1. We love them…we really do!
Chemical Bonds
We love them…we really do!

2. Exactly what are chemical bonds???
Defined as: a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together
The reason bonds are formed is because they are more stable bonded together than as individual atoms.

3. Types of Chemical Bonds
1) Ionic bonds
2) Covalent bonds
i. polar covalent
ii. non-polar covalent
3) Metallic bonds

4. Ionic Bonds Characteristics:
Results when electrons are completely given from one atom to another
Occurs between metals and non-metals

5. Covalent Bonds Characteristics:
1) Results from the sharing of electron pairs between two atoms
2) Occurs when nonmetals bond to each other

6. Ionic or Covalent?
The most important thing to remember about bonds is that they are on a continuum from ionic to covalent.
This means that you can have more ionic or more covalent character. That is why we use % ionic character.

7. Electronegativity Differences
To use the EN chart on page 162 of your text:
Find the difference in EN of the two atoms involved in the bond. (you can find EN values on the new periodic table I gave you!)
If it falls the in the range of then it is non-polar covalent.
If it falls in the range of , it is polar covalent.
If it falls in the range of , it is ionic.

8. Non-polar Covalent Bonds
These are a type of covalent bond in which the electrons are shared equally between the two atoms in the bond.
This means that the negative charge is distributed evenly!
Examples of compounds that are non-polar: hexane and oil

9. Polar Covalent Bonds
Polar covalent bonds do not share electrons equally between the two atoms involved.
This means that “poles” of positive and negative charges begin to form at each atom.
Examples of polar compounds: water, ethanol

10. Molecular compounds
A molecule is a neutral group of atoms that are held together by covalent bonds.
Examples: H2O and CO2
A chemical compound whose simplest units are molecules is called a molecular compound.

11. Chemical vs. Molecular Formulas
Really, chemical formulas encompasses all the formulas involving chemicals (for all types of bonding).
Molecular formulas are types of chemical formulas that describe molecules held together by covalent bonds.

12. Diatomic Molecules Molecules that contain only two atoms are diatomic.
There are several molecules that exist as diatomics in nature.
HOFBrINCl - (pronounced hoffbrinkle)

13. Formation of a covalent bond
When atoms at a distance are first attracted to one another, the attraction of their nucleus with the other atoms electrons is STRONGER than the repulsion between the nuclei and electrons.
Therefore, they will come close enough together to bond. When they reach the optimal distance, their potential energy is at its lowest.

14. Characteristics of a Covalent Bond
The distance between two bonded atoms at their minimum potential energy is called the bond length.
Thus each atom releases energy as they change from individual atoms to a molecule.
The same amount of energy must be supplied to separate the newly formed bond. This energy is called the bond energy.

15. Bond energies Scientists report these energies in kJ/mol.
To break a H-H bond, 436 kJ of energy is needed.
Bond energies and bond lengths are different depending on which two atoms are involved in the bond.

16. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.

17. Exceptions to the Rule
Hydrogen- can only bond to one other atom…wants 2 valence electrons.
Boron wants 6 valence electrons. (It already has 3!)
Some elements can be surrounded by MORE than 8 electrons (called expanded valence) when bonded to halogens or other highly EN elements. Examples: P, As, S

18. Electron Dot Notation Shows the valence electrons of an element
Uses the element symbol surrounded by up to 8 dots.
The order in which to place the dots:
Generally, one dot is placed on each side first before pairing the dots.
Let’s do some examples!

19. Lewis Structures
Electron dot notation can also be used to represent molecules.
In this case, the valence electrons also show us how the atoms are bonded.

20. Lewis Structures
Unshared pairs of electrons are called lone pairs. These are electrons that are not involved in bonding.
A pair of dots involved in bonding may also be represented as a dash.

21. Structural Formulas
A structural formula indicates the kind, number, arrangement and bonds but does not include the unshared or lone pairs of electrons.

