Solution Types & Classifications

Types of Solutions

• Solutions are homogenous mixtures in one phase with solute particles dissolved in a solvent.
• Solutions are classifiable by:
  • Phases
  • Saturation
  • Concentration

Solutions

• This is a homogenous mixture of two or more pure substances.
• They contain a solute and a solvent.

Classifications of Solutions

• Solutions can be classified on the solvent phase, saturation, and concentration.

Types of Solutions Based on the Solvent Phase

• There are three types: liquid, solid, and gaseous

Liquid Solutions

• These solutions have a liquid solvent and are the most common type.
• Examples are solid-liquid, liquid-liquid, and gas-liquid solutions (Ex: seawater).

Solid Solutions

• These solutions have a solid solvent and can be solid-solid, liquid-solid, or gas-solid state.
• An example is hydrogen gas in palladium metal.

Gaseous Solutions

• These solutions have a gas solvent and can be solid-gas, liquid-gas, or gas-gas.
• Ex: air

Types of Solutions Based on Saturation

• Solutions are classifiable based on the amount of solute dissolved in a certain amount of solvent at a specific temperature
• Solubility must also be considered.

Saturated Solution

• The amount of solute is equal to the solute's solubility at a certain volume and temperature.

Unsaturated Solution

• The amount of solute is less than the solute's solubility at a given volume and temperature.

Supersaturated Solution

• The amount of solute is greater than the solute's solubility at a given volume and temperature.
• Achieved by dissolving a solute at high temperature then cooling.
• Agitation causes crystallization because they are unstable

Types of Solutions Based on Concentration

• Solutions can be dilute, concentrated, or saturated.

Concentrated Solution

• This contains an excessively large amount of solute.
• Molarity (M) is the most common unit of concentration and molarities are >1M
• Concentrated sulfuric acid = 18 M, and glacial acetic acid = 17 M.
• They are often used as stock solutions when preparing diluted solutions
• Highly concentrated solutions emit fumes

Diluted Solution

• Solutions with low concentrations.
• Prepared by dilution: taking a portion of the stock solution or using an aliquot, and adding more solvent to lower the concentration.

Crystallization

• This occurs in a structured, organized pattern.
• This process occurs when solute particles arrange themselves in a repeating, orderly structure.

Precipitation

• The solute particles arrange themselves in an unordered solid

Terms Related to Solutions

• Solutions can be classified based on the solvent phase, saturation, and concentration.
• Liquid solutions have a liquid solvent whereas solid solutions have a solid solvent, and gaseous solutions have a gas solvent.
• Unsaturated solutions are solutions in which the quantity of solute is lower than the solute's solubility at a given temperature and volume.
• Saturated solutions are solutions in which the quantity of solute is equal to the solute's solubility at a given temperature and volume.
• Supersaturated solutions are solutions in which the quantity of solute is greater than the solute's solubility at a given temperature and volume.
• Solutions may be either diluted or concentrated.
• Concentrated solutions have too much solute.
• Dilute solutions have lower concentrations than stock solutions.

Energy of Solutions Formation

• Almost all substances dissolve in water under observable temperature and pressure but some do not.
• Substances that mix easily possess compatible intermolecular forces of attraction.

Solution Process

• Solute particles are dissolved in solvent particles.
• It causes a change in the intermolecular forces of attraction (IMFAs) between them
• Interactions include solvent-solvent, solute-solute, and solute-solvent.

Heat of Solution (ΔHmixing)

• Also called ΔHsoln
• Total of all enthalpy changes in the process
• ΔH mixing = ΔH1 + ΔH2 + ΔH3, where ΔH1 + ΔH2 is always positive.
• The three enthalpy changes determine the sign of ΔH mixing
• If ΔH mixing > 0 it is endothermic, occurs when ΔH1 + ΔH2 is greater than ΔH3, and solvent-solvent and solute-solute interactions are stronger than solute-solvent interactions.
• If ΔH mixing < 0 it is exothermic, occurs when ΔH1 + ΔH2 is less than ΔH3, and solvent-solvent and solute-solute interactions are weaker than solute-solvent interactions.
• Enthalpy is just one factor in solution formation.

Enthalpy (ΔS)

• Even if ΔH mixing > 0, the solute still dissolves due to the importance of entropy (S).
• Entropy is the inherent tendency toward disorder in highly favorable processes.
• The amount of chaos caused when solute and solvent particles mix determines their initial ordered states.

Solubility

• A measure of the extent to which a solute dissolves in a solvent at a particular temperature.
• Determined by ΔH mixing and ΔS mixing
• Solute and solvent with similar IMFAs tend to mix but those with differing IMFAs typically do not mix.
• Remember the phrase "like dissolves like".

