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:

• 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.

Compression involves decreasing amount of empty space in container.

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.

Within molecules,chem bonds Between molecules

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.