1. Chpt. 5: Chemical Bonding – Chemical Formulas

2. This topic will be investigated under five main headings:
Chemical Compounds
Ionic Bonding
Covalent Bonding
Electronegativity
Shape of Molecules and Intermolecular forces

3. Chemical Compounds Table Salt Water Dry Ice
In terms of elements can you think of a more scientific name for each of the above substances????

4. Chemical Compounds:
A compound is a substance that is made up of two or more different elements chemically combined.
2H2(G) + O2 (G) → 2H20(L)
The atoms of the elements in a compound are held
together by attractive forces called chemical bonds
Compounds can be broken down into their elements.
If an electric current is passed through water
(electrolysis) the compound splits into its elements of
hydrogen and oxygen

5. *Note: Valence electrons – are the outer electrons
Bonding involves only the valence electrons

6. The Octet Rule:
The Octet Rule states that when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell
Gilbert Lewis & Irving
Langmuir proposed the
Octet Rule
Elements will try to loose,
gain or share electrons to
achieve eight electrons in
their outer shell
This outermost energy level
is also known as the valence
shell

7. Exceptions to the Octet Rule:
Transition Metals - can have more OR less than eight electrons in their outermost energy level
Elements near Helium e.g. Hydrogen, Lithium etc. tend to achieve electronic arrangement of Helium -
2 electrons in outer shell!!!!

8. Loss & Gain of electrons

9. Valency Group number No. of Electrons in outer shell
Electrons needed/to be lost for octet rule
Valency 1
1 e-
1 e- (lost) +1 2
2 e-
2 e- (lost) +2 3
3 e-
3 e- (lost) +3 4
4 e-
4 e- (shared)
+/-4 5
5 e-
3 e- (gained) -3 6
6 e-
2 e- (gained) -2 7
7 e-
1 e- (gained) -1 8
8 e-
0 e-

10. An ion is a charged atom or group of atoms
Ions
An ion is a charged atom or group of atoms
Metals (left side of periodic table)
- lose valence electrons - achieve a stable valence shell (usually 8 e-) - gains a positive charge, i.e. a positive ion.

11. Loss of electrons (formation of cations):
Na 2,8,1 -> Na+ + e  2,8
Li 2,1 -> Li+ + e  2
K 2,8,8,1 -> K+ + e  2,8,8
Mg 2,8, > Mg2+ + 2e-  2,8
Ca 2,8,8, > Ca e-  2,8,8
Al 2,8, > Al e-  2,8.

12. Non-metals (right side of periodic table):
- gain valence electrons - achieve a stable valence shell (usually 8 e-)
- gains a negative charge i.e. a negative ion

13. Gain outer electrons (formation of anions):
F , e- -> F ,8
Cl ,8, e- -> Cl ,8,8
Br ,8,8,7 + e- -> Br- 2,8,8,8
I ,8,18,18,7 + e- -> I ,8,18,18,8
O , e- -> O2- 2,8
S ,8, e- -> S2- 2,8,8
N ,5 + 3e- -> N3- 2,8
H e- -> H- 2.

14. Uses of unreactive group – Noble Gases
Helium is a much safer alternative to hydrogen as it is stable and is therefore used in weather balloons and blimps.
Argon is used in electric light bulbs to prevent the tungsten filament from evaporating or reacting.

15. Lewis Symbols – ‘Dot and Cross Diagrams’
Lewis symbols show the valence electrons as dots arranged around the atomic symbol.
hydrogen:
sodium: Na
chlorine: Cl H
· ·
· ·

16. Sodium & Chlorine

17. 2 Ionic Bonding
Ionic bonding is the force of attraction between oppositely charged ions in a compound.
Remember:
Ions are elements which have a positive or negative charge e.g.
Na has 11 e- (E.C.= 2,8,1) when Na gives away this one e- it now has more protons than electrons so it has an overall positive charge.
Ionic bonds generally form between metals and non-metals.

