STATES OF MATTER

STATES OF MATTER 

Most of the observable characteristics of chemical systems are represented as bulk properties of matter, i.e., the properties associated with a collection of a large number of atoms, ions or molecules.

For example1. An individual molecule of a liquid does not boil but the bulk boils.

2. Collection of water molecules have wetting properties; individual molecules do not wet.

3. Water can exist as ice, which is a solid; it can exist as liquid; or it can exist in the gaseous state as water vapour or steam. Physical properties of ice, water and steam are very different. In all the three states of water chemical composition of water remains the same i.e., H₂O. Characteristics of the three states of water depend on the energies of molecules and on the manner in which water molecules aggregate. Same is true for other substances also.

Chemical properties of a substance do not change with the change of its physical state; but rate of chemical reactions do depend upon the physical state. Many times in calculations while dealing with data of experiments we require knowledge of the state of matter. Therefore, it becomes necessary for a chemist to know the physical laws which govern the behaviour of matter in different states.

Here we study more about these three physical states of matter particularly liquid and gaseous states.It is necessary to under stand the nature of inter molecular forces, molecular interactions and effect of thermal energy on the motion of particles because a balance between these determines the state of a substance.

INTERMOLECULAR FORCES

Inter molecular forces are the forces of attraction and repulsion between interacting particles (atoms and molecules). (This term does not include the electrostatic forces that exist between the two oppositely charged ions and the forces that hold atoms of a molecule together i.e., co valent bonds.) Attractive inter molecular forces are known as van der Waals forces, in honour of Dutch scientist Johannes van der Waals (1837- 1923)

Vander Waals forces vary considerably in magnitude and include 
1. dispersion forces or London forces,

2. dipole-dipole forces, and 
 3. dipole-induced dipole forces.

hydrogen bonding is the strongest type of dipole-dipole interaction. Only a few elements i.e. F, O and N can participate in hydrogen bond formation, therefore it is treated as a separate category.

Note-- attractive forces between an ion and a dipole are known as ion-dipole forces and these are not vander Waals forces.

Dispersion Forces or London Forces---Atoms and non polar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed. But a dipole may develop momentarily in such atoms and molecules. Suppose we have two atoms ‘A’ and ‘B’ in the close vicinity of each other 

It may so happen that momentarily electronic charge distribution in one of the atoms, say ‘A’, becomes unsymmetrical i.e., the charge cloud is more on one side than the other  This results in the development of instantaneous dipole on the atom ‘A’ for a very short time. This instantaneous or transient dipole distorts the electron density of the other atom ‘B’, which is close to it and as a consequence a dipole is induced in the atom ‘B’. The temporary dipoles of atom ‘A’ and ‘B’ attract each other. Similarly temporary dipoles are induced in molecules also. This force of attraction was first proposed by the German physicist Fritz

London, and for this reason force of attraction between two temporary dipoles is known as London force or dispersion force. These forces are always attractive and interaction energy is inversely proportional to the sixth power of the distance between two interacting particles(i.e., 1/r⁶ where r is the distance between two particles). These forces work at short distances (~500 pm) and their magnitude depends on the portability of the particle.

Dipole - Dipole Forces---Dipole-dipole forces act between the molecules possessing permanent dipole. Ends of the dipoles possess “partial charges” shown by δ (delta).(Partial charges are always less than the unit electronic charge (1.6×10–19 C). The polar molecules interact with neighboring molecules. Fig (a) shows electron cloud

distribution in the dipole of hydrogen chloride and Fig. (b) shows dipole-dipole interaction between two HCl molecules. This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved.

(b)

The attractive force decreases with the increase of distance between the dipoles. The interaction energy is inversely proportional to distance between polar molecules. Dipole-dipole interaction energy between stationary polar molecules (as in solids) is proportional to 1/r³and that between rotating polar molecules is proportional to 1/r ⁶, where r is the distance between polar molecules. Besides dipole dipole interaction, polar molecules can interact by London forces also. Thus cumulative effect is that the total of inter molecular forces in polar molecules increase.

Dipole  Induced Dipole Forces---- This type of attractive forces operate between the polar molecules having permanent dipole and the molecules lacking permanent dipole.Permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electronic cloud 

Thus an induced dipole is developed in the other molecule. In this case also interaction energy is proportional to 1/r 6 where r is the distance between two molecules. Induced dipole moment depends upon the dipole moment present in the permanent dipole and the polarisability of the electrically neutral molecule. We know that molecules of larger size can be easily polarised. High polarisability increases the strength of attractive interactions. In this case also cumulative effect of dispersion forces and dipole-induced dipole interactions exists.

