Saturday 26 January 2013

Simple Kinetic Molecular Model of Matter

IGCSE Cambridge Syllabus Area 2.1

.1       States of Matter – Solids, Liquids and Gases

State the distinguishing properties of solids,
liquids and gases (in terms of simple kinetic molecular model of matter)


Matter can be in 3 possible states: solids, liquids or gaseous states.

In solids: the constituent particles are closely packed in regular / lattice arrangement and vibrating in fixed position - the higher the temperature, the more vigorous the vibration and vice versa. Solids have fixed volumes and fixed shapes

In liquids: the constituent particles are no longer in  regular arrangement - they are constantly translating past each other. Liquids have fixed volume and take the same of their containers.

In gases: the constituent particles are far apart and they move randomly in straight lines in all direction, colliding with each other and with the walls of their containers in the process. The higher the temperature, the higher the kinetic energy of the particles.

.2 Molecular Model (of Matter)

1.    Matter is made up of tiny particles which can be: mono-atomic (as in noble gases like helium, He), molecular (as in water, H2O; oxygen, O2; nitrogen, N2;  etc) or (as in common salt, sodium chloride, NaCl). The atoms that make up the particles of each substance are unique in terms of: the type of atoms and the ratio they bear to each other - eg. the particles of water are made up of hydrogen and oxygen atoms in the ration of 2:1 whatever the states (solid, liquid or gaseous) that water is in.

2.    The particles of matter are too tiny to be seen with our naked eyes - however, they can only be seen with the aid of instruments such as electron microscope. Nevertheless, their random nature of motion can be indirectly seen with normal microscope: When large smoke particles are in suspension in air in a glass cell or smoke chamber or larger pollen grains are in suspension in liquid water and seen through normal microscope, these larger particles can be seen as moving around haphazardly - in what is known as Brownian motion - due to random bombardments by the tiny but fast-moving energetic particles of the air or liquid water, whichever is applicable. 

3.    Depending on its temperature, a substance can exist either in solid, liquid and/or gaseous states. In solid state, its temperature is lower than that in liquid state; and in gaseous state, its temperature is even higher.

4.    The melting or freezing point of a substance refers to the temperature (under normal atmospheric pressure) below which the substance exists as a solid; and at which, the substance changes state from solid to liquid or vice versa. This temperature is a unique physical property of the substance. For pure water, it is 0o C or 273o K and is also referred to as the ice point.

5.    The boiling or condensation point of a substance refers to the temperature below which the substance exists as a liquid or solid; and at which the substances changes state from liquid to gaseous state or vice versa. Again, this temperature is unique to each substance. For pure water, this temperature is 100o C or 373o K and is also known as the steam point.

6.    The distinguishing properties of a substance in solid, liquid or gaseous states may be described in terms of the following characteristics of its constituent particles:
i)      distance between them – closely-packed, further apart or furthest apart;
ii)    their arrangement - in fixed position, in loose-attachment or without any attachment;
iii)   their movement – to and fro vibrations about fixed position; moving and changing position while loosely attached to one another; or dashing randomly and independently from one another;
iv)  forces of attraction and repulsion between them – strongest to hold the particles in fixed positions, present to hold them loosely together or negligible.

7.    Hence, in terms of kinetic molecular model of matter, a substance in solid, liquid or gaseous state may be described as follows (syllabus area 2.1 (b)):
a.       In Solid State:
Its constituent particles are closely-packed, vibrating to and fro in fixed positions experiencing the strongest forces of attraction and repulsion between them much like having springs holding them together in fixed shape and volume.
Temperature and Expansion: When the temperature of the solid rises (without exceeding its melting point), the vibrations become more vigorous resulting in the distance between neighbouring particles getting bigger and hence the expansion of the solid.
b.      In Liquid State:
The constituent particles are further apart and in loose attachments, vibrating vigourously and changing positions and slipping past each other easily – resulting liquid having fixed volume with shape that varies with the shape of its container.
Convection and Expansion: When the temperature of the liquid rises (without exceeding its boiling point), the vibrations and movement of the particles become even more vigorous resulting in the distance between neighbouring particles getting bigger - hence, the hotter part of the liquid becomes less dense and the particles rise as in convection; and, hotter liquid expands.
Evaporation (syllabus area 2.1 (c)):
Energetic liquid particles at the surface of the liquid, though still below its boiling point, may have enough kinetic energy to escape from the main body of the liquid in what is known as evaporation.
Evaporation results in cooling because the evaporated particles leave with higher energy level leaving behind particles with lower thermal energy.
Factors influencing rate of evaporation: The higher the room temperature, the bigger the surface area for evaporation, the higher the movement of air (draught) and the lower the humidity of the surrounding, the higher the rate of evaporation.
Evaporation, unlike boiling that only occurs at boiling point, can occur at any temperature below boiling point as long as the surface particles have enough energy to break away from the main body of the liquid.
c.       In Gaseous State:
The constituent particles are furthest apart (least dense of all), no longer experiencing any significant force of attraction and repulsion between them and dashing around at high speed (about 500 m/s for air molecules at 0o C), randomly and independently of one another in all directions. The higher the temperature of gas particles, the faster and more vigourous the random motion of the particles and the higher the force and rate of bombardment on the inner walls of its container and hence, the higher the gas pressure exerted. 
Brownian Motion (in gas):
When randomly-moving gas particles collide with other more massive particles such as smoke particles, they have enough energy to exert a net force on the more massive particles in what is seen as Brownian motion – the random bombardments of massive particles by high-energy tiny particles.
Pressure Changes in Gas (syllabus area 2.1 (d)):
When gas particles collide with the internal walls of its container, the particles exert a force and therefore a pressure on the inner walls of its container.
At constant temperature, the pressure exerted by a gas is directly proportional to its density. The higher its density ρ, the higher the pressure exerted P due to the higher rate of collisions of the particles on the inner walls of its container:
Thus, P α ρ    P α mass (m)/volume (V) ≡ P α m/V
·        when mass m increases (such as by pumping in more air into a tyre) with temperature and volume remaining constant, pressure P increases because density ρ of the gas has increased.
Thus, P1/m1 = P2/m2  = constant

·        when volume V decreases (such as by squeezing a balloon) with temperature and mass remaining constant, pressure P also increases because lower volume means greater density ρ.
Thus, P1V1 = P2V2 = constant (Boyle’s Law)
At constant volume, the pressure exerted by a gas is directly proportional to its temperature. The higher the temperature, the higher the pressure exerted because at higher temperature, the gas particles hit the inner walls of its container at higher frequency and with greater force.
Thus, P α T P1/T1 = P2/T2 = constant (Pressure Law)

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