1. COVALENT COMPOUNDS Chapter 6
In this chapter we will look at covalent compounds and see how they differ from ionic compounds (Chapter 5)

2. 1) HOW DO COVALENT COMPOUNDS DIFFER FROM IONIC COMPOUNDS?
Unit Essential Question:
1) HOW DO COVALENT COMPOUNDS DIFFER FROM IONIC COMPOUNDS?

3. Lesson Essential Questions:
1) WHY DO ATOMS SHARE ELECTRONS? 2) HOW DO BONDS AND PROPERTIES RELATE?

4. Section 1: Covalent Bonds
Covalent- electrons are shared.
Simplest covalent bonds are diatomic molecules.
Always 2 atoms; atoms can be different or the same (diatomic elements).
Ways to remember diatomic elements: BrINClHOF, HON & the halogens, #7
Ex: H2
Each H atom shares its valence electron so each can have 2.
Called a shared pair of electrons..
Remember your diatomic molecules – I warned you in Chapter 2 to remember them. They will come back up later in the year as well. BrINClHOF

5. Molecular Orbitals Where shared electrons are located.
Region of high probability of finding the shared electrons- recall quantum model.
Form when covalent bonds occur and atomic orbitals overlap.
The electron clouds of atoms will overlap and form a molecular orbital for the covalent compound. See the diagram to see how atoms will repel and attract each other. In order to bond, the attraction must be greater than the repulsion.

6. Energy and Stability Non-bonded atoms have high PE and low stability.
Become more stable when bonded.
E is released when bonds form, PE decreases.
This determines the bond length.
Ideal length is when the two bonded atoms are at their lowest PE.
Atoms become more stable if they bond to achieve the octet rule. Energy is released in this process. The potential energy of the atoms decreases when the bond forms. The bond length is where the lowest potential energy of the atoms occurs. Covalent bonds are not as strong as ionic, but are more flexible.

7. Energy and Stability Cont.
Covalent bonds are flexible.
Bond length is an average distance.
Bond length oscillates like a spring between the two nuclei.
Bond energy – energy required to break a bond.
Different from lattice energy!
More E is needed to break a shorter bond.
Shorter bond length = stronger bond.

8. Re-visiting Electronegativity
How do we know if atoms will transfer or share electrons in a bond?
Check electronegativity values!
If the difference is greater than 2.1 it is considered ionic.
Ex: KCl
Electronegativities: K= 0.8 and Cl= 3.2
Subtract: = 2.4, so the bond is ionic.
You will not have to memorize the electronegativity chart but you must know how to use it. You will need to subtract the values to find the difference (you will not have a calculator). You will also need to know the “cut-off” numbers to determine which type of bond is formed.

9. Re-visiting Electronegativity
If the difference is less than 2.1 it is considered covalent.
Ex: CO2
Electronegativities: C=2.6 and O=3.4
Subtract: 3.4 – 2.6 = 0.8, so the bond is
covalent.
You do not need to memorize electronegativity values! They will be given to you on tests and quizzes.
pg.194 Figure 6 shows electronegativity values
You will not have to memorize the electronegativity chart but you must know how to use it. You will need to subtract the values to find the difference (you will not have a calculator). You will also need to know the “cut-off” numbers to determine which type of bond is formed.

10. Re-visiting Electronegativity
It is important to note that this cut-off value of 2.1 is arbitrary! Therefore, it is not 100% accurate.
Properties of the compound must be investigated for better classification.
Ex: magnesium chloride
Electronegativities: Mg=1.3 and Cl=3.2
Subtract: 3.2 – 1.3 = 1.9
It looks like it should be a covalent bond, but properties indicate it is actually ionic.

11. Polar vs. Nonpolar Covalent
If the electronegativity difference is less than 0.5 we will consider the bond to be nonpolar covalent.
Means the atoms are essentially sharing the electrons equally.
If the difference is between 0.5 and 2.1, we will consider it to be polar covalent.
Means one atom attracts the electrons more towards itself; there is unequal sharing of electrons.
We can also classify covalent bonds into 2 types. Polar molecules will have special properties we will discuss later in the year.

12. Determining Bond Type

13. Practice Classifying Bonds
A) Classify the following bonds as ionic, polar covalent, or nonpolar covalent:
1) Li – Cl
2) B – C
3) N – O
4) Mg – Br
5) C – F
B) Since F has the highest
electronegativity value, can F
ever form a nonpolar covalent
bond? Why or why not?
2.2 ionic
0.6 polar covalent
0.4 nonpolar covalent
1.7 ionic
1.4 polar covalent

14. Simple Polar Molecules
Ions have full charges (electrons are transferred).
Polar molecules have partial charges (unequally shared electrons).
Use Greek symbol delta δ (and +/-) to indicate partial charges.
Dipole – molecule that contains both positive and negative partial charges.
Ex: HCl
We will show partial charges using the Greek symbol delta. I think it looks like a Hershey kiss (with the little flag out of the top).