18. Bohr Type Circle Diagram
Formation of Ionic Compounds Using Ionic Bonding
Formation of Table Salt (Sodium Chloride):
Bohr Type Circle Diagram
Dot and Cross Diagram

19. Example 1: Formation of Sodium Chloride – Bohr-type diagram
Step 1: Draw Bohr diagrams of each atom
11p
11n
18n
17p
Na atom
2,8,1
Cl atom
2,8,7

20. Step 2: Determine ions formed
AND
17p
18n
Na+ ion
2,8
Cl- ion
2,8,8

21. Step 3: The formation of an Ionic Bond
The attraction between the positive sodium ion and the negative chlorine ion results in the formation of an ionic bond to give the ionic compound Sodium Chloride:

22. Na + Cl Na Cl Example 1: Formation of Sodium Chloride
– Dot-and-Cross Diagrams
Step 1: Represent atoms using dot-and-cross diagram
Step 2: Determine ions formed: xx Na + Cl xx x xx
Na atom
2,8,1
Cl atom
2,8,7 __ xx + Na
and Cl xx x xx
Na+ ion
2,8
Cl- ion
2,8,8

23. [ ] ® Na+ Cl Na + Cl Step 3: Formation of Ionic Bond: Sodium Chloridexx
Na+ [ ] Cl xx Na + Cl xx x x xx xx xx
Sodium Chloride

24. Example 2:
Show the formation of the ionic bond in magnesium fluoride, MgF2, by means of dot-and-cross diagrams
Try:
Show the formation of the ionic bond in magnesium chloride by means of dot-and-cross diagrams

25. Writing Formulas of Ionic Compounds
A chemical formula is a way of representing a compound using symbols for the atoms present and numbers to show how many atoms of each element are present.
You must know how to write the formulas of ionic compounds of the first 36 elements.
When writing formula remember: an ionic compound has no net charge, overall it is neutral
*Note: transition metals will be discussed separately

26. *Note:
A compound which contains just two elements always
ends in - ide
A compound which contains oxygen as well as two
other elements always ends in -ate

27. Working with Simple Ions:
Write the formula of:
i) Potassium bromide
ii) Calcium Chloride
iii) Sodium Sulphide
iv) Aluminium Oxide
H+, H-
Na+ ,K+
Be2+, Mg2+, Ca2+
Al3+
O2- ,S2-
F-, Cl-, Br-, I-,

28. Potassium Bromide 1- + K Br
*Note:
The potassium ion is K+.
The bromide ion is Br-
Same number of positive and negative charges

29. Calcium Chloride 1- Ca 2+ Cl
*Note:
The calcium ion is Ca2+.
The chloride ion is Cl-
In order to have the same number of negative
charges as positive must use two Cl- ions

30. Sodium Sulphide 2- + S Na
*Note:
The sodium ion is Na+.
The sulphide ion is S2-
In order to have the same number of positive
charges as negative must use two Na+ ions

31. Aluminium Oxide
3 +
2 - Al O
*Note:
Bring all charges up to their lowest common
denominator – 6
To get six positive charges we need 2 Aluminium
ions
To get six negative charges we need 3 Oxide ions

32. Working With Complex Ions:
Must learn off by heart!!!!
Name
Formula
Hydroxide Ion
Nitrate Ion
Hydrogen carbonate Ion
OH-
NO3-
HCO3-
Carbonate Ion
Sulfate Ion
Sulfite Ion
CO32-
SO42-
SO32-
Ammonium Ion
Phosphate Ion
NH4+
PO43-

33. Write the formula of the following:
Potassium Hydroxide
Calcium Hydroxide
Sodium Sulphate
Ammonium Phosphate

34. Potassium Hydroxide + 1- K OH
Note:
The potassium ion is K+.
The hydroxide ion is OH-
Same number of positive and negative charges

