10.3 Phase Transitions
3 min read•june 25, 2024
Matter changes states through phase transitions. These shifts between solid, liquid, and gas are driven by temperature and pressure changes. Understanding these transitions helps explain everyday phenomena like boiling water or why ice cubes melt.
play a crucial role in phase transitions. Stronger forces require more energy to break, leading to higher and boiling points. This explains why water, with its strong hydrogen bonds, boils at a higher temperature than many other substances.
Phase Transitions
Phase transitions and key temperatures
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- Phase transitions change matter between solid, liquid, and gas phases
- converts solid to liquid by overcoming
- Key melting points at 1 atm pressure: water (0°C), ethanol (-114°C), gold (1064°C)
- converts liquid to gas when equals atmospheric pressure
- Boiling vaporizes throughout the liquid at the
- Key boiling points at 1 atm pressure: water (100°C), ethanol (78°C), gold (2856°C)
- converts solid directly to gas, skipping the liquid phase
- Occurs with substances such as dry ice (solid CO2), iodine, naphthalene
- converts gas to liquid by removing heat energy
- converts liquid to solid at the same temperature as the melting point for a given pressure
- converts gas directly to solid, the reverse of
- converts solid to liquid by overcoming
Intermolecular forces in phase transitions
- Intermolecular forces (IMFs) attract molecules to each other
- Stronger IMFs lead to higher temperatures needed for melting and boiling
- IMFs increase in strength from to to
- London dispersion forces (LDFs) exist between all molecules due to temporary dipoles, increasing with molecular size and surface area
- Dipole-dipole forces occur between polar molecules with permanent dipoles from uneven charge distribution
- Hydrogen bonding, the strongest IMF, occurs when H bonds to highly electronegative N, O, or F
- Substances with stronger IMFs require more energy to overcome attractions and change phase
- Water has strong hydrogen bonding, resulting in high melting (0°C) and boiling (100°C) points compared to similar-sized nonpolar methane (-182°C and -161°C)
Interpretation of heating and cooling curves
- Heating curves show temperature change as heat is added at a constant rate
- Plateaus represent phase transitions where heat is added but temperature remains constant
- Heat added during a equals the enthalpy (H) of that transition
- represents the for melting/
- represents the for boiling/condensing
- represents the
- Cooling curves show temperature change as heat is removed at a constant rate, the reverse of heating curves
- Calculate heat flow (q) and enthalpy changes (ΔH) using:
- for temperature changes (m = mass, c = specific heat capacity, ΔT = temperature change)
- or for phase transitions (m = mass, ΔH = enthalpy of fusion or )
- Enthalpy change (ΔH) equals the heat flow (q) at constant pressure
- The energy absorbed or released during a phase change without temperature change is called
Phase diagrams and critical points
- Phase diagrams graphically represent the relationship between temperature, pressure, and physical state of a substance
- Key features of a include:
- : where solid, liquid, and gas phases coexist in equilibrium
- : the highest temperature and pressure at which liquid and gas phases can coexist
- Vapor pressure curve: shows the pressure at which a liquid and its vapor are in equilibrium at a given temperature
Key Terms to Review (36)
Boiling point: The boiling point is the temperature at which a liquid's vapor pressure equals the external pressure surrounding the liquid. At this temperature, the liquid transitions to a gas phase.
Boiling Point: The boiling point is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid, and bubbles of vapor form inside the liquid. It is the point at which the liquid phase transitions to the gas phase.
Clausius-Clapeyron equation: The Clausius-Clapeyron equation describes the relationship between the vapor pressure and temperature of a substance. It is used to characterize phase transitions, such as from liquid to gas.
Condensation: Condensation is the phase transition from gas to liquid. It occurs when a vapor cools and loses enough thermal energy to change state.
Condensation: Condensation is the process by which a gas or vapor transitions into a liquid state. It is a fundamental phase transition that occurs when a substance cools and the kinetic energy of its molecules decreases, allowing them to form a more condensed liquid phase.
Cooling Curve: A cooling curve is a graphical representation that illustrates the temperature change of a substance as it transitions from a higher temperature to a lower temperature, typically as it cools and solidifies. This curve provides insight into phase transitions, showing how the substance behaves during melting, freezing, and other changes in state, which are crucial for understanding the physical properties of materials.
Critical point: The critical point is the end point of a phase equilibrium curve, where the properties of gas and liquid phases become indistinguishable. It represents the highest temperature and pressure at which a substance can exist as a liquid and gas in equilibrium.
Critical Point: The critical point is a unique point on a phase diagram where the distinct liquid and gas phases of a substance merge into a single, homogeneous supercritical fluid phase. At the critical point, the properties of the liquid and gas phases become indistinguishable, marking the end of the phase transition between the two states.
Deposition: Deposition is the phase transition in which a gas turns directly into a solid, bypassing the liquid state. This process releases energy and occurs under specific temperature and pressure conditions.
Deposition: Deposition is the process by which a substance is deposited or laid down, typically from a state of suspension or vapor, onto a surface or into a solution. This term is particularly relevant in the context of phase transitions and phase diagrams, as it describes the direct transition from a gaseous state to a solid state without an intermediate liquid phase.
Dipole-Dipole Forces: Dipole-dipole forces are a type of intermolecular force that occurs between polar molecules, where there is an unequal sharing of electrons resulting in a partial positive and partial negative charge on opposite ends of the molecule. These attractive forces arise from the electrostatic interactions between the partially charged regions of neighboring polar molecules.
