Thermodynamic of Solutions
.phasecomposed of only onemixturehomogeneousis asolution, achemistryIn
ubstance, knownin another sdissolvedis a substancesoluteIn such a mixture, a
. The solvent does the dissolving. The solution more or less takes onsolventas a
characteristics of the solvent including its phase, and the solvent is commonlythe
of a solute in a solution is aconcentrationthe major fraction of the mixture. The
much of that solute is dissolved in the solventmeasure of how
A solution is a homogeneous mixture
The particles of solute in solution cannot
be seen by naked eye.
The solution does not allow beam
of light to scatter.
A solution is stable.
The solute from the solution cannot be
separated by filtration (or mechanically).
Homogeneous means that the components of the mixture form a single phase.
The properties of the mixture (such as concentration, temperature, and density)
can be uniformly distributed through the volume but only in absence of diffusion
phenomena or after their completion. Usually, the substance present in the
greatest amount is considered the solvent. Solvents can be gases, liquids or
solids. One or more components present in the solution other than the solvent
are called solutes. The solution has the same physical state as the solvent
If the solvent is a gas, only gases 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 molecules play almost no role, dilute gases
form rather trivial solutions. In part of the literature, they are not even classified as
solutions, but addressed as mixtures.
If the solvent is a liquid, then gases, liquids, and solids can be dissolved. Here
are some examples:
Gas in liquid:
Oxygen in water
Carbon dioxide in water – a less simple example, because the solution is
accompanied by a chemical reaction (formation of ions). Note also that
the visible bubbles in carbonated water are not the dissolved gas, but only
aneffervescence of carbon dioxide that has come out of solution; the
dissolved gas itself is not visible since it is dissolved on a molecular level.
Liquid in liquid:
The mixing of two or more substances of the same chemistry but different
concentrations to form a constant. (Homogenization of solutions)
Alcoholic beverages are basically solutions of ethanol in water.
Solid in liquid:
Sucrose (table sugar) in water
Sodium chloride (table salt) or any other salt in water, which forms
an electrolyte: When dissolving, salt dissociates into ions.
Counterexamples are provided by liquid mixtures that are
not homogeneous: colloids, suspensions, emulsions are not considered
Body fluids are examples for 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
If the solvent is a solid, then gases, liquids and solids can be dissolved.
Gas in solids:
Hydrogen dissolves rather well in metals, especially in palladium; this is
studied as a means of hydrogen storage.
Liquid in solid:
Mercury in gold, forming an amalgam
Hexane in paraffin wax
Solid in solid:
Steel, basically a solution of carbon atoms in a crystalline matrix of iron
Alloys like bronze and many others.
Polymers containing plasticizers.
Main article: Solubility
Main article: Solvation
The ability of
one compound to dissolve in
another compound is
called solubility. 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 called immiscible.
All solutions have a positive entropy of mixing. The interactions between different
molecules or ions may be energetically favored or not. If interactions are
unfavorable, then the free energy decreases with increasing solute concentration.
At some point the energy loss outweighs the entropy gain, and no more solute
particles can be dissolved; the solution is said to besaturated. 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
anexothermic 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.
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
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).
See also: Solvent § Solvent classifications
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
Water is a good solvent because the molecules are polar and capable of forming hydrogen
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.
For non-ionic solutes, the general rule is: like dissolves like.
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 andgrease that easily mix with each other, while being
incompatible with water.
An example for 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
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 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 know, same equation as case 1. VA=Vt; mB =
C VA /100
Example: Make 2 g/100mL of NaCl solution with 1 L water Water (properties).
The density of resulting solution is considered to be equal to that of water,
statement holding especially for dilute solutions, so the density information is not
Freezing and Boiling Points
For a solution with a liquid as solvent, the temperature at which it
freezes to a solid is slightly lower than the freezing point of the pure
solvent. This phenomenon is known as freezing point depression and is
related in a simple manner to the concentration of the solute. The
lowering of the freezing point is given by
ΔT 1 = K fm
where K f is a constant that depends on the specific solvent and m is
the molality of the molecules or ions solute. Table 1 gives data for
several common solvents.
In physics, Henry's law is one of the gas laws formulated by William Henry in
1803. It states:
"At a constant temperature, the amount of a given gas that dissolves in a
given type and volume of liquid is directly proportional to the partial
pressure of that gas in equilibrium with that liquid."
An equivalent way of stating the law is that the solubility of a gas in a liquid is
directly proportional to the partial pressure of the gas above the liquid.
An everyday example of Henry's law is given by carbonated soft drinks.
Before the bottle or can of carbonated drink is opened, the gas above the
drink is almost pure carbon dioxide at a pressure slightly higher
than atmospheric pressure. The drink itself contains dissolved carbon dioxide.
When the bottle or can is opened, some of this gas escapes, giving the
characteristic hiss. Because the partial pressure of carbon dioxide above the
liquid is now lower, some of the dissolved carbon dioxide comes out of
solution as bubbles. If a glass of the drink is left in the open, the concentration
of carbon dioxide in solution will come into equilibrium with the carbon dioxide
in the air, and the drink will go "flat".
Temperature dependence of the Henry constant
When the temperature of a system changes, the Henry constant will also
change. This is why some people prefer to name it Henry coefficient. Multiple
equations assess the effect of temperature on the constant. These forms of
the van 't Hoff equation are examples:
kH for a given temperature is Henry's constant (as defined in this article's
first section). Note that the sign of C depends on whether kH,pc or kH,cp is
T is any given temperature, in K