User:Xiang Li(Gary)/Solution (chemistry)

Lead
In chemistry, a solution is a special type of homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent. If the attractive forces between the solvent and solute particles are greater than the attractive forces holding the solute particles together, the solvent particles pull the solute particles apart and surround them. These surrounded solute particles then move away from the solid solute and out into the solution. The mixing process of a solution happens at a scale where the effects of chemical polarity are involved, resulting in interactions that are specific to solvation. The solution usually has the state of the solvent when the solvent is the larger fraction of the mixture, as is commonly the case. One important parameter of a solution is the concentration, which is a measure of the amount of solute in a given amount of solution or solvent. The term "aqueous solution" is used when one of the solvents is water.

Characteristics

 * A solution is a homogeneous mixture of two or more substances.
 * The particles of solute in a solution cannot be seen by the naked eye. By contrast, particles may be visible in a suspension.
 * A solution does not cause beams of light to scatter. By contrast, the particles in a suspension or colloid can cause Tyndall scattering or Rayleigh scattering.
 * A solution is stable, and solutes will not precipitate unless added in excess of the mixture's solubility, at which point the excess would remain in its solid phase. A solution containing more dissolved solutes than at equilibrium is referred to as supersaturated.
 * The solutes and solvents in a solution cannot be separated by filtration (or mechanically).
 * It is composed of only one phase.


 * Solutions are homogeneous mixtures where the solute is completely dissolved in the solvent, forming a clear and uniform phase. In contrast, suspensions and colloids are heterogeneous mixtures with larger dispersed phase particles, displaying distinct optical properties like light scattering. Particles in suspensions are visible and large enough to be separated by filtration, whereas colloid particles, though smaller, still scatter light, giving the mixture a Tyndall effect.

Types
In discussions about mixtures, it is essential to differentiate between homogeneous and heterogeneous mixtures. Homogeneous mixtures consist of components that blend into a single phase, whereas heterogeneous mixtures contain components that exist in distinct phases. The properties such as concentration, temperature, and density of a mixture can be uniform throughout its volume, but this uniformity is contingent on the absence of diffusion phenomena or after their completion. Generally, the component that is present in the largest quantity is considered the solvent, which may be in the form of gases, liquids, or solids. Any other component in the solution is termed a solute. The solution will maintain the same physical state as the solvent.

If the solvent is a gas, only gases (non-condensable) or vapors (condensable) are dissolved under a given set of conditions. An example of a gaseous solution is air (oxygen and other gases dissolved in nitrogen). Since interactions between gaseous molecules play almost no role, non-condensable gases form rather trivial solutions. In the literature, they are not even classified as solutions, but simply addressed as homogeneous mixtures of gases. The Brownian motion and the permanent molecular agitation of gas molecules guarantee the homogeneity of the gaseous systems. Non-condensable gaseous mixtures (e.g., air/CO2, or air/xenon) do not spontaneously demix, nor sediment, as distinctly stratified and separate gas layers as a function of their relative density. Diffusion forces efficiently counteract gravitation forces under normal conditions prevailing on Earth. The case of condensable vapors is different: once the saturation vapor pressure at a given temperature is reached, vapor excess condenses into the liquid state.

If the solvent is a liquid, it can dissolve nearly all gases, liquids, and solids. The dissolution of oxygen in water and carbon dioxide in water are examples. Carbon dioxide in water is more complex because it involves a chemical reaction (formation of ions). Notably, the visible bubbles in carbonated water are not the dissolved gas but are merely carbon dioxide effervescing from the solution; the dissolved gas itself is invisible at the molecular level. The blending of liquids, such as achieving homogeneous solutions from two substances with the same chemical composition but different concentrations, and alcoholic beverages, which are solutions of ethanol in water, exemplify liquid-liquid solutions. Solids in liquid solutions include sucrose in water and sodium chloride in water. NaCl or any other salt dissociates into ions to form an electrolyte. Solutions in water are particularly common and are known as aqueous solutions, whereas non-aqueous solutions occur when the solvent is not water. Examples of non-homogeneous liquid mixtures include colloids, suspensions, and emulsions, all of them are not classified as solutions. At same time, body fluids are examples of complex liquid solutions, containing many solutes. Many of these are electrolytes since they contain solute ions, such as potassium. Furthermore, they contain solute molecules like sugar and urea. Oxygen and carbon dioxide are also essential components of blood chemistry, where significant changes in their concentrations may be a sign of severe illness or injury.

If the solvent is a solid, it can dissolve gases, liquids, and other solids. The absorption of gases in solids, like hydrogen in metals such as palladium, is researched for hydrogen storage applications. Examples of liquids in solids include mercury in gold, forming an amalgam, and water in solid salts or sugars to form moist solids. Hexane in paraffin wax and polymers with added plasticizers, like phthalate in PVC, are instances of liquid in solid. Solids dissolving in solids, such as steel, which is a solid solution consisting of carbon atoms dispersed within the crystal structure of iron atoms. And alloys like bronze also belongs to solid-solid solutions. An example of a true solid solution is radium sulfate dissolved in barium sulfate.

