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Chapter 11. States of Matter. Physical states. Property differences among Physical states. Compressibility: measure of volume change resulting from pressure change. Thermal Expansion: Measure of volume change resulting from temperature change . The KMT of Matter:.

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Chapter 11

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chapter 11

Chapter 11

States of Matter

property differences among physical states
Property differences among Physical states
  • Compressibility: measure of volume change resulting from pressure change.
  • Thermal Expansion: Measure of volume change resulting from temperature change
the kmt of matter
The KMT of Matter:
  • Five statements used to explain three states.
  • Matter is composed of discrete tiny particles.
  • Particles are in constant motion and possess K.E (energy because of motion).
  • Particles interact through attractions and repulsions and possess P.E(stored energy).
the kmt of matter5
The KMT of Matter:
  • Particles of opposite charges attract, like charges repel.
  • Electrostatic force: An attractive force or repulsive force that occurs between charged particles.
  • The K.E (velocity) increases as temp increases.
  • Particles in a system transfer energy through elastic (total KE remains constant) collisions.
the solid state
The Solid State:
  • Physical state characterized by a dominance of P.E (cohesive forces) over K.E (disruptive forces).
  • Strong cohesive forces hold particles in fixed positions, def S and def V.
  • Large number of particles in a unit volume, high density.
the solid state7
The Solid State:
  • Very little space in-between particles, small compressibility.
  • Increase in temp causes K.E to increase , very small thermal expansion.
the liquid state
The Liquid state:
  • Physical state characterized by P.E and K.E of same magnitude.
  • Definite volume and indefinite shape.
  • High density: Particles not widely spread.
  • Small compressibility: Very less empty space.
  • Small thermal expansion
the gaseous state
The Gaseous state:
  • Physical state characterized by a complete dominance of K.E over P.E
  • Indefinite V and Shape.
  • Low density: Particles are widely separated and few of them present in a given volume.
  • Large compressibility: Particles widely separated.
  • Moderate thermal expansion: Volume increases with increase in temp.
a comparison of solids liquids and gases
A Comparison of Solids, Liquids and Gases
  • In gases particles are far away from each other compared to solids and liquids.
  • The distance ratio between particles of s, l, g: 1 to 1.1 to 10
endothermic exothermic
  • Endothermic: System absorbs energy. Ex: melting, sublimation, evaporation.
  • Exothermic: System releases (exits) energy. Ex: deposition, condensation, freezing.
heat energy and specific heat
Heat energy and Specific Heat:
  • The SI unit for heat energy is the joule (pronounced “jool”).
  • Another unit is the calorie.
  • 1 Joule of energy is required to raise the temperature of 1 g of water by 1 C.
  • 1Calorie= 4.184 J.
  • 1kcal= 4.184 kJ.
  • 1)Convert 55.2 kJ into joules, kilocalories and calories.
  • 2)Convert 11,900 calories into joules, kilocalories and kilojoules.
specific heat
Specific Heat:
  • The specific heat of a substance is the quantity of heat required to change the temperature of 1 g of that substance by 1oC.
  • The units of specific heat in joules are: J/g C
  • Q= mc∆t
  • Q= heat in J, m= mass in g, c= specific heat in J/g C, ∆t= change in temperature in C
  • 3)Calculate the specific heat of a solid in J/goC and in cal/ goC if 1638 J raise the temperature of 125 g of the solid from 25.0oC to 52.6oC.
  • 4)Calculate the number of Joules of heat energy needed to increase the temperature of 50.0 g of Cu from 21.0 C to 80.0 C. c of Cu= 0.382 J/g C.
evaporation of liquids
Evaporation of Liquids
  • Evaporation: Process by which molecules escape from a liquid to a gaseous phase.
  • Vapor: Gaseous sate of a substance at a temperature and pressure at which substance is normally a liquid or solid.
  • Equilibrium state: Two opposite processes take place at same rate.
vapor pressure of liquids
Vapor Pressure of liquids
  • Vapor pressure: Pressure exerted by a vapor above a liquid when liquid and vapor are at equilibrium.
  • Volatile: Readily evaporates at RT.
  • Boiling: Conversion from liquid to vapor (evaporation) occurs within the body through bubble formation.
vapor pressure of liquids22
Vapor Pressure of liquids
  • Boiling point: Temperature of a liquid at which the vapor pressure of the liquid becomes equal to the external atmospheric pressure exerted on the liquid.
  • Normal boiling point: Temperature of liquid at which it boils under a pressure of 760 mm Hg.
intermolecular forces in liquids
Intermolecular forces in Liquids
  • Intermolecular force: Attractive forces between molecules.
  • Weak forces compared to intra molecular.
dipole dipole interactions
Dipole-Dipole interactions
  • Dipole-Dipole interactions: occurs between polar molecules.
hydrogen bond
Hydrogen bond

H is covalently bonded to a highly electronegative element of small size(F, O , N). It is a very strong dipole-dipole interaction. Ex: water.

london forces
London Forces

Weakest type of all intermolecular. Occurs between an atom and a molecule.

ion dipole interactions
Ion dipole interactions
  • Occurs between an ion and a polar molecule
ion ion interactions
Ion-Ion interactions

Ionic compounds dissolved in water.