Physical Pharmacy: States of Matter, Intermolecular Forces, and Solutions

Vocabulary practice flashcards covering states of matter, intermolecular forces, vapour pressure, solutions, suspensions, solubility manipulation, and pharmaceutical formulation concepts.

Condensed phase

A state of matter, such as a solid, in which particles are restricted to vibrations around a mean position and possess the least amount of energy among all phases.

Liquid phase

A state of matter that possesses a defined volume and more energy than a solid, where particles are in motion while still undergoing collisions.

Gas phase

The state of matter in which particles possess the most energy, move independently, and experience random but rare collisions.

Van der Waals forces

Weak intermolecular forces consisting of four main types ranked from highest to lowest strength: ion-dipole, dipole-dipole (including hydrogen bonding), dipole-induced dipole, and induced dipole-induced dipole (dispersion or London forces).

Ion-dipole forces

Intermolecular electrostatic attractions occurring between an ion and a polar molecule when the charge of the ion attracts the opposite partial charge of the polar molecule, such as $\text{Na}^+$ attracting the electronegative oxygen and $\text{Cl}^-$ attracting the electropositive hydrogen of $\text{H}_2\text{O}$ in salt water.

Dipole-dipole forces

Forces of attraction occurring between polar molecules where two permanent dipoles attract each other; this category includes hydrogen bonding.

Hydrogen bonding

A strong attractive intermolecular force occurring between a hydrogen atom covalently bonded to a highly electronegative atom such as $\text{N}$, $\text{O}$, or $\text{F}$.

Dipole-induced dipole forces

Interactions occurring when a polar molecule with a permanent dipole distorts the electron cloud of a nonpolar molecule to make it temporarily polar, explaining why non-polar substances like oxygen dissolve in water.

Induced dipole-induced dipole forces

Really weak intermolecular forces, also known as dispersion or London forces, caused by fluctuations in the electron cloud; they occur in all molecules and explain why non-polar molecules like pentane exist as liquids.

Phase diagram of water

A phase boundary plot illustrating that higher pressure leads to a lower melting point (mp) and a higher boiling point (bp), displaying a triple point at $0.01^\text{ o}\text{C}$ and $0.006\text{ atm}$ and a critical point at $374^\text{ o}\text{C}$ and $218\text{ atm}$.

Triple point

The specific temperature and pressure condition ($0.01^\text{ o}\text{C}$ and $0.006\text{ atm}$ for water) at which solid, liquid, and gas phases coexist in equilibrium, below which sublimation occurs.

Critical point

The point of equilibrium between the liquid and water vapor phase ($374^\text{ o}\text{C}$ and $218\text{ atm}$ for water) at which supercritical fluids are achieved.

Vapour pressure

The pressure exerted by molecules that possess sufficient kinetic energy to escape across the interface into the gas phase; it increases as temperature rises and decreases when intermolecular bonds are stronger.

Solution

A clear, single-phase preparation containing individual solute molecules or ions dispersed homogeneously (evenly) within a solvent, resulting in no scattering of light.

Colloid

A dispersion containing solute particles ranging in size from $1\text{ nm}$ to $1\text{ }\beta\text{m}$ ($1\text{ }\beta\text{m}$ / 1\text{ }\text{\textmu m}) that is heterogeneous on the molecular level and causes scattering of light.

Suspension

A heterogeneous mixture containing dispersed solute particles larger than $1\text{ }\beta\text{m}$ (1\text{ }\text{\textmu m}), which causes scattering of light and is generated when solute is added beyond its solubility limit.

Molality

A concentration unit defined as the number of moles of solute present within $1000\text{ g}$ ($1\text{ kg}$) of solvent ($\text{mol}/1000\text{ g}$).

Molarity

A concentration unit defined as the number of moles of solute contained within one litre of solution.

Mole fraction

A concentration ratio expressing the number of moles of solute divided by the total number of moles of all species present in the mixture (solute plus solvent).

Plasma concentration

The measurement of a drug's concentration in the blood over time, represented graphically by a plasma concentration–time profile.

Solubility

The concentration of a saturated solution at constant temperature and pressure, marking the maximum point where no more solute can dissolve.

Supersaturation

An unstable phenomenon in which the solute concentration exceeds normal solubility, commonly used to generate crystals as excess solute precipitates out to leave a saturated solution.

Henry's law

A gas solubility principle represented by $C = kP$, where $C$ is the concentration of dissolved gas, $k$ is a constant characteristic of the gas, and $P$ is the partial pressure.

Miscible

The property of two or more liquids to mix together in all proportions to form a single homogeneous mixture.

Dissolution of a solid

A three-step process consisting of: 1) removal of a solute molecule from the crystal (requires energy), 2) creation of a cavity in the solvent by breaking solvent-solvent bonds (requires energy), and 3) incorporation of the solute into the cavity with solvent shell formation (releases energy and increases entropy).

Dissolution

The kinetic rate at which a solute dissolves into solution, which increases with higher thermodynamic solubility and decreases with larger particle size (lower surface area).

Cosolvents

Secondary solvents (such as ethanol) added to water to alter its solvent properties and significantly increase the solubility of a poorly soluble drug.

Surfactants

Amphiphilic chemicals that form and maintain emulsions; they form a surface monolayer at low concentrations and aggregate into micelles at high concentrations to entrap lipophilic drugs.

Henderson-Hasselbalch equation

The equation $\text{pH} = \text{p}K_a + \text{log}\frac{[\text{A}^-]}{[\text{HA}]}$, which relates solution pH, acid dissociation constant, and the ratio of ionised to un-ionised species, demonstrating that degree of ionisation is $50\%$ when $\text{pH} = \text{p}K_a$.

Cyclodextrins

Complexing agents with a hydrophobic inner cavity and a hydrophilic exterior that increase the aqueous solubility of non-polar drug substances.

Methylcellulose

A polymeric material added to liquid formulations such as eyedrops to increase their viscosity.

Trace elements (chemical stabilisers)

Metals incorporated into pharmaceutical solutions that catalyze oxidation reactions preferentially so that dissolved oxygen reacts with the metal rather than destroying the drug.