35. Calcium Hydroxide 2+ Ca OH 1-
*Note:
The calcium ion is Ca2+.
The hydroxide ion is OH-
In order to have the same number of negative charges
as positive we must use two OH- ions

36. In order to have the same number of positive charges
Sodium Sulphate +
SO42- Na
*Note:
The sodium ion is Na+.
The sulphate ion is SO42-
In order to have the same number of positive charges
as negative charges we must use two Na+ ions

37. The phosphate ion is PO43-
Ammonium Phosphate 3- 1+
PO4
NH4
*Note:
The ammonium ion is NH4+.
The phosphate ion is PO43-
In order to have the same number of positive charges
as negative charges we must use three NH4+ ions

38. Valency of Transition Elements
Transition elements have a variable valency (they loose
electrons)
The valency depends on the other elements the
transition element is bonding to:
Transition Metal Ion
Example
Iron (II)
FeCl2
Iron (III)
FeCl3
Copper (I)
Cu2O
Copper (II)
CuO
Chromium (III)
CrCl3
Chromium (VI)
Na2Cr2O7
Manganese (IV)
MnO2
Manganese (II)
MnSO4
Manganese (VII) +7
KMnO4

39. The reason for this variable valency is because there is
such a small energy difference between the 4s and 3d
sublevels
*Note: You do not need to remember the actual valancies of any of the transition metals but you should be able to interpret the names as shown in the following examples

40. Working with Transition Metals
1) Write the formula of:
i) Iron (II) carbonate
ii) Chromium (III) Chloride
2) Name the compound Cr2(SO4)3

41. Iron (II) Carbonate
CO3 2+ 2- Fe
*Note:
The Roman number (II) indicates Fe2+
The carbonate ion is CO32-
Valencies are balanced

42. Chromium (III) Chloride
Cl1- 3+ Cr
*Note:
The Roman number (III) indicates Cr3+.
The chloride ion is Cl-
In order to have the same number of negative charges
as positive charges we must use three Cl- ions

43. Name the compound Cr2(SO4)3
One sulphate ion has a negative charge of 2- so
three sulphate ions together must have a total
negative charge of 6-
Since overall charge must be zero the two
chromium ions together must carry a charge of 6+
Therefore each chromium ion must be Cr3+
Compound Name:
Chromium (III) Sulphate

44. D- block elements and transition elements
All of the elements from scandium to zinc are the first row of the d-block
All of the elements from titanium to copper incl. are referred to as transition elements

45. Transition elements have the following characteristics:
In d-block of table
Variable valency
Note: Sc only forms Sc3+ ions and Zn only forms Zn2+
ions
Usually form coloured compounds
Note: Sc and Zn only form white compounds
Widely used as catalysts
Note: Sc and Zn show little catalytic activity
Have incomplete d sublevel
Note: Sc3+ has empty d-sublevel and Zn2+ has full
d-sublevel

46. A transition metal is one that forms at least one ion with a partly filled d-sublevel
Lets take a closer look:
Even though the 4s sublevel is filled before the 3d sublevel (as it is lower in energy) electrons are lost from the 4s first as location wise it is further from the nucleus and are therefore more easily removed:
4s electrons are lost before 3d electrons
Investigate e.c. of Sc, Sc3+, Fe, Fe2+ , Ti and Ti3+

47. Ionic Compounds and their Crystal Lattice Structure
Ionic bonds result in a three dimensional crystal lattice structure.
Unit Cell
Sodium Chloride has a cubic structure

48. Uses of Ionic Substances in Everyday Life
(Know at least two)
Too little salt in diet = muscle cramps, so many athletes
take salt tablets to replace salt lost in sweating.
Food preservative (Brine)
Salt used in manufacture of soap, leather, detergents
Salt is spread on roads during winter to help melt frost
and snow
Fluoridation of water supplies to prevent tooth
decay

49. Properties of Ionic Compounds
Contain positive and negative ions (Na+Cl-)
Solids such as table salt (NaCl(s))
High melting and boiling points
Strong force of attraction between particles
Separate into charged particles in water to give a solution that conducts electricity.

