Solid
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Particles are packed very closely
Cannot be compressed
Fixed in volume
Held together by a very strong force of attraction (Eg. Electrostatic forces, intermolecular forces)
2 types: crystalline solid n amorphous solid
Types of Solids
Crystalline
Amorphous
Particles
Regular repeating 3D geometric
Randomly arranged and not in
arrangement
structure
order
Shapes
Well-defined edges and faces
Irregular or curved surfaces
Melting point
Sharp melting and boiling points
Melt over a wide range of
temperature
Examples
Salts (NaCl), quartz (SiO2),
Rubber, plastics, glove
metals, diamonds
Ionic solids
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Applies to ionic cmpd
Giant ionic structure
Alternating positive (cations) and negative (anions)
Held together by strong electrostatic forces (btw cations & anions)
1. NaCl
o Na+ ion (cation)  donate 1 ion to anions
o Cl- (anion)  receive 1 extra ions, electron cloud become bigger compared to the metal cation
o Unit cell: Face centered cubic (FCC)
 Each Na+ ions surrounded by 6 Cl- ions
 Each Cl- ions surrounded by 6 Na+ ions
o Coordination number = 6 (both Na+ ions & Cl- ions)
2. MgO
o Mg2+ (cation); O2- (anion)
o Unit cell: face centered cubic (FCC)
Physical Properties :
*Anion (bigger size) : occupy the corner & face center
*Cation (smaller size) : fill the voids
o Hard
 Strong electrostatic forces
 Surface can’t be scratched easily
o Brittle
 When a strong force is applied, layers of ion will be displacing, ions of the same charge will
become adjacent. Thus, create a repulsion that cause crystal to shatter and cracks
o Solubility
 Soluble in water but not in organic solvent
 Polar water molecule attracted to ions in NaCl
 Na+ (cations) attracted to partially negative region (𝛿 − ) of water molecule (O)
 Cl- (anions) attracted to partially positive region (𝛿 + ) of water molecule (H)
 Attraction between polar water molecules and ions in NaCl pull the ions away from the crystal
lattice, hence collapsing crystal lattice in NaCl
o
o
High melting points and boiling points
 Due to the strong electrostatic force of attraction between the cations and anions
 Requires a lot of heat energy to overcome them
 MgO has higher boiling point compare to NaCl
 Factors affecting:
 Charge of the ions
- Higher charge, stronger electrostatic attractions, stronger bonds
- Eg. MgO (2+/2-) attract each other more strongly than NaCl (1+/1-)
 Size of the ions
- Smaller ions are arranged closer together, stronger electrostatic attractions
- Eg. NaCl melts at 801°C and Rubidium iodine melts at 674°C (NaCl has smaller
size compared to RbI)
Electrical conductivity
 Does not conduct electricity in solid form
 Ions are strongly held by electrostatic force and
 Ions in a fixed position in crystal lattice.
 Ions are not free to moves
 Conduct electricity in aqueous or molten form
 Ions are free to move. Electrolytes are one of the examples.
 Ions will be carrying electrical current
Simple molecular / covalent solids
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Btw the molecules : IMF (VdW forces)
Simple covalent mlcs
 In the molecules : Covalent bond
Eg. CO2, I2, P4, S8
Held together by weak intermolecular forces (IMF) btw the molecule (Van der Waals attraction / Hydrogen
bonding)
1.