Hydrogen bond---- This is special case of dipole-dipole interaction.This is found in the molecules in which highly polar N–H, O–H or H–F bonds are present. Although hydrogen bonding is regarded as being limited

to N, O and F; but species such as Cl may also participate in hydrogen bonding. Energy of hydrogen bond varies between 10 to 100 kJ mol–1; therefore, hydrogen bonds are powerful force in determining the structure and properties of many compounds, for example proteins and nucleic acids. Strength of the hydrogen bond is determined by the coulombic interaction between the lone-pair electrons of the electronegative atom of one molecule and the hydrogen atom of other molecule. 

Molecules also exert repulsive forces on one another. When two molecules are brought into close contact with each other,the repulsion between the electron clouds and that between the nuclei of two molecules comes into play. Magnitude of the repulsion rises very rapidly as the distance separating the molecules decreases. This is the reason that liquids and solids are hard to compress. In these states molecules are already in close contact; therefore they resist further compression; as that would result in the increase of repulsive interactions.

THERMAL ENERGY--- is the energy of a body arising from motion of its atoms or molecules. It is directly proportional to the temperature of the substance. It is the measure of average kinetic energy of the particles of the matter and is thus responsible for movement of particles.This movement of particles is called thermal motion.

INTER MOLECULAR FORCES vs THERMAL INTERACTIONS

We know inter molecular forces tend to keep the molecules together but thermal energy of the molecules tends to keep them apart. Three states of matter are the result of balance between inter molecular forces and the thermal energy of the molecules.

When molecular interactions are very weak, molecules do not cling together to make liquid or solid unless thermal energy is reduced by lowering the temperature. Gases do not liquefy on compression only, although molecules come very close to each other and inter molecular forces operate to the maximum. However, when thermal energy of molecules is reduced by lowering the temperature; the gases can be very easily liquefied.

THE GASEOUS STATE

This is the simplest state of matter, we remain immersed in the ocean of air which is a mixture of gases. We spend our life in the lowermost layer of the atmosphere called troposphere, which is held to the surface of the earth by gravitational force. The thin layer of atmosphere is vital to our life. It shields us from harmful radiations

and contains substances like dioxygen, dinitrogen, carbon dioxide, water vapour, etc.

Let us study the behaviour of substances which exist in the gaseous state under normal conditions of temperature and pressure. In the periodic table only eleven elements exist as gases under normal conditions

The gaseous state is characterised by the following physical properties.

• Gases are highly compressible.

• Gases exert pressure equally in all directions.

• Gases have much lower density than the solids and liquids.

• The volume and the shape of gases are not fixed. These assume volume and shape of the container.

• Gases mix evenly and completely in all proportions without any mechanical aid.

Simplicity of gases is due to the fact that the forces of interaction between their molecules are negligible. 
Their behaviour is governed by same general laws. These laws are relationships between measurable properties of gases. These properties are pressure, volume, temperature and mass because relationships between these variables describe state of the gas. Interdependence of these variables leads to the formulation of gas laws.

THE GAS LAWS

Boyle’s Law (Pressure – Volume Relationship)

According to Robert Boyle,”at constant temperature, the pressure of a fixed amount (i.e., number of moles n) of gas  varies inversely with its volume. This is known as Boyle  law. Mathematically, it can be written as……..

where k₁is the proportionality constant. The

value of constant k1 depends upon the amount of the gas, temperature of the gas and the units in which p and V are expressed. On rearranging the above equation we obtain

pV = k₁

It means that at constant temperature, product of pressure and volume of a fixed amount of gas is constant.

If a fixed amount of gas at constant temperature T occupying volume V₁at pressure p₁undergoes expansion, so that volume becomes V₂and pressure becomes p₂, then 

Above figures are showing two conventional ways of graphical presentation of Boyle’s law.

Fig.(a) is the graph of equation (pV=k₁) at different temperatures. The value of k1 for each curve is different because for a given mass of gas, it varies only with temperature.

Each curve corresponds to a different constant temperature and is known as an isotherm (plot/graph at constant temperature).

Fig. (b) represents the graph between p and1/V

Experiments of Boyle, in a quantitative manner prove that gases are highly compressible because when a given mass of a gas is compressed, the same number of molecules occupy a smaller space. This means that gases become denser at high pressure.

