Understanding Solutions and Solutes

Solutions

• A solution is a homogenous mixture of two or more components at a molecular level.

Components of a Solution

• The component which determines the phase of the solution and accounts for the largest proportion of the system is the solvent.
• Solutes are dispersed as individual molecules or ions throughout the solvent, and may be solid, liquid, or gas.

Dissolution of Crystalline Solute

• The dissolution of a crystalline solute involves the removal of a solute molecule, creation of a cavity in the solvent, and the solute molecule filling the solvent cavity.

Types of Solutes

• Solutes can be divided into two main classes: electrolytes and non-electrolytes.
• Non-electrolytes don't form ions when dissolved in water; examples of non-electrolytes are sucrose, glycerin, naphthalene and urea.
• Electrolytes form ions when dissolved in water, strength varies depending on ionization; examples of strong electrolytes are HCI and Na sulfate, examples of weak electrolytes are ephedrine and phenobarbital.

States of Solutions

• Solutions are classified based on the states of the solute and solvent.
• With three states of matter (gas, liquid, solid), there are nine possible types of homogenous mixtures of solute and solvent.

Kinetic Theory of Matter

• Gaseous State: High thermal motion overcomes attractive forces, molecules move randomly within the container.
• Liquid State: Van der Waals forces lead to coherence, liquids maintain a definite volume.
• Solid State: Strong intermolecular forces cause a high degree of order, thermal motions have little influence.

Binding Forces Between Molecules

• There are two main types of forces: intramolecular and intermolecular.

Intramolecular Forces

• Intramolecular forces are chemical bonds that hold atoms together within a compound.
• Covalent bonds holding hydrogen and oxygen atoms in water molecules are examples of intramolecular forces.

Intermolecular Forces

• Intermolecular forces hold two or more molecules to one another.
• The attraction between nearby water molecules are examples of intermolecular forces.
• Intermolecular forces are based on magnetic attractions related to polarity, not chemical bonds.

Manifestations of Intermolecular Forces

• Intermolecular forces are a factor behind various phenomena, such as cohesion, adhesion, gas, liquid and solid properties, interfacial phenomena, flocculation, stabilization of emulsions, and powder compaction.
• Cohesion is the attraction between like molecules.
• Adhesion is the attraction between unlike molecules.

Intermolecular Force Types

• Van der Waals forces, dispersion forces (London forces), and hydrogen bonding constitute three main types of intermolecular forces.

Van der Waals Forces (Dipole-Dipole Forces)

• Van der Waals forces occur between molecules with permanent dipoles (polar molecules).
• Dipole-dipole interactions occur between H2S molecules.
• The strength of the dipole-dipole force is determined by molecular polarity.
• Polar compounds stick together due to dipole-dipole forces.

Dipole-Induced Dipole Forces

• Occur between polar and non-polar molecules.
• Permanent dipoles can induce electrical dipoles in easily polarized, non-polar molecules.

London Dispersion Forces

• Arise between non-polar molecules that can induce polarity in one another.
• Discovered in 1930, these are the weak electrostatic forces that cause attraction between molecules like f2, hydrogen gas, carbon tetrachloride, and benzene.

Significance of London Dispersion Forces

• London forces are sufficient to condense non-polar gas molecules into liquids and solids when they are brought closer together.

Ion-Dipole and Ion-Induced Dipole Forces

• Occur between polar/non-polar molecules and ions.
• These forces are thought to be involved in the formation of the iodide complex.
• I2 + K+I- forms K+I3- , which explains the solubility of iodine in potassium iodide solution.

Hydrogen Bonds

• Hydrogen bonds, considered as a special case of dipole forces, occur between molecules with a permanent net dipole due to hydrogen being covalently bonded to fluorine, oxygen, or nitrogen.
• Water (H2O), ammonia (NH3), hydrogen fluoride (HF), and hydrogen peroxide (H2O2) molecules exhibit hydrogen bonds.

Relative Strength of Intermolecular Forces

• Intermolecular forces are weaker than intramolecular forces.
• Dispersion forces are the weakest, followed by dipole-dipole interactions, with hydrogen bonds being the strongest.
• Resulting force order is dispersion forces < dipole-dipole interactions < hydrogen bonds.

Solvent-Solute Interaction

• Solubility is enhanced when adhesive forces between solute and solvent molecules are greater than cohesive forces between solute-solute or solvent-solvent molecules.
• Polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes, described as "like dissolves like".
• Polar solvents have high dipole moments and dielectric constants, giving them a greater ability to separate oppositely charged bodies.
• Crystalline solids have low solubility because of their stable crystalline structure and low intermolecular forces between solvent and solute.