Iodine (I2)
o Unit cell: Face center cubic (FCC)
o Molecules are closely packed but they are still separate molecule
o Forces: weak intermolecular forces – Van der Waals forces
2. Ice
o Oxygen atom of each water molecule forms hydrogen bonds with two hydrogen atoms of
nearby water molecule -> 3D tetrahedral arrangement
o One water molecule able to form 4 hydrogen bonds
o Extensive hydrogen bonding  very open structure  molecule further apart from each other
compared to liquid form  less dense in water
Physical properties:
o Low melting point
 Weak intermolecular forces
 Tend to be gases, liquid at room temperature
 Larger molecules increase more surface area between molecules  more IMF attractions 
more energy is needed to break down the bond  increase the melting n boiling point
 Presence of hydrogen bonding, melting n boiling point increase too
o Sparingly soluble in water
 Soluble in organic solvents
o Non-conductors of heat or electricity
 No free ions / don’t have free delocalised e when they are in water
Giant covalent solid structure
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Non-metallic atoms
Joined by extended network of covalent bonds (intramolecular forces)
No discrete molecule is present
Higher melting and boiling points than simple molecular solids
Eg. Allotropes of carbon: Diamond or graphite / quartz (silicon (IV) oxide)
[ Allotropes= carbon which can exist more than 1 form by having difference in arrangement of atoms ]
1.
Diamond
o Sp3 hybridization and 109.5° bond angle
o Coordination number= 4 (4 single covalent bonds)
o Each carbon is bonded to 4 neighboring carbon
o Extended network of covalent bond
o Physical properties
 Very high melting and boiling point
 Due to the very strong carbon-carbon covalent bonds must be broken throughout the
structure before melting occurs (Extended network of covalent bond)
 Very hard
 The atoms are held by strong covalent bond & hard to break
 Doesn’t conduct electricity
 All the electrons are held tightly between the atoms, and aren’t free to move
 Insoluble in water and organic solvents
 Attractions between solvent molecule and carbon atoms will never be strong enough to
overcome the strong covalent bonds in diamond
2. Graphite
o 2d layers of hexagonal rings
o Planar structure and sp2 hybridization and 120° bond angle
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o
o
Unhybridized p orbital – 1 free electron
Can be delocalized throughout the lattice 
account for the electrical conductivity
Coordination number = 3 (can form 3 single covalent bonds)
Physical properties
 High melting and boiling points
 Due to strong covalent bonds (between carbons) in each layer. High energy is needed to
overcome this strong bond
 Soft and slippery
 Easily scratched. Layers of graphite can slide over each when a force is applied due to
weak intermolecular force (between layers) pencil lead
 Good electrical conductors
 Due to the presence of pi electrons that are freely delocalised through the entire sheet.
/ The delocalised electrons can move along the layer (when voltage is applied)
 Insoluble in water and organic solvents
 Attractions between solvent molecule and carbon atoms will never be strong enough to
overcome the strong covalent bonds in graphite
Diamond
Transparent (able to see through)
High density
Non-electrical conductor  no delocalize electron
Graphite
Black and opaque (can’t see through)
Lower density than diamond  between the sheets has
wasted relatively large amount of space
Electrical conductor  delocalize electron in the π bond
3. Silicon dioxide
o Tetrahedral arrangement
 Each silicon atom is bonded to 4 oxygen atoms
 Each oxygen atom is bonded to 2 silicon atoms
o Has strong network of covalent bond
o Sp3 hybridization and 109.5° bond angle
o Coordination number= 4
o Physical properties
 Very high melting and boiling point
 Due to the very strong carbon-carbon covalent bonds must be broken
throughout the structure before melting occurs
 Very hard
 Due to need to break very strong covalent bonds operating in 3D
 Doesn’t conduct electricity
 All the electrons are held tightly between the atoms, and aren’t free to move
 Insoluble in water and organic solvents
 Attractions between solvent molecule and silicon atoms nor oxygen atoms will
never stronger to overcome the strong covalent bonds in silicon dioxide
Metallic solid
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Bonding : Metallic bond [ Strong electrostatic attraction btw metallic cations (+ve) & ‘sea’ of electrons (-ve) )
1. Copper
o Unit cell: Face-centered cubic (FCC)
o Coordination number = 12
o Physical properties
 Hard, high melting and boiling points
 Strong metallic bond
Size cations  bigger the size of atoms, larger distance between nucleus
and valence electrons, decrease in strength of metallic bond
No. of valence electrons  increase in no. of valence electron, increase in
strength of metallic bond
 Good electrical conductors (Conduct electricity in both solid & molten state)
 Due to delocalised sea of electron that are free to move
 Good thermal/heat conductors
 Heat energy picked up by electrons (additional kinetic energy)  move faster
 Good strength / Not brittle
 Does not shatter when force is applied (not brittle)
 The metallic bonds do not break as electrons are free to move
 High flexibility
 Malleability (deformed under compression) Eg. Flatten in a sheet
 Ductility (deformed under tension) Eg. Pulled out into wires
The presence of layers in crystal lattice that can slide over on another
Metallic bonds are non-directional : electrons can take up new positions &
reform after deformation
When a force is applied, one plane of metal ions will be slides over another.
The delocalized electrons take up new position and the metallic bond can
be maintained. (Due to non-directional)
Gas
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Particles far apart and collide between each other
Can be compressed  small particles + far apart
Mix evenly  confined into the same container [Effusion]; Manage to escape through tiny hole of container
[Diffusion]
 Exert pressure uniformly on all sides of a container
 Relatively low densities
Force
= Pressure
1 atm= 760 mmHg = 760 torr = 101325 Pa = 101325 Nm-2
Area
[STP]
𝐺𝑖𝑣𝑒𝑛 𝑢𝑛𝑖𝑡 ×
𝐷𝑒𝑠𝑖𝑟𝑒𝑑 𝑢𝑛𝑖𝑡
= 𝐷𝑒𝑠𝑖𝑟𝑒𝑑 𝑢𝑛𝑖𝑡
𝐺𝑖𝑣𝑒𝑛 𝑢𝑛𝑖𝑡
Gas Law
Boyle’s Law
1
[ 𝑃𝑉 = 𝑘]
𝑃∝
𝑉
 T (fixed)
Charles’ Law
𝑉
𝑉 ∝ 𝑇 [ = 𝑘]
𝑇
 P (fixed)
Gay-Lussac’s Law
𝑃
𝑃 ∝ 𝑇 [ = 𝑘]
𝑇
 V (fixed)
Avogadro’s Law
𝑉
𝑉 ∝ 𝑛 [ = 𝑘]
𝑛
 T & P (fixed)
[ 𝑃 ↑, 𝑉 ↓ ]
[ 𝑇 ↑, 𝑉 ↑ ]
[ 𝑇 ↑, 𝑃 ↑ ]
[ 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 ↑, 𝑉 ↑ ]
𝑃1 𝑉1 = 𝑃2 𝑉2
𝑉1 𝑉2
=
𝑇1 𝑇2
𝑃1 𝑃2
=
𝑇1 𝑇2
𝑉1 𝑉2
=
𝑛1 𝑛2
n= No. of moles
Combine Gas Law
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
𝑃1 𝑉1 𝑃2 𝑉2
=
𝑇1
𝑇2
Molar Volume
STP: 1 mol of gas is 22.4L
Room temperature : 1 mole of gas is 24L
Volume (L/dm3) = mol × molar volume
*Temperature must be in kelvin (K)
1°C = 273K
**STP = 1atm and 0°C
Charles’ Law
Gay-Lussac’s Law
Ideal Gas Law

