1. Chemical Bonding

2. Background Review
Nuclear Charge – This is the total charge present on the nucleus (always positive) It is equal to the number of protons present.
Electrons – In a neutral species, the # of electrons is equal to the number of protons
Ions – species where # of protons is not equal to the # of electrons. (All ions are formed by adding or removing electrons.)
Valence Electrons – All the electrons in the outermost energy level for a given atom or ion. These are involved in chemical bonding. For A families the group number is equal to the number of valence electrons.

3. Lewis Dot Diagrams (Electron Dot Diagrams)
These are used to represent the valence electrons.
Examples – using the elements in period 2

4. TERMS
Bond capacity – Is equal to the number of bond pairs. This tells us how many other atoms a particular atom can bond to at one time.

5. Chemical Bonding Bond – any force that holds atoms (or ions) together.
Compounds – Two or more elements chemically combined
Stable chemical bonds will form if one or both of the following conditions are met.
A) The bonded atoms have less energy than the individual atoms.
B) Atoms gain or lose electrons to get a configuration like an inert gas.

6. Two models are used to explain chemical bonding:
(a) Ionic Bond – this model results from the complete transfer of one or more electrons from one atom to another.
(b) Covalent Bond – in this model, the electrons are shared between the bonded atoms.
Note: Real bonds have characteristics of both models…We will explain bonding in terms of one or the other.
Electronegativity – tells us how strongly a given atom attracts electrons. The higher the number, the stronger the attraction.

7. Part I
The Ionic Bond

8. Predicting when to use this model?
In this model, electrons are completely transferred from one atom to another.
This results in 2 oppositely charged ions.
Predicting when to use this model?
Two rules allow us to predict this bond type:
(A) If the electronegativity difference is 1.7 or greater…..we will use the ionic model.
(B) If the elements involved are from group IA and IIA (groups 1 & 2)matched with (bonded to) elements from groups VIA and VIIA (groups 16 & 17), the bond is usually ionic.

9. Empirical Formula
Ionic compounds can be represented by a chemical formula known as an empirical formula.
The empirical formula is based on measurement and observation.
It tells us the ratio of positive ions to negative ions in the compound sample.
Example: NaCl(s) There is one positive sodium ion (Na+) for each negative chloride ion (Cl-).
There is no such thing as a molecule of NaCl(s) .

10. Lewis Diagrams for Ionic Compounds

11. Sheet 20/8…8

12. REVIEW
IONIZATION ENERGY
This is the amount of energy required to remove an electron(s) from an atom or ion. (See formation of a positive ion..)
IN ANY FAMILY
As atomic number increases – ionization energy decreases
This is a function of size. The outermost electrons are further and further from the nucleus and are easier to remove.
IN ANY PERIOD
As atomic number increases – ionization energy increases
This is due to increased nuclear attraction on electrons in the same energy level.

13. 3 Step Reaction
The formation of an ionic compound can be represented with 3 steps.
Example: NaCl (aq)
The first step is ENDOTHERMIC. Energy must be supplied in order to remove the electron from the sodium atom.
That is: Na kJ → Na+ + e-
Step 2 is EXOTHERMIC. Energy is released as the electron is accepted by the chlorine atom.
That is: Cl + e- → Cl kJ

14. 3. Step 3 is also EXOTHERMIC
3. Step 3 is also EXOTHERMIC. Energy is again released as the ions move together and take up positions in the forming crystal.
That is: Na+ + Cl- → NaCl kJ

15. Adding the reactions to get a net ionic equation

17. Characteristics of Ionic Compounds
Solids with high melting points and boiling points.
Non conductors in solid state.
Conductors in aqueous solution.
Conductors in molten (melted) state.

18. Ionic Properties (Ionic Character)
Relative ionic character can be determined by comparing electronegativity differences.
The greater the difference – the greater the ionic character.
Example; which of the following has the greatest ionic character? Li2O(s) , KBr(s) , H2O(l), AlCl3 (s)

19. Types of Covalent Bonds

20. Predicting Molecular ShapeV S E P R T
Valence
Shell
Electron
We use this theory to predict the shape of molecules.
Pair
Repulsion
Theory
The shape that we are describing may be the shape around a specific atom or it may be the overall shape of the molecule.

