CHEM111: Module 10 Acid-Base Equilibrium

how can you measure the strength of acids & bases (other than pH)

• use the pKa/pKb scale

• this = -log10(Ka), providing values on a 1-10 (Strong-weak) scale

• Ka/Kb is the equilibrium constant, so higher Ka=more products=more dissociation=stronger acid/base (so -log of it gives a pKa scale of strong-weak)

how does the equilibrium constant (Ka/Kb) take form for acids/bases

why is this relevant

what can we use this to find

• K = [products] / [reactants]

• Ka = [H+] [A-] / [HA]

• Kb = [OH-] [B+] / [BOH]

• this is relevant as high Ka=lots of products=lots of dissociation=strong acid/base

• we can use this to find pKa (-log) & pH (rearrange)

how do you calculate pH from Ka (For both weak & strong acids)

• pH is -log[H+], so we want to use this to find [H+]

• Ka = [H+] [A-] / [HA]

• Weak acids =

  • we know [A-]=[H+] (donating 1H+ forms 1A-), so Ka becomes = [H+]² / [HA]

  • we can assume [HA]_equil=[HA]_{time 0}

  • then rearrange and [H+] = sqroot Ka x [HA]

• Strong acids =

  • complete dissociation of HA into [H+] & [A-]

  • so [H+]=[HA]

• then just take the -log, and pH is provided

who came up with pH and when?

what is pH

• Soren Sorensen (1800s) said that [H3O+] = 10^pH

• he said a pH=7 is neutral, while <7=acidic, and >7=basic

• therefore pH is the -log of [H3O]+, providing a way to measure acidity / basic of a solution

what did Arrhenius suggest about acids / bases

and when?

• 1884

• he suggested acids deprotonate to produce H+ in aqueous solutions

• and bases protonate, to leave behind OH- in solutions

• and these can be complete or incomplete (equilibrium) dissociates (weak/strong bases / acids)

what did Bronsted & Lowry suggest about acids / bases

and when?

• 1923

• they went a step further from Arrhenius and suggested acids are proton donors, while bases are proton acceptors

• thus broadening the definition, suggesting acid/base reactions involve transferring a proton (acid→base transfer)

what did Lewis suggest about acids / bases

and when?

• 1923

• he suggested acids accept electron pairs, and that bases donate electron pairs (opposite to proton transfer) - producing a base adduct (species formed from this)

• this extended the definition to look at electron movement instead of proton transfer

• acids / bases described this way are ‘lewis acid’ and ‘lewis base’

whos definition of acids / bases is most commonly used today

• the Bronstead & Lowry definition, where acids are proton donors while bases are proton acceptors, and acid/base reactions involve proton transfer from acid→base

what are acids

how do they behave in aqueous (w/ Water) solutions)

• acids are species that deprotonate, they dissociate to produce H+

• this is usually from a breaking of an -O-H bond (except for H halides, where this H-halide bond is broken) - leaving behind an -O- on the molecule

• this H+ is attracted by water’s negative dipole (on O), forming H3O+ (hydronium ion)

what is the difference between strong and weak acids

give examples of each type

• strong acids completely dissociate / are completely ionised, into the acid ion and H+, leaving behind no original acid

• e.g. HCl, HBr, HI, HNO3, HClO4, H2SO4

• weak acids only partially dissociate / are ionised into the acid ion and H+, so some original acid remains, and an equilibrium mixture of species is formed

• e.g. HF, most organic acids, H3PO4

what is an example of a molecule that behaves like an acid, but is not truly an acid

• metal ions in aqueous solutions are surrounded by hydration shells (associations with water molecules)

• these can produce acidic solutions, as additional water molecules may interact with the hydration shells to form stronger H interactions

• therefore accepting an H+ and the solution becomes more acidic (more H3O+ produced)

what are bases

give examples of weak and strong bases

• bases are species that protonate, they dissociate to produce an ion and OH-, as they accept an H+ from water

• strong bases include: NaOH, KOH, LiOH, Ca(OH)2, Ba(OH)2, Sr(oh)2

• weak bases include: NH3, amines (R-NH2), CO3²-)

what are conjugate acid/base pairs

• these are the species formed when an acid or base dissociates, which can then go on to act as the opposite (acid → conjugate base, base → conjugate acid)

• this is due to acids deprotonating, so they can then protonate, and bases protonate, so they can then deprotonate

• e.g. H2O (base) protonates to form H3O+, which can then act as an acid and deprotonate

• strong acids / bases form weak conjugates and vice versa, due to Ka increasing as Kb decreases and vice versa

what is an amphiprotic species

• a species that can either donate or accept a proton

• e.g. H2O, can donate a proton and become OH-, or accept and become H3O+

how does the example of corals dissolving due to climate change, relate to acids?