𝑃𝑉 = 𝑛𝑅𝑇
1. Molar mass
o 𝑀=
𝑚𝑅𝑇
𝑃𝑉
2. Density
o 𝜌=
𝑀𝑃
𝑅𝑇
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P  Pressure (atm)
V  Volume (L)
n  no. of mole (mole)
R  Universal gas content (0.08206 atm.L/mol.K)
T  Temperature (K)
g  Mass (g)
M  Molar mass (g mol-1)
Dalton’s Law of partial pressure
 The total pressure of a mixture of gases in one container is the sum of the pressure of individual

𝑃𝑡𝑜𝑡𝑎𝑙 = 𝑃1 + 𝑃2 + 𝑃3 …
o 𝑃𝑇 = 𝑃𝐴 + 𝑃𝐵
𝑛 𝑅𝑇 𝑛 𝑅𝑇
= 𝐴𝑉 + 𝐵𝑉
𝑅𝑇
= (𝑛𝐴 + 𝑛𝐵 )
𝑉
𝑅𝑇
= (∑𝑛)
𝑉
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𝑛𝐴 → 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑔𝑎𝑠 𝐴
𝑛𝐵 → 𝑛𝑜. 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑔𝑎𝑠 𝐵
If total pressure is given, then use the formula
𝑛𝐴
+𝑛
𝐴
𝐵 (𝑛𝑇 )
𝑛𝐵
𝑛𝑇