Solutes and Solubility

• Solutes possessing polar groups (-OH) exhibit high solubility by establishing hydrogen bonds with water.
• Solutes possessing non-polar groups (-CH3, -Cl) are hydrophobic and exhibit low solubility.
• Polar solvents such as water mix in all proportions with alcohols, and dissolve sugars and polyhydroxy compounds, as well as ionic and polar solutes.

Polarity of Solvents and Solutes

• Water has a dielectric constant of 80 and dissolves inorganic and organic salts.
• Glycols have a dielectric constant of 50 and dissolve sugars and tannins.
• Methyl and Ethyl Alcohols have a dielectric constant of 30 and dissolve castor oil and waxes.
• Aldehydes, Ketones, Higher Alcohols, Ethers, Esters and Oxides have a dielectric constant of 20 and dissolve resins, volatile oils, weak electrolytes including barbiturates, alkaloids and phenols.
• Hexane, Benzene, Carbon tetrachloride, Ethyl Ether, have a dielectric constant of 5 and dissolve fixed oils, fats, petrolatum, paraffin, other hydrocarbons.
• Petroleum Ether, Mineral Oil and Fixed Vegetable Oils have a dielectric constant of 0.

Solubility Defined

• Solubility is the concentration of a substance in a saturated solution at a certain temperature.

Saturated Solutions

• A saturated solution contains the maximum amount of solute that a solvent can dissolve at a given temperature.

Unsaturated Solutions

• An unsaturated (or sub-saturated) solution contains a solute concentration below that for complete saturation at a specific temperature.

Supersaturated Solutions

• Supersaturated solutions contain a higher solute concentration than necessary for saturation at a given temperature and are not very stable.

Creating and Converting Supersaturated Solutions

• Supersaturated solutions may be made by cooling, heating, or adding a third substance.
• These solutions can be converted to stable saturated solutions through seeding with a solute crystal, vigorous agitation, or scratching the container walls.

Solubility Expressions

• Most pharmacopoeias define drug solubility by the number of solvent parts needed to dissolve one drug part.
• The USP and National Formulary express solubility by the number of mL of solvent needed to dissolve 1 g of solute.
• Boric acid USP solubility: 1 g dissolves in 18 mL water, 18 mL alcohol, and 4 mL glycerin.

Terms of Approximate Solubility

• Very soluble: less than 1 part of solvent required.
• Freely soluble: 1 to 10 parts of solvent required.
• Soluble: 10 to 30 parts of solvent required.
• Sparingly soluble: 30 to 100 parts of solvent required.
• Slightly soluble: 100 to 1000 parts of solvent required.
• Very slightly soluble: 1000 to 10,000 parts of solvent required.
• Practically insoluble/insoluble: more than 10,000 parts of solvent required.

Quantitative Expressions of Solubility

• Solubility is expressed quantitatively in terms of percentage (%)
• (% w/w) as grams of solute per 100 grams of solvent.
• (% v/v) as mL of solute per 100 mL of solvent.
• (% w/v) as grams of solute per 100 mL of solvent.

Molarity (M)

• Molarity expresses concentration as the number of moles of solute per liter of solvent.
• Molarity is measured in mol/L.
• The Number of moles of solute = weight / molecular weight.

Molality (m)

• Molality describes the concentration as the number of moles of solute per 1000 grams of solvent.
• The unit of molality is mol/g.

Normality (N)

• Normality describes how many grams of equivalent weight there are of solute per liter of solution.
• 1 N = (mol.wt. / valency) / liter.

Mole Fraction (x)

• Mole fraction is the ratio of the moles of one component (e.g., solute) to the total moles in the solution.
• Mole fraction of solute = n1 / (n1 + n2), where n1 and n2 are the number of moles of solute and solvent, respectively.

Determination of Solubility

• Add an excess substance to the solvent and stir it for several hours until equilibrium is achieved to prepare a saturated solution.
• Separate any undissolved substance from the saturated solution through filtration or extraction.
• Analyze the prepared saturated solution.

Solubility of Gases in Liquids

• The solubility of a gas in liquid is the concentration of dissolved gas in equilibrium with the gas above the solution.
• The solubility of gases in liquids is dependent upon pressure, temperature, presence of salts and chemical reaction with the solvent.

Effect of Pressure on Gas Solubility

• Henry’s Law states that, in a dilute solution, the mass of gas dissolved in a volume of liquid at a constant temperature is directly proportional to the partial pressure of the gas.

