10.5 Changes of State

Overview of States of Matter

• In the study of chemistry, specifically in CHM130 Fundamental Chemistry, states of matter are a primary focus.

• While many different states of matter exist, three are considered the most visible and prevalent in this course:

  • Solids

  • Liquids

  • Gases

The Six Major Phase Changes

• Phase changes are the transitions between the physical states of matter. There are six primary phase changes identified:

  • Melting (fusion): The transition from a Solid to a Liquid.

  • Freezing (crystallization): The transition from a Liquid to a Solid.

  • Vaporization: The transition from a Liquid to a Gas.

  • Condensation: The transition from a Gas to a Liquid.

  • Sublimation: The transition from a Solid to a Gas.

  • Deposition: The transition from a Gas to a Solid.

Scientific Principles of Phase Changes

• The Role of Heat: One of the primary factors causing a phase change is either the addition or removal of heat.

• Attractive Forces and Properties: The differences observed in the properties of solids, liquids, and gases result from the strengths of the attractive forces between the atoms, molecules, or ions within each phase.

• Intermolecular Forces (IMFs):

  • These are various forces of attraction that exist between the atoms and molecules of a substance.

  • They are caused by electrostatic phenomena.

  • IMFs serve to hold particles close together.

• Kinetic Energy (KE):

  • The kinetic energy of molecules provides the energy necessary to overcome attractive forces.

  • High KE allows for an increase in the distance between particles.

• Determinants of Substance Phase: The specific phase of a substance is dependent on the balance between its IMFs and the kinetic energies of its molecules.

• Temperature Effects: Changing the temperature of a substance changes its average KE, which can induce a change in its physical state.

Dynamics of Melting and Freezing

• Heating Crystalline Solids: When heat is applied to a crystalline solid, the average energy of its constituent atoms, molecules, or ions increases, and the solid's temperature rises.

• The Process of Melting:

  • At a certain point, the energy added becomes sufficient to partially overcome the forces keeping the molecules or ions in fixed positions.

  • At this stage, the solid begins to transition to a liquid (melting).

• Temperature Plateaus:

  • During the melting process, the temperature of the solid stops rising even though heat is continually being added.

  • The temperature remains constant until every part of the solid has transitioned into a liquid.

  • Further heating will only increase the temperature of the liquid after the melting process is entirely complete.

Dynamics of Vaporization and Condensation

• Influence of Chemical Identity: The types and strengths of intermolecular attractions possible in a liquid are determined by the chemical identities of the molecules.

• IMFs and Boiling Points:

  • Liquids with stronger attractive forces between molecules require more energy to overcome those forces.

  • Stronger IMFs result in a higher boiling point because a higher temperature is needed to vaporize the liquid.

  • Liquids with weak intermolecular attractions allow molecules to vaporize more easily, leading to a lower boiling point.

Specialized Phase Changes: Sublimation and Deposition

• Sublimation: This is the process where a solid transitions directly into a gaseous state, completely bypassing the liquid state.

  • Example 1 (Dry Ice): A piece of dry ice (solid $CO_2$) sublimes at room temperature and standard pressure, appearing to disappear without forming liquid.

  • Example 2 (Snow and Ice): Snow and ice can sublime at temperatures below the melting point of water. This is usually a slow process but can be accelerated by winds and the reduced atmospheric pressures found at high altitudes.

  • Example 3 (Iodine): When solid iodine is warmed, it sublimes and creates a vivid purple vapor.

• Deposition: This is the reverse of sublimation; a process where gaseous substances condense directly into a solid state, bypassing the liquid state.

  • Example: The formation of frost is a primary example of deposition.

Detailed Physical Characteristics of the Three States

• Solids:

  • Characterized by very strong attractive forces that hold particles tightly together.

  • Particles do not move freely but vibrate in a fixed position.

  • These fixed positions provide solids with a definite shape and a definite volume.

• Liquids:

  • Particles possess enough energy to move in random directions.

  • Particles remain close together with strong attractions.

  • These strong attractions provide liquids with a definite volume.

• Gases:

  • Molecules possess a significant amount of energy and move around freely.

  • The forces between gas molecules are not very strong.

  • Gases will fill the entirety of whatever container they are placed in.

Questions & Discussion

• Practice Question 1: What is the relationship between the intermolecular forces in a liquid and its boiling point?

  • Concept Context: Stronger IMFs require more energy to break, thus increasing the boiling point.

• Practice Question 2: Why does spilled gasoline evaporate more rapidly on a hot day than on a cold day?

  • Concept Context: Heat increases the kinetic energy of the molecules, allowing them to overcome IMFs and escape into the gas phase more quickly.

• Practice Question 3: The molecular weights of $CH_4$ and $NH_3$ are nearly the same. Account for the fact that the melting and boiling points of $NH_3$ ($-77.7\,^{\circ}C$ and $-33.3\,^{\circ}C$) are higher than those of $CH_4$ ($-184\,^{\circ}C$ and $-161\,^{\circ}C$).

  • Concept Context: Even with similar weights, the chemical identities lead to different strengths in intermolecular forces (e.g., hydrogen bonding in $NH_3$ vs. London dispersion forces in $CH_4$).