1. Unit 3: Chemical Bonding and Nomenclature Part 2

2. Electronegativity
The ability of an atom in a molecule to attract electrons to itself

3. Electronegativity
Since atoms have varying electronegativities they are pulling electrons towards them with different amounts of force
Polar bond occurs when one atom is pulling electrons towards it with more force than the other atom involved in the bond

4. Electronegativity This value can be read from the periodic table
We use the electronegativity (ΔEN) difference to determine if a bond is polar or not
Find the EN values in the table then subtract the smallest from the largest number
Compare this value to the ranges for different types of bonds

5. Polar and Non-polar Bonds

6. Polar and Non-polar Bonds
Electronegativity Difference (ΔEN)
Type of Bond
>2.0
Ionic
Polar Covalent
<0.4
Purely (non-polar) Covalent

7. Polar and Non-polar Bonds
Determine whether the following bonds are polar, non-polar, or ionic
K-Cl
H-P
B-F

8. Partial Charges
In polar covalent bonds, one atom is pulling electrons closer to them
Higher electronegativity
This causes the electrons to spend more time closer to that atom
The atom with higher electronegativity has a partially negative charge
The atom with the lower electronegativity has a partially positive charge

9. Partial Charges Partial charges are represented by: δ+ and δ-
Polar bonds are said to have a dipole
Shown by an arrow with a positive (+) end and the arrow head pointing to the more electronegative atom

10. Partial Charges

11. Partial Charges

12. Polarity of Molecules Determined by:
Presence of lone pairs of electrons on the central atom in a molecule
Presence of different elements surrounding the central atom
Binary molecules involving two different elements

13. VSEPR Theory Valence Shell Electron Pair Repulsion
Based on idea of like charges repelling each other
Electron domains (bonds or lone pairs) are negatively charged and therefore repel one another

14. VSEPR Theory
We can predict the 3D shape a molecule will take on by looking at how many electron domains the central atom has
Remember the atom that needs the most electrons is placed in the middle of our Lewis dot structures

15. VSEPR Theory There are 6 shapes you are responsible for
Found on page 2 of your Data Sheets booklet
First we must figure out how many electron domains surround the central atom
How many bonding pairs?
How many lone pairs?
From here we can figure out the shape the molecule will take on

16. VSPER Shapes
Linear
2 electron domains
2 bonding pairs
0 lone pairs

17. VSEPR Shapes Trigonal planar 3 electron domains 3 bonding pairs
0 lone pairs

18. VSPER Shapes Bent or V-shaped 3 electron domains 2 bonding pairs
1 lone pair

19. VSPER Shapes Tetrahedral 4 electron domains 4 bonding pairs
0 lone pairs

20. VSEPR Shapes Trigonal pyramidal 4 electron domains 3 bonding pairs
1 lone pair

21. VSEPR Shapes Bent or V-shaped 4 electron domains 2 bonding pairs
2 lone pairs

22. Molecular Polarity We know how to find out if a bond is polar
For molecules we must look at all bonds and then the molecule as a whole
To find out if something is polar we must figure out if all sides are the same or different

23. Molecular Polarity
If all sides of the shape are the same partial charge (all negative or all positive) then the molecule is non-polar
If two or more sides have different partial charges (some positive some negative) then the molecule is polar

24. Molecular Polarity
Molecules that contain polar bonds can be non-polar as a whole
Examples:
CO2
CH4
CH3Cl
NH3

25. Inter and Intra Molecular Forces
Inter = between
Intermolecular forces are forces between molecules
Influence boiling and melting points
Intra = within
Intramolecular forces are forces within a molecule
Ionic, covalent, metallic

26. Intermolecular Forces
London dispersion forces
Weakest
Dipole-dipole forces
Hydrogen bonding
Strongest

27. Intermolecular Forces
Found in molecular (NOT ionic) compounds
Allows us to organize compounds based on melting and boiling points
Tells how tightly molecules of a substance are held together
How much energy needs to be added to change states

