ionic equlibria, pH and buffers Flashcards
what is an electrolyte solution
an electrolyte solution is one which conducts electricity, therefore there are ions present, often aqueous solutions of ionic compounds.
what happens when ionic compounds dissolve
he ions become separated from one another by water molecules, ions are not formed when an ionic solid dissolves but become free to move apart in the presence of water
solute in strong electrolyte solution
the solute is present entirely as ions- difficult to measure equilibrium constant
solute in weak electrolysed solution
solute is only incompletely ionised in solution, i.e. some molecules survive- able to express dissociation by the law of mass action in terms of the equilibrium constant
equilibrium constant
if equilibrium lies to RHS, the numerator will be grater than 1, K will be greater than 1
if equilibrium lies to LHS, K will be less than 1
what is autoprotolysis
a reaction where one molecule transfers a proton to another molecule of the same kind is called autoprotolysis
autoprotolysis constant
Kw=ionic product or autoprotolysis constant for water
= 55.51M
autoprotolysis constant in relation to acids and alkalis
in pure water [H]=[OH]
[H]=square root Kw which is 1x10-7mole l-1
solutions with [H] more than 10-7 mol l-1 are acidic and less than this value are alkaline
what is pH defined as
the negative log (to the base of 10) of the hydrogen ion concentration
the higher the [H] the lower the pH
ionic strength
concentration of other ionic species ((and magnitude of their change) in the solution is called the ionic strength
presence of these ions tend to limit mobility of H thereby decreasing activity of H
how to measure pH
standard hydrogen electrode
Nernst equation
Nernst factor=slope factor
Nernst factor will change with temp. must calibrate pH meter before doing analysis to allow for any changes, the slope will therefore also change with temperature
pH sensor
a pH electrode is made up of; measuring electrode, reference electrode
pH is a measurement of two half cell electrode potentials; one form pH sensitive electrode and other from reference electrode
reference electrode
at the reference electrode, there is a solid/solution interface
this allows chemical reaction to occur, enables electrical current to flow through and measures pH
measuring electrode
voltage produced here depends on ionic activity of the ion
at pH 7, output in 0mV- this is called the isopotential point for a perfect electrode
point at which both electrodes give the same reading
as pH increases mV becomes more negative
as pH decreases mV becomes more positive
reference electrode
both electrodes will produce a pH measurement when connected to a measuring device
to achieve an accurate value from the measurement electrode, the reference electrode must have a constant and stable potential
consists of a silver wire coated with silver chloride immersed in an electrolyte solution
also has a salt bridge junction which enables a physical and electrical interface with solution being measured, this completes the path from the glass measuring electrode to the reference electrode
calibration of pH sensor
must calibrate pH sensor to maintain accurate readings, compensates for changes in potential within and between the measuring and reference electrodes
electrodes are matched at the factory to produce a 0mV reading in zero solution
to calibrate- check pH of pH 7 buffer solution
check slope of measuring electrode
to do this always use a buffer value at least 3pH units from zero buffer e.g pH4 should pick buffers close tot that of solution being measured
Bronsted-lowry theory
acid=proton donor
base=proton acceptor
weak acids
weak acids in aqueous solutions are not completely ionised, they are in equilibrium with the undissociated acid
as is the case for water which is a very weak acid.
dissociation of a weak acid
HA(reversible arrow) H+A
A=conjugate base
Ka=ionisation or dissociation constant
Ka=[H][A]/[HA]
If numerator > denominator, Ka is large, acid is strong
if numerator < denominator, Ka is small, acid is weak
dissociation of weak base
for any weak base A, A9reversible arrows) HA+OH
HA=conjugate acid
Kb=ionisation or dissociation constant
Kb=[HA][OH]/[A]
If numerator > denominator, Kb is large, base is strong.