A relationship can be obtained between density and pressure of a gas by using Boyle’s law :

By definition, density ‘d’ is related to the mass ‘m’ and the volume ‘V’ by the relation d=m/v. If we put value of V in this equation from Boyle’s law equation, we obtain the relationship.

This shows that at a constant temperature, pressure is directly proportional to the density of a fixed mass ofthe gas.

Charles’s Law (Temperature – VolumeRelationship)

According to this law, ‘for a fixed mass of a gas at constant pressure, volume of a gas increases on increasing temperature and decreases on cooling. He found that for each degree rise in temperature, volume of a gas increases by 1/273.15 of the original volume of the gas at 0 °C. Thus if volumes of the gas at 0 °C and at t °C are V₀and V_trespectively, then

At this stage, we define a new scale of temperature such that t °C on new scale is given by T = 273.15 + t and 0 °C will be given by T₀= 273.15. This new temperature scale is called the Kelvin temperature scale or

Absolute temperature scale.

Thus 0°C on the celsius scale is equal to 273.15 K at the absolute scale.

(Note that degree sign is not used while writing the temperature in absolute temperature scale, i.e., Kelvin scale.)

Kelvin scale of temperature is also called Thermodynamic scale of temperature and is used in all scientific

works.

Thus we add 273 (more precisely 273.15) to the celsius temperature to obtain temperature at Kelvin scale.

If we write T_t= 273.15 + t and T₀= 273.15 in the above equation we obtain the relationship

Thus we can write a general equation as follows.

Thus V = k₂ T

The value of constant k₂is determined by the pressure of the gas, its amount and the units in which volume V is expressed. Above equation is the mathematical expression for Charlesí law, which states that at constant pressure, the volume of a fixed mass of a gas is directlyproportional to its absolute temperature.

Charles found that for all gases, at any given pressure, graph of volume vs temperature (in celsius) is a straight line and on extending to zero volume, each line intercepts the temperature axis at – 273.15 °C. Slopes of lines obtained at different pressure are different but at zero volume all the lines meet the temperature axis at

– 273.15 °C (Fig.).

Fig. Volume vs Temperature ( ⁰C) graph

Each line of the volume vs temperature graph is called isobar.

Observations of Charles can be interpreted if we put the value of t in above equation as – 273.15 °C. We see that the volume of the gas at – 273.15 °C will be zero. This meansthat gas will not exist. All the gases get

liquified before this temperature is reached. The lowest hypothetical or imaginary temperature at which gases are supposed to occupy zero volume is called Absolute zero.

All gases obey Charles’ law at very low pressures and high temperatures.

Gay Lussacís Law (Pressure- Temperature Relationship)

The mathematical relationship between pressure and temperature was given by Joseph Gay Lussac and is known as Gay Lussac’s law. It states that at constant volume, pressure of a fixedamount of a gas varies directly with the temperature. Mathematically,

This relationship can be derived from Boyle’s law and Charles’ Law. Pressure vstemperature (Kelvin) graph at constant molar volume is shown in Fig.

Fig. Pressure vs temperature (K) graph (Isochores) of a gas.

Each line of thisgraph is called isochore.

Avogadro Law (Volume – AmountRelationship)

It states‘equal volumes of all gases under the same conditions of temperature and pressure contain equal number of molecules’.

Thismeans that as long as the temperature and pressure remain constant, the volume depends upon number of molecules of the gas/amount of the gas.Mathematically ……..

here n is the number of moles of the gas.

The number of molecules in one mole of a gas has been determined to be 6.022 ×10²³and is known as Avogadro constant.

Since volume of a gas is directly proportional to the number of moles; one mole of each gas at standard temperature andpressure (STP) will have same volume.

Standard temperature and pressure means 273.15 K (0 ⁰C) temperature and 1 bar (i.e., exactly 10⁵pascal) pressure. These values approximate freezing temperature of water and atmospheric pressure at sea

level. At STP molar volume of an ideal gas or a combination of ideal gases is 22.71098 L mol^ñ1.

Number of moles of a gas can be calculated as follows….

n = M/m here m = mass of the gas under investigation and M = molar mass

Thus, V = k₄M/m

Above equation can be rearranged as follows :

M = k₄m/V = k₄d

Here ‘d’ is the density of the gas.

It means from above equation that the density of a gas is directly proportional to its molar mass.

