MATTER AND ITS PROPERTIES III
This is a basic article primarily meant for grade 9 and 10 level students of various educational boards like CBSE, ICSE, IGCSE, and other state educational boards. It can be used as introductory material for higher levels
THE CHANGE OF STATE
Prior Knowledge:
Key Terms or concepts discussed in the article
Matter exists in three clearly defined states, viz, solid, liquid and gas. All of the three states differ from each other with regard to the arrangement of particles at the micro-level and some bulk-scale properties like shape, volume, compressibility and their ability to flow or not. The knowledge of the kinetic theory of matter helps us understand these differences in a clearer way.
The above three forms of matter can inter-change their state from one to another depending on the conditions of temperature and pressure. How does this happen? Can such changes of state be understood in terms of what happens at the particle (or micro) level? For this, we need to have a basic idea about the energy possessed by particles in various forms of matter and their comparisons.
ENERGY: A concept that cuts across disciplines within science
Energy is the capacity to do work. This is a cross-cutting concept in science, that is, its definition and application are central to all science. Whenever some work is done against a force, that is, an object is moved a certain distance by application of force, energy is used. In fact, the total work that can be done is equal to energy. This energy must come from some other source, it can neither be created nor be destroyed. This forms the basis of another cross-cutting concept in science, which is the ‘law of conservation of energy’. For example, you are using some of your energy to read this material and this came from the food you ate. The energy in the food ultimately came from the sun (through the process of photosynthesis).
Note: As a student of science you should develop the ability to identify the cross-cutting concepts because they essentially remain the same whether you are studying chemistry, biology, or physics. However, the context or scale at which they are applied may differ.
The total energy an object possesses can simply be taken as the sum of its potential energy and its kinetic energy. Potential energy is the energy possessed by an object due to its position as compared to other objects. For example, an object raised above the ground level (see (a) of the figure below), or a system in which an object lies at the end of an unstretched and stretched and spring (see (b)), or two charges either removed apart or brought closer to each other (see (c) and (d)). In all these examples, one of the situations has energy stored and that can be converted to work or any other form of energy. Simply, the lower energy state of a system will not be able to do work (i.e. will tend to remain as it is in the absence of an external force) while the situation in which a system is having some stored energy due to its position (i.e. potential energy) is said to have higher potential energy which tends to do some work. In this event, the system returns to a position of lower potential energy. This fact brings us to another important concept which is that a system with higher energy tends to change (i.e. it is unstable) while a system with lower energy tends to remain as it is (i.e. it is stable). In the context of chemistry, the figures (c) and (d) consisting of situations in which the potential energy of a system varies with the position of particles is encountered more. That’s why potential energy is additionally defined as the energy possessed by a chemical system due to its composition too.
Apart from potential energy, another major kind of energy is the kinetic energy which is energy due to motion and takes into account both mass and velocity of a particle. The average kinetic energy of all particles in an object gives the measure of temperature. The increase or decrease in the kinetic energy of particles of any form of matter results in an increase or decrease in its temperature respectively. But, to say that if heat is supplied to an object its temperature will necessarily rise is not correct. It may raise the average kinetic energy of particles comprising it (in which case its temperature will rise) but it may also raise potential energy of particles comprising it, and in this case, its temperature will not rise.
It is easier to understand the above concept in terms of pure substances (i.e. those which are made up of the same kind of particles). For example, ice, water, and steam differ in terms of potential energy.
In the same substance, the potential energy of solid-state is always least and it is the most in the gaseous state.
The reason for the above fact is that freer the particles, the greater the distances to which they can potentially move and do greater work when they strike each other. Hence the order of potential energies in the three states of matter is shown in the figure below.
What are the names given to different processes that result in a change of state in the matter?
Solids absorb heat when they change to the liquid state of higher potential energy and generally higher total energy (potential + kinetic energy). As the solid is heated, its temperature begins to rise until it reaches a point where the rise in temperature stops but the heat is continued to be absorbed. It is in this phase that change of state happens, i.e., a solid start changing into its liquid state and until this change of state does not complete, the temperature does not rise. The heat absorbed in this phase is called ‘latent’ (meaning hidden) heat. The fixed temperature at which a solid changes into its liquid state at atmospheric pressure is called its melting point.
But where is the heat energy absorbed during the change of state going?
As the temperature stops rising once the solid reaches its melting point, it means that kinetic energy has stopped increasing and it is to increase the potential energy that all the heat absorbed during this ‘change of state’ period is going. Until all the solid turns to its liquid state, the temperature does not rise and for example, in case of water whose melting point is 00C, we have ice at 00C changing to water at 00C . Which has more energy and why?
Water at 00C has greater total energy even though the average kinetic energy of both is the same (because the temperature is the same) because its potential energy is greater than that of ice. On similar lines, the total energy of steam at 1000C is greater than that of water at 1000C.
The pure substances, which are made up of identical particles at micro-scale, have fixed temperatures at which they melt (or boil or condense or freeze), provided the atmospheric pressure remains the same. So, when it is stated that ice melts at 00C , it actually means that the melting is happening at normal or standard atmospheric pressure (1 atmosphere or 76 cm of Hg or 101,325 pascals or 14.696 psi) although this fact is sometimes not written with it. If the atmospheric pressure changes, the new values for melting, boiling, condensation, and freezing points emerge which are characteristic of that particular pressure. However, in scientific writings, the pure substances generally have the temperatures at which they melt, boil, or freeze written at standard or normal atmospheric pressure.
