chapter 5 states of matter 3 hours 5.1 gas 5.2 liquid …
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22
LECTURE 1
GASLearning Outcomes:(a) Explain qualitatively the basic assumptions of the kinetic
molecular theory of gases for an ideal gas.
(b) Define gas laws:
i.Boyle’s law
ii.Charles’s law
iii.Avogadro’s law
(c) Sketch and interpret the graphs of Boyle’s and Charles’s Laws.
44
● There are four parameters describing the
gaseous state:
Parameter Unit (SI)
● Quantity, n moles
● Volume, V litres
● Temperature, T kelvin
● Pressure, P pascal/ Nm-2
55
1 atm = 101325 Pa
= 101.325 kPa
= 101325 Nm-2
= 760 mmHg
= 760 torr
1 Pa = 1 Nm-2
1 mmHg = 1 torr
0°C = 273.15 K
Relationship between units:
Example:
Convert 749 mmHg to atm
Solution:
1 atm = 760 mmHg
0.986 atm = 749 mmHg
66
Kinetic Molecular Theory of Gases
4 basic assumption (postulate) :
1. The gas consists of tiny particles of negligible volume.
2. Intermolecular forces attraction do not existbetween gas particles.
3. The molecules of a gas are in continuous random motion. The gaseous particles are perfectly elastic.
4. The average kinetic energy of the gas molecules is directly proportional to the absolute temperature.
77
Ideal gas behaviour
1. The gas consists of tiny particles of negligible volume.
1. Intermolecular forces attraction do not exist between gas particles.
99
BOYLE’S LAW
( The pressure –volume relationship)
● For a fixed amount of gas at a constant temperature, gas volume is inversely proportional to gas pressure.
● As the pressure (P) increases, the volume (V) decreases.
● V= 1/P at constant T and n
● P1V1 = k1 (a constant)
P1V1 = P2V2
1111
Graph A : Pressure against Volume
Illustrate that : pressure is inversely proportional to
volume
12KMP/ChemUnit/SK017-09/10 12
Graph B : Pressure against 1/Volume
Illustrate that : P is directly proportional to 1/ V
1414
CHARLES’S LAW
( The Temperature – Volume relationship)
● The volume of a fixed amount of gas at constant pressure is
directly proportional to the absolute temperature
V ∝ T at constant P and n
V = constant
T
T in Kelvin (K) !!!
T(K) = T°C + 273.15
V1 = V2
T1 T2
1616
AVOGADRO’S LAW
( The Volume-amount relationship)
At constant pressure and temperature, the volume of a gas is
directly proportional to number of moles of the gas present.
V ∝ n ( P and T remain constant )
…………….. 4
1717
Avogadro’s “equal volumes-equal numbers” hypothesis can be stated in either of two ways:
1. At the same temperature and pressure, equal volume of different gases contain the same number of molecules (or atoms if the gas is monatomic).
1. Equal numbers of molecules of different gases compared at the same temperature and pressure occupy equal volumes.
At constant P and T, if VCO2 = VN2 = VAr ⇒ nCO2 = nN2 = nAr
At STP: T = 273.15K (0°C), P = 1.0 atm;
1 mol of gas = (6.023x1023 particles) = 22.4 dm3.
The volume of 22.4 dm3 gases at STP is called molar volume of a gas at STP
19KMP/ChemUnit/SK017-09/10 19
The Ideal EquationPV = nRT ……………...............
(6)
P = cRT
R= gas constant
= 0.08206 L atm K-1 mol-1
= 8.314 N m K-1 mol -1
2121
Dalton’s Law states that the total
pressure of a mixture of gases is
the sum of the partial pressures
of all the components in the
mixture
Definition:
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Partial pressures is the pressures of individual gas components in the mixture.