A gas that follows Boyle’s law, Charles’ law and Avogadro law strictly is called anideal gas.

Such a gas is hypothetical. It is assumed that intermolecular forces are not present between the molecules of an ideal gas. Real gases follow these laws only under certain specific conditions when forces of interaction are practically negligible. In all other situations these deviate from ideal behaviour.

IDEAL GAS EQUATION

The three laws which we have learnt till now can be combined together in a single equation which is known as ideal gas equation.

here R is proportionality constant. On rearranging the equation

R is called gas constant. It is same for all gases. Therefore it is also called UniversalGas Constant.

Equation is called idealgas equation.

Above equation shows that the value of R depends upon units in which p, V and T are measured.

This equation will be applicable to any gas, under those conditions when behaviour of the gas approaches ideal behaviour.

Volume of one mole of an ideal gas under STP conditions (273.15 K and 1 bar pressure) is 22.710981 L mol^–1. Value of R for one mole of an ideal gas can be calculated under these conditions as follows :

= 8.314 Pa m³K^–1mol^–1

= 8.314 × 10^–2bar L K^–1mol^–1

= 8.314 J K^–1mol^–1

At STP conditions used earlier (0 °C and 1 atm pressure), value of R is 8.20578 × 10^–2L atm K^–1mol^–1.

Ideal gas equation is a relation between four variables and it describes the state of any gas, therefore, it is also called equationof state.

Combined gas law

. If temperature, volume and pressure of a fixed amount of gas vary from T₁, V₁and p₁to T₂, V₂and p₂as..

Step-1 According to Boyle’s law…. At constant temp. T₁, By change in pressure p₁to p₂, volume changes from V₁to V_x, Hence

p₁V₁= p₂V_x or V_x= p₁V₁/p₂ ……….eq1 ( Continuous change)

Step-2 according to Charle’s law….. At constant pressure p₂, By change in temperature T₁to T₂, Volume changes from V_xto V₂, Hence

V_x = V₂ or Vx = V₂T₁/T₂ ………..eq 2

T₁ T₂

From eq 1 and 2….

p₁V₁ = V₂T₁

p₂ T₂ Or

This equation is known as Combined gas law.

Density and Molar Mass of a Gaseous Substance

Ideal gas equation can be rearranged as follows:

Replacing n by M/m, we get

(where d is the density)

On rearranging this equation we get a relation to calculate the molar mass of a gas.

Dalton’s Law of Partial Pressures

It states that the total pressure exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases

In a mixture of gases, the pressure exerted by the individual gas is called partial pressure. Mathematically,

p_Total= p₁+p₂+p₃+......(at constant T, V)

here p_Totalis the total pressure exerted by the mixture of gases and p₁, p₂, p₃etc. are partial pressures of gases.

Significance of this law; Gases are generally collected over water and therefore are moist. Pressure of dry gas

can be calculated by subtracting vapour pressure of water from the total pressure of the moist gas which contains water vapours also. Pressure exerted by saturated water vapour is called aqueous tension. (Aqueous tension of water at different temperatures may be taken by standard Table)

p _{Dry gas} = p _Total– Aqueous tension

Partial pressure in terms of mole fraction Suppose at the temperature T, three gases,enclosed in the volume V, exert partial pressure p₁, p₂and p₃respectively. then,

here n₁n₂and n₃are number of moles of these gases. Thus, expression for total pressure will be

p_Total= p₁+ p₂+ p₃

On dividing p₁by p _totalwe get

here n = n₁+n₂+n₃

x₁is called mole fraction of first gas.

Thus, p₁= x₁p _total

Similarly for other two gases we can write p₂= x₂pt_otaland p₃= x₃p_total

Thus a general equation can be written as p_i= x_ip_total

where pi and xi are partial pressure and mole fraction of i^thgas respectively.

KINETIC MOLECULAR THEORY OF GASES

Assumptions or postulates of the kinetic molecular theory of gases are given below.

• Gases consist of large number of identical particles (atoms or molecules) that are so small and so far apart on the average that the actual volume of the molecules is negligible in comparison to the empty space between them (i.e. total volume of gas). They are considered as point masses. This assumption explains the great compressibility of gases.

• There is no force of attraction between the particles of a gas at ordinary temperature and pressure. The support for this assumption comes from the fact that gases expand and occupy all the space available to them.

• Particles of a gas are always in constant and random motion. This is the reason gases do not have fixed shape and volume.