50. 3. Covalent Bonding
A covalent bond is the chemical bond formed by sharing a pair of electrons
Covalent bonds are typical of non-metal elements.
A single bond has 1 shared pair of electrons.
A double bond has 2 shared pairs of electrons.
A triple bond has 3 shared pairs of electrons.
E.G. H-H O=O N N
A molecule is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.

51. Examples of single covalent bonds: Hydrogen Molecule (H2): H H
Both hydrogen atoms want a full outer shell, but as they both only need one electron they each share an electron. Thus, each hydrogen atom achieves a stable electron structure like the noble gas Helium.
H H X X
Covalent bond

52. Examples of single covalent bonds:
1. The chlorine molecule
2. The water molecule
3. The ammonia molecule
4. The methane molecule

53. The chlorine molecule (Cl2):

54. 2.The Water Molecule (H2O):
*Note: There are two pairs of electrons on the oxygen atom that are not involved in bonding. These are called lone pairs.
Lone Pairs – not bonded to another atom
Bond Pairs – pairs of electrons that are involved in bonding
Thus, the oxygen atom in a water molecule has two lone pairs and two bond pairs

55. 3. The Ammonia Molecule (NH3):

56. 4. The Methane molecule (CH4):
* Note: a single bond is formed when one pair of electrons is shared between two atoms.

57. Valency Can predict number of covalent bonds formed around
certain atoms.
Hydrogen is said to be monovalent i.e. one hydrogen
atom can only combine with one atom of any other
element
The valency of an element is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.
e.g. Chlorine has a valency of 1 – HCl
Oxygen has a valency of 2 – H2O
Nitrogen has a valency of three – NH3

58. Double and Triple Bonds
In some molecules the atoms share two or even three pairs of electrons
When two pairs of electrons are shared (between two
atoms), we say that a double bond is formed
Examples of double covalent bonds:
1. The oxygen molecule
2. The carbon dioxide molecule

59. 1. The oxygen molecule (O2):

60. 2. The carbon dioxide molecule (CO2):

61. When three pairs of electrons are shared (between
two atoms), we say that a triple bond is formed
Example of triple covalent bonds:
1. The nitrogen molecule
The nitrogen molecule (N2):
*Note: the number of electron pairs is the bond order

62. Combining single, double and triple covalent bonds
Draw electron dot structures showing the bonding in each of the following:
a) HCHO
b) CH3OH
c) COCl2
d) HOCl
Steps:
Determine number of bonds required for each atom
Draw skeleton structure with no bonds (note: order of atoms)
Complete the octet rule for each atom by assigning appropriate bonds

63. Sigma ( σ ) and Pi ( π ) Bonding
It is also possible to describe chemical bonding in terms of atomic orbital’s
Sigma Bonding (σ):
A sigma bond is formed when electrons are shared in line with the nuclei – a head on overlap of orbitals
example H2 molecule
1s orbital 1s orbital
Molecular orbital H2

64. *ALL single bonds can also be called sigma bonds
Sigma Bonding:
Example Cl2 molecule:
E.C. : 1s2, 2s2, 2p6, 3s2, 3px2, 3py2, 3pz1
Two half filled pz atomic orbitals overlap head on to
form a sigma bond
*ALL single bonds can also be called sigma bonds

66. Pi Bonding (π):
A pi bond is formed when the shared orbitals are side on – not in line with the nuclei
example O2 molecule -
E.C. 1s2,2s2 2px2, 2py1 2pz1
Direct overlap
Between the two
2py1 of each atom
Sigma bond
Sideways overlap
between The two
2pZ1 of each atom
Pi bond

67. From our study of covalent bonding we know that the oxygen molecule is formed by a double covalent bond which we now know consists of 1 sigma and 1 pi bond
example N2 molecule: - E.C. 1s2, 2s2, 2px1, 2py1, 2pz1
- 3 half filled p orbitals
- 2px orbitals overlap head on – sigma bond
- 2py orbitals overlap sideways – pi bond
- 2pz orbitals overlap sideways – pi bond

