Exam 1 Definitions and Concepts 5-7 Flashcards
drugs can be in many different ionized forms
weak organic acids
weak bases
salts
nonelectrolytes
quaternary ammonium salts
ionization affects
absorption, distribution, elimination of drugs
*not metabolism
pH partition hypothesis
partitioning into lipophilic membrane is greatly inhibited by ionic charges
*ionized will not partition, unionized will
*pH can be altered to increase excretion or absorption
3 acid base theories
Arrhenius
Lowry-Bronsted
Lewis
Arrhenius acid and base
acid= any species that can increase the concentration of H+ in an aqueous solution
base= any species that can increase the concentration of OH- in an aqueous solution
*limiting because may not always have H+ and OH-
Lowry-Bronsted acid and base
acid donates proton
base accepts proton
water is amphiprotic
Lewis acid and base
acid= electron donating
base= electron accepting
*won’t really use because mostly applies to inorganic molecules
amphiprotic
substance that can both accept and donate a proton (acts as an acid and a base)
*ex. water and amino acids
*amphoteric substances are called ampholytes
equilibrium
established when the rate of the forward reaction equals the rate of the reverse reaction
pH
negative logarithm of the hydrogen ion concentration
pH + pOH =
14 (pKw)
mole fraction
always part over total
*for HA = Ka/([H3O+]+Ka)
*for BH+ = [H3O+]/(Ka+[H3)+])
salts of weak acids and bases
weak acid + strong base or strong acid + weak base
dissociate fully to equilibrium
ionization of ordinary ampholytes
pKa of acid group > pKa of basic group
Zwitter ionic ampholytes
pKa of acid group < pKa of basic group
rules of ionization
like dissolves (dissociates) like
*acid drugs become more NON ionized in acidic pH (will absorb in stomach)
*basic drugs become more NON ionized in basic pH (will absorb in intestine)
buffer solution
solution that changes pH only slightly when small amounts of a strong acid or strong base are added
*resists changes in pH
*contains significant concentrations of both a weak acid and its conjugate base or a weak base and its conjugate acid
blood buffer system
pH of blood is 7.35-7.45
maintained by H2CO3/HCO3
buffer capacity, B = dC/d(pH) where dC is
number of moles of alkali needed to change the pH of 1 liter of solution by an amount d(pH)
hydrogen bonding
intra or inter molecular interactions between H and an electronegative atom (O,N)
polymorph
crystalline vs. amorphous
bioavailability
amount that gets to blood/organ after first pass if applicable
important pharmaceutical buffer
PBS (phosphate buffered saline)
*NaCl and Na2PO4 or KCl and KH2PO4 or CaCl2 or CaCl2 or MgSO4
acidic buffer
combination of weak acid and its salt with a strong base
*ex. HCOOH/HCOONa (formic acid and sodium formate)
basic buffer
combination of weak base and its salt with a strong acid
buffers with 2 salts
monobasic potassium phosphate (KH2PO4)
dibasic potassium phosphate (K2HPO4)
buffer action
resistance of a buffer solution to a change in pH
MOA of acidic buffers when an acid is added ex. HCl and CH3COONa/CH3COOH
hydrogen ion yielded by HCl are quickly made into acetic acid (CH3COOH), so hydrogen ion concentration does not change
*strong electrolytes (CH3COONa and HCl) dissociate quickly
MOA of acidic buffers when a base is added Ex. NaOH and CH3COOH/CH3COONa
hydroxyl ions yielded by NaOH are removed as water
*CH3COOH dissociates to H+ and NaOH to OH-
MOA of basic buffers when an acid is added
makes water, same as MOA of acidic buffers when base is added
MOA of basic buffers when base is added ex. NH4OH/NH4CL and NaOH
hydroxyl ions quickly removed as ammonium hydroxide. NH4OH
water is also made
MOA of phosphate buffers KH2PO4/K2HPO4
H2PO4- is weak acid, HPO42- is conjugate base
water is made whether acid or base is added
common ion effect
shift in equilibrium potential that occurs because of addition of ion already involved in equilibrium reaction
a solution of HCN and NaCN is more/less acidic than HCN alone
less
buffer capacity
measure of a buffer’s magnitude of resistance to change in pH
*buffer index, buffer value, buffer efficiency, buffer coefficient
*B= delta A or B / delta pH (approximate)
*solution has a buffer capacity of 1 when 1 liter of it requires 1 g equivalent of a strong acid or base to change the pH by one unit
secondary buffers in erythrocytes
oxy-hemoglobin/hemoglobin and acid/alkali potassium salts of phosphoric acid
*primary is blood
*erythrocytes + RBCs
*carbon monoxide poisoning changes blood pH
common ophthalmic buffers
borate, carbonate, phosphate
common ointment and cream buffers
citric acid/ its salts and phosphoric acid/ its salts
use of buffers in pharmaceutical systems
- adjust pH for stability
- very important for injections (want it to be 7.4)
- acetate, phosphate, citrate, glutamate often used
effect of pH on drug solubility
weak acids best absorbed in acidic stomach
weak bases best absorbed in more basic intestine
drug classifications by solubility and permeability (BCS= biopharmaceutics classification system)
Class I- high solubility and permeability
Class II- low solubility, high permeability
Class III- high solubility, low permeability
Class IV- low solubility and permeability
pHp
for a weakly acidic drug, the pH below which the drug precipitates
for a weakly basic drug, the pH above which the drug precipitates
addition of H3O+ will increase/decrease the solubility of a salt that contains the anion of a weak acid
increase
crenation
cell shrinks, hypertonic solution
hemolysis
cell bursts, hypotonic solution
solution
mixture of 2 or more components that form a homogeneous mixture
solute
dissolved agent, less abundant in solution
solvent