Henry’s Law

• C = kp.
• C is the concentration of dissolved gas in g/L.
• P is the partial pressure in mm Hg of undissolved gas.
• k is a proportionality constant, the solubility coefficient.

Temperature Effect on Gases

• Solubility of most gases will decrease as temperature increases because gases tend to expand.
• Gaseous solutions should be handled carefully and should be opened containers in cool environments.
• Vessels with gaseous solutions or liquids with high vapor pressure (e.g. ethyl nitrate) should be cooled before opening to reduce pressure and prevent loss.

Salting Out

• The addition of electrolytes (e.g., sodium chloride) or non-electrolytes (e.g., sugar) to gas causes reduced solubility.
• This is due to higher affinity between the solvent and the added substance compared to the original gas solute.

Effect of Chemical Reaction on Gas Solubility

• Henry's law strictly applies to slightly soluble gases that do not react with the solvent.
• A chemical reaction between a gas and solvent generally increases the gas's solubility.
• Hydrogen chloride gas reacts with water via hydrogen bonding which increases its solubility.

Solubility of Liquids in Liquids

• Vapor pressure is a factor in solubility of liquids in liquids.
• Liquid surface molecules transform to gas phase by breaking intermolecular forces.
• This process is reversible; thus, gaseous molecules may be trapped in the liquid (condensation).
• Escape (vaporization) rate equals capture (condensation) rate, establishing equilibrium vapor pressure over the liquid.

Vapor Pressure

• Vapor pressure is the pressure exerted by vapor at equilibrium.
• Condensed systems (solids and liquids) give rise to vapor pressure.
• Vapor pressure exerted by solids is typically lower than that of liquids.
• This is because solids have stronger intermolecular forces and less of an escaping tendency.
• Surface loss of vapor from liquids through evaporation is more common than that of solids by sublimation.

Vapor Pressure of Solutions containing Solid Solutes

• In solutions, both solvent and solute molecule escape.
• Intermolecular forces between solid solutes are strong, thus solid solutes generally do not escape, instead contributing to vapor pressure.

Vapor Pressure of Solutions Containing Liquid Solutes

• In mixtures of miscible liquids, molecules evaporate and contribute to vapor pressure.

Ideal Solutions and Raoult's Law

• Raoult's Law says intermolecular strengths are the same, thus solvent-solvent, solute-solvent, and solute-solute interactions are equal in the same strengths.
• Solute and solvent molecules escape because only their numbers on the surface are a factor.
• The relative numbers can be expressed by the mole fractions of the components.
• Definition is mole fraction of solute= n1/n1+n2 (n1 and n2 numbers of moles of solute and solvent).

Raoult's Law

• Each volatile constituent's vapor pressure is the pure constituent's vapor pressure multiplied by its mole fraction in the solution.
• The total vapor pressure P: P = P1 + P2 = P1° X1 + P2° X2
  • X1/X2: mole fractions of solute and solvent.
  • P1/P2: partial vapor pressures from solute and solvent.
  • P1°/P2°: vapor pressures from pure solute and solvent.

Ideal Vs Real (non-ideal solution)

• Ideal solutions obey Raoult's law.
• Ideal behaviour is expected from chemically similar components, as equal intermolecular forces allow for it.
  • E.g. benze + toluene.
  • n-hexane + n-heptane.
  • ethyl bromide + ethyl iodide.
• The majority of real solutions do not exhibit ideal behavior.
• Positive deviation from Raoult's law is when the attractive forces between solute - solvent molecules are weaker than those of solute-, solute, or solvent–solvent.
  • Results indicate components have little affinity and reduce solubility.
  • P1, P2 and P are thus greater than expected.
• Negative deviations from Raoult’s law are present if the solute and solvent have affinity for each other.
• This will increase the solubility.
• Adhesive forces are greater than cohesive forces  -ve deviation i.e., P1, P2 and P are smaller than anticipated.

Miscibility of Liquids

• Liquid- liquid systems can be two categories dependent on a substance; complete and partial miscibility.
• Complete miscibility is when substances mix with each other at any concentration i.e. alcohol and water.
• Partial miscibility will only partially mix where layers of saturated fluid is made, producing complex/conjugated phases.

Miscibility with temperature

• Partial miscibility depends on temperature and can be described by use of phase diagrams.
• Phase diagrams are graphs of temp dependant on composition measured at constant pressure.
• Phenol water system will experience positive deviation due to forces that exists between molecules in a liquid mixture.

