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Solubility and ion concentration

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Intros
Lessons
  1. Updating our knowledge of solubility
  2. Recap on solutions and solubility.
  3. Electrolytes and ionic solutions.
  4. Using molecular formula to predict electrolytes.
  5. Solubility and saturation, using equilibrium.
  6. Calculating solubility: example.
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Examples
Lessons
  1. Calculate the solubility of substances given mass and volume of solvent.
    Some students tested the solubility of different substances by dissolving increasing amounts of them in water until the solutions became saturated. Some results of theirs are below:

    Chemical A: 54g was dissolved in 700 mL water.
    Chemical B: 3g was dissolved in 1 L water.
    Chemical C: 58g was dissolved in 750 mL water.

    1. Which of the salts has the highest solubility?
    2. Are these chemicals likely to be electrolytes or non-electrolytes? Explain your reasoning.
    3. Chemical B has a molar mass of 86 grams per mole. What is its molar solubility?
    Topic Notes
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    Introduction to Solubility and Ion Concentration

    Solubility and ion concentration are fundamental concepts in chemistry that play a crucial role in understanding various chemical processes. Our introduction video provides an essential foundation for grasping these topics, making it an invaluable resource for students and enthusiasts alike. Solubility refers to the ability of a substance to dissolve in a solvent, while ion concentration measures the amount of ions present in a solution. These concepts are intricately linked to the principle of equilibrium, where the rates of dissolution and crystallization balance each other. Understanding the relationship between solubility, ion concentration, and equilibrium is vital for predicting chemical reactions, analyzing solution properties, and solving real-world problems in fields such as environmental science, pharmaceuticals, and materials engineering. By exploring these interconnected concepts, we can gain deeper insights into the behavior of substances in solution and their practical applications in various scientific disciplines.

    Understanding Chemical and Physical Changes Using the Particle Model

    The particle model is a fundamental concept in chemistry that helps us understand the behavior of matter at the microscopic level. This model is particularly useful in explaining both chemical and physical changes, providing insights into how substances interact and transform. By visualizing matter as composed of tiny particles, we can better comprehend the processes occurring during various chemical reactions and physical transformations.

    Let's first explore how the particle model aids in understanding chemical changes. A chemical change involves the rearrangement of atoms to form new substances with different properties. The combustion of methane (CH4) serves as an excellent example to illustrate this concept. When methane burns in the presence of oxygen, a chemical reaction occurs, resulting in the formation of carbon dioxide (CO2) and water (H2O). At the particle level, we can visualize the methane molecules breaking apart and their atoms recombining with oxygen atoms to create new molecules.

    In the combustion of methane, the particle model helps us understand that the carbon atom from CH4 combines with two oxygen atoms to form CO2, while the hydrogen atoms join with oxygen to create H2O. This rearrangement of atoms explains why the products have different properties from the reactants. The particle model allows us to visualize how the bonds between atoms break and new bonds form, leading to the creation of entirely new substances.

    On the other hand, physical changes involve alterations in the physical state or form of a substance without changing its chemical composition. The particle model is equally valuable in explaining these transformations. Let's consider the state changes of carbon dioxide as an example. CO2 can exist as a solid (dry ice), liquid, or gas, depending on temperature and pressure conditions.

    When solid CO2 sublimes directly into a gas, the particle model helps us understand that the individual CO2 molecules remain intact, but their arrangement and behavior change. In the solid state, the CO2 molecules are closely packed and have limited movement. As energy is added, these molecules gain kinetic energy, overcoming the intermolecular forces holding them together. They begin to move more freely, transitioning directly into the gaseous state where they are far apart and move rapidly.

    The particle model illustrates that during physical changes, the fundamental particles (atoms or molecules) retain their identity. The changes occur in how these particles are arranged or how they interact with each other. This is in contrast to chemical changes, where the particles themselves are altered through the breaking and forming of chemical bonds.

