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    JEE Main Study Notes for States of Matter: Solid State Properties, Laws of Matter, Study Tips

    Anam Shams Anam Shams
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    JEE Main Physical Chemistry section carries a weightage of 30-40 % hence is considered important in the trio Physical, Organic and Inorganic chemistry. Numerical problems on topics like Chemical Kinetics, Equilibrium, Thermodynamics are asked frequently, but states of matter is one such topic in which direct theory questions are asked making it quite an easy chapter to cover from that point of view. 

    • States of matter carries a weightage of 23.2% in Chemistry section, thus 1-2 questions are necessarily put forward. 

    • Some sub topics repeatedly find place in the questions like Void, Radius-ratio, Density and Bravais Lattice. 

    • On a general note, any substance that has mass and occupies space is called matter. The matter is composed of atoms and molecules. 

    • The arrangement of these building blocks gives various states, physical and chemical properties to the matter. 

    Detailed study notes of States of Matter, based on the latest JEE Main Chemistry Syllabus have been given for reference to the students in the article. 

    General Properties of Matter

    General Properties of Matter

    PropertySolidLiquidGas
    ShapeDefinite shapeIndefinite shapeIndefinite shape
    VolumeDefinite VolumeDefinite VolumeIndefinite Volume
    Inter particular ForcesStrong Inter particular ForcesComparatively weaker Inter particular ForcesInte rparticular forces are negligible
    Inter particular SpaceNegligible inter particular spaceComparatively large inter particular spaceVery large Inter particular space
    Particular MotionParticle motion is restricted to vibratory motion.Particle motion is very slowParticle motion is very rapid and also random.
    Packing of ParticlesParticles are very Closely packedParticles are loosely packedParticles are very loosely packed
    CompressibilityIncompressibleCompressibleHighly Compressible
    DensityVery High DensityLow DensityVery low density

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    Solid                         

    In solid, constituent particles (ions, atoms, or molecules) are closely packed together. The forces between particles are so strong that the particles cannot move freely but can only vibrate. As a result, a solid has a stable, definite shape, and a definite volume. Solids can only change their shape by an outside force, as when broken or cut. 

    Liquid

    A liquid is a nearly incompressible fluid that conforms to the shape of its container but retains a (nearly) constant volume independent of pressure. The volume is definite if the temperature and pressure are constant. 

    • When a solid is heated above its melting point, it becomes liquid, given that the pressure is higher than the triple point of the substance. Intermolecular (or interatomic or interatomic) forces are still important, but the molecules have enough energy to move relative to each other and the structure is mobile. 

    • This means that the shape of a liquid is not definite but is determined by its container. The volume is usually greater than that of the corresponding solid, the best-known exception being water, H2O. The highest temperature at which a given liquid can exist is its critical temperature. 

    Gas

    Gas is a compressible fluid. In a gas, the molecules have enough kinetic energy so that the effect of intermolecular forces is small (or zero for an ideal gas), and the typical distance between neighbouring molecules is much greater than the molecular size. A gas has no definite shape or volume but occupies the entire container in which it is confined. A liquid may be converted to a gas by heating at constant pressure to the boiling point, or else by reducing the pressure at a constant temperature. 

    These notes contain all the important GAS LAWS and some important terms and relations.

    The gas laws are a set of laws that describe the relationship between THERMODYNAMIC TEMPERATURE ( T ), PRESSURE (P) AND VOLUME (V) of gases.

    • Boyle's law, 

    • Charles's law, 

    • Gay Lussac's law, 

    • Avogadro law, 

    • Ideal gas equation, 

    • Dalton's law of partial pressure, 

    • Graham’s law of diffusion 

    • Kinetic Theory of Gases 

    • Vander Waal's equation 

    Boyle’s Law

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    Boyle’s Law

    The volume of a given mass of a gas is inversely proportional to its pressure at a constant temperature. 

     ; V ∝ 1 / p or PV = K 

    K is a constant and its value depends on the mass, temperature and nature of the gas. 