68. A covalent single bond is a sigma bond
A covalent double bond consist of one sigma and one
pi bond
A covalent triple bond consists of one sigma and two
pi bonds
*Note: Since there is less overlapping of orbitals in
forming pi bonds these bonds are not as strong as
sigma bonds

69. Characteristics of Ionic & Covalent Compounds
Remember ionic compounds usually consist of a network of ions whereas covalent compounds generally consist of individual molecules. This fact helps is understand many of the properties of ionic and covalent compounds

70. Characteristics of Covalent Compounds
Property
Explanation
Contain individual molecules
Solids, liquids, or gases at room temperature
(C6H12O6(s), H2O(l), CO2(g)
Soft
Consist of molecules e.g. One crystal of iodine can be made up of millions of iodine molecules (I2) with only weak forces of attraction between one molecule and another
Low melting and boiling points
During melting/boiling, molecules become separated. Forces of attraction between molecules are weak and little energy is required to separate them.

71. Property
Explanation
Do not conduct electricity
No mobile charged particles Molecules not charged Electrons tightly bound to atoms or shared by atoms in covalent bonds.

72. 5. Shapes of Covalent Molecules
Ionic compounds consist of giant crystal lattices but covalent compounds consist of separate molecules with each individual molecule having a particular shape

73. Shapes of Covalent Molecules
The Valence Shell Electron Pair Repulsion Theory
(VSEPRT)
The shape of a molecule depends on the number of pairs of electrons in the valence shell of the central atom
(i.e. that lie around the central atom of the molecule)
Since electrons are negatively charged, the electron
pairs repel each other and arrange themselves in space
so as to be as far apart as possible.
The following shapes will arise:
- Linear - Pyramidal
- Triangular Planar - Planar/V-Shaped
- Tetrahedral

74. Linear
Beryllium Chloride (BeCl2) has 2 bond pairs of electrons around the central atom.
These bond pairs (-ive) repel each other as far as possible resulting in a linear shape.
The bond angle is 180o.

75. Triangular (Trigonal) Planar
Boron trichloride (BCl3) has 3 bond pairs of electrons around the central atom.
The three pairs of electrons repel each other as far as possible resulting in a triangular planar shaped molecule.
The bond angle is 120o.

76. Tetrahedral (Tetrahedron)
Methane (CH4) has 4 bond pairs of electrons around the central atom.
The four pairs of electrons repel each other as far as possible resulting in a tetrahedral shaped molecule.
The bond angle is 109.5o.
C-H bond in plane of paper
The wedge drawn indicates that this particular C-H bond is coming out of the plane of the paper
The dashed line indicates the C-H bond is going behind the paper

77. In three examples given all pairs of electrons have been bond pairs i
In three examples given all pairs of electrons have been bond pairs i.e. no lone pairs present!!! However, some molecules have lone pairs present which must be taken into account.
The VSEPR Theory has been adapted to account for lone pairs:
*Lone pair/lone pair > lone pair/bond pair > bond pair/bone pair*
i.e. Repulsion between two lone pairs in a molecule is greater than the repulsion between a lone pair and a bond pair etc.

78. Using the VSPER Theory to predict shapes of molecules
Determine number of electrons around central atom.
Work out the number of electrons contributed by the other atoms to the bonds around the central atom
Work out the total number of electron pairs around the central atom.
Taking lone pairs into account decide on the shape of the molecule.

