KEY IDEA: Temperature and pressure induce changes in the state of substances.
- The physical state of matter - solid, liquid, gas - is an everyday example of the organization of subatomic particles.
- Changes in state of matter (freezing, boiling, condensation, etc.) involves a rearrangement of the particles.
Intro to Kinetic Theory
The physical state of a substance - solid, liquid, gas - is a common example of the submicroscopic organization of atoms and molecules.
Solid, liquids and gases are usually described in terms of their shape and volume:
Solid, liquids and gases are usually described in terms of their shape and volume:
- Solids mostly have a definite shape and definite volume. (A solid is its own container.)
- Liquids have an indefinite shape and a definite volume. (Liquids conform to the shape of their container.)
- A gas has indefinite shape and indefinite volume. (A gas conforms to the shape and volume of its container.)
Kinetic-molecular theory describes the motion of the particles, and attempts to explain the behavior of states of matter. Most of the ideas from the kinetic theory came from a study of gases.
According to kinetic theory, all matter consists of tiny particles that are in continual motion, and that the temperature of an object is a measure of the intensity of this motion. The physical state of a substance depends upon the kinetic energy of its particles.
According to kinetic theory, all matter consists of tiny particles that are in continual motion, and that the temperature of an object is a measure of the intensity of this motion. The physical state of a substance depends upon the kinetic energy of its particles.
Gas particles have very high kinetic energy; particles of a solid have very low kinetic energy:
- Particles of a solid are tightly packed, arranged in an orderly pattern, and vibrate in a fixed location.
- Particles of a liquid are close together, disorderly arranged, and able to flow past one another (fluidity).
- Under normal conditions, gas particles are spaced very far apart from each other, and spread out to occupy all available volume.
Daniel Bernoulli, Rudolf Clausius, James Clerk Maxwell, and Ludwig Boltzmann were among scientists contributing ideas to kinetic theory.
PLASMA
- Super-heated gas
- Electrically charged
2 more states of matter:
- Bose-Einstein condensate consists of a gas cooled near absolute zero; atoms lose their individual identity and begin to coalesce into a "super-atom." Quantum effects are observed on near-macroscopic scale.
- When a gas is subjected to extremely high temperatures, the collisions become so energetic that some electrons are stripped off. The atoms and molecules become ionized, and the resulting state of matter is called a plasma. Stars are made of plasma.
Since the gas state is the "simplest" state of matter, it's the easiest to understand and explain theoretically.
Behavior of gases
Kinetic theory models gas particles as point masses (no volume) in constant, rapid, random linear motion, moving independent of each other. Collisions between particles are perfectly elastic (no net loss in kinetic energy). The avg. kinetic energy of the particles is directly proportional to its temperature.
Microscopic behaviors of the particles reflect macroscopic properties of gases:
gas pressure
Pressure is defined as the force per unit area. Moving bodies exert a pressure when they collide. Gas pressure results from the simultaneous collisions of billions of particles with a surface.
Air exerts a pressure on Earth because gravity holds particles in the atmosphere. Atmospheric pressure can be approximated as the weight of a column of air extending above the surface:
On average, one square inch of atmospheric air 62 miles high has a weight of about 14.7 pounds (14.7 pounds per square inch, or "psi"). The standard atmosphere (atm) is the unit of this pressure.
Atmospheric pressure is measured using a barometer: the height of mercury in a tube reflects the atmospheric pressure exerted by the particles in air.
At sea-level, the pressure is sufficient to support a mercury column about 760 mm high. (760 mm Hg) The SI unit of pressure, the Pascal (Pa), is defined as 1 Newton per square meter, a very small amount of pressure. One "standard" atmosphere is equivalent to 101,325 Pa, or 101.325 kilopascals (kPa). 1.00 atm = 760. mm Hg = 101.325 kPa
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kinetic theory and temperature
The particles in any collection of gas have a wide range of kinetic energies; so consider an "average" kinetic energy. At any given temperature, all particles of a substance have the same average kinetic energy.
An increase in kinetic energy results in an increase in temperature. The red and blue curves on the graph below show the kinetic energy of a typical collection of molecules at two different temperatures:
An increase in kinetic energy results in an increase in temperature. The red and blue curves on the graph below show the kinetic energy of a typical collection of molecules at two different temperatures:
phase diagrams
A phase diagram is a chart showing the relationship between the effects of pressure and temperature on the phase (solid, liquid, gas) of a substance. The figure is an example of a phase diagram:
NOTE: The boundary lines between phases indicate the conditions of temperature and pressure under which equilibrium of different phases of a substance can exist; i.e. BOTH phases exist on these lines.
Melting/Freezing
Vaporization/Condensation
Sublimation/Deposition
Critical Point
Triple Point
Below is the phase diagram for water (left) and carbon dioxide (right):
Melting/Freezing
- The line between the solid and liquid phases (Line AD) is the equilibrium of solid and liquid phases at that specific pressure and temperature, i.e. a curve of all the freezing/melting points.
- Any point on this line the substance is both solid and liquid.
- Usually, melting point is only slightly affected by pressure. For this reason, the melting point line is nearly vertical.
- “Normal” melting/freezing point – the temperature at which the solid changes to a liquid at STP.
Vaporization/Condensation
- The line between the liquid and gas phases (Line AB) is the equilibrium of liquid and gas phases at that specific pressure and temperature, i.e. a curve of all the vaporization/condensation points.
- Any point on this line the substance is both liquid and gas.
- “Normal” boiling/condensation point – the temperature at which the liquid changes to a gas at STP.
Sublimation/Deposition
- The line between the solid and gas phases (Line CA) is the equilibrium of solid and gas phases at that specific pressure and temperature, i.e. a curve of all the deposition/sublimation points.
- Any point on this line the substance is both solid and gas.
Critical Point
- The vapor pressure curve ends at the critical point, the temperature above which the gas cannot be liquefied no matter how much pressure is applied (the kinetic energy simply is too great for attractive forces to overcome). Any substance beyond its critical point is called a super-critical fluid. Above this pressure and temperature, the gas and liquid phases are indistinguishable.
Triple Point
- The triple point is the condition of temperature and pressure where ALL THREE phases exist in equilibrium (solid, liquid, gas).
Below is the phase diagram for water (left) and carbon dioxide (right):
Note how the line solid-liquid phase boundaries are leaning in opposite directions:
For most substances, the solid form is more dense than the liquid. However, for water, solid ice is less dense than liquid water. This is quite unusual, especially for a substance as "small" as water. If this were not the case, and solid ice was more dense than liquid water, lakes and ponds would freeze from the bottom-up, and life as we know it could not have evolved.
Some substances exist in various phases (called "allotropes") besides just solid, liquid, and gas. Allotropes are pure forms of the same element that differ in crystal structure. The different structural forms give rise to different physical and chemical properties.
- If a liquid is more dense than its solid, the melting point line leans to the left, causing the melting point to decrease with pressure.
- If a liquid is less dense than its solid, the melting point line leans to the right, causing the melting point to increase with temperature.
For most substances, the solid form is more dense than the liquid. However, for water, solid ice is less dense than liquid water. This is quite unusual, especially for a substance as "small" as water. If this were not the case, and solid ice was more dense than liquid water, lakes and ponds would freeze from the bottom-up, and life as we know it could not have evolved.
Some substances exist in various phases (called "allotropes") besides just solid, liquid, and gas. Allotropes are pure forms of the same element that differ in crystal structure. The different structural forms give rise to different physical and chemical properties.