Enthalpy of Fusion: Enthalpy of fusion, also known as latent heat of fusion, is the amount of energy (heat) required to transform a substance from a solid state to a liquid state at its melting point, without changing the temperature of the substance. It represents the energy required to overcome the intermolecular forces that hold the solid structure together, allowing the transition to a more disordered liquid phase.
Enthalpy of Sublimation: The enthalpy of sublimation is the amount of energy required to transform a substance directly from the solid phase to the gas phase, without passing through the liquid phase. It represents the energy change associated with the sublimation process, which is the direct transition from solid to vapor state.
Enthalpy of Vaporization: The enthalpy of vaporization is the amount of energy required to transform a substance from a liquid state to a gaseous state at a constant temperature and pressure. It represents the energy needed to overcome the intermolecular forces that hold the molecules together in the liquid phase, allowing them to transition into the gas phase.
Freezing: Freezing is the phase transition from the liquid state to the solid state. It occurs when a liquid's temperature is lowered below its freezing point.
Freezing: Freezing is the physical process by which a liquid transitions into a solid state due to the removal of thermal energy. It is a crucial phase transition that is central to understanding the behavior of substances in various contexts, particularly in the study of phase diagrams.
Freezing point: The freezing point is the temperature at which a liquid turns into a solid. It is the same temperature as the melting point for a given substance.
Heating Curve: A heating curve is a graphical representation that depicts the changes in the physical state of a substance as it is heated or cooled. It illustrates the relationship between temperature and the phase transitions that occur during the heating or cooling process.
Hydrogen Bonding: Hydrogen bonding is a type of dipole-dipole intermolecular force that occurs when a hydrogen atom covalently bonded to a highly electronegative element, such as nitrogen, oxygen, or fluorine, experiences an attractive force with another nearby highly electronegative element. This attractive force is significantly stronger than a typical dipole-dipole interaction and has a significant impact on the physical and chemical properties of various compounds.
Intermolecular forces: Intermolecular forces are the forces of attraction and repulsion between molecules that influence the physical properties of substances. These forces are weaker than intramolecular forces, which hold atoms together within a molecule.
Intermolecular Forces: Intermolecular forces are the attractive or repulsive forces that exist between molecules, as opposed to the intramolecular forces that hold atoms together within a molecule. These forces play a crucial role in determining the physical properties and behavior of substances across various topics in chemistry, including non-ideal gas behavior, the properties of liquids, phase transitions, and the dissolution process.
Latent Heat: Latent heat is the amount of energy released or absorbed by a substance during a phase change, such as the transition from a solid to a liquid or from a liquid to a gas, without a change in temperature. It is a crucial concept in understanding the behavior of matter and the energy transformations that occur during phase transitions.
London Dispersion Forces: London dispersion forces are a type of intermolecular force that arises from the temporary, spontaneous polarization of atoms or molecules. These forces are the weakest of the intermolecular forces, but they play a crucial role in the properties and behavior of many substances, including liquids, gases, and the noble gases.
Melting: Melting is the phase transition from a solid to a liquid. It occurs when a substance absorbs sufficient heat to overcome its molecular forces.
Melting: Melting is the phase transition that occurs when a solid substance is heated, causing it to transform from a rigid, crystalline state into a liquid state. This change in phase is driven by the increase in kinetic energy of the atoms or molecules within the material, overcoming the intermolecular forces that hold the solid structure together.
Normal boiling point: The normal boiling point is the temperature at which a liquid boils under standard atmospheric pressure (1 atm or 101.3 kPa). At this temperature, the vapor pressure of the liquid equals the surrounding atmospheric pressure.
Phase diagram: A phase diagram is a graphical representation that shows the conditions of temperature and pressure under which distinct phases (solid, liquid, gas) of a substance exist. It illustrates the equilibrium between different states of matter.
Phase Diagram: A phase diagram is a graphical representation that shows the relationship between the physical states or phases of a substance, such as solid, liquid, and gas, as a function of variables like temperature and pressure. It provides a comprehensive overview of the conditions under which a substance can exist in different phases and the boundaries between these phases.
Phase Transition: A phase transition is a transformation of a substance from one physical state or phase to another, such as the transition from a solid to a liquid or from a liquid to a gas. These changes in phase are driven by changes in temperature, pressure, or other external conditions and involve the rearrangement of the molecular structure of the substance.
Sublimation: Sublimation is the phase transition in which a substance changes directly from a solid to a gas without passing through the liquid state. This process occurs under specific conditions of temperature and pressure.
Sublimation: Sublimation is the direct transition of a substance from the solid phase to the gas phase without passing through the liquid phase. This phase change occurs when the vapor pressure of the solid is equal to the surrounding pressure, allowing the solid to transform directly into a gas.
Triple point: The triple point is the unique set of conditions at which all three phases (solid, liquid, and gas) of a substance coexist in thermodynamic equilibrium. It is specific to each substance and occurs at a precise temperature and pressure.
Triple Point: The triple point is the unique temperature and pressure at which the solid, liquid, and gaseous phases of a substance can coexist in equilibrium. It is a critical point on a substance's phase diagram, where the three phases converge.
Vapor pressure: Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid phase at a given temperature. It indicates how volatile a substance is.
Vaporization: Vaporization is the phase transition from the liquid phase to the gas phase. It occurs when molecules in a liquid gain enough energy to overcome intermolecular forces and enter the vapor phase.
Vaporization: Vaporization is the process by which a substance in the liquid or solid state transitions to the gaseous state. It is a fundamental phase transition that occurs when the vapor pressure of a substance exceeds the surrounding atmospheric pressure, allowing the molecules to escape the liquid or solid and enter the gas phase.