topic 12 COPY COPY Flashcards
what strength are organic acids
organic acids- weak
what is the concentration of weak acids
[HA]start = [HA]equilibrium
what is a bronsted-lowry acid
proton donor
acid donates protons to bases
what is a bronsted-lowry base
Bronsted- Lowry base- proton acceptor
bases accept protons from acid
what does acid base equilibria involve
involves transfer of protons between substances
what is the equation for acid-base equilibria
what is a conjugate acid
used to be a base but accepted protons to become an acid
what is a conjugate base
used to be an acid but donated protons to become a base
what is a strong acid
completely dissociates in aqueous solutions
what is the equation for the disocciation of an acid
HA ⇌ H+ + A-
what is a weak acid
partially dissociates in aqueous solutions
what is pH
potential hydrogen, measures concentration of H+ ions
how do you work out pH when given the conc. of H+ ions
pH= -log[H+]
how do you work out the conc. of H+ ions when given pH
what is the Ka expression
how do you calculate the Ka of a weak acid given the pH and molarity of the solution
1) write the ionic equation for the dissociation of the acid
2) write the equilibrium expression
3) determine the conc. of H+, HA, A-
4) calculate Ka by putting the numbers into the equilibrium expression
what does it mean when the Ka value is smaller
weaker acid
what does it mean if the KA value is larger
stronger acid as the conc. of H+ ions is larger
how to calculate molarity (HA) when given Ka and pH of a solution
HA= [H+] squared/ Ka
how do you calculate the pH of a weak acid given the molarity and Ka solution
[H+] = square root Ka x molarity
then do -log[H+] to find pH
how to calculate Ka from mass and pH
1) find moles of acid using mass and mr
2) find conc of acid using moles and vol
3) 10 to the power of -pH to find [H+]
4) plug values into Ka expression
what is pKa
pKa is the negative base 10 logarithm of the acid dissociation constant (Ka)
how do you find pKa when given Ka
-log(Ka)
how to find Ka when given pKa
10 to the power of -pKa
how to calculate pH when given the molarity and pKa of the solution
1) find Ka by doing 10 to the power of -pKa
2) then do square root of Ka x molarity to find conc. of H+ ions
3) then do -log[H+]
how to calculate molarity when given the pH and pKa of the solution
find Ka using pKa
find conc of H+ using pH
HA= [H+]2/Ka
what is the expression for Kw
[H+][OH-]/ [H2O]=1
therefore
Kw= [H+][OH-]
whats the units for Kw
mol2 dm-6
what is Kw constant dependent on
- Kw constant is dependent on a constant temp
- as temp increases, Kw increases
- this is because there is more energy to split the molecule apart into H+ and OH- ions
how do you find H+ from Kw
[H+]= Kw/ [OH-]
why is the pH of pure water 7
how do you find pH from Kw
[H+]= square root Kw
then do -log[H+] to find pH
why is the pH of water not acidic
neither H+ not OH- ions are in excess
what does the conc. of a strong base equal
conc. of a strong base= conc. of OH- ions
what happens when you dilute acids
reduce conc of H+ ions, increasing pH
what does the conc of H+ ions equal to
conc of H+= conc of acid
how can you compare solutions through pH measurements
use acids with the same number of moles and compare their pHs to determine the strength of the acid compared to to other acids
same can be done to determine relative strength of bases
what happens when pH increases by 1 for a strong acid
the concentration decreases by a factor of 10
what happens for every 10 fold decrease in concentration
with a weak acid the pH value increases by a factor of 0.5 for every 10 fold decrease in concentration
how to determine the Ka of an acid experimentally
- accurately weigh the acid and dissolve it in a small volume of deionised water in a beaker
- transfer the solution to a 250cm3 volumetric flask. wash the beaker several times and pour the washings into the flask. make up to the mark with deionised water
- invert the flask several times
- take a sample from the solution and place it in a small beaker
- measure the pH of the solution using a calibrated pH metre
what are the errors with the experimental procedure to determine Ka
- assumed that the equilibrium of the acid is identical to the original concentration of the acid
- incorrect measurement of pH value
- transfer errors when weighing out small amounts
what can titrations be used for
- titrations can be used to find conc. of acid or base
- in burette- acid or base has known conc.
