Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Solutions and Mixtures ↓


Solubility and Solubility Rules


Solubility is an important concept in chemistry that relates to the ability of a substance to dissolve in a solvent, forming a solution. Understanding solubility is important in a variety of chemical processes and applications, such as pharmaceuticals, environmental science, and materials engineering. This lesson will explore the fundamentals of solubility, the concept of solutions, and solubility rules that help predict whether certain substances will dissolve in others.

What is solubility?

Solubility is defined as the maximum amount of solute that can dissolve in a given volume of solvent at a specified temperature and pressure, resulting in a saturated solution. Solubility is usually expressed in terms of concentration, such as grams of solute per 100 grams of solvent or moles of solute per liter of solution.

Solutions and their types

A solution is a homogeneous mixture made up of two or more substances. The solute in a solution is the substance that dissolves, while the solvent is the substance that is being dissolved. For example, when you mix sugar in water, the sugar is the solute and the water is the solvent, creating a sugar solution.

Solutions may be classified based on the state of matter of the solvent:

  • Gaseous solutions: Here the solvent is a gas, such as air, which is a solution of oxygen, nitrogen and other gases.
  • Liquid solution: The solvent is a liquid, such as tea or juice.
  • Solid solution: The solvent is a solid, such as an alloy like brass, which is a solution of zinc in copper.

Visual representation of solutions

Solvent solute

In this illustration, the larger circle represents the solvent molecules, while the smaller circles represent solute molecules dispersed evenly throughout the solvent.

Factors affecting solubility

The solubility of a substance depends on several factors:

Nature of solute and solvent

The chemical nature of the solute and the solvent has a significant effect on solubility. The "like dissolves like" principle states that polar solutes generally dissolve well in polar solvents, and nonpolar solutes dissolve in nonpolar solvents. For example, NaCl is very soluble in water because both are polar, while oil (nonpolar) does not dissolve in water.

Temperature

Temperature plays an important role in determining solubility. For most solid solutes, solubility increases with increasing temperature. However, this is not a universal rule; some solutes may become less soluble as the temperature increases. In contrast, the solubility of gases in liquids generally decreases with increasing temperature.

Pressure

Pressure mainly affects the solubility of gases. According to Henry's Law, the solubility of a gas in a liquid is directly proportional to the pressure of that gas above the liquid. This is why carbonated beverages fizz; they contain dissolved carbon dioxide under high pressure.

Henry's law can be represented by this formula:

S = kH × P

where S is the solubility, kH is the Henry's law constant, and P is the partial pressure of the gas.

Solubility rules

Solubility rules are guidelines that help predict whether an ionic compound will dissolve in water. These rules are based on empirical observations and provide a quick reference for predicting the solubility behavior of various compounds.

Here are some common solubility rules:

Rule 1

  • Compounds containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) and ammonium ions (NH4+) are soluble.

Rule 2

  • Nitrates (NO3-), bicarbonates (HCO3-), and chlorates (ClO3- are usually soluble.

Rule 3

  • Chloride (Cl -), bromide (Br -), and iodide (I -) are soluble, except when they are combined with silver (Ag+), mercury (Hg22+), and lead (Pb2+).

Rule 4

  • Sulfates (SO42-) are soluble, with a few exceptions such as barium sulfate (BaSO4), lead sulfate (PbSO4), and calcium sulfate (CaSO4).

Rule 5

  • Carbonates (CO32-), phosphates (PO43-), chromates (CrO42-), and sulfides (S2-) are usually insoluble, except when they are combined with alkali metal ions or the ammonium ion.

Rule 6

  • Hydroxides (OH -) are insoluble, with the exception of those combined with alkali metals and the barium ion (Ba2+).

Practical applications of solubility

Knowledge of solubility and solubility rules is important in many industries and natural processes. Here are some examples:

Medicines

Solubility is an important factor in drug design and delivery. The efficacy of a drug often depends on its ability to dissolve in bodily fluids. Poorly soluble drugs may be difficult to absorb, reducing their effectiveness.

Environmental science

Solubility plays an important role in the transport and distribution of pollutants. For example, the solubility of some pollutants in water can affect their movement in an ecosystem.

Food industry

The solubility of various ingredients affects food processing and production. For example, the solubility of sugar is important in the production of beverages and confectionery.

Visual example of solubility in action

Water Salt

In the above diagram, salt is shown in a green circle. When salt is added to water, the salt dissolves and spreads throughout the solvent, forming a homogeneous solution.

Conclusion

Solubility and solubility rules provide an essential foundation for understanding chemical reactions and processes that occur in solutions. By knowing which substances are likely to dissolve in one another, scientists and engineers can better design experiments, develop new products, and manage environmental challenges.

Whether investigating the behavior of a pharmaceutical compound or assessing the effects of pollutants in nature, the principles of solubility help us understand the behavior and interactions of substances at the molecular level. Solubility laws serve as a guide in predicting and controlling the outcomes of mixing different compounds, thereby expanding our understanding and application of chemistry.


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Solubility and Solubility Rules

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