1. Noble Gases

2. Discovered He 1868; Ar 1894 Ne, Kr and Xe in 1898 Rn 1900
Unreactive because they have a full outer shell of electrons
Called Noble then Inert now Noble

3. Uses Helium Airships as it is lighter than air
Not as light as hydrogen [twice as heavy per volume] but does not burn [Hindenberg]
Used in atmosphere of deep sea divers
Less likely to cause the “bends”
Gives the diver “Mickey Mouse” voice because it has such a low density compared to air.

4. Argon The most common Noble Gas.
Used to fill normal incandescent light bulbs to stop them imploding, to reduce evaporation of the filament and because it is unreactive
Used as inert atmosphere for welding

5. Trends in Noble gases
Don’t react in general because they have a full outer shell.
Down the Group – molecules get larger
Makes it easier to form temporary dipoles due to larger electron clouds
So van der Waal’s forces are greater
More energy required to break these.
So MP and BP get higher
Same trend in halogens

6. Octet Rule A chemical rule of thumb
Octet Rule says atoms with 8 electrons in their outer shell are stable
Atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electron configuration as the nearest Noble gas

7. The rule applies to the main-group elements, especially carbon, nitrogen, oxygen, the halogens, and also to metals such as sodium or magnesium.
In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain 8 electrons.

8. Limitations Doesn’t allow for H, He or Li
Stable with 2 e- in their outer shells - Duet Rule
Transition elements - 18 electron rule
BF3 which only has 6 e- in its outer shell

9. How can atoms get a full outer shell?
Give away electrons
Take in electrons
Share electrons
Don’t bother getting a full outer shell at all – (Rare but does happen [BF3])

10. Ionic Bonding

11. Ar Noble Gas - Also Ne, He, Xe, Kr, Rn Stable or Unreactive40 18
Protons 18
18+
Electrons 18
18 -
Overall Charge
ATOM
Noble Gas -
Also Ne, He, Xe, Kr, Rn
Stable or Unreactive
Full outer shell
18+
All atoms want to be like this i.e. have a full outer shell - Octet Rule

12. Na Protons 11 11+ Electrons 11 11- Overall One electron in outer shell23 11
Protons 11
11+
Electrons 11
11-
Overall
ATOM
One electron in outer shell
Wants to get a full outer shell
11+
To have the same pattern as a Noble Gas
Loses electron from outer shell

13. Na 23 11
Protons 11
11+
Electrons 10 11
11-
10-
Overall 1+
ATOM
11+

14. Na+ Where has the electron gone? ION Protons 11 11+ Electrons 10 10-
Overall 1+
ION
Full outer shell
Same pattern as Neon
11+
Stable
Now has a positive charge
Called an ION
Where has the electron gone?

15. Cl Protons 17 = 17+ Electrons 17 = 17 - Overall = 035 17
Protons 17
= 17+
Electrons 17
= 17 -
Overall
= 0
ATOM
7 electrons in outer shell
Wants to get a full outer shell
17+
Takes an electron into its outer shell

16. Cl- ION Full outer shell Same pattern as Argon 17+ Stable
Protons 17
17+
Electrons 18
18 -
Overall
1 -
ION
Full outer shell
Same pattern as Argon
17+
Stable
Now has a - charge
Called an ION

17. Unlike charges attract
Cl-
Na+
Na+
Cl-

18. Summary Sodium atoms lose an electron to become sodium ions
Sodium ions have a +ve charge
Chlorine atoms gain an electron to become Chloride ions
Chloride ions have a − ve charge
Opposite charges attract so the Na+ and Cl- come together and stick to each other
This is called an ionic bond

20. General Points Ionic Bonding involves a transfer of electrons
It happens when the difference in electronegativity between atoms > 1.7
Metals lose electrons to form positive ions
They gain one plus charge for each electron lost
Non-metals gain electrons to form negative ions
They gain one minus charge for each electron gained
Total electron loss must equal electron gain

21. Keeping Track of Electrons
Atoms in the same group [column]
Have the same outer electron configuration.
Have the same valence electrons.
Easily found by looking up the group number on the periodic table.
Group 2A - Be, Mg, Ca, etc.-
2 valence electrons

22. General Points Group I lose 1 e- to become M+
Group II lose 2 e- to become M2+
Group III lose 3 e- to become M3+
Group V gain 3 e- to become X3-
Group VI gain 2 e- to become X2-
Group VII gain 1 e- to become X-
There are exceptions