- conical flask contains acid/base with an unknown conc. but known vol.
what could be an error for titrations
adding too much past the end point
what does it mean when the indicator changes colour
- acid and base are mixed in the right proportions
- when indicator changes colour- end point
what is the equivalence point
when the acid and base have reacted together in the exact proportions dictated by the stoichiometric equation
what does the pH of the equivalence point depend on
- depends on the combination of acid and base used
- e.g if a strong acid is titrated against a weak base, the pH will be less than 7
draw a titration pH curve for strong acid- strong base titration
- the pH falls a small amount until near the equivalence point
- after the equivalence point there is a big drop
draw a pH curve for a weak acid-strong base titration
- the curve is the same as that for a strong acid-strong base up to the equivalence point
- acid is in excess
draw the pH curve for a strong acid-weak base titration
- when acid is first added, pH starts to fall sharply
- but then the curve levels out quickly because a buffer solution has been formed
draw a pH curve for a weak acid-weak base titration
no steep section to the graph which means that its difficult to do a titration of a weak acid against a weak base using an indicator
what are the features of titration curve
- vertical part= end point/ equivalent point
- change in pH is smallest when using a weak acid and weak base together in a titration
- horizontal part below 7- pH of the acid
- horizontal part above 7- pH of base
what is the half neutralisation point
the point halfway between zero and the equivalence point
what can the half neutralisation point be used to calculate
can be used to calculate pKa of a weak acid using pH at that point
how would you calculate pKa using the half neutralisation point
at this point [HA]=[A-]
Ka= [H+]
pKa=pH
what equation can be used to represent the dissociation of indicators that are a weak acid in aqueous solution
Hln and its conjugate base ln- have different colours in aqueous solutions
how does methyl orange work with reference to the indicator dissociation equation
methyl orange
Hln= red In- = yellow
- when H+(aq) is large, the equilibrium will shift far enough to the left so that the solution will be red
- if H+(aq) is low. equilibrium will shift to the right and the yellow colour will predominate
- thats why the indicator colour changes according to the pH of the solution
when does methyl orange appear orange
there will be a stage when Hln= In- and the indicator will appear orange
how can the exact stage where HIn=In- be determined
KIn= [H+][In-] / [HIn]
how can the pH at which different indicators change colour be determined
the pH at which different indicators will change colour can be determined using their pKHln values
what does a good indicator do
- a good indicator shows a full colour change from one drop of acid from the burette
- necessary to determine accurate end point of titration
where does the pH range of indicators fall into
the pH range of each indicator falls within the steep section of the curve where a large pH change is occurring upon the addition of one drop of acid from the burette
what is the pH range of the colour change phenolphthalein
roughly pH 8-10
what is the pH range of the colour change of methyl orange
roughly pH 3-5
what colour is methyl orange in acid, base and neutral
acid- red
base- yellow
neutral- yellow
what colour is phenolphthalein in acid, base and neutral
acid- colourless
base- purple
neutral- colourless
how do you know which indicator is suitable to use
when their pH range falls within the steep section of the curve
what is a buffer solution
- a buffer solution is one who’s pH remains almost unchanged when small amounts of acid or base are added
- resists change in pH
what two types of buffers are there
acidic and basic
what do acidic buffers do
acidic buffers resist changes in pH to keep solution below pH 7
what is an acidic buffer made up of
- made from a weak acid and a salt of its conjugate base
- salt must be formed from the acid that youre using
what are the two equilibrium equations that exist in a buffer solution
- weak acid dissociates weakly so equilibrium lies to the left as the conc. of [H+] and [A-] is low (which is what the acid dissociates into
- salt dissociates fully into its positive and negative ion so equilibrium lies to the right
what happens when you add some acid to an acidic buffer
- the H+ ions react with the negative ions in solution
- there is a high concentration of these negative ions from the dissociation of the
- more acid is produced which means equilibrium shifts to the left due to the H+ ions reacting with the negative ions
what happens when you add base to an acidic buffer
- OH- ions react with the H+ ions in solution
- there is a low conc. of these OH- ions however they can be reproduced from a high concentration of acid to counteract the change (Le Chatiliers principle)
- if you dec. conc. of H+ ions, equilibrium will shift to the right to replace them
whats another way to make an acidic buffer
- acidic buffers can also be made from excess weak acid and a strong base
- all of the base reacts with the acid
- since there is an excess of acid there will still be some left even after all of the OH- ions have reacted
- in the beaker there will be a mixture of weak acid, its salt and water
- because there is some weak acid remaining, it partially dissociates to form H+ and A- ions
what do basic buffer do
resist the change in pH in order to keep the pH above 7
how can you make a basic buffer
make from weak base and its salt
how does a basic buffer work
- weak base dissociates weakly and produces very little OH- ions
- equilibrium lies to the left