Partial pressure of gas A in the mixture is given as:
PA = XA . Ptotal
XA = mole fraction of gas A in the mixture= nA/ntotal
The total mole fraction of all gasses in the mixture is equal to 1
XA + XB + XC = 1
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In a mixture of gases A, B and C, the total pressure PT is the result of collisions of 3 types of molecules, A, B and C, with the wall of container. Thus, according to the Dalton’s law,
PTotal = PA + PB + PC
= nART + nBRT + nCRTV V V
= RT ( nA + nB + nC )
V
= nTotalRT
V
A B C
Combining
the gases
PT = PA + PB + PC
nTotal = nA + nB + nC
Volume and temperature are constant
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The ratio number of moles of gas A to the total number of moles,
XA = nA = PA ( because PA = nART/V)
ntotal Ptotal PT nTRT/V
Mole fraction of gas A PA = XA PT
Where :
XA + XB + XC + ….. = 1
2525
Learning Outcomes:
(f) Explain the ideal and non-ideal behaviors of gases
in terms of intermolecular forces and molecular
volume.
(g) Explain the conditions at which real gases approach
the ideal behavior.
2626
Real Gases- Deviation From Ideality
● Boyle’s Law and Charles Law have led to the derivation of an ideal gas equation, PV=nRT.
● Ideal gas obey the equation and fits the assumption of the Kinetic Molecular Theory.
● However, real gases showed deviation from ideal behavior.
● There are two main reason:● Volume of molecules
● Intermolecular forces
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Volume of molecules
•Ideal gas assume that gases consist of tiny molecules that does not occupy any space.
•However, for a real gas, the molecules have a certain volume.
•When a gas is compressed, the volume of the gas is decreased.
•Thus, the volume of molecules begins to occupy a sizable portion of the container.
28KMP/ChemUnit/SK017-09/10 28
Intermolecular forces
•Ideal gas assume that, there are no attractive force or repulsive force between gaseous particles.
•However,•When the volume of a container becomes smaller (by increasing the pressure)
•The distance between molecules decrease•The force of attraction between the molecules increase.
•These caused the behavior of real gases to deviate from ideal behavior.
3030
● For 1 mole of ideal gas, the value of PV/RT is equal to 1.
● The lines plotted display the deviation of real gases from the ideal behavior.
● However, all the lines converge to 1.0 when P is near zero.
● Thus, real gas behave ideally at very low pressure.
31
● When gases are compressed, the molecules are closed enough to experience the attractive force among them.(curve below 1.0, PV<RT)
● NH3 (polar molecule) shows the largest deviation because it has strongest attractive force.
● At higher pressure, the molecules are pushed too close to each other that cause the repulsive forces among them.
● This repulsion makes them less compressible, hence the line above 1.0, PV>RT.
31
3333
● Figure B shows the real gases display deviation from ideal behavior.
● However, when the temperatures increases, the line PV/RT against P for N2 approaches the dotted line for ideal gas.
● Thus, real gases behave almost ideally at high temperature.
3535
● The ideal gas law equation, PV=nRT shows the variables of gas (P, V and T) in a single equation.
● But, it does work in a real gases, especially at high temperature and low pressure.
● In a real gas:● The molecules have a finite volume.
● The molecules experience intermolecular forces
● It fails to obey the ideal gas equation and gas laws, but obeys Van der Waals equation better.
● So that, for the real gases the equation needs to be adjusted.
● The 2 parameters that need to be adjusted are P and V
3636
PRESSURE
● Attractive forces that act between molecules have an effect on the speed of the moving molecules.
● The molecules that experience this attractive will move slower.
● As a result they will give less impact to the wall of the container when collided.
● The pressure exerted by the real gases therefore is less compared to the ideal gas.
3737
● Since Preal < Pideal, the term P need to be
corrected by adding a coefficient,
to fit the ideal gas situation.
● ☞ Pideal = P real + n2a
V2
Where:
● n = number of moles
● a = positive constant
(intermolecular forces)
Depends on the strength of the attractive forces
acting between molecules.
The higher the value of a, the stronger forces among the molecules.
n2a
V2
3838
VOLUME
● Volume within which the molecules cannot move because of their own finite volume (the volume of the gas molecules).
● Therefore the effective volume available is not V (volume of a container) but a new value, V-the excluded volume.
V = Vcontainer – nb
Where,
n= moles of molecules
b= a constant representing the volume occupied by a molecule
40
Difference between Ideal Gas Equation
and non-Ideal/Real Gas Equation
Ideal Gas Non-Ideal/Real Gas
PV = nRT (P + an2) ( V – nb) = nRT
V2
Volume of gas molecules is
negligible.
