Phases of matter

Phases of Matter

Matter exists in different physical forms called phases (or states of matter). The phase of a substance is determined by the balance between two competing factors:

• The kinetic energy of its particles (their energy of motion)

• The intermolecular forces holding those particles together

When kinetic energy dominates, particles move more freely and the substance exists in a less organized phase. When intermolecular forces dominate, particles are held more tightly together and the substance exists in a more organized phase. Four phases of matter are formally recognized, and a fifth phase has been produced under specialized laboratory conditions.

Figure 1: The Phases of Matter

Explore the phases of matter—solid, liquid, gas, plasma, and Bose-Einstein condensate—illustrated by their particle arrangements, properties, and real-world examples, highlighting how temperature and pressure influence state changes.

The Four Main Phases of Matter

The four commonly recognized phases of matter are:

1. Solid

2. Liquid

3. Gas

4. Plasma

Each phase exhibits unique physical properties because of differences in particle arrangement, particle motion, and intermolecular forces.

Solid

In a solid, particles (atoms, molecules, or ions) are packed closely together in a highly ordered arrangement.

The intermolecular forces are strong relative to the particles' kinetic energy. As a result, particles can vibrate about fixed positions but cannot move freely from place to place.

Because of this restricted motion, solids possess several characteristic properties:

• Definite shape

• Definite volume

• High density

• Very low compressibility

Examples include:

• Ice (H₂O(s))

• Sodium chloride (NaCl)

• Iron (Fe)

• Diamond (C)

• Quartz (SiO₂)

Figure 2: Particle Arrangement in a Solid

Illustrating the solid state, particles are depicted in a fixed, regular lattice structure where they remain closely packed and vibrate minimally, emphasizing the solid's definite shape and low compressibility.

Crystalline Solids

In a crystalline solid, particles are arranged in a repeating, long-range ordered structure called a:

Crystal Lattice

Examples include:

• Sodium chloride

• Diamond

• Quartz

Crystalline solids possess:

• Highly ordered structures

• Sharp melting points

• Distinct crystal shapes

Amorphous Solids

In an amorphous solid, particles lack long-range order.

Examples include:

• Glass

• Rubber

• Many plastics

Unlike crystalline solids, amorphous solids:

• Do not possess repeating crystal structures

• Soften gradually over a range of temperatures

• Do not exhibit sharp melting points

Figure 3: Crystalline versus Amorphous Solids

Exploring the differences between crystalline and amorphous solids, this diagram highlights their particle arrangements, key characteristics, and examples, emphasizing how structure influences properties and uses.

Liquid

In a liquid, particles remain close together but are no longer fixed in position.

The particles possess enough kinetic energy to move past one another while remaining in close contact.

This mobility gives liquids their ability to flow.

Liquids have:

• No definite shape

• Definite volume

• Moderate density

• Very low compressibility

Liquids conform to the shape of their container while maintaining essentially constant volume.

Examples include:

• Water (H₂O(l))

• Mercury (Hg(l))

• Ethanol (C₂H₅OH)

• Molten iron

Figure 4: Particle Arrangement in a Liquid

Illustrating the fluid nature of liquids, this diagram shows how particles are closely packed yet move freely, allowing them to flow and adapt to the shape of their container.

Surface Tension

Particles at the surface of a liquid experience unequal intermolecular attractions.

This creates a phenomenon known as:

Surface Tension

Surface tension causes liquids to minimize surface area and explains behaviors such as:

• Water droplets forming spheres

• Insects walking on water

• Capillary action

Figure 5: Surface Tension

A water strider effortlessly skims the surface of a pond, demonstrating its unique ability to walk on water due to surface tension.

Viscosity

Viscosity is a measure of a liquid's resistance to flow.

Examples:

• Honey has high viscosity.

• Water has low viscosity.

Viscosity depends on:

• Intermolecular forces

• Molecular size

• Temperature

Figure 6: Viscosity Comparison

A visual comparison of viscosity: honey exhibits high viscosity, flowing slowly, while water demonstrates low viscosity, flowing quickly.

Gas

In a gas, particles possess enough kinetic energy to overcome intermolecular attractions almost completely.

Gas particles:

• Move rapidly

• Move randomly

• Are widely separated

• Interact primarily through collisions

As a result, gases exhibit very different properties from solids and liquids.

Gases have:

• No definite shape

• No definite volume

• Very low density

• High compressibility

A gas expands to fill any container available to it.

Examples include:

• Water vapor (H₂O(g))

• Oxygen (O₂)

• Nitrogen (N₂)

• Carbon dioxide (CO₂)

Figure 7: Particle Arrangement in a Gas

Illustration of gas particles moving rapidly and randomly, highlighting their dispersed arrangement and tendency to fill their container completely.

Plasma

Plasma is often called the:

Fourth State of Matter

Plasma forms when a gas is heated to extremely high temperatures or subjected to intense electromagnetic fields.

Under these conditions, particles gain enough energy to remove electrons from atoms.

The result is a mixture of:

• Positive ions

• Free electrons

Because plasma contains charged particles, it behaves very differently from ordinary gases.

Plasma:

• Conducts electricity

• Responds strongly to magnetic fields

• Responds strongly to electric fields

Examples include:

• The Sun

• Stars

• Lightning

• Fluorescent lights

• Plasma globes

Figure 8: Plasma

Plasma, the fourth state of matter, consists of ionized gas with free electrons and ions, exemplified by natural phenomena like lightning and the Sun.

Plasma in the Universe

Although plasma is relatively uncommon in everyday life, it is actually the most abundant form of visible matter in the universe.

Nearly all stars consist primarily of plasma.

This means most visible matter in the universe exists in the plasma state.

