Chapter 11

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# Chapter 11 - PowerPoint PPT Presentation

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

States of Matter

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:
• 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 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:
• 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 State:
• Very little space in-between particles, small compressibility.
• Increase in temp causes K.E to increase , very small thermal expansion.
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:
• 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
• 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:
• 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.
Problems:
• 1)Convert 55.2 kJ into joules, kilocalories and calories.
• 2)Convert 11,900 calories into joules, kilocalories and kilojoules.
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
Problems
• 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: 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: 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 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 force: Attractive forces between molecules.
• Weak forces compared to intra molecular.
Dipole-Dipole interactions
• Dipole-Dipole interactions: occurs between polar molecules.
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

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

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

Ionic compounds dissolved in water.