Solubility
Solubility can be defined as the proportion of a specific solute in a specific solvent in saturated solution. This proportion can be expressed in various ways, such as concentration, molality, mole fraction, or mole ratio. These terms describe the amount of solute dissolved in a solvent at equilibrium in a saturated solution. When a liquid can completely dissolve in another liquid the two liquids are miscible. Two substances that can never mix to form a solution are said to be immiscible.

All solutions have a positive entropy of mixing. The interactions between different molecules or ions can be energetically favored or not. If the interactions are unfavorable, then the free energy of the system actually increases, not decreases, with an increase in solute concentration. This is because unfavorable interactions require additional energy to overcome the repulsion forces between molecules, to disperse the solute molecules or ions in the solvent. At some point, the energy loss will surpass the gain in entropy, hence no more solute molecules or ions can be dissolved; at this point, the solution is said to be saturated. However, the point at which a solution can become saturated can change significantly with different environmental factors, such as temperature, pressure, and contamination. For some solute-solvent combinations, a supersaturated solution can be prepared by raising the solubility (for example by increasing the temperature) to dissolve more solute and then lowering it (for example by cooling).

Usually, the greater the temperature of the solvent, the more of a given solid solute it can dissolve. However, most gases and some compounds exhibit solubilities that decrease with increased temperature. Such behavior is a result of an exothermic enthalpy of solution. Some surfactants exhibit this behaviour. The solubility of liquids in liquids is generally less temperature-sensitive than that of solids or gases.

Properties
The physical properties of compounds such as melting point and boiling point change when other compounds are added. Together they are called colligative properties. There are several ways to quantify the amount of one compound dissolved in the other compounds collectively called concentration. Examples include molarity, volume fraction, and mole fraction.

The properties of ideal solutions can be calculated by the linear combination of the properties of its components. If both solute and solvent exist in equal quantities (such as in a 50% ethanol, 50% water solution), the concepts of "solute" and "solvent" become less relevant, but the substance that is more often used as a solvent is normally designated as the solvent (in this example, water).

Liquid solution characteristics
In principle, all types of liquids can behave as solvents: liquid noble gases, molten metals, molten salts, molten covalent networks, and molecular liquids. In the practice of chemistry and biochemistry, most solvents are molecular liquids. They can be classified into polar and non-polar, according to whether their molecules possess a permanent electric dipole moment. Another distinction is whether their molecules can form hydrogen bonds (protic and aprotic solvents). Water, the most commonly used solvent, is both polar and sustains hydrogen bonds.

Salts dissolve in polar solvents, forming positive and negative ions that are attracted to the negative and positive ends of the solvent molecule, respectively. If the solvent is water, hydration occurs when the charged solute ions become surrounded by water molecules. A standard example is aqueous saltwater. Such solutions are called electrolytes. Whenever salt dissolves in water ion association has to be taken into account.

Polar solutes dissolve in polar solvents, forming polar bonds or hydrogen bonds. As an example, all alcoholic beverages are aqueous solutions of ethanol. On the other hand, non-polar solutes dissolve better in non-polar solvents. Examples are hydrocarbons such as oil and grease that easily mix, while being incompatible with water.

An example of the immiscibility of oil and water is a leak of petroleum from a damaged tanker, that does not dissolve in the ocean water but rather floats on the surface.

Preparation from constituent ingredients
It is common practice in laboratories to make a solution directly from its constituent ingredients. There are three cases in practical calculation:


 * Case 1: amount of solvent volume is given.
 * Case 2: amount of solute mass is given.
 * Case 3: amount of final solution volume is given.

In the following equations, A is solvent, B is solute, and C is concentration. Solute volume contribution is considered through the ideal solution model.


 * Case 1: amount (mL) of solvent volume VA is given. Solute mass mB = C VA dA /(100-C/dB)
 * Case 2: amount of solute mass mB is given. Solvent volume VA = mB (100/C-1/ dB )
 * Case 3: amount (mL) of final solution volume Vt is given. Solute mass mB = C Vt /100; Solvent volume VA=(100/C-1/ dB) mB
 * Case 2: solute mass is known, VA = mB 100/C
 * Case 3: total solution volume is known, same equation as case 1. VA=Vt; mB = C VA /100

Example: Make 2 g/100mL of NaCl solution with 1 L water. The density of the resulting solution is considered to be equal to that of water, statement holding especially for dilute solutions, so the density information is not required.


 * mB = C VA = ( 2 / 100 ) g/mL × 1000 mL = 20 g

Chemists often make concentrated stock solutions that may then be diluted as needed for laboratory applications. Standard solutions are those where concentrations of solutes are accurately and precisely known.