    Understanding the differences between chemical and physical changes at the particle level is crucial. In chemical changes, we observe:

    • Rearrangement of atoms
    • Formation or breaking of chemical bonds
    • Creation of new substances with different properties
    • Changes in the electronic structure of atoms

    Conversely, in physical changes, we see:

    • No alteration of the fundamental particles
    • Changes in the arrangement or state of particles
    • Modifications in intermolecular forces
    • Retention of the chemical identity of the substance

    The particle model helps us visualize these differences. For instance, when water freezes (a physical change), the H2O molecules remain intact but arrange themselves in a crystalline structure. In contrast, when water undergoes electrolysis (a chemical change), the H2O molecules are split into hydrogen and oxygen gases, fundamentally altering the substance at the particle level.

    Intermolecular forces play a crucial role in both types of changes. In physical changes, these forces are often overcome or altered without breaking the bonds within molecules. For example, when ice melts, the hydrogen bonds between water molecules are weakened, allowing for greater molecular movement. In chemical changes, intermolecular forces may influence reaction rates or product formation, but the primary focus is on the breaking and forming of chemical bonds within molecules.

    The particle model's strength lies in its ability to provide a visual and conceptual framework for understanding these complex processes. It allows us to mentally picture the behavior of atoms and molecules during various transformations, making abstract concepts more tangible. This model is particularly valuable in educational settings, helping students grasp the fundamental concepts of carbon dioxide state changes and kinetic energy.

    Solubility as a Chemical and Physical Process

    The process of dissolving is a fascinating phenomenon that can be observed in everyday life, and it serves as an excellent example of how chemical and physical changes intertwine. To understand this process better, let's explore the dissolution of sodium chloride (table salt) in water, a common occurrence in kitchens worldwide.

    Sodium chloride is an ionic compound, consisting of positively charged sodium ions (Na+) and negatively charged chloride ions (Cl-) held together by strong ionic bonds. When we add salt to water, we initiate a complex series of events at the molecular level that ultimately lead to the salt dissolving.

    The dissolving process begins as water molecules, which are polar, surround the salt crystals. The positive end of the water molecules is attracted to the negatively charged chloride ions, while the negative end is drawn to the positively charged sodium ions. This attraction creates a force that begins to overcome the ionic bonds holding the salt crystal together.

    As more water molecules interact with the ions on the surface of the salt crystal, they gradually pull these ions away from the crystal structure. This process is known as hydration. The sodium and chloride ions become surrounded by water molecules, forming hydration shells. These hydrated ions then move freely throughout the solution, dispersing evenly.

    From a chemical perspective, the dissolving of sodium chloride involves the breaking of ionic bonds within the crystal lattice. This bond breaking could be considered a chemical change, as it alters the arrangement of atoms and ions. However, no new substances are formed during this process, and the chemical identity of sodium and chloride ions remains unchanged.

    Simultaneously, from a physical standpoint, the dissolving process results in a change of state. The solid salt transitions into a homogeneous solution, where the ions are evenly distributed throughout the liquid. This change in physical state and the resulting uniform mixture are characteristics of a physical change.

    The particle model provides an excellent visual representation of this process. Imagine the salt crystal as a well-organized structure of alternating sodium and chloride ions. As water molecules surround the crystal, they begin to pull individual ions away from the surface. These ions become enveloped by water molecules, forming a new arrangement where the ions are separated and dispersed throughout the solution.

    One of the most intriguing aspects of the dissolving process is its reversibility. If we were to evaporate the water from our salt solution, we would once again obtain solid sodium chloride crystals. This ability to reverse the process further supports the classification of dissolution as primarily a physical change, despite the breaking and forming of interactions at the molecular level.

    The reversibility of dissolving is crucial in many natural and industrial processes. For example, the formation of stalactites and stalagmites in caves occurs when minerals dissolved in water precipitate out as the water evaporates. In industrial settings, this principle is used in purification processes, where substances are dissolved and then recrystallized to remove impurities.

