1. The covalent bond
Chapter VIII

2. Why do atoms bond?
When atoms form bonds they don’t do it just for the sake of doing it there has to be a reason.
Atoms want to be like the noble gases
They want stability
They want low potential energy
When atoms form bonds they are getting these things from the union.
They obtain noble gas notation
They have full outer shells
or they have a lower potential energy
Making them more stable

3. What is a covalent bond? Atoms share a valence electron.
Different than an ionic which an electron is given up of accepted
When atoms share electrons in a covalent bond a molecule is formed.
In these molecules the shared electron is considered to belong to both electrons highest energy level
Ionic only the accepter gets credit for this electron.
These bonds usually happen with non-metals
Ionic is a metal and a non metal usually
The atoms in these bonds tend to be close together on the periodic table.
Ionic tend to be far apart

4. How are covalent bonds formed
Lets look a Fluorine
Fluorine has seven valence electrons
If two fluorine atoms approach each other several forces act.
Two repulsive forces act on the atoms
One form like-charged electrons
One from like-charged protons
A force of attraction also acts
One atoms protons attractions the other atoms electrons
As the atoms get closer the attraction of the protons to the electrons increases till it reaches a maximum.
When this maximum is achieved the atoms covalently bond.
However if the atoms move closer
The attraction force is less then the repulsive force
The most stable position for these atoms is when the attraction force exceed the net repulsion.

5. Looks like this

6. The single covalent bond
In the fluorine atom a single covalent bond is formed
This can be represented in a Lewis structure with a line or vertical dots
Again they do this to be at low energy.

7. Group 17 and single bonds These atoms have seven valence electrons
They will form one single bond to they can achieve a full octet
They will form these bonds with other non metals such as carbon
Some of these will form bonds with themselves
F2, Cl2

8. Group 16 and single bonds These atoms have six valence electrons
They will share two electrons forming to covalent bonds to form an octet
For example oxygen will form a covalent bond with two hydrogen to make H2O
In this example all of the atoms have noble gas configurations and are stable.

9. Group 15 and single bonds
These atoms have 5 valence electrons and are able to form three single covalent bonds
When they form these bonds the complete the octet and are stable
Nitrogen for example is able to bond with other non metals to form new molecules
NH3 (ammonia), NF3, NCl3, NBr3

10. Group 14 and single bonds
These atoms have 4 valence electrons and therefore need 4 electrons to complete the octet
They can make 4 single bonds
Carbon for example bond with everything it seems because it needs to make the bonds
CH4 (methane)

11. The Sigma (σ) bond The sigma bond is just a single covalent bond
These bonds overlap orbitals between atoms
S-S overlap
S-P overlap
P-P overlap
This tells us where the electron is most likely to be found.

12. Multiple covalent bonds
In some molecules, atoms have noble-gas configurations when they share more than one pair of electrons with on or more atoms.
Sharing multiple pairs of electron forms multiple covalent bonds
Double bonds and triple bonds are examples of multiple bonds.

13. Multiple covalent bond types
Double bonds
Two pairs of electrons are shared between atoms
Example is O2
each Oxygen has six valence so needs to share two pairs to reach an octet.
Triple bonds
Three pairs of electrons are shared between atoms
N2 is an example

14. The Pi bond (a fat kids dream)
A multiple covalent bond consists of one sigma bond and at least on pi bond
The pi bond (π) forms when parallel orbitals overlap and share electrons
The shared electron pair a pi bond occupies the space above and below the line that represents where the two atoms are joined together
Any multiple covalent bond has both sigma and pi bonds.
Double one sigma and one pi
Triple one sigma and two pi bonds

15. Pi bond

16. Bond strength of covalent bonds
Bonds between different atoms have different strengths based upon multiple factors.
Distance between the nuclei of the bonding atoms is the biggest factor
The shorter the distance the stronger the bond
Multiple bonds tend to have shorter bond lengths making them stronger bonds.

17. Bonds and energy
An energy change occurs when a bond between atoms in a molecule forms or breaks.
Energy is released (exothermic) when bonds are formed
Energy must be added (endothermic) to break a bond.
The amount of energy required to break a specific covalent bond is called bond-dissociation energy and is always a positive value.
Bond dissociation energy also indicates the strength of the bond
Remember bond length is inversely related to bond energy
The smaller the bond length the higher the energy.
To determine the potential energy of a molecule all one does is add up the bond dissociation energies for all the bonds in said molecule.

18. Energy in vs. Energy Out
Ok so we know that when a bond is formed it releases energy and when a bond is broken it requires or takes in energy, but what happens in a chemical reaction
Remember that in a chemical reaction bonds are being broken and made most of the time
To figure out total energy one must calculate the difference between the energy taken in vs the energy given out
In the energy taken in is more than its an endothermic reaction
Think of instant ice packs- take in energy that’s what make them cold
If the energy give out is more than its an exothermic reaction
Think of hand warmers or MREs or camping food
End 8-1

19. Naming molecules
We learned how to name ionic compounds already. Covalent not much harder.
Binary compounds
1. the first element in the formula is always named first, using the entire element name.
2. the second element in the formula is named using its root and adding the suffix –ide
3. prefixes are used to indicate the number of atoms of each element that present in the compound
Never use mono for the first element
If double vowels dropped on one
Not monooxide
Monoxide

20. Common prefixes Number of atoms Prefix 1 Mono- 6 Hexa- 2 Di- 7 Hepta-3
Tri- 8
Octa- 4
Tetra- 9
Nona- 5
Penta- 10
Deca-