79. Pyramidal (Distorted tetrahedral)
Using the VSEPR Theory deduce the shape of the ammonia molecule
Step 1: Nitrogen has 5 electrons in its outer shell.
Step 2: Each hydrogen atom contributes one electron so a total of 3 electrons will be contributed to the bond pairs
Step 3:
N = 5 electrons
3H = 3 electrons
Total = 8 electrons
= 4 pairs of e-
suggests tetrahedral shape but ammonia has three bond pairs and 1 lone pair

80. Step 4: However, the lone pair of electrons in ammonia repel the bond pairs more than the bond pairs repel the lone pair, resulting in a pyramid shaped molecule.
The bond angle is 107o.
Remember:
LP/LP>LP/BP>BP/BP

81. V Shaped (Distorted tetrahedral)
Using the VSEPR Theory deduce the shape of the water molecule
Step 1: Oxygen has 6 electrons in its outer shell.
Step 2: Each hydrogen atom contributes one electron so a total of 2 electrons will be contributed to the bond pairs
Step 3:
O = 6 electrons
2H = 2 electrons
Total = 8 electrons
= 4 pairs of e-
suggests tetrahedral shape but the water molecule has 2 bond pairs and 2 lone pairs

82. Since LP/LP>LP/BP>BP/BP the pairs of electron repel each other to result in a V-shaped/planar molecule.
The bond angle is 104.5o

83. Use the VSEPR Theory to deduce the shape of the following molecules:
i) SiCl4
ii) PBr3

84. Shapes of Molecules Summary
Number of Bond Pairs
Number of Lone Pairs
Shape
Bond Angle
Example 2
Linear
1800
BeCl2 3
Triangular Planar
1200
BCl3 4
Tetrahedral
109.50
CH4 1
Pyramidal
1070
NH3
V-Shaped
104.50
H2O

85. 4. Electronegativity
Dividing chemical compounds into two groups, ionic and covalent, simplifies the subject of bonding to a considerable extent. Many compounds have bonding which is neither fully one nor the other, their properties suggest that they are partially ionic and partially covalent.
Electronegativity is a property which can be used to predict the type of bond formed when two elements combine

86. In a covalent bond between identical atoms e.g. H2, Cl2
the pair of electrons is shared equally between the
two atoms in the molecule.
In bonds between different atoms, the pair of electrons
is often more attracted to one of the atoms than the
other e.g. HCl.
In HCl the two electrons in the bond are more attracted
to the chlorine than to the hydrogen. As a result this
gives chlorine a slightly negative charge and leaves the
hydrogen atom with a slightly positive charge.
H ---- Cl δ+ δ-

87. Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond i.e. chlorine is more electronegative than hydrogen
Linus Pauling developed a scale of relative values of
electronegativity:

88. From Pauling’s table of electronegativity values (pg 61/log tables) it can be seen that:
- the more electronegative elements are the very reactive non-metals – fluorine, chlorine etc.
- the least electronegative elements are the very reactive metals like potassium and sodium. These are said to be electropositive

89. Polarity:
Polarity means that there is a positive and negative charge on the molecule and these two charges are separated by some distance.
In a polar bond, electrons are shared unequally
because of the difference in electron density:
e.g. HCl is a polar covalent molecule
HCl:

90. In a non polar bond (purely covalent) the bonding electrons are shared equally:
*Note: in a polar covalent substance:
- the element with the lower electronegativity is δ+ and is written first
- the element with the higher electronegativity is δ- and is written second
e.g. HCl H --- Cl
H2,
Cl2: δ+ δ-

91. Uses of electronegativity values:
The difference in the electronegativity values of two
joined atoms can be used to tell
1. whether the bond between them is ionic or
covalent
2. if covalent how polar that bond is

92. Is bond ionic or covalent???
If predicting whether a bond is ionic or covalent the
following rules should be applied:
E.N. Difference = → non polar bond (purely covalent)
E.N. Difference ≤ → bond is slightly polar
E.N. Difference > 0.4 < 1.7 → bond is polar (polar covalent)
E.N Difference > → bond is ionic
*Note:
Polar covalent bonds are intermediate between pure
covalent bonds and ionic bonds and have therefore a
certain amount of ionic character
When amount of ionic character is more than 50 % bond is ionic