• Particles of a gas move in all possible directions in straight lines. During their random motion, they collide with each other and with the walls of the container. Pressure is exerted by the gas as a result of collision of the particles with the walls of the container.

• Collisions of gas molecules are perfectly elastic. This means that total energy of molecules before and after the collision remains same. (There may be exchange of energy between colliding molecules, their individual energies may change, but the sum of their energies remains constant.)

Note: If there were loss of kinetic energy, the motion of molecules will stop and gases will settle down.

• At any particular time, different particles in the gas have different speeds and hence different kinetic energies, but the distribution of speedsremains constant at a particulartemperature.

• If a molecule has variable speed, then it must have a variable kinetic energy. Hence we can calculate only average kinetic energy.

The average kinetic energy of the gas molecules is directly proportional to the absolute temperature.

It is observed that on heating a gas at constant volume, the pressure increases. Because on heating the gas, kinetic

energy of the particles increases and these strike the walls of the container more frequently thus exerting more pressure.

BEHAVIOUR OF REAL GASES: DEVIATION FROM IDEAL GAS BEHAVIOUR

To test behaviour of real gases or their deviation from ideal behaviour we plot pV vs p plot of gases because at constant temperature, pV will be constant (Boyle’s law) and pV vs p graph at all pressures will be a straight line

parallel to x-axis. Fig.

Fig. Plot of pV vs p for real gas and ideal gas Fig. Plot of pressure vs volume for real gas and ideal gas

Plot shows that at constant temperature pV vs p plot for real gases is not a straight line. There is a significant deviation from ideal behaviour. Two types of curves are seen…..

1. In the curves for dihydrogen and helium, as the pressure increases the value of pV also increases.

b)The second type of plot is seen in the case of other gases like carbon monoxide and methane. In these plots first there is a negative deviation from ideal behaviour, the pV value decreases with increase in pressure and reaches to a minimum value characteristic of a gas. After that pV value starts increasing. The curve then crosses the line for ideal gas and after that shows positive deviation continuously.

It is thus, found that real gases do not follow ideal gas equation perfectly under all conditions.

Real Gases; They do not follow, Boyle’s law, Charles law and Avogadro law perfectly under all conditions.

Q i. Why do gases deviate from the ideal behaviour?

Ans. We find that two assumptions of the kinetic theory do not hold good. These are

(a) There is no force of attraction between the molecules of a gas.

(b) Volume of the molecules of a gas is negligibly small in comparison to the space occupied by the i.e. volume of the gas.

Note: If assumption (a) is correct, the gas will never liquify. However, we know that gases do liquify when cooled and compressed. Also, liquids formed are very difficult to compress. This means that forces of repulsion are powerful enough and prevent squashing of molecules in tiny volume. If assumption (b) is correct, the pressure vs volume graph of experimental data (real gas) and that theoritically calculated from Boyles law (ideal gas) should coincide.

Real gases show deviations from ideal gas law because i) molecules interact with eachother. At high pressures molecules of gases are very close to each other so they can not strike the walls of thecontainer with full impact because these are dragged back by other molecules due tomolecular attractive forces. This affects the pressure exerted by the molecules on the wallsof the container. Thus, the pressure exerted by the gas is lower than the pressure exertedby the ideal gas.

Here, ‘a’ is a constant. Repulsive forces also become significant.Repulsive interactions are short-range interactions and are significant when molecules are almost in contact. This is the

situation at high pressure. The repulsive forces cause the molecules to behave as small but impenetrable spheres.

ii) The volume occupied by the molecules also becomes significant because instead of moving in volume V, these are now restricted to volume (V–nb) where nb is approximately the total volume occupied by the molecules themselves. Here, ‘b’ is a constant.

Having taken into account the corrections for pressure and volume, we can rewrite ideal gas equation as..

This equation is known as van der Waals equation.In this equation n is number of moles of the gas. Constants a and bare called van der Waals constantsand their value depends on the characteristic of a gas.

Significance of vanderwall’s constants ‘a’ and ‘b’----

Valueof ‘a’ is measure of magnitude of intermolecular attractive forces within the gas and is independent of temperature andpressure.

And b is the measurement of actual volume of 1 mole gas molecules.

Note; Also, at very low temperature, intermolecular forces become significant. As the molecules travel with low average speed, these can be captured by one another due to attractive forces.

Q.(ii) What are the conditions under which gases deviate from ideality?