    Understanding solubility as both a chemical and physical process has significant implications across various fields. In chemistry, it forms the basis for many reactions and separations. In environmental science, it helps explain the transport of pollutants in water systems. In biology, the solubility of different substances in cell membranes is crucial for many life processes.

    The dissolving process of sodium chloride exemplifies the intricate interplay between chemical and physical changes in nature. While the breaking of ionic bonds suggests a chemical aspect, the overall process - including the change in physical state and the ability to reverse the dissolution - aligns more closely with a physical change. This dual nature of solubility highlights the complexity of molecular interactions and the sometimes blurry line between chemical and physical phenomena in the natural world.

    Updated Definitions of Solubility and Saturation

    In the realm of chemistry, understanding solubility and saturation is crucial for various applications, from pharmaceutical development to environmental science. The modern definitions of these concepts are intrinsically linked to the principle of equilibrium, providing a more comprehensive understanding of solution dynamics.

    Solubility, in its updated definition, refers to the maximum amount of a solute that can dissolve in a given quantity of solvent at equilibrium under specific conditions. This equilibrium-based definition emphasizes that solubility is not just a static property but a dynamic state where the rate of dissolution equals the rate of crystallization. When a solution reaches this point, it is said to be saturated.

    Saturation, therefore, is the state of a solution that exists in equilibrium with undissolved solute. At saturation, the solution contains the maximum concentration of dissolved solute for the given conditions. It's important to note that a saturated solution can still be in equilibrium with excess undissolved solute, maintaining a constant concentration of dissolved particles.

    The concept of molar solubility provides a quantitative measure of solubility. Molar solubility is defined as the number of moles of solute that can dissolve in one liter of solvent to form a saturated solution at a specific temperature. This definition directly relates to the equilibrium concentration of the solute in the saturated solution. For instance, if the molar solubility of a compound is 0.1 mol/L, it means that at equilibrium, a saturated solution contains 0.1 moles of the dissolved compound per liter of solution.

    The relationship between molar solubility and equilibrium concentration is straightforward for simple dissociation reactions. For a salt that dissociates into two ions, like NaCl, the equilibrium concentration of each ion in a saturated solution is equal to the molar solubility. However, for more complex dissociations, the relationship becomes more intricate, requiring consideration of stoichiometry and equilibrium constants.

    Several factors significantly affect solubility, with temperature and solvent type being among the most influential. Temperature generally increases the solubility of solid solutes in liquid solvents. This is because higher temperatures provide more energy for molecular motion, enhancing the ability of solvent molecules to overcome the intermolecular forces in the solid solute. However, it's worth noting that for gases dissolved in liquids, the relationship is often inverse, with solubility decreasing as temperature rises.

    The nature of the solvent plays a crucial role in determining solubility. The general rule "like dissolves like" applies here, meaning polar solutes tend to dissolve better in polar solvents, while nonpolar solutes dissolve better in nonpolar solvents. Water, being a polar solvent, is excellent at dissolving ionic compounds and polar molecules but poor at dissolving nonpolar substances. Organic solvents like hexane, on the other hand, are better suited for dissolving nonpolar compounds.

    Pressure, while not as universally impactful as temperature, significantly affects the solubility of gases in liquids. Increasing pressure typically increases the solubility of gases, as described by Henry's law. This principle is vital in understanding phenomena like the formation of gas bubbles when a carbonated drink is opened, reducing the pressure and thus the solubility of dissolved CO2.

    The presence of other solutes can also influence solubility through the common ion effect or by altering the properties of the solvent. For instance, adding a common ion effect to a solution can decrease the solubility of a sparingly soluble salt containing that ion, a principle utilized in various precipitation reactions.