• more CO2 in atmos → more dissolves into carbonic acid (Weak) in the sea → this dissociates to form hydrogen carbonate + H+

• this hydrogen carbonate reacts with CaCO3 structures of coral → CaCO3 dissolves as it protonates to form aqueous calcium hydrogen carbonate

• this also forms CO2 as a byproduct, which then adds to the increasing CO2 and the cycle continues

what is the Henderson Hasselbalch equation?

what can we use this to find?

• pH = pKa + log ([A-] / [HA])

• we can use this to find pH, pKa, etc

• depending on the information provided to us

what is a polyprotic acid

provide an example

• these can deprotonate multiple times, donate multiple protons

• e.g. H2SO4 → HSO4- → SO4²-

what is a buffer solution

• what is it made of

• what does it do

• a solution that contains a weak acid, and its conjugate base

• it is resistant to pH change when a strong acid / base is added, therefore acting as a pH buffer

how do buffer solutions work?

• these are solutions of a weak acid and its conjugate base, in equilibrium

• adding a strong acid (adding H+) will shift equilibrium to the left, to use up H+ and counteract the addition, therefore resisting a pH change as H+ still doesn’t increase

• adding a strong base (adding OH- and decreasing H+ as it reacts to form H2O) will shift equilibrium to the right, to make more H+ and counteract the removal, thus resisting a pH change as H+ still doesnt decrease

• eventually the added strong acid / base will change pH (buffer action stops) when equilibrium shifts too far (Weak acid conc. decreases too much)

how do you find the pH of a buffer solution

how do we determine a buffer to maintain a certain pH

• the pH will be close to the pKa of the weak acid

• so use the Ka for the buffer system to calculate the pKa, and this will resist the pH change of pH close to this pKa

• to find the exact pH, use the Henderson Hasselbalch Equation (pH=pKa+log( [A-]/[HA] )

what happens if you dilute a buffer

• diluting a buffer does not change the ratio of [H+][A-] : [HA] (used to determine the Ka) so Ka won’t change, therefore pH won’t change

• eg water entering cells (body’s buffer)

how are buffer systems involved with red blood cells in the body

• red blood cells contain haemoglobin, which pick up O2 at the lungs, transfer it to tissue, then remove the CO2 from the tissue

• the bicarbonate buffer regulates this transfer, as CO2 dissolves in the cell’s cytoplasm → forms carbonic acid → releases H+ → associated with haemoglobin & 2O → cause shape change and O2 release

• at the lungs, the reverse occurs, depleted haemoglobin associates with O2 → CO2 pops off the carbonic acid → CO2 release to be breathed out the lungs

• therefore CO2 removal enabling O2 release

what is an indicator

describe an example in the natural world

• solutions that change colour due to changes in pH, each with different pH ranges (ranges where they change colour)

  • e.g. hydrangeas change colour based on the pH of the soil they are in, acidity (and Al present in soil) will turn them blue

how does phenolphthalein indicator work

• this is a weak acid, which is diprotic (can release 2H+ to form a dibasic ion), so in equilibrium between the two

• is colourless when protonated, is bright magenta when deprotonated (when dibasic)

  • adding a base (OH-) will remove H+ from the acid to form water, causing equilibrium to shift to the right (form more H+ & base) so magenta indicates basic

  • adding an acid (H+) will favor the reverse to use up the H+ (form more acid) so colourless indicates acidity

• 8.3-10 pH range

how does 2-methyl Orange indicator work?