23. Group I with Group VII Li F Na Cl K Br Rb I Cs At Fr Formula MX M+1 M

24. Group II with Group VI Be O Mg S Ca Se Te Sr Ba Po Ra Formula MX M2+ M

25. Group III with Group V B N Al P Ga As In Sb Tl Bi Formula MX M3+ M X3-

26. Group II with Group VII Be F Mg Cl Ca Br Sr I Ba At Ra Formula MX2 X

27. Group III with Group VIIX
X1- B Al Ga In Tl F Cl Br I At
X1- X M+
M2+
M3+ M
X1- X
Formula MX3

28. Group I with Group VI Be O Mg S Ca Se Sr Te Ba Po Ra Formula M2X M M+

29. Group I with Group V Li Na K Rb Cs Fr N P As Sb Bi Formula M3X M+ M X

30. Group III with Group VI O B S Al Se Ga Te In Po Tl Formula M2X3 X2- X

31. Group II with Group V Be Mg Ca Sr Ba Ra N P As Sb Bi Formula M3X2 M

32. Covalent Bonding
Peter Jackson

33. General Rules Involves a sharing of electrons
Electrons are shared in pairs
In each shared pair one electron comes from each atom
This type of bond occurs between elements of Groups IV, V, VI and VII
If there is a choice between ionic and covalent - covalent is preferred
There are exceptions

34. H2 Hydrogen H — H — Represents a shared pair of electrons.1+ 1+
H — H —
Represents a shared pair of electrons.
Called a Single Bond

35. Hydrogen - Alternative View
H — H H2
Peter Jackson

36. Chlorine
Only draw the outer shell electrons Cl Cl
Cl — Cl
Cl2

37. Chlorine – Alternative ViewCl Cl
Cl — Cl

38. Oxygen O O
Peter Jackson

39. Oxygen O O
O = O O2
Two pairs of electrons shared called a Double Bond

40. Oxygen – Alternative View
O = O O2

41. N ≡ N Nitrogen N2 Three pairs of electrons shared called a Triple Bond

42. Nitrogen – Alternative View
N ≡ N N2
Peter Jackson

43. Methane H
CH4 H C H H C H H H H
4 single C- H Bonds
Peter Jackson

44. Peter Jackson
Ozone O3 O O O O O O

45. Bonding

46. Electronegativity
SCC Science Dept

47. Is the relative attraction that an atom in a molecule has for the shared pair of electrons in a single covalent bond between two atoms of the same element
In a bond between two identical atoms the pair of electrons are shared equally
Chemists have found that in many bonds the pair of electrons are attracted to one of the atoms more than to the other.
SCC Science Dept

48. Hydrogen and Chlorine
Electrons attracted to chlorine more than to hydrogen [bigger, but more +ve nucleus]
Electrons spend more time near the chlorine than near the hydrogen
This gives the chlorine a slightly negative charge δ- delta minus
This gives the hydrogen a slightly positive charge δ+ delta plus
SCC Science Dept

49. H Cl
SCC Science Dept

50. Linus Pauling measured the electronegativity of each element and put them in a table
Noble gases are not in the table because they do not form bonds H Cl
+
 –
SCC Science Dept

51. Differences and Bond Type
Difference 0 to 0.45
[Pure] Covalent Bond
Difference > 0.45 but < 1.7
Polar Covalent Bond
Difference is = or > 1.7 then
Ionic Bond
SCC Science Dept

52. Trends in Electronegativity
Across a period
Goes up
Bigger nuclear charge
Smaller atomic radius [distance from nucleus]
Increased effective nuclear charge
SCC Science Dept

53. Increased Atomic Radius - electron is much further from nucleus
Down a Group
Gets less
Bigger nuclear charge
But! But! But! But!
Increased Atomic Radius - electron is much further from nucleus
Increased shielding by more inner electron shells
Decreased effective nuclear charge
SCC Science Dept

54. Intra and Intermolecular Forces
Intramolecular forces are the forces within a molecule
They are essentially
Ionic Bonds and Covalent Bonds.
They are powerful forces

55. Intermolecular Forces
Forces of attraction between molecules
They are generally weak
Three types
Dipole – Dipole [permanent]
Hydrogen Bonding [special type of permanent dipole – dipole]
Van der Waals’ Forces [transient or temporary dipole]