Figure 9: Plasma in Stars

The Sun, a massive ball of plasma, consists of ionized gases with charged particles like electrons and ions. Plasma is the universe's most abundant form of visible matter, making up over 99% of known ordinary matter.

The Fifth State: Bose–Einstein Condensate

A fifth state of matter known as the:

Bose–Einstein Condensate (BEC)

was first produced experimentally in 1995 by:

• Eric Cornell

• Carl Wieman

using rubidium atoms cooled to temperatures extremely close to absolute zero.

The temperature achieved was approximately:

170 nanokelvin

which is only:

0.000000170 K

above absolute zero.

At such extreme temperatures:

• Atomic motion nearly ceases.

• Atoms lose individual identities.

• Large numbers of atoms occupy the same quantum state.

The result is a collection of atoms behaving as a single quantum entity.

Producing a Bose–Einstein condensate requires:

• Laser cooling

• Magnetic trapping

• Specialized laboratory equipment

Cornell and Wieman received the:

2001 Nobel Prize in Physics

for this achievement.

Figure 10: Bose–Einstein Condensate

A Bose-Einstein Condensate (BEC) forms when atoms are cooled to nearly absolute zero, allowing them to occupy the same quantum state and behave as a single "super-atom."

Comparing the Phases

This table compares the characteristics of the four states of matter: solid, liquid, gas, and plasma, highlighting their properties such as shape, volume, density, compressibility, particle motion, and intermolecular forces.

Figure 11: Comparison of Phases of Matter

Explore the distinct characteristics of matter phases: solids with definite shape, liquids with fluid volume, gases with free movement, and plasma with high-energy ionization.

Phase Changes

A phase change occurs when a substance transitions from one phase to another.

Phase changes are:

Physical Changes

because the chemical identity of the substance remains unchanged.

Only intermolecular forces are broken or formed.

The molecules themselves remain intact.

Phase changes are driven by changes in:

• Temperature

• Pressure

Figure 12: Phase Changes Among States

Explore the six phase changes that occur between solid, liquid, and gas states: melting, freezing, vaporization, condensation, sublimation, and deposition.

Melting (Fusion)

Melting occurs when:

Solid → Liquid

Melting is an:

Endothermic Process

because energy must be absorbed to overcome intermolecular forces.

Example:

Ice melting into liquid water.

Freezing (Solidification)

Freezing occurs when:

Liquid → Solid

Freezing is an:

Exothermic Process

because intermolecular forces form and release energy.

Example:

Liquid water freezing into ice.

Vaporization

Vaporization occurs when:

Liquid → Gas

Vaporization is:

Endothermic

because particles must absorb energy to escape intermolecular attractions.

Example:

Boiling water.

Condensation

Condensation occurs when:

Gas → Liquid

Condensation is:

Exothermic

because intermolecular attractions form.

Example:

Formation of rain droplets from water vapor.

Sublimation

Sublimation occurs when:

Solid → Gas

without passing through the liquid phase.

Sublimation is:

Endothermic

Example:

Dry ice (solid CO₂) becoming carbon dioxide gas.

Deposition

Deposition occurs when:

Gas → Solid

without passing through the liquid phase.

Deposition is:

Exothermic

Example:

Formation of frost on a cold surface.

Summary of Phase Changes

This table illustrates the energy changes associated with different phase transitions of matter, including examples of each process.

Phase Diagrams

A phase diagram is a graph that shows which phase of a substance is stable under various conditions of:

• Temperature

• Pressure

Temperature is plotted on the x-axis.

Pressure is plotted on the y-axis.

Different regions of the diagram correspond to different phases.

Figure 13: Generic Phase Diagram

A phase diagram illustrating the pressure-temperature relationship for a pure substance, highlighting regions of solid, liquid, gas, and supercritical fluid, with key points such as the triple point and critical point.

Key Features of a Phase Diagram

Solid Region

Stable at:

• Low temperature

• High pressure

Liquid Region

Stable at:

• Intermediate temperatures

• Intermediate pressures

Gas Region

Stable at:

• High temperature

• Low pressure

Triple Point

The unique temperature and pressure at which:

• Solid

• Liquid

• Gas

all coexist in equilibrium.

Critical Point

The point beyond which liquid and gas become indistinguishable.

Above this point, the substance becomes a:

Supercritical Fluid

Supercritical Fluids

A supercritical fluid exists beyond the critical point.

Supercritical fluids possess properties intermediate between:

• Liquids

• Gases

For example, supercritical carbon dioxide:

• Diffuses like a gas

• Dissolves substances like a liquid

Industrial applications include:

• Decaffeination of coffee

• Chemical extraction

• Dry cleaning

Summary

Matter exists in several distinct phases determined by the balance between particle kinetic energy and intermolecular forces. Solids possess definite shape and volume, liquids have definite volume but no fixed shape, gases have neither definite shape nor volume, and plasmas consist of ionized particles that respond strongly to electromagnetic fields. Under extremely low-temperature conditions, matter can also exist as a Bose–Einstein condensate.

Phase changes are physical changes driven by temperature and pressure changes. Phase diagrams provide a graphical representation of phase stability and identify important features such as the triple point, critical point, and supercritical fluid region.

Key Points

• Matter exists in multiple phases.

• Solids have definite shape and volume.

• Liquids have definite volume but no definite shape.

• Gases have neither definite shape nor definite volume.

• Plasma consists of ions and free electrons.

• Bose–Einstein condensates occur near absolute zero.

• Phase changes are physical changes.

• Melting and vaporization are endothermic.

• Freezing and condensation are exothermic.

• Sublimation is solid-to-gas conversion.

• Phase diagrams map phase stability as a function of temperature and pressure.

• Water has an unusual negatively sloped solid-liquid boundary.

• Supercritical fluids possess properties of both liquids and gases.