• weak acid in equilibrium with its conjugate base

• red in acidic form, orange in basic form

  • when acid is added (H+) reverse is favored, conjugate base is protonated, and solution turns red

  • when base is added (OH- reacts to remove H+ and form H2O) the forwards is favored, more conjugate base formed, so turns orange

• pH range 3.1-4.4

describe the carbonic acid / hydrogen carbonate buffer system

• CO2 + H2O <==> H2CO3 <==> H+ + HCO3-

• so adding CO2 to water forms H2CO3 in equilibrium with H+ and HCO3- (its conjugate base)

  • adding an acid (add H+) will shift to the left, using up H+ to resist pH change

  • adding a base (adding OH- to remove H+) will shift to the right to form more H+ and resist pH change

what is another name for titrations in the lab

name the equipment used, and each of their purposes

Volumetric analysis

equipment used includes

• burette (various volumes and classes for accuracy, filled with base which it allows to be gradually added into an acid beaker)

• volumetric flask (to make up a solution to a certain volume)

• pipette (glass, measures specific amounts of chemicals)

• beaker

• indicator solution (indicates near end point)

• pH meter (glass electrode in the beaker being titrated on to measure pH)

what does a titration curve display

how do these differ between titrations of (strong acid/strong base) and (weak acid/strong base)

• plotted pH values of the titration, against the volume of base added, including indication of the equivalence and end point

• strong acid / strong base:

  • low pH at the start due to strong acid, which rises slowly

  • until reaching equivalence point (Steepest part of the graph) at pH7, where it steeply rises, as it neutralises

  • then afterwards it flattens off at a high pH (strong base)

• weak acid / strong base:

  • higher pH at the start due to weak acid, which rises slowly as it resists pH change (buffer region)

  • this contains half equivalence point (Veq/2) where pH=pKa ([HA]=[A-])

  • it then rises steeply, reaching equivalence point (steepest part), at a pH higher than 7

  • it then flattens off at high pH just like strong acid curve (due to strong base)

in an acid/base titration, what is the difference between an end point and an equivalence point?

how do we get these as close as possible, and why is that important?

• an end point refers to when the indicator changes colour, to signify a certain pH has been reached within the solution

• an equivalence point refers to when number moles acid = number of moles base (or according to reaction stoichometry), where neutralisation occurs of OH-&H+ to form H2O

• it is important these are close as possible, so we can figure out the pH of the equivalence point (differs from pH7 for weak acids/strong bases), and use this to figure out unknown concentrations

• this is done by choosing an indicator with a pH colour change range, over the equivalence point, by ensuring its pKa is close (the point where indicator is in middle of colour change, [HA]=[A-]=[H+]=Ka, so pKa=pH)

how do equivalence points differ between

• strong acid / strong base titrations

• weak acid / strong base titrations

strong acid / strong pace

• equivalence point will be pH of 7

• this is because both species fully dissociate into OH- and H+, therefore at this point there is an equal balance of [OH-] and [H+], which will react and neutralise to form water

• this results in a neutral pH of 7

what is a half equivalence point

what can we use this to find

• only in weak acid/strong base titrations, the part of the curve where V = Veq/2

• this is also where pH=pKA ([HA]=[A-]), so we can use this to find these

why is the equivalence point at a higher pH for weak acid x strong base titrations (compared to strong acid x strong base Eqiv=pH7)

• the weak acid dissociates into a strong conjugate base, which forms OH- and makes the pH more basic (higher at equivalence point)

• base is stronger, so it ‘wins’ and pulls equiv pH to a more basic value

how to find pOH

how does this differ from pH

pOH is -log[OH-], so the opposite of pH

so to find, pOH = 14-pH

what is the expression for Kb

[OH-] [HB+] / [B]

what is a common approximation we make for titration calculations, for concentrations of acids / bases at equilibrium?

• for equilibriums where K is very small (equilibrium far to the left), we can assume that the [acid or base] at time 0, is the same as the [acid or base] at equilbrium

• therefore we can use it to rearrage the equilibrium expression to find unknowns

how would you find concentrations of weak acids / bases after the titration has progressed some amount (not using the approximation that the conc is the same as time 0)

• look at the amount of titrant added, and find the n for this (use V and conc), which will provide the n(H+) or n(OH-) (strong so it fully dissociates)

• at equivalence point, this will equal the number of mols of conjugate formed (as pKa = pH, as [HA]=[H+])

• then do n(weak at time 0) - n(conjugate), which will equal the n(weak) at the equivalence point

• this can also be used to find pH or pOH

what would a titration curve look like for a polyprotic acid

• will deprotonate multiple times, forming Ka1 and Ka2 (multiple equilibrium constants)

• so graph will have multiple changes

• e.g. carbonic acid

• e.g. sulphuric acid

why does pH =pKa at half equivalence point?

• because at half equivalence point, (HA) = (A-), so in the Ka expression they cancel out, and Ka simply becomes =(H+)

• pKa is -log (Ka), and pH is negative log of (H+), therefore they equal