    Understanding these updated definitions and concepts is essential for predicting and controlling solubility in various applications. In the pharmaceutical industry, for example, knowledge of solubility is crucial for drug formulation and bioavailability. Environmental scientists use these principles to understand the distribution of pollutants in water systems. Chemical engineers apply solubility concepts in designing separation and purification processes.

    In conclusion, the modern definitions of solubility and saturation, rooted in equilibrium principles, provide a dynamic view of solution behavior. Molar solubility offers a quantitative link to equilibrium concentrations, while factors like temperature, solvent nature, an

    Calculating Solubility and Molar Solubility

    Understanding how to calculate solubility in chemistry and molar solubility is crucial in chemistry. These calculations help determine the amount of solute that can dissolve in a given amount of solvent. Let's explore the step-by-step process using examples of sodium chloride (NaCl) and copper sulfate (CuSO4) to demonstrate these important solubility calculations.

    Step 1: Understanding Solubility

    Solubility in chemistry is typically expressed in grams per liter (g/L) and represents the maximum amount of solute that can dissolve in a specific volume of solvent at a given temperature. Molar solubility, on the other hand, is expressed in moles per liter (mol/L) and represents the number of moles of solute that can dissolve in one liter of solvent.

    Step 2: Calculating Solubility

    To calculate solubility, you need to know the mass of solute that dissolves in a specific volume of solvent. For example, if 36 grams of NaCl dissolves in 100 mL of water, the solubility would be:

    Solubility = (36 g NaCl) / (0.1 L water) = 360 g/L

    Step 3: Calculating Molar Solubility

    To find molar solubility, you need to convert the solubility from g/L to mol/L. This requires knowing the molar mass of the solute. For NaCl (molar mass = 58.44 g/mol):

    Molar Solubility = (360 g/L) / (58.44 g/mol) = 6.16 mol/L

    Step 4: Unit Conversion

    It's essential to be able to convert between g/L and mol/L. To convert from g/L to mol/L, divide by the molar mass. To convert from mol/L to g/L, multiply by the molar mass.

    Example with Copper Sulfate (CuSO4)

    Let's calculate the solubility and molar solubility of CuSO4. Suppose 20 grams of CuSO4 dissolves in 150 mL of water.

    Step 1: Calculate solubility Solubility = (20 g CuSO4) / (0.15 L water) = 133.33 g/L

    Step 2: Calculate molar solubility Molar mass of CuSO4 = 63.55 + 32.07 + (4 × 16) = 159.62 g/mol Molar Solubility = (133.33 g/L) / (159.62 g/mol) = 0.84 mol/L

    Importance of Units

    Using the correct units is crucial in solubility calculations. Grams per liter (g/L) provides information about the mass of solute that can dissolve, while moles per liter (mol/L) gives insight into the number of particles in solution. This is particularly important when dealing with chemical reactions or comparing solubilities of different compounds.

    Practical Applications

    Understanding solubility calculations is essential in various fields, including:

    • Pharmaceutical industry: Determining drug solubility for proper dosage forms
    • Environmental science: Assessing pollutant concentrations in water bodies
    • Chemical engineering: Designing separation processes and crystallization techniques
    • Food science: Developing food products with specific solubility properties

    Practical Applications of Solubility and Ion Concentration

    Understanding solubility and ion concentration concepts has numerous real-world applications across various fields, from environmental science to industry and everyday life. These principles play a crucial role in problem-solving and decision-making processes in many sectors. In environmental science, solubility and ion concentration are essential for water quality management and pollution control. Scientists use these concepts to assess the health of aquatic ecosystems, determine the impact of pollutants, and develop strategies for water treatment. For instance, measuring the concentration of dissolved oxygen in water bodies helps evaluate their ability to support aquatic life. Similarly, understanding the solubility of various pollutants aids in designing effective remediation techniques for contaminated sites.