𝑋𝐴 → 𝑛

𝑋𝐵 →
o 𝑃𝐴 = 𝑋𝐴 𝑃𝑇 ; 𝑃𝐵 = 𝑋𝐵 𝑃𝑇
Deviation from Ideal Gas Behaviour

Ideal gas : A gas which obeys the gas laws and gas eq (PV = nRT) strictly at all conditions of temperature
and pressure.
 Real gases : The gases that obeys the gas laws fairly well if the pressure is low OR the temperature is high.
*Deviation from ideal gas behaviour are observed particularly :
- At HIGH pressure
- At LOW temperature
𝑃𝑉
𝑉
Compressibility factor, Z [𝑍 = 𝑛𝑅𝑇 = 𝑉 𝑟𝑒𝑎𝑙 ]
𝑖𝑑𝑒𝑎𝑙
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For ideal gas, 𝑃𝑉 = 𝑛𝑅𝑇 ∴ 𝑍 = 1
For real gas, 𝑃𝑉 ≠ 𝑛𝑅𝑇 ∴ 𝑍 ≠ 1
o Therefore Z can be < 1 OR > 1 :
o When Z < 1, it is a negative deviation. It shows that the gas is more compressible than expected
from ideal behaviour. (Eg: CH4, CO2, etc)
o When Z > 1, it is a positive deviation. It shows that the gas is less compressible than expected from
ideal behaviour.
Causes of Deviation from ideal behaviour
Due to 2 wrong assumptions of Kinetic Theory of Gases :
1. When compared to the total volume of the gas, the volume occupied by gas molecules is negligible
2. The force of attraction (IMF) btw gas molecules are negligible
[The assumptions ONLY true at LOW pressure & HIGH temperature - as the distance btw the molecules is large]
Ideal Gas
 Assumptions fits Kinetic Theory of Gases
o Are in random motion in straight line
o Have perfect elastic collision. No loss in kinetic energy during Collision.
o The average kinetic energy is directly proportional to Kelvin temperature of the gas
o The volume occupied by the molecule themselves is entirely negligible relative to
the volume of the container.
o Negligible intermolecular forces between molecules.
 Obeys ideal gas laws [ 𝑃𝑉 = 𝑛𝑅𝑇 ]

False!
!
Compressibility factor, Z = 1 (if n=1) – A straight line
Gas will behave ideally when :
Low pressure  Volume of gas particles is
negligible and relatively small compared to
the volume of container. IMF negligible due
to bigger volume of the container
High temperature  molecules higher
kinetic energy and moving at high speed.
They can easily overcome the IMF attraction
between molecule
Real Gas
 Only obeys ideal gas laws where the pressure is low, and temperature is high
 Do not always the kinetic theory
o The volume occupied by the molecule themselves is entirely may not
negligible relative to the volume of the container.
 When the pressure is high, and temperature is low  molecules being incompressible 
volume of molecules is no more negligible as compared to the total volume of the gas
o The force of attraction btw gas molecules are negligible TRUE when at low pressure & high
temperature
 When the pressure is high, and temperature is low  total volume of gas is small  forces of
attraction become can’t negligible
High pressure
Low temperature
-More molecules / High density (small
-kinetic energy relatively smaller 
volume)
molecules move in low speed
-Closer from each other  Increase in IMF -IMF increase  less collision to the wall 
(no more negligible)
lower pressure is observed
 Compressibility factor, Z ≠ 1
o Positive deviation, Z > 1 [Due to volume]  Above the line of ideal gas (gas is less compressible
expected than from ideal behavior)
o Negative deviation, Z < 1 [Due to IMF]  Below the line of ideal gas (gas is more compressible than
expected from ideal behavior)
Higher temperature,
gas behaves more
ideally (less deviation)