1. Properties of Solids
SCH4U1

2. Intra vs. Intermolecular Bonds
The properties of a substance are influenced by the force of attraction within and between the molecules.

3. Intra vs. Intermolecular Bonds
Intramolecular Bonds: Bonds within a molecule (covalent or polar covalent) Intermolecular Bonds: Bonds between molecules.

4. Intermolecular Forces
The physical properties of a molecule (e.g. melting point) are mainly due to the strength of intermolecular bonding:
H2O (s)  H2O (l)  H2O (g)  2H + O
MP = 0oC BP =1000C Decomposes: >2000oC
intermolecular bonds breaking intramolecular bonds breaking

5. 1) Atomic Solids
Noble gases form liquids and solids at very low temperature due to the very weak bonds between the atoms.
Since the attraction is so weak, they are weakened and broken at very low temperature.
E.g. argon (Ar): mp = -189oC; bp = -186oC

6. van der Waals (London) Forces
Since they do form solids, a very weak attraction must exist between the atoms.
These are explained by weak attractions between molecules called van der Waals forces.
London forces are the weakest type of van der Waals attraction.

7. London Forces
London forces form due to the attraction between instantaneous dipoles (charge imbalances) that form in the atoms.
At low temperature can induce dipoles in other atoms, causing solidification of helium:

8. London Forces and Electrons
As the number of electrons in an atomic solid increases, the mp/bp also increases.
Group Electrons Boiling Point ('C) He Ne Ar Kr Xe Rn

9. Summary: Properties of Atomic Solids
very low melting points/ boiling point
do not conduct electricity
mp/bp increase down the group (increasing London forces)
Liquid helium is a very strange substance.

10. Molecular Solids
Substances with covalent bonds or polar covalent bonds (N2, CH4, H2O, C6H12O6 etc.).
Can exist in all states at room temperature.
Check out liquid nitrogen!

11. 2) Non-Polar Molecular Compounds
Compounds without bond dipoles only have London Forces between molecules.
This results in LOW bp/mp
Group 17
Boiling Point ('C) F2
-188.1
Cl2
-34.6
Br2
58.8 I2
184.4

12. BOILING POINTS OF RELATED MOLECULAR COMPOUNDS
Formula
Number of Electrons
Boiling Point (oC)
CH4 10
-161
SiH4 18
-112
GeH4 36
-90
SnH4 54
-52

13. Comparing Larger Compounds
When comparing non-polar compounds, the forces of attraction are greater between molecules with the greatest number of atoms.
There are more locations for London (van der Waals) forces to occur between adjacent molecules.

14. Boiling Points of Hydrocarbons
Molecular
Formula
Boiling Point (oC)
State at STP
CH4
-161.5
gas
C2H6
-88.6
C3H8
-42.1
 gas
C4H10
-0.5
C5H12
36.1
liquid
C6H14
68.7
 liquid
C10H22
174.1
C22H46
327
solid

15. 3) Polar Molecular Compounds
Compounds with bond dipoles AND molecular dipoles (e.g. HCl, H2S, CF2H2) have higher boiling points.
This is due to intermolecular forces between permanent dipoles.
These are called dipole-dipole forces or bonds.
Dipole-Dipole Force (Bond)

16. Boiling Points of Some Polar and Nonpolar Substances
Boiling Point (oC)
Molar Mass (g/mol)
Number of Electrons
HCl
polar
(molecular dipole)
-84.9 36 18
H2S
-60.7 34 F2
nonpolar
(NO molecular dipole)
-188.1 38 Ar
-185.7 40

17. 4) Polar Molecules: Hydrogen Bonding
If hydrogen is bonded to a VERY electronegative atom (F, O or N), a very strong dipole forms.
These atoms are also very small, concentrating this positive and negative charge.
Dipole-dipole bonds between molecules containing O-H, N-H or H-F bonds form “hydrogen bonds”.

18. Water “bends” near a charged object.

19. Properties of Hydrogen-bonded Molecules
A hydrogen bond is about 10x weaker than a covalent bond BUT 10x stronger than a normal dipole-dipole bond
Thus H-bonded molecules have the highest mp/bp of the molecular compounds:
mp (oC) bp (oC)
Propane (C3H8)
Propanol (C3H7OH)
Glycerol (C3H6(OH)3)

20. General Properties of Molecular Compounds
Molecular compounds do not conduct electricity sine their electrons cannot move between molecules.
They have relatively low bp/mp due to the existence of weaker intermolecular forces
As the strength of these intermolecular forces increase, so does the mp and bp.

21. Unusual Properties of Water
Water is called the “universal solvent” since it dissolves both polar molecules (e.g. sugar) and ionic compounds (e.g. NaCl).
Water expands when it freezes due to the organization of the many hydrogen bonds in the solid.

22. 5) Metallic Solids General Properties: Few valence electrons
Low ionization energies
Malleable, ductile and shiny
Moderate mp/bp
Good conductors of heat and electricity in the solid and liquid states.