- salt dissociates strongly so equilibrium lies to the right
what happens when you add a base to a basic buffer
- OH- ions react with the positive ions from the salt
- high concentration of these positive ions from the salt
- more water product is formed (water and names molecule) which means equilibrium shifts to the left
what happens when you add acid to a basic buffer
- the H+ ions react with the OH- ions in solution
- low conc. of H+ ions but they can be reproduced from a high conc. of named molecule and H2O to counteract the change (Le Chatiliers principle)
- equilibrium shifts to the right to replace the reacted OH- ions
how can buffer action be seen on a titration curve
initially pH changes quickly and there is plenty of OH- ions from the strong base to react with the H+ ions from the weak acid
- the curve then becomes more level
- a buffer solution exists between the weak acid and the salt of a conjugate base
- this resists changes in pH as more base is added so the curve is more level
what happens when all of the weak acid has been used up in a buffer
- when all of the weak acid has been used up, the equivalence point is reached and the curve rises steeply
- the buffer system doesnt exist now
- if more is added after the end point where all of the acid and base has reacted this would lead to an imbalance
how do you calculate the pH of a buffer
how can you calculate the concentrations in a buffer
- calculate pKa
- substitute the figures into the Henderson-Hasselbalch equation to find the ratio of [A]:[HA]
- to simplify we have to remove log10 by doing the inverse→ 10^ratio
- just press shift then ‘log’
- we know salts fully dissociate so [salt]=[A-]. we can calculate [HA] at equilibrium as its equal to [HA] at the start
what are the uses of buffers
maintain pH in blood
soap
how does the buffering system work in blood
- carbonic acid is controlled by respiration in cells
- when you breathe out CO2 the level of carbonic acid reduces as equilibrium shifts to the right to replace the carbon dioxide molecules
- hydrogen carbonate is controlled by kidneys and excess is removed by the them
- this allows us to eat acidic foods
how do you calculate the new pH of an already made buffer that has has aid or base added to it
Find the initial moles of HA and A- in the buffer solution
Find the moles of H+ added or removed (if OH- is added)
Using an ICE table, we can find the moles of HA and A- at equilibrium
calculate [HA] and [A-] by dividing by the total volume of the solution
calculate [H+] using [HA], [A-] and Ka
where is the half neutralisation point on a titration curve
half way between origin and equivalent point
conc of weak acid = conc of conjugate base
therefore pH= pKa
what happens when you mix an acid with water
- H= dont exist on their own in water → form hydronium ion (H3O+) but can still use H+ in equations
- water acts as a base in this reaction as it accepts a proton
what happens when you mix a base with water + give the equation
H+ ions react with bases to form OH- ions → to produce a basic solution
how is a salt made
salt is made from the metal from the base or NH4+ and the non metal from the acid (excluding hydrogen)
definition of enthalpy change of neutralisation
enthalpy change of neutralisation → the enthalpy change when acid and base solutions react together under standard conditions to form 1 mole of water
in neutralisation reactions involving weak acids and bases why is the acid and base constantly dissociating
since weak acids and bases dissociate weakly and the OH- and H+ ions are used up quickly in a neutralisation reaction, theres only a small number in solution
because of this the acid and base are constantly dissociating to replace the H+ and OH- ions that have reacted
what are the 2 types of enthalpy involved in neutralisation reactions with weak acids/bases
- enthalpy of dissociation
- enthalpy when OH- and H+ react
why does enthalpy of neutralisation vary with weak acid/base neutralisation reactions
enthalpy of dissociation varies depending on the acid/base being used up meaning that enthalpy of neut varies too
why is standard enthalpy of neutralisation similar for strong acids/ bases
- in neutralisation reactions in weak acids and bases there is no enthalpy of dissociation as they fully dissociate
- therefore standard enthalpy of neutralisation is similar for all reactions of strong acids and bases
what is the conc. of strong acids equal to
conc of H+ ions= conc of strong acid (for monoprotic acids)
conc of the acid is x2 the conc of H+ ions (multiply conc. of H+ by 2)→ diprotic. expect really low numbers
- this can then be used to calculate pH of strong acids using -log
what is a polyprotic acid
some acids can donate more than one proton → called polyprotic or polybasic
- depends on how many hydrogen atoms the molecule of acid contains
how would you calculate the pH of a base
to calculate the pH of a base use the Kw expression (still need the H+ conc)
strong base→ assume it strongly dissociates to conc. of base=conc. of OH-
what is Ka used for
used for weak acids to calculate pH
what do we assume when using Ka to determine the pH of an acid
- only a small amount of the weak acid (HA) dissociates so we can assume:
[HA]equilibrium=[HA]start
- the dissociation of acid is greater than the dissociation of water present in the solution so we can assume :
- all of the H+ ions come from the acid
- [H+]=[A-]
what are the units for Kw
mol2 dm-6
why do we assume water has a constant conc. value
- water dissociates into its ions very weakly
- we assume that water has a constant concentration value because theres so little OH- and H+ ions compared to H2O molecules
ow do you calculate pKw from Kw
-log[Kw]
how do you calibrate pH meter
put it into distilled water first. the reading should be 7.0
