you must have 1 pair or 2 electrons in the vin
diagram
F
F

31. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
1)
Draw the Lewis structures for 1 carbon & 2 oxygen atoms.
Place carbon as the central atom since there is only 1 carbon
and there are 2 oxygens.
Since carbon is in group 4, show 4 valence electrons, and since
oxygen is in group 6 show 6 valence electrons for each atom.
Next place the single electrons between the oxygen and carbon
atoms, (there are 4 singles in carbon and 2 singles in oxygen).
Place the single electrons between the atoms because that is
where the bonds will be and because all atoms need to make pairs
Which will be shared to make an octet (8) of valence electrons.
O
C
O

32. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
1)
Draw the Lewis structures for 1 carbon & 2 oxygen atoms.
Place carbon as the central atom since there is only 1 carbon
and there are 2 oxygens.
Since carbon is in group 4, show 4 valence electrons, and since
oxygen is in group 6 show 6 valence electrons for each atom.
Next place the single electrons between the oxygen and carbon
atoms, (there are 4 singles in carbon and 2 singles in oxygen).
Place the single electrons between the atoms because that is
where the bonds will be and because all atoms need to make pairs
Which will be shared to make an octet (8) of valence electrons.
O
C
O

33. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
1)
Draw the Lewis structures for 1 carbon & 2 oxygen atoms.
Place carbon as the central atom since there is only 1 carbon
and there are 2 oxygens.
Since carbon is in group 4, show 4 valence electrons, and since
oxygen is in group 6 show 6 valence electrons for each atom.
Next place the single electrons between the oxygen and carbon
atoms, (there are 4 singles in carbon and 2 singles in oxygen).
Place the single electrons between the atoms because that is
where the bonds will be and because all atoms need to make pairs
Which will be shared to make an octet (8) of valence electrons.
O
C
O

34. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
1)
Draw the Lewis structures for 1 carbon & 2 oxygen atoms.
Place carbon as the central atom since there is only 1 carbon
and there are 2 oxygens.
Since carbon is in group 4, show 4 valence electrons, and since
oxygen is in group 6 show 6 valence electrons for each atom.
Next place the single electrons between the oxygen and carbon
atoms, (there are 4 singles in carbon and 2 singles in oxygen).
Place the single electrons between the atoms because that is
where the bonds will be and because all atoms need to make pairs
Which will be shared to make an octet (8) of valence electrons.
O
C
O

35. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
1)
Draw the Lewis structures for 1 carbon & 2 oxygen atoms.
Place carbon as the central atom since there is only 1 carbon
and there are 2 oxygens.
Since carbon is in group 4, show 4 valence electrons, and since
oxygen is in group 6 show 6 valence electrons for each atom.
Next place the single electrons between the oxygen and carbon
atoms, (there are 4 singles in carbon and 2 singles in oxygen).
Place the single electrons between the atoms because that is
where the bonds will be and because all atoms need to make pairs
Which will be shared to make an octet (8) of valence electrons.
O
C
O

36. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
Now circle an octect for each atom.
O
C
O

37. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
Now circle an octect for each atom.
O
C
O

38. Problem: Draw the covalent bonds for a molecule of Carbon Dioxide
Now circle an octect for each atom.
Notice that there are 2 pairs (4) of electrons in the vin diagram
between each oxygen and the carbon atom.
This means that there are 2 double bonds holding the carbon
dioxide molecule together.
O
C
O

39. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
1) Draw the Lewis structures for 1 carbon & 4
chlorine atoms.Carbon is in group 4 so it has 4
valence electrons and chlorine is in group 7, so it has
7 valence electrons.
C
Cl
Cl
Cl
Cl

40. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom.
C
Cl
Cl
Cl
Cl

41. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom.
Cl
Cl
Cl
Cl
C

42. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom.
Cl
Cl
C
Cl
Cl

43. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom.
Cl
Cl
Cl
C
Cl

44. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom.
Cl
Cl
Cl
C
Cl

45. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Since there is only 1 carbon and there are 4 chlorine atoms
Now place each single electron from chlorine between
make the carbon atom the that each single
carbon and the chlorine so central atom. with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

46. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Now place each single electron from chlorine between
carbon and the chlorine so that each single with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

47. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Now place each single electron from chlorine between
carbon and the chlorine so that each single with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

48. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Now place each single electron from chlorine between
carbon and the chlorine so that each single with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

49. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Now place each single electron from chlorine between
carbon and the chlorine so that each single with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

50. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Now place each single electron from chlorine between
carbon and the chlorine so that each single with chlorine
and carbon can make a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

51. Draw the Lewis Structure showing the covalent
bonding in Carbon Tetrachloride (CCl4)
Notice that each single singleatom of chlorine and for the
Now circle there arefor electron from chlorine between
place an octet 4 each bonds holding the molecule
carbon atom. chlorine so shared pair is from carbon and
and the
together. 1 electron in eachthat each single with chlorine
and carbon can make
1 is from chlorine. a pair. Place the singles making pairs
between the atoms because that is where the bond will be.
Cl
Cl
C
Cl
Cl

52. Lewis Structures in Ionic Bonding
Remember in ionic bonds the metal gives
its electrons to the non-metal, and the
non-metal takes electrons from the metal
so that each is able to have an octect of
electrons in its outer most energy level.
Q: I don’t understand, how does a metal get 8 electrons
in its valence if it gives electrons away?
A: When it gives its few electrons in its valence away
a new energy level or layer of 8 electrons lies underneath
the electrons that were given away.

53. Lets show the Lewis structure for
NaCl (Sodium Chloride)
Na
+
Cl
+
-
Na Cl

54. Ionic Bonding for Aluminum
Bromide (AlBr3)
1st Draw 1 Al & 3 Br
Electron Dot Structures
-
Br
+1
Al
+1
+
-
Br
+1
Br
-
Next give 1 electron
From Al to each of
The Br atoms; notice
The change in charge
Of the atoms to form
Ions.
The overall (-) and (+)
Charges balance for a
Net charge of 0