Dissolution of Ionic Solids in Liquids

• Polar solvents are necessary during this process due to the differences in IMFAs

Solvation

• Solvation is the process in which solvent molecules surround an ion or a molecule.

Hydration

• Hydration occurs when water is the solvent and ions are attracted to the dipole by ion-dipole interaction.

Steps During Dissolution of Ionic Solids: Sodium Chloride dissolving in water example

• Step 1: Ions separate in the NaCl crystal lattice, forming Na+ and Cl- ions
  • Endothermic process
  • The crystal lattice energy causes ΔH1 > 0 .
• Step 2: Solvent molecules separate within the H2O molecules through H-bonding
  • This is endothermic with ΔH2 > 0.
• Step 3: Solute and solvent particles interact with Na+ interacting with partially negative O atoms, and Cl- interacting with partially positive H atoms.

Crystal Lattice Energy (CLE)

• The energy released when a mole of a solid's formula units is formed from its constituent ions.
• Formula: M+(g) + X-(g) → M+X(s) + CLE.
• CLE reflects the overall strength of IMFA in a solid, so a stronger IMFA holding the particles means higher CLE, meaning a higher amount of energy is needed to break the solid
• The constant ΔH1 = -CLE is always positive.

Solvation Shell

• This forms when solvent molecules surround ions, atoms, or molecules of the solute in a specific arrangement.

Hydration shell

• This occurs when water is the solvent and the solvation shell is then called a hydration shell.
• IMFAs form between ions and H2O molecules with ΔH3 < 0 because of the hydration energy.

Hydration Energy

• This is the energy released when one mole of formula units becomes hydrated.
• Hydration energy increases with charge density (ratio of charge to size/volume of the ion).
• Higher charge density means a greater magnitude of hydration energy.

Dissolution of Liquids in Liquids

• Similar to the process for liquid solutes in liquid solvents.
• The solvation process depends heavily on IMFAs present in both the solute and the solvent.

Immiscible Liquids

• Liquids that don't mix well and form distinct layers, exhibiting distinct IMFAs.

Ideal Solutions

• IMFAs of the solute and solvent are "compatible."
• ΔHmixing = 0

Dissolution of Gases in Liquids

• Like the general processes of dissolution in liquids
• "Like dissolves like" applies: polar gases are soluble in polar solvents, while nonpolar gases are soluble in nonpolar solvents.
• Some polar gases enhance solubility by reacting with water.

Properties of Gases in Liquids

• Few nonpolar gases are soluble in water because few can react.
• Gases have very weak IMFAs, so the energy needed to break the IMFAs is often negligible.
• Dissolution of gases in liquid is always exothermic.

Solution Terms

• The diagram on the next slide allows the molecular view of the solution process to be split into three steps.
• The overall heat of the solution process, ΔHmixing (also known as ΔHsoln), is calculable as the sum of all the enthalpy changes involved with each step. When ΔHmixing < 0, mixing is preferable because the interaction between solute and solvent overcomes the interactions between same molecule-type particles(solute-solute/solvent-solvent).
• Entropy (S) is the inherent tendency toward disorder in highly favorable processes. - Entropy of mixing is always positive.
• Solubility is the limit in which a solute dissolves in a solvent at a precise temperature.
• Solvation is when solvent molecules encompass an ion or molecule. If the solvent is water, then solvation is called hydration.
• Crystal lattice energy (CLE) is energy discharged when one mole of a formula unit in a solid is formed from its components.
• Hydration energy (HE) is energy discharged when one mole of formula units becomes hydrated and rises with the charge density of the ion.
• Solvation shells are formed when solvent molecules surround particles of a solute; if water is the solvent, they are called hydration shells.
• Two liquids are miscible if they exhibit similar IMFAs and dissolve into each other in all proportions.
• "Like dissolves like:" governs the rule that liquids and gases follow.

Affecting Solubility

• Dissolved oxygen amount rises as the temperature lowers, so fish are abundant during times of low temperature

Macroscopic View of Solution Formation

• Dynamic equilibrium is found through dissolution and crystallization processes
• The forward reaction is also the dissolution process
• The reverse reaction is also the crystallization process
• Solubility depends on the amount of solute and solvent.

Factors Affecting Solubility

• Nature of solute and solvent:
  • Temperature
  • Pressure or volume
  • Mechanical changes

Solute-Solvent Interaction

• ' Like dissolves like'
• When solute and solvent have same IMFAs they are soluble to one another
• Stronger interaction between solute and solvent with similar IMFAs causes more solubility
• Ionic solids dissolve in polar solvents where hydrogen bonds hold the solution together better than ion-induced dipole

Temperature

• The most obvious factor is temperature
• most solids are soluble at high temperatures
• temperature increase increases the solubility
• solubility of gases decreases as temperature increases

Henry's Law

• Henry's Law defines the relationship between the pressure of the gas over the solution and solubility of the gases