Why a change in surrounding or atmospheric pressure changes melting or boiling points?
Simply, melting and boiling points generally increase with an increase in surrounding pressure
For the melting process, most solids expand on becoming liquids and an increase in pressure tends to oppose this expansion. So, a solid has to be taken to a higher temperature to melt and that is why melting point increases with an increase in pressure. Similarly, vapors occupy more volume and an increase in pressure opposes an increase in volume. That is why the boiling point also increases with an increase in atmospheric pressure.
There are, however, exceptions to the above general trends. One of the major exceptions is the melting process in the case of ice, which is less dense than water. So, when ice changes to water, it actually contracts. An increase in pressure, therefore, decreases the melting point in case of water
Just like melting and boiling point, the heat absorbed by pure substances for every 1-degree rise in temperature and latent heat absorbed during the ‘change of state’ processes are also characteristic of substances. To be compared to each other, they must be for fixed masses (conventionally 1g or 1kg). The associated terms and their definitions are presented below:
Specific Heat Capacity (c): The heat absorbed by a unit mass of substance for 1-degree rise in temperature
If ‘Q’ is the heat absorbed by a mass ‘m’ for ‘t’ degree rise in temperature, then by definition:
Note that heat absorbed is divided by mass and rise in temperature because by definition heat absorbed by a unit mass for a unit degree rise in temperature is needed.
This is an application of the ‘unitary method’ of basic mathematics
Can you write units of specific heat capacity provided the above formula? What if you know that units of heat (Q) are joules and that of the mass (m) and the temperature (t) are kg and degree Celcius respectively. If these units are put in the formula above, the units of specific heat capacity (c) come out to be J/kg.0C or Jkg-1.0C-1.
Another set of important terms associated with ‘change of state’ are definitions of ‘latent heat’. An example is latent heat of vaporization. It is defined as:
The latent heat of vaporization is defined as the heat absorbed by liquid at its boiling point to change into its vapor state without a rise in temperature at standard atmospheric pressure.
It can be observed from the above definition that it is important to mention that the temperature does not change when the latent heat is absorbed and the pressure at which change of state takes place. Further, the above definition can be changed to another ‘latent heat’ definition by changing some of the words marked in red in the above definition. This is illustrated below:
The latent heat of fusion is defined as the heat absorbed by solid at its melting point to change into its liquid state without a rise in temperature at standard atmospheric pressure.
The latent heat of condensation is defined as the heat given out by gas at its condensation point to change into its liquid state without a rise in temperature at standard atmospheric pressure.
On similar lines try writing the other definitions of latent heats for different processes of the change of state.
Two other processes that require special mention in the change of state processes are sublimation and evaporation. Their main points are discussed in the sections that follow.
Sublimation
Normally, the majority of substances pass through the liquid state on their way to become a gas from their solid-state on being heated. But in some substances, the kinetic energy of the particles suddenly increases so much that they turn from solid to gaseous state straightaway skipping the liquid state altogether.
When a solid changes directly to its gaseous state on being heated, the process is called sublimation and the solid which undergoes this change is called a sublimate.
The reason for sublimation is the nature of the particles comprising the solids that undergo this process. Due to their nature, the inter-particle forces decrease suddenly.
Solid carbon dioxide, camphor, iodine, and naphthalene are some of the common examples of solids that undergo sublimation at different temperatures under standard atmospheric pressure. Note that, again, mention of pressure is important. From the above sublimates, solid carbon dioxide and naphthalene sublime at room temperature (250C) under standard atmospheric pressure while the others have to be heated.
The reverse of sublimation, that is, direct change from gaseous to solid state is called deposition.
Evaporation
When a liquid changes to its vapors at the temperature of its surroundings rather than upon heating, the change is called evaporation. As opposed to the process of boiling (which happens at a fixed temperature called a boiling point), evaporation is a continuous process in any liquid and keeps on happening irrespective of the temperature. The rate of evaporation, however, differs depending upon the conditions. Boiling happens throughout the liquid (bulk phenomenon) while evaporation is a surface phenomenon, that is, only the liquid particles on the surface will change to vapors. There is always a chance of a vapor returning back to the liquid state (condense) due to inter-particle attractions, especially, when vapors in the medium surrounding the surface of the liquid are in high concentration (the number of vapors per unit volume) or temperature is lowered or pressure is increased.
Can you relate the above phenomenon to everyday life experiences like the drying of clothes on a rainy day or a dry day, which happens quickly? or the water droplets formed outside a glass containing a cold beverage, where do those water-droplets come from? or how the water in an earthen pot becomes cold?
All the above three examples can be explained at the level of particles, in this case, water molecules. The air on a rainy day is heavily laden with water vapors which slows down the escape of water vapors from the surface of clothes. The temperature of the glass tumbler containing cold water is less than the temperature of surroundings and this makes the water vapors condense on the outside of the tumbler. Similarly, the water in the earthen pot gets cool due to evaporation of water through the pores of the earthen pot and the heat required for the process is taken majorly from the water inside it. When heat is drawn out from something, it gets cold. As a rule, evaporation causes cooling.
Test your knowledge
Change of State of Matter
This quiz tests your knowledge about change of state of matter at the level of grade 9 and 10
