Chapter 11 States of Matter Physical states Property

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 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. • 1 Calorie= 4. 184 J. • 1 kcal= 4. 184 k. J.

Problems: • 1)Convert 55. 2 k. J 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 1 o. C. • 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/go. C and in cal/ go. C if 1638 J raise the temperature of 125 g of the solid from 25. 0 o. C to 52. 6 o. C. • 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.