9.3 Activity and activity coefficients
3 min read•august 6, 2024
and activity coefficients are key concepts in thermodynamics, helping us understand non-ideal behavior in mixtures. They quantify how components interact, affecting vapor pressure and equilibrium in solutions.
Standard states provide reference points for thermodynamic properties, while Raoult's and Henry's laws describe behavior. Understanding these concepts is crucial for analyzing real-world mixtures and their deviations from ideality.
Activity and Standard States
Activity and Activity Coefficients
Top images from around the web for Activity and Activity Coefficients
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
Non-Ideal Gas Behavior | Chemistry View original
Is this image relevant?
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
Non-Ideal Gas Behavior | Chemistry View original
Is this image relevant?
1 of 3
Top images from around the web for Activity and Activity Coefficients
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
Non-Ideal Gas Behavior | Chemistry View original
Is this image relevant?
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
8.6: Non-Ideal Gas Behavior | General College Chemistry I View original
Is this image relevant?
Non-Ideal Gas Behavior | Chemistry View original
Is this image relevant?
1 of 3
- Activity quantifies the effective concentration of a species in a mixture
- Accounts for non-ideal behavior and interactions between species
- relates activity to mole fraction:
- For ideal solutions, and activity equals mole fraction
- Activity coefficients depend on composition, temperature, and pressure
Standard States and Conventions
- Standard state is a reference condition for expressing thermodynamic properties
- For pure substances, standard state is the pure component at the system temperature and pressure
- For mixtures, standard state depends on the convention chosen
- Symmetrical convention defines standard state as the pure component at system temperature and pressure for all components
- Unsymmetrical convention defines standard state differently for solvent (pure liquid) and solutes (hypothetical ideal solution at unit molality)
Ideal Solution Laws
Raoult's Law for Ideal Solutions
- describes vapor-liquid equilibrium for ideal solutions
- Vapor pressure of each component is proportional to its mole fraction in the liquid phase:
- is the vapor pressure of pure component at the system temperature
- Applies to components that are liquid at the system temperature (typically the solvent)
- Examples: Benzene-toluene, ethanol-water (at low concentrations)
Henry's Law for Dilute Solutions
- describes vapor-liquid equilibrium for dilute solutions
- Partial pressure of a solute is proportional to its mole fraction in the liquid phase:
- is the Henry's law constant, which depends on temperature and solute-solvent pair
- Applies to components that are gases at the system temperature (typically solutes)
- Examples: Oxygen in water, carbon dioxide in beverages
Non-Ideal Solutions
Excess Gibbs Energy and Non-Ideality
- Non-ideal solutions deviate from ideal solution laws due to interactions between species
- quantifies the deviation from ideal behavior:
- is the actual Gibbs energy of the mixture, and is the Gibbs energy for an ideal mixture
- Activity coefficients relate excess Gibbs energy to composition:
- Positive indicates from ideality (repulsive interactions), while negative indicates (attractive interactions)
Examples of Non-Ideal Solutions
- Ethanol-water exhibits positive deviations from ideality at high concentrations due to hydrogen bonding
- Chloroform-acetone shows negative deviations from ideality due to strong intermolecular attractions
- are non-ideal mixtures with a composition where the vapor and liquid phases have the same composition at equilibrium
- Examples of azeotropes: Ethanol-water (95.6% ethanol), nitric acid-water (68% nitric acid)
- Non-ideal behavior is important in separation processes, as it affects vapor-liquid equilibrium and can limit separation efficiency
Key Terms to Review (20)
Activity: Activity is a measure of the effective concentration of a species in a mixture, which reflects its behavior compared to an ideal solution. It plays a crucial role in understanding how substances interact in real solutions, taking into account non-idealities like intermolecular forces. This concept is tied closely to activity coefficients, which quantify how much a species deviates from ideal behavior, and is foundational in laws that govern vapor-liquid equilibria.
Activity coefficient: The activity coefficient is a factor used in thermodynamics to account for the deviation of a substance's behavior from that of an ideal solution. It quantifies how interactions between particles in a mixture affect the effective concentration of a component, thereby influencing properties like vapor pressure and chemical potential. Understanding this concept is crucial for analyzing non-ideal solutions and mixtures, especially in determining phase equilibria and excess properties.
Azeotropes: Azeotropes are mixtures of two or more liquids that exhibit the same composition in both the liquid and vapor phases at a given temperature and pressure, leading to unique boiling behaviors. This characteristic means that when an azeotropic mixture is boiled, the vapor has the same composition as the liquid, making it impossible to separate the components by simple distillation. Understanding azeotropes is crucial for processes like distillation and is deeply connected to the concepts of activity and activity coefficients.
Chemical potential: Chemical potential is the change in the Gibbs free energy of a system when an infinitesimal amount of substance is added, reflecting how the potential energy of a chemical species can influence its behavior in a mixture. It is crucial for understanding how substances interact, particularly when it comes to phase changes, equilibrium, and the thermodynamic properties of mixtures.
Dilution: Dilution refers to the process of reducing the concentration of a solute in a solution, typically by adding more solvent. This concept is important as it affects the activity and behavior of solutes in various chemical and thermodynamic systems. Understanding dilution helps in calculating how changes in concentration can influence the properties of solutions, such as their reactivity and equilibrium states.