component in which solute is dissolved, more abundant part of solution
unsaturated solution
contains dissolved solute in a concentration below that necessary for complete saturation at a definite temp
saturated solution
one in which an equilibrium is established between dissolved and undissolved solute at a definite temperature
supersaturated solution
contains more of the dissolved solute than it would normally contain in a saturated state at a definite temperature
solubility quantitatively vs. qualitatively
quant- concentration of solute in a saturated solution at a certain temperature
qual- spontaneous interaction of 2 or more substances to form a homogeneous molecular dispersion
degree of saturation
how much solute can be dissolved in a certain amount of water at a given temp
solubility curve
any solution can be made more saturated, unsaturated, or supersaturated by changing the temperature
thermodynamic solubility of a drug in a solvent
max amount of the most stable crystalline form that remains in solution in a given volume of the solvent at a given temp and pressure under equilibrium conditions
*balance of solvent with solvent, solute with solute, solvent with solute
solubility process
mechanistic perspective of solubilization process for organic solute in water involving these steps
1. break up of solute-solute bonds
2. break up of solvent-solvent bonds
3. formation of cavity in solvent phase that can accommodate solute molecule
4. transfer of solute into the cavity
5. formation of solute-solvent intermolecular bonds
3 types of interaction in solution process
solvent-solvent, solute-solute, solvent-solute
delta Hsol = delta H1 + delta H2 + delta H3
enthalpy
overall amount of heat released or absorbed during dissolving process at constant pressure
*can be positive or negative, endothermic or exothrmic
molarity
M,c
moles of solute in 1 L of solution
molality
m
moles of solute in 1000g of solvent
normality
N
gram equivalent weights of solute in 1 L solution
mole fraction
x
ratio of moles of solute to total moles solute + solvent
percentage by weight
%w/w
g solute in 100g solution
percentage by volume
%v/v
mL solute in 100mL of solution
percentage weight in volume
%w/v
g solute in 100mL solution
USP solubility
number of mL of solvent in which 1g of solute will dissolve
polar solvents
polarity measured as dipole moment
ability to form hydrogen bonds is generally more important (and will be reflected by high dipole moment)
water dissolves phenols, alcohols, other nitrogen and oxygen containing compounds
as length of non-polar chain of an aliphatic alcohol increases
solubility in water decreases
branching increases/decreases solubility in water
increases
non-polar solvents
unable to form hydrogen bonds
dissolve non-polar solutes through weak van der Waals forces
semi-polar solvents
ex. ketones, propylene glycol
can induce a certain degree of polarity in non-polar solvents
intermediate solvents- bring about miscibility of polar and non-polar liquids
solubility product principle
Ksp, mathematical product of its dissolved ion concentrations raised to the power of their stoichiometric coefficients
barium sulfate
radiopaque contrast media
coats the esophagus, stomach, intestine with material not absorbed by body so it can be seen on an x-ray
dielectric constant
property of solvent relating to the amount of energy required to separate 2 oppositely charged bodies in a solvent as compared with the energy required to separate the same 2 oppositely charged bodies in a vacuum
*highest: water then glycols
*drugs can be polar, non-polar, semipolar
solubility of gases in liquids
when pressure above solution is released, solubility of gas decreases
as temperature increases, solubility of gas decreases
2 types of liquid-liquid systems
ones with complete miscibility (alcohol and water)
ones with partial miscibility (phenol and water)
miscibility
mutual solubility of components in liquid-liquid systems
complete miscibility
adhesive forces between different molecules (A-B) > cohesive forces between like molecules (A-A or B-B)
partial miscibility
cohesive forces of constituents of a mixture are different
ex. water (A) and hexane (B), A-A > B-B
*think oil and water
decreasing particle size _ surface area and _ solubility
increases, increases
molecular size effect on solubility of solids in liquids
larger molecules more difficult to solvate
*carbon branching increases solubility
boiling point of liquids and melting point of solids effect on solubility
aqueous solubility decreases with increasing boiling and melting point
influence of substituents on solubility
polar groups capable of H bonding impart high solubility
non-polar groups impart low solubility
ionization of substituent increases solubility
influence of temperature on solubility
if solution process absorbs energy, solubility will increase as temperature increases
if solution process releases energy, solubility will decrease with increasing temperature
influence of crystal properties on solubility
polymorphs- same chemical structure but different physical properties
amorphous- noncrystalline form
amorphous dissolves more rapidly than same drug in crystalline form
influence of pH on solubility
very important
solubility depends on degree of ionization, and this changes with pH
carboxylic acids with more than _ carbons are insoluble in water
5
*react with sodium hydroxide, carbonates, bicarbonates to form salts
fatty acids
carboxylic acids with 12-20 carbons
*soluble in solvents with low dielectric constants
benzoic acid soluble in
sodium hydroxide solution
phenol soluble in
NaOH (not water, weak acid)
most weak electrolytes not very soluble in water, but are soluble in
dilute solutions of acids
