Classwork Series and Exercises {Basic Science – JSS2}: Kinetic Energy

Basic Tech JSS 2 Week 1

Topic: Kinetic Energy

Introduction

Matter contains very tiny particles such as atoms, molecules or ions. These particles are always in continuous random motion. The energy which these particles possess that enable them move about is called kinetic energy. The word kinetic is derived from the Greek word “kineo” which means, “I move”. Therefore, kinetic energy is the energy possessed by any body or object as it moves from one place to another.

Statement of the Kinetic Theory

The kinetic theory of gases describes the physical behaviour of a gas in “ideal” or perfect conditions. The statements used to describe the behaviour are called “assumptions” or “postulates”.

They are as follows:

1. All gases are made up very tiny particles, which are in constant motion and possess kinetic energy.
2. These particles (atoms, molecules or ions) move in a straight line and in every direction
3. The molecules collide with each other and with the walls of the vessels in which they are contained
4. The collisions made by the gas molecules are said to be perfectly elastic. This means that the collisions do not result in any change in the kinetic energy of the gas.
5. A distance that is large compared with their sizes separates the gas molecules from each other. This means that actual volume occupied by the gas molecules themselves is negligible when compared with the volume of the vessel containing the gas.
6. An increase in temperature tends to increase the motion of the gas molecules. This implies that the average kinetic energy of the gas molecules is directly proportional to the absolute temperature of the gas
7. The increase in temperature leads to increase in volume of the gas

Explanation of Some Phenomena Using Kinetic Theory

Matter exists in three states, namely; solid, liquid and gas. These states are made up of atoms, molecules or ions. One clear difference between these states of matter is the extent (degree) of movement of their particles when heat is applied to them.

Solid State

In solids, the particles are very closely or tightly packed and held firmly by attractive forces. These forces are called forces of cohesion. Examples of cohesive forces are electrovalent, covalent, metallic and even weak ones like Vander Waal forces. It is these forces that hold the particles of a solid together in a fixed position and so restrict their movement from one point to another.

The particles however vibrate about their fixed position. The fixed position of particles of a solid is the reason solids have definite shapes and fixed volumes and are difficult to compress.

Liquids State

Ina liquid, the particles are not so closely packed as in a solid, but they are not confined to a fixed position. They have sufficient kinetic energy to move (vibrate, rotate, translate and even slide) about randomly. Hence, the particles have some freedom of movement although they are still held together by cohesive forces (covalent, Vander  Waals Forces) into a fixed volume. So, when a liquid is poured into a vessel, it immediately takes the shape of the container.

Gaseous State

In a gas, the particles are very widely spread and so are easily compressed. They have much more kinetic energy than those of a liquid. The cohesive forces holding the particles are so weak and so small (negligible) that the particles are free to move about in all directions at great speed. This is the reason a gas occupies the whole volume of its container.

A good aroma (odour) of your mother’s food, which reaches you in the sitting room, clearly demonstrates how easily gas particles mix with air and move about and far too. The same is true when a classmate of yours in one corner of your classroom passes out a foul gas (i.e. carbon (IV) oxide), which quickly spreads to other parts of the classroom.

Summary of properties of states of matter

Property	Solid	Liquid	Gas
Motion	There is no movement of particles	There is a slight movement of particles	There is very rapid movement of particles
Volume	It has a fixed volume	It has a fixed volume	It occupies the whole volume of the vessel
Shape	It has a fixed shape	Takes the shape of container	It takes the shape of vessel and fills it completely
Compressibility	It  incompressible	It is incompressible	It is highly compressible

Using Molecular Structure to Explain Kinetic Theory

Kinetic theory can be used to explain certain phenomena in nature to show the differences in them. Such phenomena are evaporation, boiling and vapour pressure.

Evaporation

Evaporation is a process whereby molecules of a liquid with higher kinetic energy escape through the liquid surface into the space above the liquid. Such molecules are said to vaporize.

In any given liquid sample, some particles possess more kinetic energy than others. So, when such energetic particles come near the liquid surface, they break away from the attractive forces of other nearby liquid particles and escape into the space above. When this happens, we say the liquid is evaporating.

Boiling

Boiling is said to take place when a liquid freely changes into vapour when it is heated. The temperature at which this happens is called the boiling point.

For a liquid to boil, it must be heated. When this happens, the rate of evaporation increases. Boiling takes place because liquid molecules acquire more kinetic energy when the liquid is heated. The molecules then collide with each other and with the walls of the vessels to build up pressure in the liquid. At this time, a saturated vapour pressure for the liquid results.

This pressure builds up and increases until a temperature is reached at which the saturated vapour pressure of the liquid is equal to the atmospheric pressure at that time. At this stage, bubbles of vapour form in the liquid and rise to the surface. The liquid is now said to be boiling.

Distinction between Evaporation and Boiling

S/N	Evaporation	Boiling
	Lowers kinetic energy of molecules	Increases kinetic energy of molecules
	Does not need heating to place	It needs heating before it could take place
	Can take place at any temperature	It takes place at finite temperature

Vapour Pressure

This is the pressure that is built up when escaping molecules of a liquid collide with each other and the walls of the containing vessel. It is therefore formed by evaporation in a closed system. At this time, some vapour molecules hit the liquid surface and re-enter the liquid. This is called condensation. Therefore, two forces or processes (i.e. evaporation and condensation) are in operation here.

As the above processes continue, there comes a time when the number of molecules condensing into liquid is equal to the number of molecules evaporating from the liquid. At this point, equilibrium is set up between evaporation and condensation making the vapour pressure to remain constant. This vapour pressure is called the saturated vapour pressure of the liquid at that temperature.

Factors That Affect Evaporation

The following factors affect evaporation of liquids, namely:

1. Wind speed: How slow or fast the wind is at a particular time
2. Humidity: The amount of water vapour in the atmosphere
3. Temperature: How hot or cold the liquid and the atmosphere are
4. Vapour pressure of liquid: The degree of saturation of liquid vapour
5. Viscosity of liquid: How thick or thin a liquid phase is