93. Example :
Determine if potassium fluoride and methane are ionic or covalent compounds KF
E.N. Difference = 4.0 – 0.8 = 3.2 i.e. >1.7 → ionic bonding therefore this is an ionic compound
CH4
E.N. Difference between carbon and hydrogen atom = 2.5 – 2.1 = 0.4 i.e. < 1.7 → covalent bonding therefore this is a covalent compound

94. Exceptions to the rule:
LiH
NaH KH
CaH2
Using rule would indicate these molecules all contain covalent bonding when in fact they are all ionic compounds.
Reason for this is the presence of the hydride ion H-

95. Remember 2. Predicting polarity of covalent bonds
The greater the electronegativity difference the more
polar the bond

96. (Note please leave space of half a page to complete this exercise)
Investigate the polarity of the following covalent compounds:
H2O
HCl
CH4 H2
(Note please leave space of half a page to complete this exercise)
The first 3 examples above have polar bonds and can therefore be considered polar molecules, however, some molecules have polar bonds but are considered non polar molecules

97. Remember
Polarity means that there is a positive and negative charge on the molecule and these two charges are separated by some distance.
Let us look at the polarity of covalent compounds when taking their shape into account!!!!!
Water:
H2O - E.N. Difference 1.4 = polar bonds therefore a polar molecule
- Shape V-shaped = centre of + and – charges are in different places.
The centre of + is between 2 H’s while centre of – is in the O
Centre of -ve
δ2- δ+ H H δ+
Centre of +ve

98. Carbon Dioxide:
CO2 - E.N. Difference: 1.0 = polar bonds therefore you would expect polar
molecule
- Shape: Linear = centre of + and – charges are both in centre of C atom (i.e not separated by a distance )
- Since the centre of gravity of the partial + and – charges coincide carbon dioxide is considered non polar
Centre of +ve
δ2+ O O δ- δ-
Centre of -ve

99. By considering the shape and the polarity of the bonds in the following covalent molecules decide whether they are polar molecules or not: HF
NH3
BCl3
CCl4

100. Mandatory Demonstration: Polarity Test for liquids
(Water & Cyclohexane)
Bring a charged rod towards a thin stream of liquid
from a burette
If polar the stream will be attracted and bend towards
the rod
If polarity of rod changed- stream will still bend
towards it, but opposite side of molecule will be
attracted
If non polar liquid used stream will not bend towards
liquid
**Note: Diagram required
** Do on board

101. Importance of polarity:
One of the most important properties of water is that it
is an excellent solvent – this property depends on the
fact that it is a polar molecule
Most ionic substances and most polar covalent
substances dissolve in water e.g. NaCl (ionic substance)
The +ve sodium ion in NaCl is attracted to the –ve partial charge of oxygen in water molecule. Also, the –ve chlorine ion is attracted to the partial +ve charge of hydrogen in water molecule. As a result the crystal lattice structure of NaCl breaks up i.e. salt dissolves

102. Rule for solubility: ‘like dissolves like’
Polar substances like HCl and NH3 readily dissolve in water (polar)
Most non polar substances do not dissolve in water
Non polar liquids such as methyl benzene and cyclopropane dissolve non polar substances such as wax and oil

103. (Salt, potssium iodide and iodine in cyclohexane and water)
Mandatory Demonstration:
Testing solubility of ionic and covalent substances in different solvents
(Salt, potssium iodide and iodine in cyclohexane and water)

104. 5. Intramoloecular and Intermolecular Bonding
Intramolecular Bonding
- forces/bonds within molecules
- ionic and covalent bonding are examples of intramolecular bonding
Intermolecular Bonding
- forces/bonds that exist between one molecule and another
- 3 types of forces between molecules:
- van der Waals forces
- Dipole-Dipole
- Hydrogen Bonding

105. Dipole: is a pair of separated opposite charges
*Note:
Dipole: is a pair of separated opposite charges

106. Van der Waal’s:
These are the weakest forces caused by the movement of e- within a molecule. (The electrons move randomly within the bond so at 1 point in time they are nearer to 1 atom than the other i.e. one side of molecule slightly +ive leaving other side slightly –ive.) This induces a temporary dipole force.
This temporary dipole could induce a similar dipole in a nearby molecule.
There is then an attraction between the opposite charges – van der Waals force.

