1. Ionic and Covalent Bonding
Atoms rarely exist in nature, in their pure un- bonded 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 when atoms share
cations attract by valence electrons.
virtue of their
opposite charges
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 1st
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.
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. N

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.
Did you notice that there are 2 pairs and
2 singles? This means oxygen will form
2 bonds in most cases. Why? O
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.
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. Cl

14. .Ra. Examples Alkali Metals Li. Na. K. Rb. Fr.
Alkaline Earth .Be. .Mg. .Ca. .Sr. Metals
.Ra.
Cations they form:
Alkali Metals Li+ Na+ K+ Rb+ Fr+
Alkaline Earth Be Mg Ca Sr Metals 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 (+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,
covalent bonds can be classified
as nonpolar covalent bonds
(with electronegative
differences of 0 to 0.3) or
polar covalent bonds (with
electronegativity differences
above 0.3 to below 1.7)
Ionic to > 1.7
Polar
Covalent 1.7 to > .3
Covalant .3 to 0
Difference in
electronegativity

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
Put the electron dots between the 2 atoms,
Since that is where the bond will be! H H
The number of pairs of electrons in the vin of
The vin diagram shows the type of bond
Single 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 H-H
Triatomic Linear example: carbon dioxide O C O
Triatomic Bent example: Water H O
Pyramidal example: Ammonia N
Tetrahedral example: Carbon Tetrachloride H H H H C 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

27. Some More Covalent Bonding
Problem: Draw the covalent bonds in diatomic Nitrogen
1) Draw the Lewis structure for 2 Nitrogen atoms.
Finally circle 8 electrons for each nitrogen atom.
Since nitrogen
is in group 5 of the periodic table show 5 valence electrons for
each atom.
Notice that you must have 3 pairs or 6 electrons in the vin 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

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)
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 C Cl Cl

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 C Cl 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 C Cl Cl

45. 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.
Since there is only 1 carbon and there are 4 chlorine atoms
make the carbon atom the central atom. Cl C Cl 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 C Cl 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 there are 4 single bonds holding the molecule
together. 1 electron in each shared pair is from carbon and
1 is from chlorine.
Now circle an octet for each atom of chlorine and for the
carbon atom.
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

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
Next give 1 electron
From Al to each of
The Br atoms; notice
The change in charge
Of the atoms to form
Ions. - Al +1 Br + +1 - Br
The overall (-) and (+)
Charges balance for a
Net charge of 0