Volume of gas molecules is
significant. Gas volume for
real gas is corrected using
nb to take into account
volume of gas molecules.
Attractive forces between
gas molecules is negligible.
Attractive forces between
gas molecules is significant.
The n2a/V2 term is used to
correct real gas pressure to
include attractive forces. 40
4141
✰ Real gases tend to behave ideally at:
✰ low pressures
✰ high temperatures
At low pressure.
● The molecules in a gas are far apart
● the attractive forces are negligible
● the volume of the molecules is almost zero compared
to the average intermolecular distance (the gas will
behave ideally ).
At higher temperature
● The kinetic energy of the molecules is high.
● therefore, the intermolecular forces between them
can be ignored.
4242
Learning Outcomes:
(a) Explain the properties of liquid : shape, volume,
surface tension , viscosity, compressibility and
diffusion.
LECTURE 2
LIQUID
4343
PROPERTIES OF LIQUID
● SHAPE AND VOLUME
• A liquid has a definite/fixed volume but not a definite/fixed shape.
• The particles are arranged closely but not rigidly.
• held together by a strong intermolecular forces but not
strong enough to hold the particles firmly in place
• ∴ particles able to move freely
Thus, a liquid flows to fit the shape of its container and is
confined to a certain volume.
4545
COMPRESSIBILITY
● In liquid, the particles are so closed with one
another.
● Thus, there is very little empty space.
● Therefore, liquid are much more difficult to
compress than gases (nearly incompressible) and
also much denser under normal condition.
4646
● Diffusion
● Diffusion rate of a liquid is much slower than gas
but faster than solids.
● Due to :
● molecules are closely packed compared to gases
● lower kinetic energy than gases
● stronger intermolecular attractive forces between the
molecules compared to gases
47KMP/ChemUnit/SK017-09/10 47
● SURFACE TENSION, γ
● The surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area.
● Molecules within a liquid are pulled in all directions by intermolecular forces.
● However, molecules at the surface are pulled downward and sideways by neighbouring molecules, but not upward away from the surface.
4848
• These intermolecular attractions thus tend to pull the
molecules into the liquid
• Caused the surface to stretch and tighten.
• The surface tension decreases with increased
temperature.
• The intensity of molecular motions increase,
• The intermolecular forces become less
effective.
4949
VISCOSITY
● Viscosity is the internal resistance of liquid to flow.
● It opposite with fluidity which is the ability of the liquid to flow
● It will reduce the flow of liquid.
● Increases with strength of intermolecular forces but decreases with temperature.
5050
The factors affecting the viscosity are:
● The size of the molecules
➢Molar mass (Mr) increase, resistance increase, more
viscous the liquid.
● The intermolecular forces acting between molecules.
➢The stronger the attractive forces, the viscosity
increase.
● The temperature of the liquid.
➢Viscosity decreases with increased temperature.
5151
Learning Outcomes:
(b) Explain vaporisation and condensation
processes based on kinetic molecular theory and
intermolecular forces.
(c) Define vapour pressure and boiling point.
(d) Explain boiling process.
5252
Vaporisation
● The process of changing the state from liquid to vapour.
● Molecules in liquid moves quite freely.
● Some molecules have higher kinetic energy (move faster) than some other molecules (move slower).
● Molecules at the surface that posses enough kinetic energy to overcome the attractive intermolecular forces can escape as vapour in the gas phase.
5353
● The rate of vaporisation increases with:
● A rise in temperature.
● An increase in the surface area of the liquid.
● A decrease in the intermolecular forces of
attraction in the liquid.
● A decrease in external pressure.
5454
Condensation
● The process of changing from vapour to liquid.
● Occurs when the vapour molecules are cooled ( lose kinetic energy) or when pressure is exerted ( molecules are closer together).
● If a gas is cooled, ➢ The molecules are moving slowly and closer together.