21. *** Shape is always determined by drawing lines from the center of one specific atom to the center of all atoms bonded to it.***

22. Exercise: Bond each element in period II with as many atoms as possible of an element (‘X’). Element X has a bond capacity of one.
Group IA Group IIA Group IIIA Group IVA

23. Group IA Group IIA Group IIIA Group IVA
Green (1) black (4)
Group VA Group VIA Group VIIA Group VIIIA

24. Molecular Polarity Some molecules have a polar nature.
This means that one part of the molecule, often an end, is positively charged and one part is negatively charged.
This polarity is the result of uneven sharing of electrons – that is the presence of polar bonds.
Examples: H – Cl Polar molecule
H – H Non – Polar molecule

25. We use vectors to determine if a molecule is polar or not.
Procedure:
Replace all polar bonds with vectors. (point the arrow toward the atom with the higher elecronegativity)
Starting at a given point, place all the vectors end to end.
Now: If you end up where you started – the molecule is non – polar.
If you end up anywhere else – the molecule is polar.
Examples:

26. Intermolecular forces are attractions between molecules.
Types of Forces
Intermolecular forces are attractions between molecules.
Intramolecular forces are attractions within the molecule. (Covalent Bond)

27. Intermolecular attractions
These are forces that exist between covalently bonded molecules.
Example: from one water molecule to another
We will look at three of these intermolecular forces.
In effect, these are a weak, a medium and a stronger force.

28. I. Van der Waals (London Dispersion)
These attractive forces are present between the molecules of ALL samples of matter.
They result from the attraction between the electrons of one molecule and the nuclei of another.
Example:
RULE: the greater the number of electrons, the greater the Van der Waals forces and therefore the higher the melting point.

30. II. Dipole – Dipole These forces exist, in addition to Van der Waals.
Generally, they are a little stronger.
They consist of the attraction between the positive end of one polar molecule and the negative end of another.
Example:

31. III. Hydrogen Bonding This is a special type of dipole – dipole.
Generally these are stronger forms.
This will ALWAYS involve molecules that have a Hydrogen atom covalently bonded to a Nitrogen, N, Oxygen, O, or a Fluorine, F.
Hydrogen bonds are generally responsible for the unique characteristics of water.

33. Put the following in order of increasing melting points
Put the following in order of increasing melting points. CH4, NH3, NCl3, BeO, H2O, HBr, CCl4, LiF,

34. NaCl, C3H8, H2S, OF2, PI3, N2, CaCl2, H2, O2

35. C2H6, MgO, HBr, HF, Cu, NCl3, CO2, CaS, C4H10, PBr3

36. Typical Metallic Solid and Its “Sea of Electrons”

38. Trends in atomic size

40. Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell electrons
1A A A A A A A A
H He:
          
Li  Be   B   C   N   O  : F  : Ne :
       
          
Na  Mg  Al   Si   P   S  :Cl  :Ar:
       

41. Lewis Diagram for Ionic Compounds
 
 
Na + Cl  Na[ Cl ]-
 
 
  
 
 

42. Oxidation/Reduction Half Reactions
Oxidation  the loss of electrons by metal.
Reduction  the gain of electrons by non-metals.
LEO says GER

43. Loss of electrons (LEO)
Oxidation
Loss of electrons (LEO)
Na(s)  Na+ + 1e-

44. Gain of electrons (GER)
Reduction
Gain of electrons (GER)
Cl2(g) + 2e-  2 Cl-

45. 2Na(s) + Cl2(g)  2NaCl(aq)
Balance
The balanced equation must show a balance between the two half reactions of the equation.
Na(s)  Na+ + 1e-
Cl2(g) + 2e-  2 Cl-
( ) x 2
2Na(s) + Cl2(g)  2NaCl(aq)

46. Forces

47. Van Der Waals Forces
These are the weakest of the intermolecular forces. There are two types:
London Dispersion
Dipole Interactions

48. London Dispersion
Attraction of electrons to the other molecules positive nucleus.
Weakest of all
the greater the number of electrons in the atoms of a molecule the greater are the dispersion forces

49. Dipole Interactions When polar molecules are attracted to each other
Electrostatic attractions occur between the opposite charged regions of molecular dipoles.
Similar to ionic bonding but much weaker

50. Hydrogen Bonding Strongest intermolecular forces.
Weak bond of hydrogen atoms to the unshared electron pair of an electronegative atom in a nearby molecule.
Due to the fact that the valence electrons are not shielded from the nucleus.