    ∴ p1V1 = p2V2 ......=K

    ALSO, density d ∝ 1/V 

    Hence, P ∝ D 

    Example

    We all have seen a syringe while visiting a doctor. It is a medical device used to inject or withdraw fluid. It consists of a hollow cylinder called a barrel and a sliding plunger attached to it. The working principle of a syringe is like a reciprocating pump. When the plunger is pushed, the fluid will inject, and when the plunger is pulled, the fluid will withdraw the pushing of the plunger reduces the volume of the fluid in the barrel. This reduction in the volume causes a momentary increase in the pressure of the fluid, and the fluid is injected into the patient's body. In a similar way, the pulling of the plunger increases the volume of the fluid. It results in a momentary decrease in the pressure of the fluid, and external fluid is withdrawn.

    Charles’ Law

    Charles’ Law

    The volume of the given mass of a gas increases or decrease by 1 / 273 of its volume for each degree rise or fall of temperature respectively at constant pressure. 

    Vt = Vo (1 + t / 273) t constant p 

    or 

    The volume of a given mass of a gas is directly proportional to the absolute temperature at constant pressure. 

    V ∝ T (at constant p), V / T = constant or V1 / T1 = V2 / T2 

    Absolute zero is the theoretically possible temperature at which the volume of the gas becomes zero. It is equal to O°C or 273.15K. 

    Isobars A graph of V vs T at constant pressure is known as isobar 

    Charles’law explains that gases expand on heating, so hot air is less dense than cold air. 

    Example

    The human lungs are spongy air-filled organs play an important role in respiration. Air flows in when the lungs expand and flow out when they contract. In winters, the temperature of air decreases. As a consequent, the temperature of the air inside the body also decreases. According to Charles's law states volume is directly proportional to temperature. Hence, the volume of the air decreases with temperature. It shrinks the lungs and physical activities like jogging become difficult in freezing winter days.

    Gay Lussac’s Law

    Gay Lussac’s Law

    The pressure of a given mass of gas increases or decreases by 1 /273 of its pressure for each degree rise or fall of temperature respectively at constant volume. 

    pt = po (1 + t / 273) at constant V and n 

    or 

    The pressure of a given mass of a gas at constant volume is directly proportional to absolute temperature. 

    p ∝ T or p = KT or p / T = K at constant V and n or P1 / T1 = P2 / T2 

    Isochores A graph of p vs T at constant volume is known as isochore 

    Avogadro’s Law

    Example

    In hot summer days, the inflated tyres of vehicles may burst. The bursting of tyres is caused by Gay-Lussac's law. The inflated tyres are under high pressure. When the temperature of the air rises, the pressure of the gas in the tubes increases. After an unbearable point, the tyres fracture.

    Avogadro’s Law 

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

    Mathematically 

    V infi; n (at constant T and p) or V / n = K 

    Molar gas volume The volume of one mole of a gas, i.e., 224 Lat STP(0°C, 1 atm) i_!S known as molar gas volume 


    Ideal Gas Equation

    Ideal Gas Equation 

    V ∝1 / p, T and n constant (Boyle’s law) 

    V ∝ T, p and n constant (Charles’ law) 

    V ∝ n, p and T constant (Avogadro’s law) 

    ⇒ V ∝ nT / p or pV ∝ nT or pV = nRT. 

    This is known as an ideal gas equation. R is known as the universal gas constant. 

    From the ideal gas equation, density. 

    d = pM / RT (where, M = molecular mass) 

    Example

    Balloon filled with helium weighs much less than an identical balloon filled with air. Both balloons contain the same number of molecules. Helium atoms have a lower mass than either oxygen mPal Gas Equation

    Graham’s Law of Diffusion

    Similar conditions of temperature and pressure, the rates of diffusion of gases are inversely proportional to the square root of their densities. 

    Mathematically, r1 / r2 = √d2 / √d1 = √M2 / √M1 

    [Diffusion is the tendency of gases to distribute itself uniformly throughout the available space while effusion is the movement of gas through a small hole when it is subjected to pressure]. 

    Example

    The gases with different densities can be separated using Graham's law. It is also helpful in determining the molar mass of unknown gases by comparing the rate of diffusion of unknown gas to known gas. We can separate the isotopes of an element using Graham's law. A common example is enriching uranium from its isotope

    Dalton’s Law of Partial Pressure

    Dalton’s Law of Partial Pressure 

    At constant temperature. the total pressure. exerted by a mixture of non-reacting gases. is the sum of partial pressures of different gases present in the mixture. 

    p = p1 + p2 + p3 + …. 