22. How to Write a Lewis Structure
1. Determine the type and number of atoms in the molecule.
2. Write the electron dot notation for each type of atom in the molecule.
3. Add together the total number of valence electrons involved.
4. Arrange the atoms to form a skeletal structure of the molecule. If carbon is present, it is always in the center!

23. Continued 5. Add unshared pairs of electrons where appropriate.
6. Count the number of electrons to be sure that the number used equals the number available.
Practice:
CH3I
CO3-2
NH3
H2S

24. Multiple covalent bonds
Some elements can share more than one pair of electrons, especially carbon, nitrogen, and oxygen.
A double bond is a covalent bond produced by sharing two pairs of electrons.
A triple bond is a covalent bond produced by sharing 3 pairs of electrons!

25. About Multiple Bonds
Double and triple bonds have higher bond energies and have shorter bond lengths than single bonds.
Practice: N2
HCN
CO2

26. Resonance Structures
Some molecules and ions cannot accurately be represented with one Lewis structure.
This occurs when a molecule is asymmetrical with respect to bonds of the same type.
Example:
Ozone
When writing resonance structures you must include all the possibilities.

27. Covalent Network Bonding
All the covalent molecules you have learned about to this point are molecular.
Some covalent molecules do not exist as individual molecules. They are bound together by forces acting between them.
Continuous 3-D networks of bonded atoms are referred to as a covalent network. You will learn more about these later!

28. More About Ionic Compounds
Ionic compounds are composed of positive and negative ions that are combined so that the number of positive and negative charges are equal.
Most ionic compounds exist as crystalline solids. Many minerals are ionic compounds.

29. Formula Units
The chemical formulas of ionic compounds show the ratio of ions present in any size sample.
The ratio of ions depends on the charges of each.
Example: Ca and F

30. Characteristics of Ionic Bonding
Ions in ionic compounds often form a crystal structure of repeating units.
The 3D arrangement of ions depends on the strength of attraction between them and their sizes.
To compare bond strengths, chemists use lattice energy.

31. Lattice Energy
The energy released from one mole of an ionic crystalline compound as it turns to a gas is called the lattice energy.
The values are negative indicating that the energy is being released from the compound.

32. Intramolecular vs. Intermolecular Forces
Intramolecular forces are forces IN a molecule that keep them together.
Intermolecular forces are those BETWEEN molecules.
The difference between these two forces is why ionic and molecular compounds have such different physical properties.

33. Physical properties
The physical properties measured in lab were: melting point, solubility and electrical conductivity.
What did you notice about the compounds that we tested?

34. Melting point For ionic compounds, melting points are high.
This is because the ions are so strongly attracted to each other in a closely compact crystalline solid.
For molecular compounds, the force between molecules is not as high. This is why they melt so much easier.

35. Electrical Conductivity
In the solid state, the ions cannot move in an ionic compound, so they are not good conductors of electricity.
When they are dissolved in water, they conduct electricity very well because the charges are able to move around.

36. More Electrical Conductivity
If the solid is molecular (covalent bonds), it does not develop strong ions as it dissolves so it would not be a good conductor.

37. Polyatomic ions
This group of molecules have atoms that bond together covalently with each other to form a group that acts like an ion.
Examples: OH- ClO3- PO4-3
Here is a list of Polyatomic Ions that you must memorize for the next unit.

38. Metallic Bonding
Bonding is WAY different in metals than in ionic and molecular compounds.
The vacant orbitals in the valence energy levels overlap with other metal atoms.
This allows electrons to roam freely between the atoms.
The electrons are said to be delocalized.
The attraction that results from the attraction of metals and their sea of electrons is called metallic bonding.

39. Characteristics of Metal Bonds
The freedom of electrons is what makes metals a good conductor of electricity!
The reason that metals are shiny and reflect light is because the metals can absorb a wide range of frequencies. The valence electrons get excited and move to a higher energy level. When they return to their ground state light is emitted!

40. The Strength of Metallic Bonds
Metallic bond strength is measured by the amount of heat required to vaporize the metal (called the heat of vaporization).
It depends on the nuclear charge and the number of electrons!