Phase Seperation

• When temperature decreases it will leads to decrased misciblility and cauase the liquid mixture top seperate phases.
• Conjugate solutions is when every phase that is present is a saturated soluble of component with other liquiods.

Tie Lines

• Within two phase region; tie lines run across region is refered to as a system with two systems.
• All systems prepared to be drawn by tie lines will have consistant 50c and separate phases with that phase.

Critical Solution Temperatures

• The critical solution in any system is the two phase region at maximum temperature.
• phenol and water is at 66.8°C.
• Above this temp will occur and combine water mixture for one phase.

Binary Liquid System and Ternary

• When added there produce an extra ternary as the critical temp are sensative to impurities.

Three Components in Phase Equilibrium

• Consists of one partially mixed liquid
  • Water, actetic acid, CCl4.
  • Water, acetone, water.
  • Water, PEG, peppermint oil.
  • Water, alcohol, castor oil, -Loran and Guth have used a titrition to measure water alcholol and caster oil.

Area inside Triangle

• This has represented combination of components required to produce components of an equilibrium. The line that goes with AC means that an A and C contains the system with the abscense. The horizontal lines that go across the triangles will increase % of B to 100, Point x on horizontal liens indicate about % of b as the system also contain to % of A and C together.
• applying simlar arguments to similar system will get other conbines. if proceding diagram of % increases unitl u get AC percent. The system lies for AC and means 30% of it.
• With a concentration of A being, so 100 - (B + C) = 100 - (15 + 30) = 55%.

Types Water/ Benze/Alch

• Contain partial amounts of misable water where the benze and alcohol conatin a mixture with 2 phase, water when mixed causes single phase which the components mix perfectly miscible.

Diagram in Termary used

• For a study of A in four, a system contains water with glyercose A making that a one phase system.
• For hair lotions with water and ethenal and hair.

Solubility in Liquids

• It has solids are the most importants for solutions for medicine.

Factor in the Sollid in Liquids

• The use of heat causes it to absorb or dissolve, when heat increases will improve dissolving the solution with heat.
  • when heat is made it is hard to dissolve as the compound decreases with heat and is exoteric. most compounds are able the absorb the right heat it can dissolve what will be added.
  • an exhample use for it when using nacll wont increase it will dissolve in cold water.

Molecule structure within the solution

• Adding a small change for a compound to dissolve, hydrophyillc group is aded that a large imprivement dissolves for water --Ex: phenol is 100xfolded with benzene when dissolved

Weak structure

• If the solution is not a hard acid causes degree to conatin dissoiation on what dissolves with some of this: -- aque solubility of salicilic acid is 1:55 and 1.1 with the right addition. The smaller soluble it is means the larger amount is dissolved.

Particle size

• Decreasing it increase ability of dissolving means more area when added
• It also occurs when the smallest size it stops what it needs to dissolve that why -It is caused by creating electro static charger to dissolve
• A natural with solvent and solute to disolve,

Dissolvability of Polar and apolar

• When increasing the ability of soluablity helps reduce to be know by the use, also with ethanol help to dissolv,

Co-Solutes

• Using ethanol and polygene glycol helps help with mix solvent with best solutes Ex: a soluble that isn that is electro and soluble causes water from glycol

Solution with Ph scale

• most druys are not strong can can rarely disolving most water based

-Ph for sollids is what causes the druys or salts to decrease

• increase will decrease amount off ph, this use with the equations of soild 3,4, with drugs that are the base.

Mimum value

what required and is added what are there

complex foramtions

The form to add and decease addition to make the solute correct. Has been know for origin.

For Example origin

--Beta adding nico increase solubabilotu of poorly water of drugs.

• • teatra forms insoluable mix with calcium that contains for.

Solute Agents

A mixed soluables used that increase agents to dissolve small

• • surfactant are in with CMCS In aqueaous its centure it take organic and then takes mix to produce

. -in organic it is made if increase has a soluble organ

• if uses high value than what used like .3 .ex fat mixed are and mixed with some solublitiy

Common ion and electrolyte effect

• for a bit to cause electrolit the electrolyte cause by it will discribe at a certain to get products right.
• chloride with the reaction with the euations.

If to come Ag in common than it will alter .

The aciliton will increase amount the chloride with less amounts for a better result. The end add more soluable so it is a very light is good bit or the end result a complex for it to increase

Electrylyts in it as no ion that come in is created solution with no ion with the use

• it create a solute the that solution has the use of no molecule.

10 Crystal

This crystal cause other form with what has the with what is used These crystals that form are know as different with the use to change the right shape on storage