States of matter
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  • CHEMISTRY
  • END

States of matter States of matter Presentation Transcript

  • STATES OF MATTER - Presented by SHRABANTI HALDER Class-XI A CHEMISTRY
    • THREE STATES OF MATTER
    does not flow easily  rigid - particles cannot move/slide past one another flows easily  particles can move/slide past one another flows easily  particles can move past one another not easily compressible  little free space between particles not easily compressible  little free space between particles compressible  lots of free space between particles retains a fixed volume and shape  rigid - particles locked into place assumes the shape of the part of the container which it occupies  particles can move/slide past one another assumes the shape and volume of its container particles can move past one another Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior Solids Liquids Gas
  •  
  •  
    • INTERMOLECULAR FORCES
    • Intramolecular forces are attractive forces between molecules.
    • They hold atoms together in a molecule.
    • INTERMOLECULAR Vs INTRAMOLECULAR
    • 41 KJ to vaporize 1 mole of water (inter)
    • 930 KJ to break all O-H bonds in 1 mole of water (intra)
    • Measure of intermolecular force
    • boiling point
    • melting point
    • ΔHvap
    • ΔHsub
    • ΔHfus
    Generally, Intermolecular forces are much weaker than Intramolecular forces.
  •  
    • TYPES OF INTERMOLECULAR FORCES
    • DISPERSION FORCES OR LONDON FORCES
    • The London dispersion force is the weakest intermolecular force. The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. This force is sometimes called an induced dipole-induced dipole attraction. London forces are the attractive forces that cause nonpolar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently..
    • DIPOLE - DIPOLE FORCES
    • Dipole-dipole forces are attractive forces between the positive end of one polar
    • molecule and the negative end of another polar molecule. Dipole-dipole forces
    • have strengths that range from 5 kJ to 20 kJ per mole. They are much weaker
    • than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).
    • INDUCED - DIPOLE FORCES
    • Ion – induced dipole forces
    • An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.
    • Dipole – Induced Dipole Forces
    • A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.
    • HYDROGEN BOND
    • The hydrogen bond is a special dipole-dipole interaction between the
    • hydrogen atom in a polar N-H, O-H or F-H bond and an electronegative
    • O, N or F atom.
    • A – H ----B or A – H ----A
    • A and B are O, N and F
    • BOYLE’S LAW (Pressure – Volume Relationship)
    • According to Robert Boyle’s Law, at constant temperature, the
    • pressure of a fixed amount of gas varies inversely with its volume.
    • GAS LAWS
    pv = constant pv = k
    • CHARLE’S LAW (Temperature – Volume Relationship)
    • Charle’s Law states that pressure remaining constant, the
    • volume of a fixed mass of a gas is directly proportional
    • to it’s absolute temperature.
    • GAY LUSSAC’S LAW (Pressure – Temperature Relationship)
    • Gay Lussac’s law states that at constant volume, pressure of a
    • fixed amount of a gas varies directly with the temperature.
    • AVAGADRO LAW (Volume – Amount Relationship)
    • It states that equal volumes of all gases under the same conditions
    • of temperature and pressure contain equal number of molecules.
    • Ideal gas equation is a relation between four variables and it describes the
    • state of any gas, therefore, it is also called equation of state.
    • IDEAL GAS EQUATION
    Combined gas law
    • This law 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.
    • DALTON’S LAW OF PARTIAL PRESSURE
    • A gas is composed of particles in constant motion.
    • The average kinetic energy depends on temperature, the higher the temperature, the higher the kinetic energy and the faster the particles are moving.
    • Compared to the space through which they travel, the particles that make up the gas are so small that their volume can be ignored.
    • The individual particles are neither attracted to one another nor do they repel one another.
    • When particles collide with one another (or the walls of the container) they bounce rather than stick. These collisions are elastic; if one particle gains kinetic energy another loses kinetic energy so that the average remains constant.
    • KINETIC MOLECULAR THEORY ON GASES
    The Assumptions of Kinetic Molecular Theory:
  •  
    • A gas which obeys the gas laws and the gas equation PV = nRT strictly at all
    • temperatures and pressures is said to be an ideal gas. The molecules of ideal gases
    • are assumed to be volume less points with no attractive forces between one another.
    • But no real gas strictly obeys the gas equation at all temperatures and pressures.
    • Deviations from ideal behaviour are observed particularly at high pressures or low
    • temperatures. The deviation from ideal behaviour is expressed by introducing a factor
    • Z known as compressibility factor in the ideal gas equation. Z may be expressed as
    • Z = PV / nRT
    •  
    • •       In case of ideal gas, PV = nRT so, Z = 1
    •  
    • •        In case of real gas, PV ≠ nRt  so,Z ≠ 1
    • Thus in case of real gases Z can be < 1 or > 1
    •  
    • (i)   When Z < 1, it is a negative deviation. It shows that the gas is more compressible than expected from ideal behaviour.
    • (ii)   When Z > 1, it is a positive deviation. It shows that the gas is less compressible than expected from ideal behaviour.
    • BEHAVIOUR OF REAL GASES : DEVIATION FROM IDEAL GAS BEHAVIOUR
    • Causes of deviation from ideal behaviour
    •  
    • The causes of deviations from ideal behaviour may be due to the
    • following two assumptions of kinetic theory of gases. There are
    •  
    • The volume occupied by gas molecules is negligibly small as compared to the volume occupied by the gas.
    • The forces of attraction between gas molecules are negligible.
    • Atoms and molecules are never truly ideal because they all interact with other gas particles; weak attractions between separate gas particles are known as intermolecular attractions or van der Waals forces after the chemist who proposed a correction to the ideal gas law to calculate pressure of a real gas. Van der Waals proposed that the ideal gas equation could be corrected for real gas behavior by subtracting the effective gas particle volume from the volume of the container and by correcting for intermolecular attractions:
    • VAN DER WAALS EQUATION
    • VAPOUR PRESSURE
    • The pressure exerted by the vapours in equillibrium wiyh liquids
    • form of the same substance is called vapour pressure of the liquid.
    • PROPERTIES OF LIQUIDS
    • SURFACE TENSION
    • Surface tension is the force acting per unit length perpendicular
    • to the line drawn on the surface of the liquid.
    • VISCOSITY
    • Viscosity of a liquid is measure of its resistance to flow. It is related to
    • inter molecular force of attraction ; stronger these forces, greater the viscosity.
  • CHAPTER COMPLETED END