Chapter 11.States of Matter

States of Matter

State is Determined by:Chemical IdentityTemperaturePressure

States of Elements

Kinetic Molecular Theory of Matter

The Kinetic Molecular Theory of Matter is an explanation of the behavior of matter, based on the idea that the particles that make up the matter are always in motion.

The particles can be atoms, molecules, or ions.

Kinetic Molecular Theory of Matter

1. Matter is composed of tiny particles; the size of the particles is fixed for each sub-stance.

2. The particles are in constant random mo-tion and therefore possess kinetic energy (energy of motion, which can be trans-ferred by collisions).

Kinetic Molecular Theory of Matter

3. The particles interact with each other by means of electrostatic attractions and repulsions, and therefore possess potential energy (stored energy, pos-sessed by matter as a result of its po-sition, condition, and or composition).

Kinetic Molecular Theory of Matter

4. The kinetic energy (velocity) of the particles increases as temperature is increased.

5. The particles in a system transfer energy by means of elastic collisions (collisions in which all energy transfer results in motion, not deformation).

Kinetic Molecular Theory of Matter

No kinetic energy is lost in elastic collisions.Kinetic energy is transformed to work and/or

heat in inelastic collisions.Kinetic energy is a disruptive force between

particles; it makes them more independent of each other.

Potential energy is a cohesive or attractive force between particles.

Comparison of States

Some Definitions:

Density is the ratio of mass to volume.Compressibility is a measure of volume

change resulting from a pressure change.

Thermal Expansion is a measure of volume change resulting from temperature change.

Properties of StatesSolids

Definite Volume, Definite Shape

Density is High: 1.0 – 20 g/cm3

Compressibility is Low

Thermal Expansion is Low: 0.01% per C

The Solid StateIn solids, cohesive forces predominate over

kinetic energy. Particles are usually held together in a regular array, and vibrate about fixed positions.

Electrostatic attractions between particles keep them close together in fixed positions.The particles fill 50-70% of the space available, the rest is “void volume.”

Properties of StatesLiquids

Definite Volume, Indefinite Shape

Density is Fairly High: 0.5 – 15 g/mL

Compressibility is Low

Thermal Expansion is Fairly Low: 0.1% per C

The Liquid StateIn liquids, cohesive forces are balanced by

kinetic energy. Particles move freely about each other but do not separate.

Electrostatic attractions between particles keep them close together, but able to move rel-ative to one another.The particles fill about 50% of the available space.

Properties of StatesGases

Indefinite Volume, Indefinite Shape

Density is Low: 0.2 – 10 g/L

Compressibility is High

Thermal Expansion is Moderate: 0.3% per C

The Gaseous StateIn gases, cohesive forces are overcome by

kinetic energy. Particles move indepen-dently of each other.

Electrostatic attractions between particles are very weak, do not hold them together.The particles fill about 1% of the space available.

Changes in StateMost substances can exist in any phase,

solid, liquid, or gas.

The phase at which the substance exists depends on its temperature and the applied pressure.

A phase diagram is a graph showing the phase behavior of a given substance.

Phase Diagram for Water

Phase Diagram for CO2

Changes of StateExothermic: Endothermic

Release Heat Energy Absorb Heat EnergyH is negative H is positive

Changes of StateFreezing point the temperature at which the

liquid and solid phases of a substance are in

equilibrium

Boiling point the temperature at which the liquid and vapor phases of

a substance are in equilibrium.

Equilibrium is a state in which two oppos-ing processes occur at

equal rates.

Energy and Heat

Energy is the capacity to do work.

Energy can exist in different forms, and can change in form:

HeatLightElectricalMechanical

Heat Energy

First Law of Thermodynamics

Energy is neither created nor destroyed, just changed in form and/or transferred.

Energy Units

The joule (J) is the base unit for energy or work (force x distance).

1 J = 1 kgm2/sec2

4.18 J = 1 calorie

1 calorie is the amount of heat energy required to raise the temperature of 1 g of water by 1C.

Specific Heat

Specific heat is the amount of heat energy required to raise the temperature of 1.00 gram of a substance by 1.00C.

It takes 4.18 J to raise the temperature of 1.00 g of water by 1.00C.

4.18 J/gC is the specific heat of water.

Specific Heat

4.18 J/gC very high!

Metals have specific heats of 0.1 to 1.0 J/gC;They conduct heat.

Most nonmetallic materials are insulators, with specific heats of 1 to 2 J/gC.

Specific Heat

How much heat is absorbed if I heat 100 g of water from 25.0C (room temperature) to 100.0C (boiling point of water)?

How much heat is given off if I cool 1.00 lb (454 g) of iron metal from 100.0C to 25.0C? Specific heat of iron is 0.444 J/gC.

Specific Heat

What is the specific heat of a rock? Its mass is 125 g. I heat it to 100.0C in a boiling water bath, then drop it into 100.0 g of water that's at 20.0C. The water temperature rises to 30.0C.

Energy and Changes of State

A heating or cooling curve is a graph showing the amount of energy required to change the temperature or phase of a given amount of a substance.

Heating and Cooling Curves

Energy and Changes of State

In an endothermic phase change: heat energy is absorbed by a substanceits particles gain kinetic energyforces between particles are overcome

The substance goes into a less ordered state.

It melts, boils, or sublimes.

