Transcript Properties of Solids

PROPERTIES OF
SOLIDS
SCH4U1 - DVORSKY
Intra vs. Intermolecular Bonds
• The properties of a substance are influenced by the force
of attraction within and between the molecules.
Intra vs. Intermolecular Bonds
Intramolecular Bonds: Bonds within a molecule (covalent
or polar covalent)
Intermolecular Bonds: Bonds between molecules.
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) 
MP = 0oC
H2O (g)  2H + O
BP =1000C
intermolecular bonds breaking
Decomposes: >2000oC
intramolecular bonds breaking
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
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.
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:
London Forces and Electrons
• As the number of electrons in an atomic solid increases,
the mp/bp also increases.
Group 18
Electrons
Boiling Point ('C)
He
2
-268.6
Ne
10
-245.9
Ar
18
-185.7
Kr
36
-152.3
Xe
54
-107.1
Rn
86
-61.8
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.
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!
2) Non-Polar Molecular Compounds
• Compounds without bond dipoles only have London
Forces between molecules.
• This results in LOW bp/mp
Group 17
F2
Cl2
Br2
I2
Boiling Point
('C)
-188.1
-34.6
58.8
184.4
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
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.
Boiling Points of Hydrocarbons
Molecular Boiling Point (oC)
Formula
State at STP
CH4
-161.5
gas
C2H6
-88.6
gas
C3H8
-42.1
gas
C4H10
-0.5
gas
C5H12
36.1
liquid
C6H14
68.7
liquid
liquid
C10H22
174.1
liquid
liquid
C22H46
327
solid
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)
Boiling Points of Some Polar and Nonpolar Substances
Substance
Boiling Point
(oC)
Molar Mass
(g/mol)
Number of
Electrons
HCl
-84.9
36
18
-60.7
34
18
-188.1
38
18
-185.7
40
18
H2S
F2
Ar
polar
(molecular dipole)
nonpolar
(NO molecular dipole)
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”.
Water “bends” near a charged object.
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:
• Propane (C3H8)
• Propanol (C3H7OH)
• Glycerol (C3H6(OH)3)
mp (oC)
bp (oC)
-188
-126
-42
97
18
290
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.
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.
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.
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.
The Electron Sea Model of a Metallic Crystal
Positive Metal Ion
Delocalized
Electron “Cloud”
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.
• 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.
• 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) + 77 kJ
K (g)
Sc (l) + 333 kJ
Sc (g)
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.
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.
Crystal Packing
• Properties of ionic compounds are related to the packing
of the crystals:
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.
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 H2O
@ 0oC)
Melting Point
(oC)
Solubility
(g/100g H2O
@ 0oC)
CsCl
646
161
MgO 2800
0.0006
NaCl
800
35.7
NaCl
35.7
800
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)
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
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
Polar covalent
Allotropes of Carbon: Diamond
• Covalent network of sp3 hybridized carbon (tetrahedral).
• Very hard; very high sublimation point (3642 oC)
• Does not conduct electricity.
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.
Summary: Types of Bonds
Intramolecular
1. Ionic bonds
2. Covalent
(Polar and Non-polar)
Metallic
Intermolecular
Metallic bonds
1. London forces
2. Dipole-dipole forces
3. Hydrogen bonds
strong bonds
weak bonds
increasing bond strength
Types of Solids formed by Elements
Atomic Solids
Non-polar Molecular Solids
Metallic Solids
Network Covalent Solids
H2
He
Li
Be
B
C
N2
O2
F2
Ne
Na
Mg
Al
Si
P8
S4
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br2
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I2
Xe
Cs
Ba
LaLu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At2
Rn
Fr
Ra
AcLw
Summary: Types of Solids
Type
Examples
Intramolecular
Bonds
Atomic
He, Ar
Molecular
Cl2, HCl, H2O covalent bonds
(non-metals) (polar or nonpolar)
Metallic
Cu, Mg, Fe
(metals)
Ionic
Network
NaCl
NaNO3
(metal + nonmetal)
quartz (SiO2)
diamond (C)
Silicon (Si)
(non-metals)
Intermolecular
Bonds
van der Waals (London forces)
van der Waals,
dipole-dipole and
hydrogen bonds
Relative Melting
Point
very low
low
moderate-high
metallic bonds
high
ionic bonds
very high
covalent bonds