    The Partial pressure of a gas = mole fraction of the gas * total pressure. 

    If n1, n2 and n3 are moles of non-reacting gases filled in a vessel of volume V at temperature T, the total pressure, p is given by 

    pV = (n1 + n2 + n3)RT 

    This is the equation of state of a gaseous mixture, 

    [Aqueous tension It is the pressure exerted by water vapours at a particular temperature. It depends upon temperature.] 

    The Pressure of a dry gas can be determined by Dalton’s law. When a gas is collected over water, its observed pressure is equal to the sum of the pressure of dry gas and the pressure of water vapour (aqueous tension) then 

    The Pressure of dry gas = pressure of moist gas – aqueous tension. 

    Example

    It refers to the effects of which partial pressure might have on scuba divers. While the total gas pressure increases as a diver increase their descent, the partial pressure of each gas involved increases as well which might cause harm to the diver's body if proper actions are not carried out.

    Kinetic Theory of Gases 

    Based on the basic postulates regarding the structure of gases, it is possible to derive a theoretical expression for the pressure of an ideal gas. The expression is

    pV = 1/3mNun2 

    m is the mass of each molecule of a gaseous system occupying volume V and exerting pressure p, N is the number of molecules contained in it and un2 is the mean square speed, defined as

    u2= (u12+u22+u32+......+un2)/N 

    From the above equation, it is possible to derive the various experimentally observed gaseous laws 


    Van der Waals’ Equation

    Van der Waals’ Equation 

    Van der Waals’ obtained the following equation for n. moles of a gas. 

    (p + n2a/Vm2) (Vm-nb) = RT 

    can also be written as 

     (p + n2a/V2) (V-nb) = nRT 

    where, 

    b = excluded volume or co-volume = 4 * actual volume of gas molecules 

    a = magnitude of attractive forces between gas molecules. 

    The greater the value of ‘a’, the greater the strength of van der Waals’ forces and greater is the ease with which a gas can be liquefied. 


    Vapour pressure  

    The pressure exerted by the vapours above the liquid surface when these are in equilibrium with the liquid at a given temperature is known as the vapour pressure of the liquid. 

    The vapour pressure of a liquid depends on: 

    1. Nature of liquid 

    1. Temperature: Vapour pressure increases with increasing temperature. 

    Example

    The amount of water vapour will increase and the pressure will increase if a bottle of water is heated up

    Boiling point 

    The temperature, at which vapour pressure of liquids becomes equal to the atmospheric pressure, is called the boiling point. 

    Example

    Milk is a mix of butterfat and water so it is slightly heavier than water. The boiling points of liquids are due to the gravity of the liquid. Water boils at 100 degrees Celsius (212 degrees Fahrenheit). While milk boils at 212.3 degrees Fahrenheit.

    Surface Tension 

     It is the force acting per unit length perpendicular to the imaginary line drawn on the surface of the liquid. It is denoted γ (gamma); 

    SI unit: Nm-1 

    Dimensions: kgs-2 

    The magnitude of surface tension of a liquid depends on the attractive forces between the molecules. It is measured with the help of an apparatus, called stalagmometer. 

    Surface tension decreases as the temperature increases. Rise or fall of liquid in a capillary tube is due to surface tension. 

    Example

    Water insects are able to walk water because of the wax secreted on their legs combined with surface tension's elastic-like cover.

    Viscosity  

    Viscosity is a measure of resistance to flow which arises due to friction between layers of fluid. 

    When there is a regular gradation of velocity, in passing from one layer to the next, it is called laminar flow. 

    F = η Adv / dz 

    where F = forces required to maintain the flow of layers. 

    A = area of contact 

    dv/ dz = velocity gradient; (the change in velocity with distance.) 

    ‘η’ is proportionality constant and is called  the coefficient of viscosity. The Viscosity of the coefficient  is the force when the velocity gradient is unity and the area of contact is unit area. CGS unit of coefficient of viscosity is poise S.I. unit of coefficient of viscosity is Nsm-2. 

    Example

    Water has a low or "thin" viscosity, for example, while honey has a "thick" or high viscosity.

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