Energy and Changes of StateIn an exothermic phase change:

heat energy is given off by a substanceits particles lose kinetic energyforces between particles can act

The substance goes into a more ordered state.

It freezes, condenses, or deposits.

Energy and Changes of StateDuring an endothermic phase change, all the energy being supplied to the substance is used to disrupt forces between particles.

No temperature change is observed.

The temperature of a substance that is present in two phases will remain constant until all of one phase is consumed.

Energy and Changes of StateDuring an exothermic phase change, all the energy being released by the substance is allowing intermolecular forces to bring particles to a more ordered state.

No temperature change is observed.

The temperature of a substance that is present in two phases will remain constant until all of one phase is consumed.

Heats of Fusion and Vaporization

The amount of heat required to cause a phase change in a given material is a physical prop-erty of that material.

Heat of fusion, Hf = heat to convert one gram of substance from solid to liquid at its melting point. Units: J/g

Heat of vaporization, Hv = heat to convert one gram of substance from liquid to gas at its boiling point. Units: J/g

Heats of Fusion and Vaporization

How much energy will be consumed when 150 g of water at 100C is boiled to steam, also at 100C?

Hv water = 2260 J/g

How much energy will be released if 250 g of water freezes to ice, all at 0C?

Hf water = 334 J/g

Heating and Cooling Curve Calculations

How much energy will be consumed when 200 g of ice at -5.0C is converted to steam at 120.0C?

Specific Heats, J/gCIce: 2.09 Water: 4.18 Steam: 2.03

Evaporation and Boiling

What is really going on when a substance goes from the liquid to the gas phase or vice-versa?

In evaporation, some particles (atoms or molecules) have enough kinetic energy to overcome cohesive forces and escape from the surface of the liquid.

Evaporation and Boiling

What is really going on when a substance goes from the liquid to the gas phase or vice-versa?

In boiling, many particles have enough kinetic energy to enter the gas phase. Bubbles of gas form in the bulk liquid.

Evaporation and Boiling

Evaporation and Boiling

Evaporation and BoilingIf a liquid is put in a closed container, molec-

ules of the liquid will escape into the gas (or vapor) phase.

The amount of vapor will depend on:

TemperatureHigher T increases amount of gas

Cohesive forces between moleculesStronger forces decrease amount of gas

Evaporation and Boiling

Device for measuringthe vapor pressureof a liquid.

Evaporation and Boiling

Vapor Pressure of Water as a function of Temperature

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 10 20 30 40 50 60 70 80 90 100

Temperature, degrees C

Vapo

r Pre

ssur

e,

Atm

osph

eres

Evaporation and Boiling

A liquid boils when its vapor pressure equals the external pressure on the liquid.

The normal boiling point of a liquid is the temperature at which it boils under atmospheric pressure.

The boiling point of a liquid will increase or decrease with changes in applied pressure.

Evaporation and Boiling

Evaporation is conversion from liquid to vapor at temperatures below the boiling point.

Vapor is usually used for the gas phase of compounds that are liquid at room temper-ature and pressure.

Gas is the term used for compounds that are not liquid at room temperature and pressure.

Evaporation and Boiling

Evaporation and Boiling

Intermolecular ForcesIntermolecular forces are attractive forces

that act between molecules (also atoms and ions). There are three types:

London Dispersion Forces

Dipole – Dipole Interactions

Hydrogen Bonds

Intermolecular forces are hierarchical and additive.

Dipole-dipole forces are electrostatic forces that occur between polar molecules.

Hydrogen bonds are especially strong dipole-dipole forces that occur in molecules with these bonds: F-H O-H N-H

Hydrogen Bonds

London dispersion forces are weak, induced, temporary dipole-dipole interactions.

These are the only forces between nonpolar molecules.

They are strongest between large molecules and atoms.

London Dispersion Forces

Intermolecular ForcesLondon dispersion forces

1 – 10 kJ/mol

Dipole-dipole forces3 – 4 kJ/mol

Hydrogen bonds10 – 40 kJ/mol

Single covalent bonds150 – 550

kJ/mol

IM Forces and Boiling Points

Types of SolidsThere are two types of solids:

Crystalline solids are characterized by a regular three-dimensional arrangement of the atoms, molecules, or ions that are make up the substance. A crystal lattice is the regular arrangement of these particles.

Amorphous solids are characterized by a random, nonrepetitive three-dimensional arrangement of the atoms, molecules, or ions that make up the substance.

Types of Solids

Types of Crystalline solids:

Ionic MolecularNetwork Metallic

Types of Amorphous solids:

Molecular Network

Ionic solids are crystalline solids composed of ions. The ions are arranged to maximize interactions between unlike charges and minimize interactions between like charges.

Crystalline molecular solids are composed of molecules that are placed in a regular array to maximize intermolecular forces.

Crystalline network solids are crystalline solids in which the atoms are held in a regular array by covalent bonds.

Metallic solids are made of metal atoms, usually in a closely packed array. Elec-trons move freely among the atoms.

Amorphous molecular solids are composed of large molecules (polymers and plastics) that exist as random coils. Think spaghetti but longer! They will melt and dissolve.

Amorphous network solids are usually polym-ers and plastics that have been crosslinked to form covalent bonds between spaghetti strands. They will not melt or dissolve.