➢ They were able to attract each other and with molecules in the liquid state. (form back into liquid)
5555
Molecules
having high
kinetic energy
can overcome the
attractive forces
and vaporisation
A vapour
pressure
exerted by
molecules on
the surface of
the liquid
A molecule may lose
its energy during
collision and get
trapped among the
liquid molecules
leading to
condensation
Vaporisation and condensation
process
5656
VAPOUR PRESSURE
● Vapour pressure is the pressure exerted by vapour molecules while in the state of dynamic equilibrium with its liquid.
● Dynamic equilibrium: rate of forward reaction is exactly equal to the rate of the reverse reaction ( rate of vaporisation and condensation are equal)
5757
Vapour Pressure
The vapour
pressure is
constant.
Known as the
equilibrium
vapour
pressure
Dynamic Equilibrium
The rate of
vaporisation and the
rate of condensation
are equal
Number of molecules
leaving the liquid
surface is the same as
the number the of
vapour molecules
entering the liquid
state
5858
BOILING POINT
● Boiling point of a liquid is defined as the temperature
at which it’s vapour pressure equals the external
pressure.
● The normal boiling point is the temperature at which
the vapour pressure of the liquid is equal to 1 atm.
5959
SOLIDLearning outcomes:(a) Explain the fixed-shape of a solid.
(b) Apply the kinetic concepts to explain the process of :
● Freezing (solidification)
● Melting (fusion)
● Sublimation
● Deposition
(c) Differentiate between amorphous and crystalline solids.
(d) State the following types of crystalline solids with appropriate
examples:
● Metallic
● Ionic
● Molecular covalent
● Giant covalent
6060
Properties of Solid
● Particles are arranged closely together;
● Only vibrate and rotate in fixed position.
● Rigid arrangement.
● Cannot move freely.
● Has definite shape and volume.
● Has strong forces between the particles.
● Has high densities.
● Non-compressibility.
● Extremely slow diffusion.
● Less energy compared to liquids and gasses.
6262
In principle, solid, liquid and gas states are
interconvertible
solid gas
liquid
sublimation
deposition
63KMP/ChemUnit/SK017-09/10 63
Melting process
● The melting point is the temperature at which a solid
changes into a liquid.
X(s) X(l)
● When heated:
● The kinetic energy of particles increases.
● The extent/limit of rotation and vibration increases.
● A point is reached when the kinetic energy is high enough
to overcome the intermolecular forces between them.
64KMP/ChemUnit/SK017-09/10 64
Freezing (solidification) process
● A liquid is changing into a solid
X(l) X(s)
● When the temperature is lowered, the kinetic energy of a liquid particles decreases.
● The vibration decreases (kinetic energy decrease).
● A point is reached when the intermolecular forces are strong enough to hold the particles together in a fixed and orderly arrangement.
● The liquid freezes.
● The freezing point of a liquid is constant, and the value is the same as the melting point of its solid.
6565
Sublimation Process
● The process by which a substances goes directly from
solid to the gaseous state without passing through the
liquid state.
X(s) X(g)
● Examples: iodine and naphthalene
6666
Deposition
● The process where the particles in gas phase are
transformed into solid phase directly during a cooling
process.
X(g) X(s)
6767
2 types of solid
Crystalline solid☺In crystalline solid the particles (atoms, molecules or ions)
are arranged in an orderly manner.
☺It is formed when a saturated liquid is cooled slowly.
☺Its atoms, molecules or ions occupy specific position
☺Eg : salt, sugar, pyrites
Crystalline
Armophous
6868
☺Amorphous solid
☺They are called amorphous which means without
shape or form
☺Non-crystalline solid
☺Particles are randomly arranged and have no ordered
structure.
☺It is formed when a saturated liquid is cooled rapidly.
☺Eg: glass, plastics, gels
69
Differences between amorphous
and crystalline solids
Amorphous Crystalline
Non crystalline solid Crystalline solid
Random structure Ordered structure
Formed when cooled rapidly Formed when cooled slowly
No definite melting point Well defined melting point
Shatter irregularly Shatter to crystalline shaped
pieces69
7171
Ionic Crystals
● The particles are made up of cations and anions
respectively.
● There are electrostatic forces of attraction between
cations and anions.
● Example: ionic crystal of NaCl:
7272
Metallic Crystal
● The particles in metallic crystals are metal atoms of the same metal. Examples, sodium, magnesium and aluminium.