15. Simple Polar Molecules Continued
Use electronegativity values to determine partial charges.
Atom with larger e-neg value = δ-
Atom with lower e-neg value = δ+
Ex: HCl
Elecronegativities: H=2.2 and Cl=3.2
So: H is δ+ and Cl is δ-
In other words, the electrons are more likely to be found at the Cl atom than the H because it has a larger electronegativity; this makes Cl more negative.
We will show partial charges using the Greek symbol delta. I think it looks like a Hershey kiss (with the little flag out of the top).

16. Different Properties for Different Bonds
Besides electronegativity, you can predict bond type by the type of elements involved in the bond.
Metallic = metal atoms (K, Cu, etc.)
Ionic = 1 metal (typically) + 1 nonmetal
Covalent = 2 or more nonmetals
Properties of these compounds are determined by bond type.
Recall ionic compound properties and how they stem from the crystal lattice.
We will do a lab on conductivity to show some of these properties.

17. Covalent Properties Some are soluble in water and some are not.
Depends if the bonds are polar or not.
Poor conductivity.
Covalent compounds can exist as a solid, liquid, or gas.
Depends upon the polarity of bonds.
Tend to have low melting and boiling points.
*Remember: covalent compounds are made of molecules, not ions!

18. Metallic Ionic Covalent
Bond type
Metallic
Ionic
Covalent
Model
Example
Potassium
Potassium chloride
Chlorine
Melting point
63 oC
770 oC
-101 oC
Boiling point
760 oC
1500 oC (sublimes)
-34.6 oC
Properties
Soft, grey lustrous
Conductor as solid
Crystalline white solid
Conductor when molten or in solution
Green-yellow gas
Insulator
There is a similar chart in your book – this will also help with conductivity lab.

19. HOW ARE COVALENT COMPOUNDS REPRESENTED?
Lesson Essential Question:
HOW ARE COVALENT COMPOUNDS REPRESENTED?

20. Section 2: Drawing and Naming Molecules
Lewis structures (electron dot diagrams) use dots and lines to show valence electrons.
Atoms only use dots:
Each dot represents a valence e-.
Molecules use both dots and dashes:
Each line represents a bond.
Help to determine and show bonds that will form in covalent compounds.
This will be very important to both Section 2 & 3 – we will work on them almost simultaneously. We will learn to draw dot diagrams, line diagrams, 3D models, and learn about shape, angles, and polarity.

21. Lewis Structures- Atoms
The 4 sides of the element symbol represent the s and three p orbitals.
s = 2 e- and p = 6 e- so total = 8
No more than 8 dots per symbol
Put one dot on each side of the atom before pairing:
Draw the Lewis structure for the following: Al, Br, N, Ne
Remember the pattern on the Periodic Table to know the number of dots to use for each element.

22. Lewis Structures- Molecules
When drawing Lewis structures for molecules, give each atom an octet.
Except for H!
Looking at the Lewis structures for each atom can help to determine how many bonds it will form.
Ex: CH4
Needs 4 e-, so it will form 4 bonds.
Remember the pattern on the Periodic Table to know the number of dots to use for each element.
Each H needs 1 e-, so they will form 1 bond.

23. Electron Pairs
Shared pairs: e- are shared between two atoms, forming a bond.
1 shared pair = single bond or
2 shared pairs = double bond
3 shared pairs = triple bond
Unshared/lone pairs: e- are not shared between two atoms, and are not involved in a bond.
Lone pairs (lone pears – get it??) – did you ever notice the “lone pear” I have hanging in the classroom?
There is no such thing as a quadruple bond – it NEVER happens.

24. Steps for Drawing Lewis Structures
1) Find the total # of valence electrons for all atoms in the compound.
Ex: NH3 N: 5e- H: 1e- total: 5 + 1(3) = 8e-
2) Arrange atoms/determine ‘backbone’ of the compound.
H and halogens tend to be on the end.
They don’t like to share more than 1 pair of e-.
C is almost always in the center, if present.
Likes to form 4 bonds.
Except for C, the atom with the lowest electronegativity will be the center.
H will NEVER be in the center!
Ex: NH N will be in the center:
If you follow these steps, it should help you to draw the molecules correctly. It does take some trial and error, but following the steps will help you. Do not get caught up in trying to keep the initial number of valence electrons on each atom; worry about the total number for the molecule and making sure they all have 8.

25. Steps for Drawing Lewis Structures
3) Keep track of how many e- are used and how many remain.
Each bond used to build the backbone uses
2 electrons. NH3 has 3 bonds, so 6 e- used.
8 e- - 6 e- = 2 e- left.
4) Distribute remaining e- by giving each atom
an octet.
Give the most electronegative atoms e- first.
Ex: 2 e- left; N is the only one that needs an octet, so it gets the remaining 2 e-.
5) Verify the structure!
Make sure all atoms have an octet and check that all valence electrons have been used.

26. Polyatomic Ions
The bonds that hold together polyatomic ions are covalent bonds.
When drawing the Lewis structure for a polyatomic ion, you must take into account its charge when determining the total number of valence electrons you’re working with:
negative ion: add electrons to the total
positive ion: subtract electrons from the total
3) Finally, put brackets around the ion and include its charge.
Ex: Draw the Lewis structure for
the sulfate ion.