If numerator < denominator, Kb is small, base is weak
how to calculate pH
for weak acids; set ub equilibrium eq and determine [H] by using acidity constant Ka then use pH=-log[H]
for weak bases use equilibrium equation to calculate pOH from Kb then convert pOH to pH by subtracting from 14
pKa
useful to express strengths of acids and bases using same term, done by considering equilibria between an acid and its conjugate base
acid and base dissociation constant related through ionic product of water
therefore a strong acid has a weak conjugate base and a weak base has a strong conjugate acid. this relationship allows the strength of acid to be expressed in terms of the dissociation constant pKa
pKa=-logKa
relationship between pKa and acid strength
since a negative log, the lower the value of pKa the stronger the acid
reaction between an acid and a base
neutralisation reaction producing an ionic compound (salt)
what is a salt
any ionic compound formed b the neutralisation of an acid with a base
usual properties of salts
ionic, water soluble, high m.p., crystalline solids, often used to increase water solubility of a drug
hydrolysis of salts
when a salt dissolves in water the compound dissociates completely to give solvated anions and cations
cation
positively charged species
anion
negatively charged species
hydrolysis
raking a bond by action of water
effect hydrolysis can have on a strong acid/strong base solution
salt dissociates to give free ions surrounded by water molecules
salt is neutral
effect hydrolysis can have on a strong acid/weak base
acidic salt
effect hydrolysis can have on a weak acid/strong base
basic salt (alkaline solution)
effect hydrolysis can have on a weak acid/weak base
equal concentrations of ions yields a neutral salt
amphiprotic salts
more complex- can function both as an acid and a base, both can donate or accept H ions in solution
how to calculate pH of an amphiprotic salt
pH=1/2(pKa1+pKa2)
what are buffer solutions used for
calibration of pH meters, to culture bacteria, control of pH of solutions for chemical reactions, administering intravenously to critically ill patients
reason why pH in pure substances changes so easily
this occurs because pure substances have no mechanism to neutralise incoming acidic or basic impurities
buffers
solutions that resist changes in pH when small amounts of acid or base are added are called buffers
what do buffers consist of
consist of a weak acid (to supply protons to a strong base) and its conjugate base (to accept protons from a strong acid) or a weak base (to accept protons from a strong acid) and its conjugate acid (to transfer protons to a strong base)
a buffer will stabilise the pH of a solution by providing a source or sink for protons
when do buffers work best
buffers work best at a pH=pKa of acid or base from which they are made
what can Henderson hasselbalch equation be used for
can be used to calculate [acid], [base] or [salt], can also be used to determine pKa of acid or base in lab, requires potentiometric titration with pH meter
how to determine pKa of weak acid and weak base
to determine pKa of weak acid; plot pH curve during a titration with strong base, identify the pH halfway to the stoichiometric point
weak base; titrate with strong acid, find pka then use pKa+pKb=pKw
buffer capacity
number of moles of strong monobasic acid or monobasic base required to produce an increase or decrease of 1 unit of pH scale
buffer capacity=no moles/change in pH
high buffer capacity
can maintain its buffering action after addition of more strong acid or base than one with a small capacity
high capacity generally when amount of base present is aboun10% more of the amount of acid
or acid about 10% or more of the amount of base
otherwise gets used up too quickly
how is the bloodstream and intracellular fluid buffered to pH of 7.4
(biological buffers)
3 main buffers; dissolved CO2, dihydrogen phosphate, protein macromolecules
amino acids
compound with COOH and NH2 groups in same molecule. all chiral except glycine. exist as internal salt or zwitterion in solution
amphiprotic, in solution, both functional groups ionise to form salt
ionised at all values of pH
isoelectric point (pl)
pl is the average of the two pKa values on either side of the zwitterion
represents minimum solubility of protein- at which migration of protein in an electrical field is slowest, used in electrophoresis to separate mixtures of proteins
%ionised for acids and bases
weak acids- %ionised=100/1+anti log (pKa-pH)
weak bases- %ionised =1+anti log(pH-pKa)
approx. rule of thumb for if pH=pKa for acidic and basic drugs
acidic-
if pH=pKa+1, drug 90%ionised
if pH=pKa+2 drug 99%ionised
if pH=pKa+3, drug 99.9%ionised
basic-
if pH=pKa-1, drug 90%ionised
if pH=pKa-2, drug 99%ionised
if pH=pKa-3, drug 99.9%ionised
how to determine acidic or basic drug
important to remember that pKa values do not tell you if a compound is acidic or basic, it is only the negative log of the dissociation constant. same rules apply for pH
acidic drugs
acidic substance that ionises to donate protons to its surrounding, structure becomes resonance stabilised through loss if hydrogen. distributes the negative charge around the anion
majority tent to be carboxylic acids
phenol groups also tent to be acidic
basic drugs
base is a proton acceptor
pharmaceutical and biological sciences are concerned mainly with behaviour of drugs in aqueous systems. in these situations, drugs are basic if they contain; a nitrogen with atom with available lone pairs of electrons e.g. amine
but not all drug which contain nitrogen are basic- only basic if lone pairs of electrons are available for reaction e.g. amides contain nitrogen but are usually neutral
pH indicators
end point of titrations determined by colour change by indicator, indicators used in pH titrations are themselves weak acids or bases, colour depends on if they are ionised or not
Important to select an indicator with an end point where
pHin=pH9stoichiometric point) equal to or plus 1
what is the end point of a reaction defined as
point at which [acid] and [base] forms of indicator are equal
predicting end point
to predict pH at end point of titration, most acid-base reactions are over when ratio of ionised form to unionised form is 1000:1
hence when pH-pKa+3 for acid of pH=pKa-3 for base
