Transcript States of Matter
Chapter 11.
States of Matter
States of Matter
State is Determined by: Chemical Identity Temperature Pressure
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 States Solids
Definite Volume, Definite Shape Density is High: 1.0 – 20 g/cm 3 Compressibility is Low Thermal Expansion is Low: 0.01% per C
The Solid State
In 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 States Liquids
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 State
In 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 States Gases
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 State
In 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 State
Most 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 CO 2
Changes of State
Exothermic: Endothermic
Release Heat Energy
H is negative Absorb Heat Energy
H is positive
Changes of State Freezing point
the temperature at which the liquid and solid phases of a substance are in equilibrium
Boiling point Equilibrium
the temperature at which the liquid and vapor phases of a substance are in 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: Heat Light Electrical Mechanical
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 kg m
2
/sec
2
4.18 J = 1 calorie 1 calorie is the amount of heat energy required to raise the temperature of 1 g of water by 1 C.
Specific Heat Specific heat
is the amount of heat energy required to raise the temperature of 1.00 gram of a substance by 1.00
C.
It takes 4.18 J to raise the temperature of 1.00 g of water by 1.00
C.
4.18 J/g C is the specific heat of water.
Specific Heat
4.18 J/g C very high!
Metals have specific heats of 0.1 to 1.0 J/g C; They conduct heat.
Most nonmetallic materials are insulators, with specific heats of 1 to 2 J/g C.
Specific Heat
How much heat is absorbed if I heat 100 g of water from 25.0
C (room temperature) to 100.0
C (boiling point of water)?
How much heat is given off if I cool 1.00 lb (454 g) of iron metal from 100.0
C to 25.0
C? Specific heat of iron is 0.444 J/g C.
Specific Heat
What is the specific heat of a rock? Its mass is 125 g. I heat it to 100.0
C in a boiling water bath, then drop it into 100.0 g of water that's at 20.0
C. The water temperature rises to 30.0
C.
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 substance its particles gain kinetic energy forces between particles are overcome The substance goes into a less ordered state.
It melts, boils, or sublimes.
Energy and Changes of State
In an
exothermic phase change
: heat energy is given off by a substance its particles lose kinetic energy forces between particles can act The substance goes into a more ordered state.
It freezes, condenses, or deposits.
Energy and Changes of State
During 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 State
During 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
, H
f
= heat to convert one gram of substance from solid to liquid at its melting point. Units: J/g
Heat of vaporization
, H
v
= 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 100 C is boiled to steam, also at 100 C?
H
v water
= 2260 J/g How much energy will be released if 250 g of water freezes to ice, all at 0 C?
H
f water
= 334 J/g
Heating and Cooling Curve Calculations
How much energy will be consumed when 200 g of ice at -5.0
C is converted to steam at 120.0
C?
Specific Heats, J/g C Ice: 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 Boiling
If 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: Temperature Higher T increases amount of gas Cohesive forces between molecules Stronger forces decrease amount of gas
Evaporation and Boiling
Device for measuring the vapor pressure of a liquid.
Evaporation and Boiling Vapor Pressure of Water as a function of Temperature 1 0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0 0 10 20 30 40 50 60 70 80 90 Temperature, degrees C 100
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 Forces Intermolecular 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 Forces
London dispersion forces 1 – 10 kJ/mol Dipole-dipole forces 3 – 4 kJ/mol Hydrogen bonds 10 – 40 kJ/mol Single covalent bonds 150 – 550 kJ/mol
IM Forces and Boiling Points
Types of Solids
There 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 Network Molecular 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.