27. Practice Problem #2: Ions
Draw the Lewis structure for the ammonium ion.

28. Practice Problem #4: Multiple Bonds
1) It is possible to run out of remaining electrons to give to atoms as lone pairs.
But they still need an octet!
2) If this happens, you take lone pairs from a neighboring atom and form another bond to the atom that needs an octet.
The neighboring atom donates both electrons to be shared in this bond.
This is how double and triple bonds are formed.
Ex: Draw the Lewis structure for O2.

29. Practice Problems #5-6: Multiple Bonds
Draw the Lewis structure for CO2.
Draw the Lewis structure for N2.

30. Resonance Structures
1) Some compounds can be represented with more than one Lewis structure.
All Lewis structures are drawn.
A double-headed arrow is put between them to indicate all possibilities.
2) Compound is an average of the possible Lewis structures (resonance hybrid).
For the average resonance structure, we use dashed lines for the double bond that can be in either place (only 1 of the two lines for the double bond will be dashed).

31. Naming Binary Covalent Compounds
1) Name the first element listed in the formula.
The name does not change.
2) Second element always has ‘-ide’ ending.
3) Prefixes are used to indicate numbers of atoms of both elements.
Exception: do not use ‘mono-’ for the first element listed if there is only one.
This is not the same as it was for ionic. Remember to identify covalent compounds by recognizing that they are made from 2 nonmetals.

32. Naming Binary Covalent Compounds
Take off the ‘a’ at the end of the prefix if the element begins with a vowel.
Ex: pentoxide, not pentaoxide.
Ex #1: CCl4
Ex #2: P2O5
Ex #3: H2O
Ex #4: CO
carbon tetrachloride
diphosphorus pentoxide
dihydrogen monoxide
carbon monoxide

33. Lesson Essential Question:
HOW CAN THE THREE DIMENSIONAL SHAPE OF A MOLECULE BE DETERMINED? WHY ARE THE SHAPES OF MOLECULES USEFUL?

34. Section 3: Molecule Shapes
1) VSEPR theory: Valence Shell Electron Pair Repulsion
Predicts shape based on electron repulsion.
e- want to be as far apart as possible!
Look at the central atom to determine the shape.
2) Unshared pairs repel more than shared pairs.
All electrons repel each other, whether they are bonded or lone pairs.
3) In addition to bond polarity, shape also plays a role in determining the properties of a substance.
We will do a lab to build these molecules – they are easier to understand if you can see them and turn them around to experience the shapes. This is an important concept since polarity (determined by shape) will determine many properties of compounds.

35. Predicting Shapes
Count the number of shared and unshared pairs on the central atom in the Lewis structure.
Count double/triple bonds as 1 shared pair.
Total pairs
Shared pairs
Unshared pairs
Shape name 4
Tetrahedral 3 1
Trigonal pyramidal 2
Bent
Linear
Trigonal planar
I will give you a chart like this when we do the lab – it also have diagrams/pictures on it.

36. Angles Bond angles also show repulsion of electrons in molecules.
Tetrahedral angles = 109.5o
Base all other angles off of this.
From here, we can predict >, < , or = 109.5o
Depends upon shared vs. lone pairs.
(Additional factors that affect bond
angles: atom size, multiple bonds take
up more space. We won’t worry about
these.)
This is not covered in your book – I hope by doing the lab and actually looking at the models you will build, you will see how this works. Remember that there are no 90 degree angles (at least with the molecules we will look at!).

37. Angles Cont. Trigonal Pyramidal: <109.5o
Remember- tetrahedral was 109.5, now we’ve replaced a shared pair with a lone pair that repels more! So angles decrease!
Linear: > 109.5o (180o is much bigger!)
Trigonal Planar: >109.5o (120o is bigger!)
Bent – depends on # of lone pairs
Bent with one lone pair is like taking trigonal planar and replacing a shared pair with a lone pair, so angle is <120 but >109.5.
Bent with two lone pairs is like taking tetrahedral and replacing two shared pairs with two lone pairs, so angles are much < This is most common!
This is not covered in your book – I hope by doing the lab and actually looking at the models you will build, you will see how this works. Remember that there are no 90 degree angles (at least with the molecules we will look at!).

38. Polarity They can also be polar if the bonds are not all the same.
Tetrahedral, trigonal planar, and linear molecules can be nonpolar.
Polar bonds can ‘cancel’ out other polar bonds because of the ‘symmetrical’ shape.
This happens if the bonds are all the same.
Ex: CH4 & CO2
They can also be polar if the bonds are not all the same.
Bonds do not completely cancel out.
Ex: CH3F
Remember to look at both the atoms and the bonds. They must all be single bonds or double bonds, etc to be symmetrical and the atoms attached to the center must all be the same atom to be symmetrical.

39. Polarity Cont. Trigonal pyramidal and bent molecules tend to be polar.
Polar bonds are not cancelled out because of the ‘asymmetrical’ shape.
Example: H2O
Even if the bonds are not polar, the lone pair makes it polar.
Let’s relate this back to covalent compound properties from section 1.
Methanol is soluble in water.
Oil is not soluble in water.