    In industry, solubility and ion concentration concepts are widely applied in manufacturing processes, quality control, and product development. The pharmaceutical industry relies heavily on these principles to formulate medications with optimal absorption rates and bioavailability. In the food and beverage sector, manufacturers use solubility data to create stable emulsions, develop new flavors, and improve product shelf life. The mining and metallurgy industries apply these concepts in ore extraction and metal purification processes, optimizing yields and reducing waste. Additionally, the agricultural sector utilizes ion concentration knowledge to formulate balanced fertilizers and monitor soil health.

    Everyday life also presents numerous examples of solubility and ion concentration applications. Water softeners in homes work by exchanging ions to reduce water hardness, improving the effectiveness of soaps and detergents. Swimming pool maintenance relies on understanding pH levels and chlorine concentration to ensure safe and hygienic conditions. In cooking, solubility principles are applied when dissolving sugar in beverages or creating emulsions like mayonnaise. Even in personal care products, formulators use solubility data to create stable and effective cosmetics and toiletries.

    Problem-solving in various fields often involves applying solubility and ion concentration concepts. Environmental engineers use these principles to design wastewater treatment plants and develop strategies for managing acid mine drainage. In forensic science, understanding solubility helps in analyzing trace evidence and reconstructing crime scenes. Medical professionals apply these concepts in diagnosing and treating conditions related to electrolyte imbalances or mineral deficiencies. By mastering these fundamental chemical principles, professionals across diverse fields can make informed decisions, optimize processes, and develop innovative solutions to complex challenges, ultimately contributing to advancements in technology, health, and environmental protection.

    Conclusion: Mastering Solubility and Ion Concentration

    In this article, we've explored the fundamental concepts of solubility and ion concentration in chemistry. We've covered key points such as solubility product constants, common ion effect, and the relationship between solubility and temperature. Understanding these principles is crucial for success in chemistry studies and practical applications. As you continue your journey in chemistry, remember to apply these concepts in your lab work and problem-solving exercises. The introduction video on solubility provided a valuable visual representation of these ideas, helping to solidify your understanding. By mastering solubility and ion concentration, you'll be better equipped to tackle more advanced topics in chemistry and related fields. We encourage you to review these concepts regularly and seek out additional resources to deepen your knowledge. With practice and persistence, you'll develop a strong foundation in this essential area of chemistry, opening doors to exciting opportunities in research, industry, and beyond.

    Remember, the solubility product constants are essential for predicting the solubility of compounds in various conditions. Similarly, understanding the common ion effect can help you manipulate solubility in practical applications. The introduction video on solubility is a great starting point for visual learners who need to see these concepts in action. Keep these resources handy as you progress in your studies.

    Updating our knowledge of solubility

    Recap on solutions and solubility.

    Step 1: Introduction to Chemical Reactions and the Particle Model

    In this lesson, we will recap what we have learned about chemical reactions so far, using the particle model to enhance our understanding. The particle model helps us visualize and explain the relationship between macroscopic properties (observable characteristics) and microscopic interactions (interactions between particles, atoms, and molecules).

    Step 2: Objectives of Using the Particle Model

    Our objectives include representing examples of chemical and physical changes using the particle model. We aim to use this model to explain the relationship between macroscopic properties and microscopic interactions. This approach will help us better understand the properties, bonding, and structure of molecules or substances.

    Step 3: Recap of Chemical Reactions

    We have learned the basics of chemical reactions, including how they differ from physical changes. We have also learned how to represent chemical reactions using word and symbol equations, balance equations, and understand various types of chemical changes such as decompositions, combustion, and acid-base neutralization.

    Step 4: Using the Particle Model to Examine Chemical Changes

    We will now use the particle model to examine chemical changes. Chemical changes involve the rearrangement of bonds, where reactant bonds are broken, and product bonds are formed. This process can be visualized using the particle model.

