1. Ch. 12 Chemical Bonding

2. 12.1 Ionic, covalent and polar covalent bonds.
A bond is a force that holds atoms together.
Ionic Bonding
a. When a metal reacts with a nonmetal, electrons are transferred from a metal to a nonmetal and an ionic compound is made.

3. Ionic Bonding
b. In ionic bonding, electrostatic attraction holds atoms together.

4. Covalent Bonding
a. Atoms make covalent bonds by sharing electrons.
b. Electrons are attracted to the nucleus of both atoms in the bond.
c. Nonmetals make covalent bonds.
Let’s share! H H

5. Electronegativity
Electronegativity is a measure of how strongly an atom attracts electrons. Look at the electronegativity chart on p403.
Fluorine has an electronegativity value of 4.0.
Hydrogen has an electronegativity value of 2.1.
Difference in electronegativity =
4.0 – 2.1 = ________

6. Bond Types
Determined by difference in electronegativity values—absolute value of DEN
Pauling Electronegativity Values

7. Electronegativity
d. At left is a picture of hydrofluoric acid (HF). At right is a picture of “The Blob.”
e. F has a higher electronegativity than H, so electrons are closer to F.

8. In a covalent bond, atoms have a difference of electronegativity of Electrons are shared equally.
In a polar covalent bond, atoms have a difference in electronegativity of Electrons are not shared equally.

9. Electronegativity Exampe: Si—C
Electronegativity of Si is ________. Electronegativity of C is _______.
Difference in electronegativity is ____________.
Si – C bond is ( covalent /
polar covalent ) [circle one].
Which attracts more electrons? ( Si / C )
1.8
2.5
0.7

10. Covalent Polar covalent Ionic In the picture below,
Label each molecule or compound as ionic bonding, covalent bonding or polar covalent bonding.
For polar and ionic bonds, label the more electronegative atom.
Covalent
Polar covalent
Ionic

11. Identify each of the following bonds as ionic,
covalent or polar covalent.
Bond
Is there a metal and nonmetal?
Electronegativity difference
Bond Type
H – H No
2.1 – 2.1 = 0
Covalent
S – H
2.5 – 2.1 = 0.4
Polar covalent
Na – I
Yes  ionic
ionic
Rb – N
B – P
H – P

12. Dipoles
In a polar molecule, one side has a partial positive charge, and the other has a partial negative charge.
A dipole moment is represented with an arrow pointing to the negative side and the Greek letter “delta” δ to show the partial positive and negative charges:

13. Dipoles
Write the partial charges and draw the dipole moment on Cl –I

14. 12.2 Ionic Bonding Ions Metals ( lose / gain ) electrons.
Nonmetals ( lose / gain ) electrons.
Group 1 elements form ions with a charge of ____.
Group 2 elements form ions with a charge of ____.
Group 6 elements form ions with a charge of ____.
Group 7elements form ions with a charge of ____.

15. Ions That’s interesting…but WHY???
Atoms gain or lose electrons to get the electron configuration of a noble gas.
Noble gases have completely filled energy levels, so they are very stable.
He has a completely filled 2s sublevel.
Other noble gases have filled s and p sublevels.

16. Ions Example: Li The electron configuration of Li is ______
Li loses one 2s electron and becomes Li+.
The electron configuration of Li+ is _____
Li + has the same electron configuration as He.

17. Ions The electron configuration of F-
The electron configuration of F is
F gains one 2p electron and becomes F-
The electron configuration of F- is
F- has the same electron configuration as ______

18. Ions What would happen if Li reacted with F?
Li gives an electron to F
Li + F  Li+ + F-
And they form _______ (write the formula of the compound)
Ionic bonding and structures
LiF is packed together in a group in order to maximize attractions of the cations and anions.
It makes a hard, tight crystal.

19. The size of ions Which is larger, Na or Na+? Why?
Na. Na loses a 3s electron, and then only has electrons in the n=2 level. n=2 orbitals are smaller than n=3 orbitals. Na
Na+ F- F

20. The size of ions Which is larger, F or F-? Why?
F- because it gains electrons.
Which is larger F- or Na+? Why?
F-. They both have the electron configuration of Ne, but Na+ has more protons (a stronger + charge), which pulls electrons closer.
Which is larger, Ca or Ca2+? Why?

21. Ionic Size Taken from: http://www. chem. umass

22. 12.3 Lewis Structures
The octet rule: Sharing of electrons usually occurs so that atoms acquire the electron configurations of noble gases (1s2 or ns2np6)

23. Lewis dot structures The element symbol represents the core electrons.
Dots to show the
valence electrons.

24. Lewis dot structures 1. First, write the symbol for the element.
2. Imagine the molecule has four sides (but don’t draw the “x”)
3. Draw one dot at a time in each empty section.
4. You should only have pairs of e- if there are no empty sections. Cl

25. Lewis dot structures The “paired” electrons cannot usually make bonds.
The three “unpaired” electrons can make bonds.

26. Arsenic
Dot Diagram?
In order to get the configuration of a noble gas, how many bonds will arsenic form?
Dot Diagram for Hydrogen?
In order to get the configuration of a noble gas, how many bonds will H form? 3 H 1

27. Arsenic Trihydride, AsH3
Lewis dot structure
Structural formula As H H H As H H H As

28. Drawing Lewis Dot Structures
Use the molecular formula to find the total number of valence electrons
= 8
H H H As

29. Drawing Lewis Dot Structures
2. Draw the symbols of each element. Draw the backbone/skeleton structure
Terminal atoms go around
the central atoms.
Least Electronegative is usually central
Hydrogen is always terminal.
Carbon is usually central
Oxygen and Halogens are usually terminal

30. Drawing Lewis Dot Structures
3. Draw the valence electrons
4. Make single bonds with pairs

31. Drawing Lewis Dot Structures
5. Each pair represents a bond.
6. Count electrons again and make sure you get the same number (8).
Single bond

32. Structural Formula
The structural formula is drawn with a “–“ line to represent the bond.
Double and triple bonds
Double bonds can be formed by sharing two pairs of electrons
Triple bonds involve sharing three pairs of electrons. O N

33. Try the Dot Diagrams for each of the following:
O2 H2O BCl3 SiI4 Si I B Cl O H O
Each pair of shared electrons can be represented by a line. Si I B Cl O H O

34. Structural formulas for polyatomic ions
Example: NO3-
The negative charge means there is one more electron
How many total electrons are there?
5 + 3(6) + 1 = 24 electrons.
NO3-
Draw:

35. Resonance Resonance Example: ozone
Ozone’s structure can be drawn like this:
The actual bonding in ozone is not like either of these structures. The actual structure lies somewhere in between these two.
These drawings are resonance structures of ozone O