Ans. At high pressure and low temperature.

Real gases show ideal behaviour when conditions of temperature and pressure are such that the intermolecular

forces are practically negligible. Thus the real gases show ideal behaviour when pressure approaches zero.

Q. What is the compressibility factor and what is its significance?

Ans. It is the ratio of product pV and nRT. Mathematically

Significanceof compressibility factor Z, --- (i) It used to measurethe deviation of real gases from ideal behaviour.

For ideal gas Z = 1 at all temperatures and pressures because pV = n RT. The graph of Z vs p will be a straight line parallel to pressure axis (Fig).

Fig Variation of compressibility factor for some gases

For gases whichdeviate from ideality, value of Z deviates fromunity.

At very low pressures all gases shown have Z ≈1 and behave as ideal gas.

At highpressure all the gases have Z > 1. These aremore difficult to compress.

At intermediatepressures, most gases have Z < 1. Thus gasesshow ideal behaviour when the volumeoccupied is large so that the volume of the molecules can be neglected in comparisonto it.

Note; Up to what pressure a gas will follow the ideal gas law, depends upon nature of the gas and its temperature. The temperature at which a real gas obeys ideal gas law over an appreciable range of pressure is called Boyle

temperature or Boyle point. Boyle point of a gas depends upon its nature.

Above explanation shows that at low pressure and high temperature gases show ideal behaviour. These conditions are different for different gases.

(ii) To calculate the ideal volume of a real gas--- if we note the following derivation

If the gas shows ideal behaviour then

On putting this value of nRT/pin above equation we have

From above equation, ‘compressibility factor is the ratio of actual molar volume of a gas to the molar volume of

it, if it were an ideal gas at that temperature and pressure’.

LIQUIFACTION OF GASES

First of all Thomas Andrewsstudy the pressure - volume - temperature relationsof a substance in both gaseous and liquid state for carbon dioxide. He plotted isotherms (Plots at constant temperature) of carbon dioxide at various temperatures (Fig.).

Fig. Isotherms of carbon dioxide at various temperatures

He found that real gases behave as carbon dioxide. Andrews noticed that at high temperatures isotherms look like that of an ideal gas and the gas cannot be liquefied even at very high pressure. As the temperature is lowered, shape of the curve changes and data shows considerable deviation from ideal behaviour.

At 30.98 °Ccarbon dioxide remains gas upto 73 atmospheric pressure. (Point E in Fig). At 73 atmospheric pressure, liquid carbon dioxide appears for the first time. The temperature 30.98 °C is called criticaltemperature (T_C) of carbon dioxide. “This is the highest temperature at which carbon dioxide gas can be liquefy by applying pressure.” Above this temperature it is gas. Volume of one mole of the gas at critical temperature is called critical volume (V_C)and pressure at this temperature is called critical pressure (p_C).

The critical temperature, pressure and volume are called critical constants. Further increase in pressure simply compresses the liquid carbon dioxide and the curve represents the compressibility of the liquid.

The steep line represents the isotherm of liquid. Even a slight compression results in steep rise in pressure indicating very low compressibility of the liquid.

Below 30.98 °C,.i.e. At 21.5 °C, carbon dioxide remains as a gas only upto point B. At point B, liquid of a

particular volume appears. Further compression does not change the pressure. Liquid and gaseous carbon dioxide coexist and further application of pressure results in the condensation of more gas until the point C is reached. At point C, all the gas has been condensed and further application of pressure merely compresses the liquid as

shown by steep line. A slight compression from volume V2 to V3 results in steep rise in pressure from p2 to p3 (Fig.). Below 30.98 °C (critical temperature) each curve shows the similar trend. Only length of the horizontal

line increases at lower temperatures. At critical point horizontal portion of the isotherm merges into one point.

Thus we see that a point like A in the Fig. represents gaseous state. A point like D represents liquid state and a point under the dome shaped area represents existence of liquid and gaseous carbon dioxide in equilibrium. All the gases upon compression at constant temperature (isothermal compression) show the same behaviour as shown by carbon dioxide. From above it is clear that gases should be cooled below their critical temperature for liquification. Critical temperature of a gas is highest temperature at which liquifaction of the gas first occurs. Liquifaction of calledig. 5.11151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515151515permanent gases (i.e., gases which show continuous positive deviation in Z value) requires cooling as well as considerable compression. Compression brings the molecules in close vicinity and cooling slows down the movement of molecules therefore, intermolecular interactions may hold the closely and slowly moving molecules together and the gas liquifies.