28. Intermolecular Forces
If a molecule has a strong intermolecular force (hydrogen bonding) it will also have the weaker forces (dipole-dipole forces and London dispersion forces)
Substances with stronger forces will boil and melt at higher temperatures than those with weaker forces

29. London Dispersion Forces
There is always at least a slight electrostatic (+ and -) attraction between molecules
Explains why non-polar gases can be liquified
Remember electrons are floating around a positive nucleus in a cloud at all times
Instantaneous distribution of electrons can vary from average distribution

30. London Dispersion Forces
Present between all molecules
Significant only when molecules are very close together
Increases with increasing atomic and molecular size

31. London Dispersion Forces

32. Dipole-dipole Forces Permanent dipole is present
Must have polar molecules
δ+ and δ- ends of the molecules align so the opposite charges are close

33. Dipole-dipole Forces

34. Hydrogen Bonding
Attraction between H atom attached to highly electronegative atom and nearby small electronegative atom in another molecule
Strongest of the intermolecular forces
Molecules must have H bonded to: O F N

35. Hydrogen Bonding

36. Intermolecular Forces
What is the strongest intermolecular force present in the following:
NH3
HCl
CCl4

37. Intermolecular Forces
As the strength of intermolecular forces increase so do boiling and melting points
If the same intermolecular forces are present in two molecules the force is increased as the size of the molecule increases
C2H6 has a higher melting/boiling point than CH4

38. Hydrogen Bonding - Water
Unusual because solid form is less dense than liquid form
In liquid form, H-bonds between molecules
When frozen there are more H-bonds made and a lattice is formed
Mass does not change
When frozen volume increases (takes up for space) which decreases density

39. Hydrogen Bonding - Water

40. Hydrogen Bonding-Water
Since H-bonding is such a strong intermolecular force water has a high melting and boiling point
Molecules are held close together
Requires a lot of energy to make molecules more mobile

41. Acids and Bases Arrhenius concept was useful but limited
Bonsted-Lowry Acids were developed
Acid-base reactions involved the transfer of H+ ions from one substance to another

42. Bronsted-Lowry Acids and Bases
Acids donate the proton (H+)
Bases accept the proton (H+)
When identifying compounds as acids for naming in this course we will look for H at the beginning of the formula

43. Acids and Bases Acid Base Low pH High pH Sour taste Bitter taste
Turns litmus paper red
Turns litmus paper blue
Increases concentration of H+ ions in solution (Arrhenius)
Increases concentration of OH- ions in solution (Arrhenius)
Proton (H+) donors (Bronsted-Lowry)
Proton (H+) acceptors (Bronsted-Lowry)

44. Acids Formula to Name Look for the H at the beginning of the formula
It’s an acid!
Is there two elements or more than two elements in the formula?
Is it binary or polyatomic?

45. Binary Acids Formula to Name
Binary acids have H and one other element in their formula
Always “hydro______ic acid” HF
Hydrofluoric acid
HBr
Hydrobromic acid

46. Polyatomic Acids Formula to Name
Polyatomic acids contain polyatomic ions attached to H
Remember:
Ate  ic
Ite ous
Then add acid to the end
We also don’t need to show H is there, drop the hydro!

47. Polyatomic Acids Formula to Name
H3PO4
Phosphoric acid
H2SO3
Sulfurous acid
CH3COOH (also HCH3COO, CH4CO2)
Acetic acid

48. Acids Name to Formula “Acid” must start with an H
Does the name begin with hydro?
Is it binary?
Work backwards

49. Acids Name to Formula Polyatomics: Binary: Balance your charges!!
Ic  ate
Ous  ite
Binary:
Look for the root element
Balance your charges!!

50. Polyatomic Acids Formula to Name
Hydroiodic acid HI
Chloric acid
HClO3
Hydrochloric acid
HCl
Dichromic acid
H2Cr2O7