21. Name these CO₂ CO P₂O₅ N₂O SiO₂ CBr₄ SO₃ PBr₅ ICl₃ NI₃ N₂O₃ N₂O₄ SO₂
As₂O₅
PCl₃
CCl₄
H2O
SF₆
XeO₃
P₄S₃

22. Naming acids
If a compound produces hydrogen ions (H+) in solution, it is an acid.
Two common types
Binary
Oxyacids
Naming binary acids
1. the first word has the prefix hydro- to name the hydrogen part of the compound the rest of the first word consists of a form of the root of the second element plus the suffix –ic
2. the second word is always acid
HCl- hydrochloric acid

23. Naming Oxyacids
An acid that contains both a hydrogen atom and an oxyanion is referred to as an oxyacid.
Oxyanion- polyatomic ion containing on or more oxygen atoms
1. first, identify the oxyanion present. The first word of an oxyacid's name consists of the root of the oxyanion and the prefix per- or hypo- if it is part of the name, and a suffix. If the oxyanion’s name ends with the suffix –ate, replace if with the suffix –ic. If the name ends with –ite, replace it with –ous.
2. the second word in always acid

24. Lets name some acids HI
HClO3
HClO2
H2SO4

25. Go the other way
Now that we know how to name the acids and other compounds we can write the formulas from the names
Just remember the naming rules and its pretty easy

26. Lets try some Silver monochloride Dihydrogen monoxide
Chlorine triflouride
Diphosphorus trioxide
Disulfur decafluoride
End 8-2

28. Molecular structures Things we need to know Lewis structures Resonance
Structural formula
Resonance
Molecule that can’t be shown by one Lewis structure
Coordinate covalent bond
End 8.3

29. Molecular geometry

30. More shapes
BENT

31. Yet more shapes
Square Antiprismatic
T-Shaped

32. Last one

33. The VSEPR Model
VSEPR
Valence Shell Electron Pair Repulsion
This model is used to determine the shape of a molecule once the Lewis structure of that molecule has been drawn.
This model is based on an arrangement that minimizes the repulsion of shared and unshared electron pairs
Molecule shape gives molecules many of its properties.
Electron densities in overlapping orbitals determine molecular shape.

34. Bond angles
Electron pairs in a molecule arrange so that they minimize interaction between them
This creates a natural shape of molecules
The unpaired electrons take up slightly more space then the shared electron pairs.
So shared bonding orbitals are pushed together by unshared pairs.

36. Why does shape matter?
Molecular shape determines if different molecules are able to react.
If the shape isn’t right then the molecules cant get close enough to react.
Think about food
Food molecules have different shapes
The different shape give us different tastes.
If they all had the same shape everything would truly taste like chicken.

37. Hybridization
This means the same as it does in biology and engineering
the combination of two different things to form something new but has qualities of the old
Orbitals can do this also
Lets look at carbon
Carbon has 4 valence electrons
2 in the s orbital
2 in the p orbital
But when carbon bonds with hydrogen the orbitals change into a hybrid called the sp3 orbital
Its an sp3 because it forms from one s and 3 p orbitals to form 4 new orbitals
Lone pairs also occupy hybrid orbitals.
End 8.4

38. Electronegativity and Polarity
The measure of the tendency of an atom to accept an electron
No noble gases
Don’t bond
Xeon will bond with certain elements
Increase from bottom left to upper right in general
These number were assigned not measure
Refer to your flipbook.

39. Using electronegativity
Electronegativity can be used to determine what type of bond will be formed.
To do this you calculate the difference in the electronegativity of the elements that are bonding together.
Take this difference and see where it falls
The problem with bonds is that they are never truly covalent or ionic they are some of both
Which ever one is more we call it that.

40. Lets do one or two Look at Oxygen gas O2
Look up the electronegativity of oxygen and find it to be 3.44
So = zero
Look at the chart
So Oxygen is nonpolar covalent
Hint any element bonded only with itself is going to equal zero and be nonpolar covalent.

41. Cont…. Now lets look at hydrochloric acid What is HCl? HCL
Look up the electronegativity
Cl= H=2.20
Do the math
= .96
Look at the chat
What is HCl?

42. So what's a Polar covalent bond
Think about a tug of war between John Cena and Paris Hilton.
This is kind of the idea being a polar covalent bond
One atom is going to “win” the electron more often than not.
Electrons have a negative charge so that atom has a negative charge
The losing atom end up with a positive charge
The charge is represented with the Greek letter δ with a + or –
Depending on the charge

43. Dipole?????? A dipole is something that has two (di) poles.
Like a magnet
Just like the magnet a dipole molecule has a positive end and a negative end
These are polar covalent bonds
With these poles they are attracted by electric fields.
Solubility is the ability of a substance to dissolve in another substance
Polar and ionic molecules dissolve in polar substances
Non polar molecules only dissolve in nonpolar substances
Good rule-like dissolves like
This tells us why white bears are afraid of water.

44. Intermolecular forces
Nonpolar molecules
The force is weak and called a dispersion force
Or induced dipole
Polar molecules
They have charged ends so opposite charges attract
Called a dipole-dipole force
The more polar a molecule the stronger this force is
Finally we have the hydrogen bond
Especially strong
Forms between hydrogen end of one dipole and
A fluorine, oxygen or nitrogen atom on another dipole