Q. Is it possible to change a gas into liquid or a liquid into gas by a process in which always a single phase is present? If yes, Explain your answer.

Ans. For example in above Fig. we can move from point A to F vertically by increasing the temperature, then

we can reach the point G by compressing the gas at the constant temperature along this isotherm (isotherm at 31.1°C). The pressure will increase. Now we can move vertically down towards D by lowering the temperature.

As soon as we cross the point H on the critical isotherm we get liquid. We end up with liquid but in this series of

changes we do not pass through two-phase region. If process is carried out at the critical temperature, substance always remains in one phase.

Note.The term fluid is used for either a liquid or a gas to recognize this continuity. Thus a liquid can be viewed

as a very dense gas. Liquid and gas can be distinguished only when the fluid is below its critical temperature and its pressure and volume lie under the dome, since in that situation liquid and gas are in equilibrium and a surface separating the two phases is visible. In the absence of this surface there is no fundamental way of distinguishing

between two states.

A gas below the critical temperature can be liquefied by applying pressure, and is called vapour of the substance. Carbon dioxide gas below its critical temperature is called carbon dioxide vapour.

LIQUID STATE--- Important Features are as ---

1. Intermolecular forces are stronger in liquid state than in gaseous state hence Molecules in liquids are so close that there is very little empty space between them and under normal conditions liquids are denser than gases.

2. Molecules of liquids are held together by attractive intermolecular forces so they have definite volume because molecules do not separate from each other but they can move past one another freely therefore, liquids can flow, can be poured and can assume the shape of the container in which these are stored.

Now we will study some of the physical properties of the liquids such as vapour pressure, surface tension and viscosity.

Vapour Pressure

The liquid evaporates and pressure exerted by vapours on the walls of the container (vapour pressure) increases.

After some time it becomes constant, an equilibrium is established between liquid phase and vapour phase. Vapour pressure at this stage is known as equilibrium vapourpressure or saturated vapour pressure..

“The liquid evaporates and pressure exerted by vapours on the walls of the closed container at equilibrium and at constant temperature is known as vapour pressure of that liquid.”

Vapourisation can occur throughout the bulk of the liquid and vapours expand freely into the surroundings. The condition of free vapourisation throughout the liquid is called boiling.

The temperature at which vapour pressure of the liquid becomes equal to the external pressure (atmospheric pressure), is called boiling temperature or boiling point at that pressure.

At 1 atm pressure boiling temperature is called normal boiling point.

If pressure is 1 bar then the boiling point is called standard boiling point of the liquid.

Standard boiling point of the liquid is slightly lower than the normal boiling point because 1 bar pressure is slightly less than 1 atm pressure. The normal boiling point of water is 100 °C (373 K), its standard boiling point

is 99.6 °C (372.6 K).

At high altitudes atmospheric pressure is low. Therefore liquids at high altitudes boil at lower temperatures in comparison to that at sea level. Since water boils at low temperature on hills, so the pressure cooker is used for cooking food. In hospitals surgical instruments are sterilized in autoclaves in which boiling point of water is increased by increasing the pressure above the atmospheric pressure by using a weight covering the vent.

Boiling does not occur when liquid is heated in a closed vessel. On heating continuously vapour pressure increases. At first a clear boundary is visible between liquid and vapour phase because liquid is more dense than vapour. As the temperature increases more and more molecules go to vapour phase and density of vapours rises.

At the same time liquid becomes less dense. It expands because molecules move apart. When density of liquid and vapours becomes the same; the clear boundary between liquid and vapours disappears. This temperature is

called critical temperature.

Surface Tension

We know that---1. the liquids assume the shape of the container.

2. the small drops of mercury form spherical bead instead of spreading on the surface.

3. the particles of soil at the bottom of river remain separated but they stick together when taken out.

4. a liquid rise (or fall) in a thin capillary as soon as the capillary touches the surface of the liquid.

5. the falling rain drops are spherical in shape.

All these phenomena are caused due to the characteristic property of liquids, called surface tension.

A molecule in the bulk of liquid experiences equal intermolecular forces from all sides, therefore it does not experience any net force. But the molecule on the surface of liquid, feels net attractive force towards the interior of the liquid (Fig) due to themolecules below it. Since there are nomolecules above it.