Excess Gibbs energy: Excess Gibbs energy is a thermodynamic property that measures the deviation of a real solution's Gibbs energy from that of an ideal solution. It provides insights into the interactions between molecules in a mixture, highlighting how real solutions differ from ideal behavior due to factors like molecular size, shape, and polarity. Understanding excess Gibbs energy is essential for analyzing properties like activity and activity coefficients, which describe the non-ideal behavior of components in solutions.
Gibbs-Duhem Equation: The Gibbs-Duhem equation is a fundamental relationship in thermodynamics that connects the changes in chemical potential to the changes in composition of a system at constant temperature and pressure. This equation highlights how the chemical potential, which describes the potential energy change due to the addition of particles, varies with the concentration of different components in a mixture, playing a crucial role in understanding phase equilibria and thermodynamic properties of mixtures.
Henry's Law: Henry's Law states that at a constant temperature, the amount of gas that dissolves in a liquid is directly proportional to the partial pressure of that gas above the liquid. This relationship highlights how the behavior of gases in liquids is influenced by pressure, and it is crucial for understanding various phenomena related to solutions and mixtures.
Ideal solution: An ideal solution is a type of mixture that perfectly follows Raoult's law, where the interactions between different molecules in the solution are similar to those between like molecules. In such solutions, the vapor pressures of the components are proportional to their mole fractions, leading to predictable behavior during phase changes. This concept is essential for understanding how different mixtures behave when subjected to changes in temperature and pressure, especially in processes like distillation and flash calculations.
Ionic strength: Ionic strength is a measure of the concentration of ions in a solution, reflecting the total ionic charge contributed by all ions present. It plays a crucial role in determining the behavior of electrolytes in solution, impacting the activity coefficients of ions, which are essential for understanding how ions interact in various chemical and biological processes.
Mixing: Mixing refers to the process of combining two or more substances to form a homogeneous mixture, where the individual components are uniformly distributed throughout. This concept is crucial in understanding how fluids interact, influencing properties such as concentration, temperature, and energy transfer. Mixing is essential in various applications, from chemical reactions to fluid dynamics, impacting the behavior of solutions and the efficiency of processes.
Negative deviations: Negative deviations occur when the behavior of a real solution differs from that predicted by Raoult's law, resulting in a lower vapor pressure than expected. This phenomenon indicates that the interactions between unlike molecules in a mixture are stronger than those between like molecules, leading to a suppression of vapor pressure. Negative deviations highlight the importance of molecular interactions and the role they play in determining the properties of solutions.
Non-ideal solution: A non-ideal solution is a mixture of two or more substances that does not obey Raoult's law and exhibits deviations in behavior from ideal solutions due to interactions between different molecules. This concept highlights how real-world solutions behave under varying concentrations, temperatures, and pressure, influenced by factors like molecular size, polarity, and hydrogen bonding. These deviations can be quantitatively expressed using activity coefficients, which help describe the effective concentration of a component in the solution compared to its behavior in an ideal scenario.
Partial Molar Volume: Partial molar volume is the change in the volume of a solution when an additional amount of a component is added, keeping the temperature and pressure constant. This concept helps understand how individual components in a mixture affect the overall properties of the solution, including their activities and how they relate to fugacity.
Positive deviations: Positive deviations refer to the phenomenon where the behavior of a solution deviates positively from ideal behavior, usually indicated by a higher vapor pressure than expected. This occurs when the interactions between different molecules in a solution are weaker than those among the pure components, leading to greater vaporization. Positive deviations highlight the importance of non-ideal solutions and how intermolecular forces affect physical properties like vapor pressure and activity coefficients.
Pressure Dependence: Pressure dependence refers to how the properties and behaviors of a substance, especially in terms of its thermodynamic state, vary with changes in pressure. This concept is crucial when discussing phase equilibria, where pressure changes can significantly alter the activity and activity coefficients of species in a mixture, affecting their chemical potential and reactivity.
Raoult's Law: Raoult's Law states that the partial vapor pressure of each component in an ideal solution is equal to the vapor pressure of the pure component multiplied by its mole fraction in the solution. This law is crucial for understanding phase behavior, phase diagrams, and the behavior of mixtures, particularly in liquid-vapor equilibrium scenarios.
Solute-solvent interactions: Solute-solvent interactions refer to the various forces and mechanisms that occur between solute particles (the substance being dissolved) and solvent molecules (the substance doing the dissolving). These interactions are critical in determining the solubility, stability, and behavior of solutions, influencing properties such as boiling point, melting point, and vapor pressure. Understanding these interactions is essential for differentiating between ideal and non-ideal solutions, analyzing equilibria in phase transitions, and applying activity coefficients.
Temperature dependence: Temperature dependence refers to how various properties of substances, such as activity, equilibrium constants, and heat effects, change with variations in temperature. This concept is crucial because it influences the behavior of chemical reactions, the composition of mixtures, and the physical properties of fluids under different thermal conditions.
Van 't Hoff equation: The van 't Hoff equation relates the change in equilibrium constant of a reaction to changes in temperature, providing a crucial understanding of how temperature affects chemical equilibria. This equation connects the concepts of activity and activity coefficients to equilibrium constants, allowing for the assessment of how reactions shift with temperature variations, ultimately influencing reaction dynamics and thermodynamics.