    Step 5: Example of Combustion of Methane

    Let's take the example of the combustion of methane. We have already balanced the equation: one carbon in, one carbon out; four hydrogens in, four hydrogens out; and four oxygens in, four oxygens out. Using the particle model, we can represent the reactants and products, showing the conservation of mass and the rearrangement of atoms.

    Step 6: Visualizing the Particle Model

    In the particle model, we can color-code atoms to represent different elements. For example, oxygen is generally red, carbon is black, and hydrogen is white. This color-coding helps us visualize the structure of molecules and the conservation of mass during chemical reactions.

    Step 7: Conservation of Mass in the Particle Model

    The particle model allows us to show the conservation of mass by balancing the number of molecules or atoms. For instance, in the combustion of methane, we can see that the number of oxygen, hydrogen, and carbon atoms remains the same before and after the reaction.

    Step 8: Representing Chemical Bonds

    While the particle model shows atoms attached to each other, it does not necessarily depict covalent bonds. For example, an O2 molecule has a double bond, but in the particle model, we can simply show the atoms as being close to each other without explicitly representing the double bond.

    Step 9: Rearrangement of Atoms in Chemical Reactions

    During chemical reactions, atoms rearrange to form new molecules. In the combustion of methane, hydrogen atoms combine to form water molecules, and carbon atoms combine with oxygen to form carbon dioxide. The particle model helps us visualize this rearrangement.

    Step 10: Communicating Chemical Reactions

    Both chemical equations and the particle model are ways to communicate that a reaction has occurred. While chemical equations provide a symbolic representation, the particle model offers a visual representation, helping us understand the process at a molecular level.

    Step 11: Using the Particle Model for Physical and Chemical Changes

    We can use the particle model to describe both physical and chemical changes. This model helps us understand the interactions between particles and how these interactions lead to observable changes in properties. In the next section, we will explore how to use the particle model to describe physical and chemical changes in more detail.

    FAQs

    1. What determines ion concentration?

      Ion concentration is determined by several factors, including the solubility of the compound, temperature, pressure, and the presence of other ions in the solution. The dissociation of the compound in the solvent and the equilibrium between dissolved and undissolved solute also play crucial roles in determining ion concentration.

    2. How to find ionic concentration?

      To find ionic concentration, you can use the molarity of the solution and the dissociation equation of the compound. For example, for NaCl, which dissociates completely in water, the concentration of Na+ and Cl- ions will be equal to the molarity of the NaCl solution. For more complex compounds, you may need to consider the degree of dissociation and use equilibrium constants.

    3. What is meant by molar solubility?

      Molar solubility is the maximum number of moles of a solute that can dissolve in one liter of solvent to form a saturated solution at a specific temperature. It is typically expressed in mol/L and provides a quantitative measure of a substance's solubility.

    4. How are concentration and solubility related?

      Concentration and solubility are closely related but distinct concepts. Solubility is the maximum amount of solute that can dissolve in a given amount of solvent, while concentration refers to the amount of solute actually dissolved in a solution. The concentration of a solution cannot exceed its solubility limit under given conditions. As the concentration approaches the solubility limit, the solution becomes saturated.

    5. What is the difference between molarity and molar solubility?

      Molarity is a measure of concentration that expresses the number of moles of solute per liter of solution, while molar solubility specifically refers to the maximum number of moles of solute that can dissolve in one liter of solvent to form a saturated solution. Molarity can describe any concentration up to the saturation point, whereas molar solubility represents the concentration at saturation.

    Prerequisite Topics

    Understanding solubility and ion concentration is crucial in chemistry, but to fully grasp these concepts, it's essential to have a solid foundation in several prerequisite topics. These fundamental areas of study provide the necessary context and knowledge to comprehend the intricacies of solubility and ion concentration.

    First and foremost, a strong introduction to solution chemistry and solubility is vital. This foundational knowledge helps students understand the basic principles of how substances dissolve in solvents and the factors that influence solubility. By mastering these concepts, students can better analyze and predict solubility behavior in various chemical systems.