Fig.Vapour pressure vs temperature curve of Fig. Forces acting on a molecule on liquid surface and

some common liquids on a molecule inside the liquid

Liquids tend to minimize their surfacearea. The molecules on the surface experiencea net downward force and have more energythan the molecules in the bulk, which do notexperience any net force. Therefore, liquids

tend to have minimum number of molecules at their surface. If surface of the liquid is increased by pulling a molecule from the bulk, attractive forces will have to be overcome. This will require expenditure of energy.

“The energy required to increase the surface area of the liquid by one unit is defined as surface energy.

Unit is J m^–2”.

“Surface tension is defined as the force acting per unit length perpendicular to the line drawn on the surface of liquid”. It is denoted by Greek letter γ (Gamma). It has dimensions of kg s^–2and in SI unit it is expressed as N m^–1.

Note; The lowest energy state of the liquid will be when surface area is minimum. Spherical shape satisfies

this condition, because spherical shapes have minimum surface area for the given volume, that is why

1.mercury drops are spherical in shape.

2.the sharp glass edges are heated for making them smooth. (On heating, the glass melts and the surface of the liquid tends to take the rounded shape at the edges, which makes the edges smooth. This is called fire polishing of glass.

3. Liquids wet the things because they spread across their surfaces as thin film. Moist soil grains are pulled together because surface area of thin film of water is reduced.

4. It is surface tension which gives stretching property to the surface of a liquid. On flat surface, droplets are slightly flattened by the effect of gravity; but in the gravity free environments drops are perfectly spherical.

Factors affecting the magnitude of surface tension---1. The surface tension of a liquid depends on the attractive forces between the molecules. When the attractive forces are large, the surface tension is large.

2. Increase in temperature increases the kinetic energy of the molecules and effectiveness of intermolecular attraction decreases, so surface tension decreases as the temperature is raised.

Viscosity-. Viscosity is a measure of resistance to flow which arises due to the internalfriction between layers of fluid as they slip past one another while liquid flows.

Strongintermolecular forces between molecules hold them together and resist movement of layerspast one another.When a liquid flows over a fixed surface, the layer of molecules in the immediatecontact of surface is stationary. The velocity of upper layers increases as the distance oflayers from the fixed layer increases. This

type of flow in which there is a regular gradation of velocity in passing from one layer to the next is called laminar flow. If we choose any layer in the flowing liquid (Fig.5.14), the layer above it accelerates its flow and the layer below this retards its flow.

Fig. Gradation of velocity in the laminar flow

If the velocity of the layer at a distance dz is changed by a value du then velocity gradient is given by the amount

du/dz. A force is required to maintain the flow of layers. This force is proportional to the area of contact of layers and velocity gradient i.e.

(A is the area of contact)

(where, du/dz is velocity gradient; the change in velocity with distance)

‘ η ’ is proportionality constant and is called coefficient of viscosity. Viscosity coefficient is the force when velocity gradient is unity and the area of contact is unit area.

Thus ‘ η ’ is measure of viscosity. SI unit of viscosity coefficient is 1 newton second per square metre

(N s m–2) = pascal second (Pa s = 1kg m–1s–1). In cgs system the unit of coefficient of viscosity is poise (named after great scientist Jean Louise Poiseuille). 1 poise = 1 g cm–1s–1 = 10–1kg m–1s–1

Greater the viscosity, the more slowly the liquid flows. Hydrogen bonding and van der Waals forces are strong enough to cause high viscosity.

Glassis an extremely viscous liquid. It is so viscous that many of its properties resemble solids. However, property of flow of glass can be experienced by measuring the thickness of windowpanes of old buildings. These become thicker at the bottom than at the top.

Factors affecting the viscosity of liquid--- Viscosity of liquids decreases as the temperature rises because at high temperature molecules have high kinetic energy and can overcome the intermolecular forces to slip past one another between the layers




Chemistry for class 11.....
elements-of-group-13-p-block-elements
states-of-matter-liquids-and-solids
geometric-isomerism-different-geometries
chemical-thermodynamics
introducation-of-carbon-chemistry
electrons-in-atom-and-periodic-table
hybridisation
intermolecular-forces-liquid-and-solids
niels-bohr-atomic-model
iupac-nomenclature-of-organic-compounds
chemical-bonding-molecular-geometry
molecular-orbital-theory
heisenberg-uncertainty-principle
some-basic-concepts-of-chemistry
equilibrium
environmental-chemistry
hydrogen
structure-of-atom

classification-of-elements

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