41. Molecular Geometry This is the 3D arrangement of molecules in space.
The polarity and geometry of the molecule determine molecular polarity, or the uneven distribution of molecular charge.

42. 2 Theories about Geometry
1) VSEPR- (pronounced vess-per)
Stands for Valence Shell Electron Pair Repulsion
The theory is based on the idea that repulsion between sets of valence level electrons surrounding atoms causes them to be oriented as far apart as possible.

43. Types of Geometry
Linear- occurs when two atoms are bonded to a central atom
The atoms are 180° apart.
Example:

44. Types of Geometry
If there are 3 identical atoms surrounding a central atom, then to orient them as far apart as possible, the bond angle would be 120°.
The shape would be trigonal-planar.

45. Types of Geometry
Molecules that have 4 atoms bonded to the central atom always follow the octet rule.
The geometry here would require the angles to be 109.5°.
The shape would be tetrahedral as a result.

46. Do lone pairs have any effect of geometry?
Short answer…YES!
VSEPR theory says that the lone pair occupies space around the molecule just like bonded atoms do.
This means that there are more shape possibilities for molecules with lone pairs!

47. More Geometry
For example, ammonia is NH3 and has 3 bonds and 1 lone pair. It geometry is described as trigonal pyramidal.
While the lone pairs do take up space, we describe the shape of the molecule with respect to the position of the atoms ONLY!

48. Bent
A molecule that has 2 bonds and 2 lone pairs on the central atom is considered bent.
The bond angle is about 105°.
Water is an example of a molecule that is bent.

49. Trigonal Bipyramidal
This type of molecule results from 5 bonded atoms with no lone pairs.
There are bond angles of both 90° and 120°.
Example: PCl5

50. Octahedral
In this molecule 6 atoms are bonded to a central atom with no lone pairs.
Example: SF6

51. Hybridization
2)Hybridization of molecular molecules is the 2nd way to predict shape.
Hybridization is the mixing of two or more atomic orbitals of similar energies to produce new orbitals of equal energy.
Types of hybrid orbitals: sp
sp2
sp3

52. Geometry of Hybrid Orbitals
Atomic orbitals
Type of Hybrid
# of Hybrid Orbitals
Geometry
s,p sp 2
linear
s,p,p
sp2 3
Trigonal planar
s,p,p,p
sp3 4
tetrahedral

53. Intermolecular Forces
Remember…these are forces between covalent molecules!
They vary in strength but are weaker than covalent, ionic and metallic bonds.
Boiling point is a good measure of intermolecular forces!

54. Molecular Polarity
Dipoles are created by equal but opposite charges that are separated by a short distance.
The direction of a dipole is from positive to negative pole.
The forces of attraction between polar molecules are called dipole-dipole forces.

55. Molecular Polarity
Dipoles determine the polarity of a bond, and in molecules with more than one bond, ALL dipoles and their directions must be considered in determining molecular polarity.

56. Examples
CCl4
Carbon Dioxide and formaldehyde

57. Dipole-Dipole Interactions
Short range, only affect molecules that are near one another.
Results from the polarity of bonds in a molecule.
The partially positive atom attracts to the partially negative atom in another molecule.
This force of attraction keeps them together.

58. Hydrogen Bonding
A particularly strong dipole-dipole force in which the hydrogen attached to a highly electronegative atom is attracted to the electronegative atom in a nearby molecule.
Example: Water

59. London Dispersion Forces
London dispersion forces are intermolecular attractions resulting from the constant motion of electrons.
They are presents between all atoms and molecules.
London forces are the only intermolecular forces involved between noble gases and non-polar molecules.
Strength increases with increasing atomic mass of atoms involved.

60. Important Things to Remember
The important thing to remember about intermolcular forces is that the strength of the forces between molecules can help us predict physical properties like boiling point and surface tension.
Also, the difference in electronegativity of atoms involved in the intermolecular force matters immensely.