107. Diagram: Please leave space for small diagram
Van der Waals forces are weak attractive forces between molecules resulting from the formation of temporary dipoles

108. Van der Waals Forces:
are the only forces of attraction that exist between
non polar molecules e.g.H2 , O2 , N2 , I2 etc.
gives low melting and boiling points
strength of van der Waals forces increases as
molecule gets bigger – this is a result of the greater
number of electrons present which allow the
temporary dipoles to form more easily

109. At room temperature: Boiling points:
Cl2 = gas H2 = -252oC
Br2 = liquid O2 = -183oC
I2 = solid
*Note: Higher boiling point is result of more heat required to break stronger van der Waals forces between bigger molecules of oxygen
Van der Waals forces also explain difference in boiling
points of organic compounds – C3H8 (propane)
and C4H10 (butane)
Which has higher boiling point and why?????

110. 2. Dipole-Dipole Forces
The positively charged end of a polar molecule is
attracted to the negative end of another polar molecule.
The dipoles in this case are permanent and as a result they are stronger than Van der Waal’s forces.
Dipole-dipole forces are forces of attraction between
the negative pole of one molecule and the positive
pole of another.

111. Diagram: Please leave space for two diagrams
Dipole-Dipole force in HCl
Dipole-Dipole force in acetone (CH3)2 CO)

112. Dipole-Dipole Forces
Results in higher melting and boiling points than
similar mass non-polar compounds e.g.
Boiling Point
HCl = -85o C
H2 = -252oC
Similarly,
C2 H4 (ethene) = -104o C (non-polar)
HCHO (methanal) = -21o C (polar grp – C=0)
HCl has much higher boiling point as enough heat must be supplied to overcome the dipole-dipole forces present in HCl

113. 3. Hydrogen Bonding
this is a specific type of dipole-dipole force and
occurs when H is bonded to a more
electronegative element such as F, N, or O (the 3
most electronegative elements in the periodic
table)
Explanation of hydrogen bonding - example H2O
- The hydrogen atom has a strong positive charge because it is bonded to a very electronegative element – oxygen
- This causes a strong polarity between all the water molecules and they line up in strong molecular chains.

114. Diagram: Please leave space for two diagrams
Hydrogen Bonding in HF
Hydrogen Bonding in H2O

115. Causes much higher boiling points as a lot of energy
is required to break these molecular chains allowing
molecules to escape from the liquid to form a gas
H2O oC
H2S oC
Similarly,
NH oC HF oC
PH o C HCl oC
Hydrogen bonds are much stronger than Van der
Waals forces and dipole-dipole forces
H2S has a lower boiling point than water even though sulphur is in the same group as oxygen and its molecule is bigger, because the effect of polarity is not as great as it is in water.

116. Note: the boiling point of. is higher
*Note: the boiling point of is higher than that of HF - 20o C (even though hydrogen bonds in HF are stronger as fluorine is more electronegative than oxygen)
This is because there are twice as many hydrogen bonds for each water molecule as there are for every molecule of HF
H2O oC

117. Hydrogen bonds are particular types of dipole-dipole attractions between molecules in which hydrogen atoms are bonded to nitrogen, oxygen or fluorine. The hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule. Thus, it acts as a bridge between two electronegative atoms.

118. Hydrogen Bonding:
- important in synthetic clothing – Kevlar
- helps fibres absorb water – wool
- ammonia is extremely soluble in water because:
a) polar molecule like water
b) forms hydrogen bonds with water
- explains why water has high surface tension