23. Metallic Bonding
Metal properties can be explained by considering them as postivie ions in an “electron sea” or “electron cloud”
Delocalized or conduction electrons are shared among multiple cations are free to move throughout a crustal of positive ions.

24. Explaining Metallic Properties
Property
Explanation
Conductivity (Electricity / Heat)
Delocalized electrons can move between ions.
Ductility and Malleability
The plane of ions can move by distorting the electron cloud.
Lustre
Reflection is caused by loosely bonded electrons absorbing and remitting all wavelengths of light.

25. e. g. 1 Lithium is far more malleable than aluminum
e.g. 1 Lithium is far more malleable than aluminum. Propose an explanation for this observation using the model of metallic bonding.
Metallic bonding occurs since the loosely held (delocalized) electrons are mutually shared by a crystal of positive ions. Since Li has only 1 delocalized valence electron compared with aluminum which has 3 and aluminum has a greater nuclear charge, we can deduce that the additional protons & electrons strengthen the metallic bonding and make it more difficult to displace the network of atoms in the crystal.

26. e.g. 2 Which element would require the most energy to undergo vapourization, K or Sc? Explain.
Scandium.
The stronger the metallic bonding, the more energy required to change state. Similar explanation as above…..scandium has more delocalized electrons.
Kl (l) kJ K (g)
Sc (l) kJ Sc (g)

27. 6) Ionic Solids
Solids formed by ionic bonds between metal cations (+) and non-metallic anions (-).
Bonded together by a 3D array or crystal lattice without distinct molecules.

28. Properties of Ionic Solids
High melting points and boiling points (many ionic bonds that must be broken to change states).
Hard but brittle.
Many are soluble in water.
DO NOT conduct electricity in the solid state since ions cannot move.
DO conduct electricity in the liquid or aqueous states since charged ions are mobile.

30. Crystal Packing
Properties of ionic compounds are related to the packing of the crystals:

31. Factors Affecting the Strength of Ionic Bonding
1. Ionic Radius of the Cation and Anion: As the radius of the ions increases, the attraction between oppositely charged ions decreases.

32. Factors Affecting the Strength of Ionic Bonding
2. Ionic Charge: As the charge of the cation and anion increases, the attraction increases.
Melting Point (oC)
Solubility
(g/100g 0oC)
CsCl
MgO
NaCl

33. 7. Covalent Network Solids
Form a lattice of continuous covalent / polar-covalent bonds.
Do not contain molecules.
Very hard, brittle substances.
Most do not conduct since electrons are either in sigma bonds or lone pairs (filled orbitals).
Some exist as different allotropes (forms with different properties)

34. Quartz: A Common Network Solid
Quartz (SiO2) and Feldspars (KAlSi3O8 , NaAlSi3O8 & CaAl2Si2O8) make up most of the Earth`s crust.
Quartz is a continuous framework of tetrahedral SiO4

35. Comparing CO2 and SiO2 Property Carbon dioxide (CO2) Quartz (SiO2)
Type of Solid
Non-polar Molecular
Covalent Network
Melting Point (oC)
-78(sublimates at 1 atm)
1650
Boiling Point oC)
N/A
2230
Bond angle (o)
180o
109o
Geometry
Linear (sp)
Tetrahedral (sp3)
Intramolecular bond
Type(s)
Polar covalent
Intermolecular bond Types
London forces

36. Allotropes of Carbon: Diamond
Covalent network of sp3 hybridized carbon (tetrahedral).
Very hard; very high sublimation point (3642 oC)
Does not conduct electricity.

37. Allotropes of Carbon: Graphite
Network of sp2 hybridized carbon (trigonal planar)
Half-filled p-orbitals form pi bonds
Graphite conducts electricity along the plane of the layers due to the network of delocalized p-orbital electrons/
Graphite is a good lubricant since the planes can slip over each other.

38. Summary: Types of Bonds
Intramolecular
Metallic
Intermolecular
1. Ionic bonds
2. Covalent
(Polar and Non-polar)
Metallic bonds
1. London forces
2. Dipole-dipole forces
3. Hydrogen bonds
strong bonds weak bonds
increasing bond strength

39. Types of Solids formed by Elements
Metallic Solids
Network Covalent Solids
Atomic Solids Non-polar Molecular Solids

40. Summary: Types of Solids
Examples
Intramolecular
Bonds
Intermolecular
Relative Melting Point
Atomic
He, Ar
van der Waals (London forces)
very low
Molecular
Cl2, HCl, H2O
(non-metals)
covalent bonds (polar or non-polar)
van der Waals,
dipole-dipole and hydrogen bonds
low
Metallic
Cu, Mg, Fe
(metals)
metallic bonds
moderate-high
Ionic
NaCl
NaNO3
(metal + non-metal)
ionic bonds
high
Network
quartz (SiO2)
diamond (C)
Silicon (Si)
covalent bonds
very high