● They are bound together by metallic bond,
● (recall* metallic bond= the electrostatic attraction between the lattice of positive ions and this sort of ‘sea’ of electrons).
7373
Molecular Covalent Crystals
● The particles consists of simple molecules which are held together by weak van der Waals forces.
● Generally, this type of solid has very low boiling and melting point.
● Examples: iodine, I2, phosphorus,P4 and sulphur,S8
7474
Giant Covalent Crystals● Very large molecules.
● Consists of strong covalent bonds binding its particles.
● The covalent bond gives gigantic structure
● The particles are non-metal.
● Examples: diamond, graphite and fullerene.
● Generally, the giant crystal structure are hard and have high melting and boiling points.
78
LECTURE 3:
PHASE DIAGRAMLearning Outcomes:
(a) Define phase, triple point and critical point
(b) Identify triple and critical point on the phase diagram
(c) Sketch the phase diagram of H2O and CO2
(d) Compare the phase diagram of H2O with CO2 and
explain the anomalous behavior of H2O
79
Phase, Triple point and Critical point
Phase
The phases of a system are parts of it which are separated by
a distinct boundary, such as solid, liquid and gas.
Triple point
The point at which the three lines representing the
solid/liquid, liquid/vapour and solid/vapour equilibriums
meet.
Critical point
The point above which it is not possible to liquefy as gas,
however great the pressure
80
Example
The number of phases in a system is denoted P.
● A gas or gaseous mixture, is a single phase (P = 1 )
● A crystal is a single phase ( P = 1 ).
● Two totally miscible liquids form a single phase
(P = 1)
● An alloy of two metals is a two-phase system ( P = 2 ) if the metals are immiscible, but
a single-phase system ( P = 1 ) if they are miscible.
81
A phase diagram
Showing the relationship between the different
phases of a given substance under varying
conditions of temperature & pressure
Phase Diagram
82
Phase diagram of water
B
C
A
Solid
Liquid
Vapour
1
Critical
Point
Boiling
PointT
Triple Point
273.16 373.15 374
Temperature / K
218
6.0 x 10-3
Pressure/atm
ATB :Solid phaseBTC : Liquid phaseATC :Vapour phase
83
Phase Diagram for Water
B
C
A
SolidLiquid
Vapour
1
Critical
Point
Boiling
PointT
Triple Point
273.16 373.15 374Temperature / K
218
6.0 x 10-3
T-C : Represents the variation of boiling temp. with pressureT-B : Represents the variation of melting temp. with pressureT-A : Represents the variation of sublimation temp. with
pressure
84
B
C
A
SolidLiquid
Vapour
1
Critical
Point
Boiling
PointT
Triple Point
273.15 373.15 374Temperature / K
218
6.0 x 10-3
At temperatures between 273.15 K and 373.15 K under atmospheric pressure, the stable phase is liquid water
Phase Diagram for Water
85
T is an unique point at which the three lines representing the solid/liquid, liquid/vapour and solid/vapour equilibriums meet is called the triple point
The triple point for water is 273.16 K and 6.0 x 10-3 atm
At the other extreme of the liquid/vapour line is the critical point, above which it is not possible to liquefy as gas, however great the pressure
Phase Diagram for Water
86
Phase Diagram for Carbon Dioxide
B
C
A
72.9
Solid Liquid
Vapour
5.2
1
Critical
Point
T
Triple
Point
195 216.4 304 Temperature / K
T-C : Represents the variation of boiling temp. with pressureT-B : Represents the variation of melting temp. with pressureT-A : Represents the variation of sublimation temp. with pressure
ATB : Solid phase
BTC : Liquid phase
ATC : Vapour phase
87
Phase Diagram for Carbon Dioxide
-The phase diagram of carbon dioxide is more typical,
showing a rightward sloping melting temperature line
-Solid carbon dioxide is denser than liquid carbon dioxide
Triple point for carbon dioxide is 216.4 K and 5.2 atm
The triple point is above atmospheric pressure, so that at
atmospheric pressure carbon dioxide sublimes, so ‘dry ice’,
which is solid carbon dioxide, changes directly from solid
to gas
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