    While it may not seem directly related, calculating cell potential in voltaic cells is also relevant to solubility and ion concentration. This topic introduces students to the concept of electrochemical reactions, which often involve ions in solution. Understanding how ions behave in these systems can provide valuable insights into their behavior in solubility scenarios.

    Another critical prerequisite is the study of intermolecular forces. These forces play a significant role in determining solubility, as they influence how solute particles interact with solvent molecules. A thorough understanding of intermolecular forces helps explain why some substances are more soluble than others and how temperature affects solubility.

    The common ion effect is particularly important when studying solubility and ion concentration. This phenomenon occurs when a soluble compound is added to a solution containing one of its ions, affecting the solubility equilibrium. Grasping this concept is crucial for predicting and explaining changes in solubility under various conditions.

    Lastly, familiarity with the solubility constant, also known as the solubility product, is essential. This quantitative measure of solubility helps students calculate and predict the concentration of ions in saturated solutions. Understanding how to use solubility product constants enables more accurate analysis of solubility equilibria and precipitation reactions.

    By mastering these prerequisite topics, students build a strong foundation for understanding solubility and ion concentration. Each concept contributes to a more comprehensive view of how substances dissolve, how ions behave in solution, and the factors that influence these processes. This knowledge is not only crucial for academic success in chemistry but also for practical applications in fields such as environmental science, pharmaceuticals, and materials engineering. As students progress in their studies, they'll find that these fundamental concepts continually resurface, reinforcing their importance in the broader context of chemical understanding.

    In this lesson, we will learn:

    • To recall the definition of electrolyte and non-electrolyte and define solubility in terms of equilibrium.
    • How to identify electrolytes and non-electrolytes using molecular formulae and predict ionic or molecular solutions based on this.
    • How to measure solubility of substances in solution.

    Notes:

    • To begin this chapter, let’s recall our definitions and knowledge about solubility and how solutions are made:
      • A solution is made of:
        • A solvent, the liquid substance in excess that does the dissolving, such as water.
        • A solute, the substance in a smaller quantity that is being dissolved, such as table salt.
        Together, these make a solution.
      • Our definition of solubility from Solution chemistry and solubility: introduction is “the extent that a substance can dissolve in a specific amount of another substance at a specific temperature”.
      • From this we learned that like dissolves like, which means substances similar to each other will interact favorably and dissolve in one another, whichever one is the solute or solvent.
      • Saturation is when a solution has dissolved the maximum amount of solute possible, given the solute’s solubility. When more solute was added after this point it simply collected, undissolved, at the bottom of the container.
      • Polarity is a very important property of chemicals when discussing solubility. Polarity (see Polarity) in a molecule exists because electrons (and charge) is unequally distributed in a molecule. This gives molecules a permanent ‘partial charge’ as there is greater electron density in some regions of the molecule than in others.
      • When comparing chemical substances and their solubility, chemicals are studied in terms of hydrophilic ("water-loving") or lipophilic ("fat-loving") nature.  
        • Hydrophilic molecules are molecules with high polarity due to a large gap in electronegativity between atoms bonded together. Hydrophilic molecules are soluble in water, and generally insoluble in organic compounds like simple hydrocarbons. Hydrophilic compounds are polar.
        • Lipophilic molecules are molecules with low polarity in the molecule – there is no significant gap in electronegativity between the atoms that form the molecule. This means the molecule has little to no partial charge, which makes it insoluble in water, and generally soluble in organic compounds like simple hydrocarbons. Lipophilic compounds are non-polar.