56. What is dipole?
Charge caused by the separation of the centres of + and - charge within an atom or molecule
Dipole-Dipole [Polar Bonding] caused by difference in electronegativity. Dipole permanent e.g. δ+ C=Oδ-
Hydrogen Bonding special dipole-dipole where H bonded to N, O or F. e.g. δ- O-H δ+ in water [1/10th]
Van der Waals’ Forces occurs in atoms or molecules where the difference in electronegativity is 0 and there should be no polarity. They are caused by temporary dipoles produced by random movement of electrons . E.g. H-H or He [about 1/1000 of covalent]

57. Why is NH3 soluble in water but PH3 is not?
NH3 is polar, PH3 is not. NH3 forms hydrogen bonds with the water, PH3 does not.
Metallic Bonds a lattice of +vely charged nuclei in a sea of electrons

58. Let us look at a Helium atom+
1. Centre of + and – charges are in the same place [centre] so there is an even distribution of charge

59. 2. Random movement of electrons results in two electrons being on one side of the atom. This side is now slightly negative δ- and the other side is therefore δ + + δ+
δ -

60. 3. Random movement of electrons results in two electrons being on another side of the atom. This side is now slightly negative δ- and the other side is therefore δ + + δ+
δ - + +
4. It is unlikely that the two electrons will be exactly opposite each other so the atom spends most of it time with one part slightly negative δ-and the other side slightly positive δ +
This is called a Temporary Dipole

61. Let us now look at a Neon atom
The bigger the atom the more electrons it has so there is more scope for a bigger charge differential between the two sides.
This is due to the possibility of more electrons being at one side.

62. +
1. Centre of + and – charges are in the same place [centre] so there is an even distribution of charge

63. +
2. Random movement of electrons results in more electrons being on one side of the atom. This side is now more slightly negative δ- and the other side is more δ +
δ +
δ -
This is a bigger Temporary Dipole
This is why neon has a higher boiling and melting point than helium

64. What is the effect of these Temporary Dipoles?
These atoms haves NO temporary dipoles at the moment.
They have a limited attraction to each other

65. δ - δ + This atom has developed a temporary dipole
This atom still has NO temporary dipole
δ -
δ +

66. δ - δ + δ - δ + This atom NOW has an induced temporary dipole
The attractive force between the two atoms is now greater due to the two dipoles

67. Since the force of attraction is greater it is more difficult to remove one of the atoms from the other.
The heat needed to give one enough energy to evaporate is greater
Thus the boiling point is greater than if there were no dipoles
Thus neon has a higher boiling point than helium.
Melting points are also greater for the same reason

68. Neon has a higher boiling point than helium because the atoms are bigger and therefore can produce greater temporary dipoles.
The boiling point of O2 is higher than H2 for the same reason
Molecules with dipoles tend to be soluble in water – like dissolves like

69. Everday experience of Van der Walls’ Forces

70. Graphite is a form of carbon [allotrope]
Graphite is pencil lead
Made up of sheets of carbon atoms in hexagons
Carbons are joined to each other by covalent bonds
Sheets are held to each other by van der Walls’ forces.

71. When you put the pencil on the paper the force of attraction between the paper and the graphite is greater than the attraction between the grapite sheets.
Thus when you move the pencil the bottom graphite sheet sticks to the paper
The pencils slides along and drops to the paper to repeat the process.
Pencils can write underwater
The american spent millions developing a pen that could write in space
The Russians were smart !!!!!!!!!!!!!!!
They brought pencils

72. Paper / graphite attraction
Pencil moves
Van der Waals’ Forces
Since the paper/grapite attraction is greater than the graphite/graphite attraction the bottom layer stays put as the upper layers moves to the side and there is now a black mark on the paper
Paper / graphite attraction

73. [ ] Dative Bonding + H3O+ H2O
Special case where one atom supplies both electrons
Two examples H2O → H3O+ and NH3 → NH4+ H+
H3O+
H2O
Proton [ ] + H O H O H
Hydroxonium Ion

74. [ ] Dative Bonding + NH4+ NH3 Ammonia Ammonium Ion
Special case where one atom supplies both electrons
Two examples H2O → H3O+ and NH3 → NH4+
Ammonia H+
NH3
NH4+
Proton [ ] + H N H N H H
Ammonium Ion