    • The ‘like dissolves like’ rule is important to remember for solubility. Hydrophilic compounds dissolve in polar solvents and lipophilic compounds dissolve in non-polar solvents.
      However, these two cases will make solutions with different properties:
      • Some compounds are electrolytes – substances that produce ions when dissolving to make a solution that conducts electricity. Electrolytes make ionic solutions.
        • Salt compounds made of metal and non-metal atoms are electrolytes and become ionic in solution. This is because salts have atoms or groups that have a very large difference in their electronegativity or electron withdrawing/donating properties. Salts therefore form oppositely charged ions in solution.
      • Some compounds are non-electrolytes – these compounds remain neutral when dissolved and the resulting solution does not conduct electricity.
        • Lipophilic compounds are generally non-electrolytes because the molecule likely doesn’t contain partial charges between the atoms bonded together, so ions would not form when it is dissolved in water, for example.

    • By studying the molecular formula of a compound, you can normally predict whether a compound is an electrolyte or not. Equations can be written to describe this.
      • Compounds that have both metal and non-metal atoms will likely be ionic in solution and are therefore electrolytes.
        • Salts fall into this category as mentioned already.
        • For example iron (iii) chloride, FeCl3, has a molecular formula that shows an iron (metal) atom and three chlorine (non-metal) atoms. This suggests it would be an electrolyte and form an ionic solution. In an equation:

          FeCl3 \enspace→ \enspace Fe3+(aq) + 3Cl-(aq)
      • Covalent compounds, made of only non-metal atoms, are generally non-electrolytes and will form molecular solutions.
        • Most organic compounds fall into this category unless it is an organic acid, which may be ionic to some degree.
        • For example hexane, C6H14, is made of just carbon and hydrogen. Both are non-metals and it is an organic compound. Both of these facts are strong evidence that it is a non-electrolyte and forms a molecular solution.
        • For non-electrolytes pay attention to state symbols. If it is not dissolved in water, do not call it aqueous or label the compound (aq)!

    • Now that we have covered dynamic equilibrium, our definitions of solubility and saturation can be adjusted to be more accurate:
      • Solubility is the equilibrium concentration of a substance dissolved in solution. This is sometimes measured in grams per liter (g / L) but also often measured in moles per liter or moles per cubic decimeter (mol dm-3) so may be called the molar solubility.
      • Saturation is when a substance that is dissolved is in equilibrium with its undissolved state.

    • Using these two ‘updated’ definitions we say that saturation occurs when two things are happening:
      • There is some undissolved material.
      • The undissolved material is in equilibrium with its dissolved state.

    • Take an example of a salt e.g. sodium chloride, NaCl, in equilibrium between the solid and aqueous phases:

      NaCl(s) ⇌\enspace\rightleftharpoons \enspace Na+(aq) + Cl-(aq)

      This can be separated into its two individual reactions – the dissolving (forward) process and the crystallization (backward) process:

      Dissolving: NaCl(s) \enspace→ \enspace Na+(aq) + Cl-(aq)

      Crystallization: Na+(aq) + Cl-(aq) \enspace→ \enspace NaCl(s)

    • Calculating solubility is usually done by adding the solute to 1 liter of solvent (assumed to be at 25oC) , and with the mass of solute recorded it is fairly simple to calculate:
      • If a saturated 1 L solution of NaCl (aq) contains 360g of NaCl. Then the solubility of the solution is 360 g / L. This can be expressed as molar solubility:

        [NaCl]=360gNaCl1L∗1 molNaCl58.4gNaCl=6.16M[NaCl] = \frac{360g \enspace NaCl}{1 \enspace L} * \frac{1\, mol \enspace NaCl}{58.4g \enspace NaCl} = 6.16 M

      • Converting from molar solubility to solubility in grams per liter is possible too, for example with copper sulfate, CuSO 4 which has molar solubility of 0.877 M.

        Solubility=0.877mol1L∗159.6g1mol=140gLSolubility = \frac{0.877 \enspace mol}{1 \enspace L} * \frac{159.6g}{1 \enspace mol} = 140 \frac{g}{L}


        When finding solubility of a substance, you can assume the solution is at 25oC, but pay attention to what volume of solution is quoted – if it is not in liters, convert it into liters first.