Unit 3.9 - Acid-base equilibria Flashcards
What is the Lowry-Bronsted theory of acids limited to?
Aqueous solutions
Acid
An ion or molecule which can donate a proton (H+ donor)
Base
An ion or molecule which can accept a proton (H+ acceptor)
What do acids do during a reaction?
Donate a proton to become a base
What do bases do during a reaction?
Accept a proton to become an acid
Equation to represent an acid-base reaction
A —><— B+ + H+
In the reaction between water and hydrochloric acid, what acts as the acid and what acts as the base and why?
Water has 2 lone pairs of electrons that can easily accept the proton from HCl to form a coordinate bond in H3O+ = behaves as a base
HCl behaves as an acid since it’s losing its proton to the H2O
Neutralisation reaction
One substance donates a proton to another substance - often water
Equation to represent a neutralisation reaction
Acid 1 + Base 2 —><— Base 1 + Acid 2
What type of pairs form during neutralisation reactions?
Conjugate acid-base pairs
Conjugate acid-base pairs
An acid and a base which differ only by he presence or absence of a proton
An acid and a base which differ only by he presence or absence of a proton
Conjugate acid-base pairs
Alkali
A base that’s soluble in water (most contain OH-)
A base that’s soluble in water
Alkali
Show the equation for the neutralisation reaction between hydrochloric acid and sodium hydroxide + explain
HCl (aq) + NaOH (aq) —> NaCl (aq) + H2O (l)
H+ + Cl- + OH- —> Na+ + Cl- + H2O
(Cancelling similar elements)
H+ (aq) + OH- (aq) —> H2O (l)
What do we assume with strong acids and bases?
That they completely dissociated in water into their ions
Problem with using Ka to indicate the acidic strength + what is used instead
Numbers are difficult to handle
pH (easier to follow)
Ka
Acid dissociation constant
Acid dissociation constant
Ka
What is pH a measure of?
The [H+(aq)] concentration
pH equation
pH = -log10[H+(aq)]
Hydrogen ion concentration equation
[H+] = 10^-pH
How do we use [H+] = 10^-pH ?
Shift log
pH scale?
Simple and widely applicable method for measuring the acidity/alkalinity of an aqueous solution
Can we get negative pH’s?
Yes
What would a negative pH indicate?
Highly strong acid, high concentration of H+ ions
Ways of measuring pH
pH meter
Universal indicator
Pros and cons of using a pH meter to measure pH
Pros = +-0.01, accurate
Cons = needs to be calibrated against a solution of known pH
Pros and cons of using a universal indicator to measure pH
Pros = quick, convenient, paper or solution, cheap
Cons = less accurate
Good use for universal indicators to measure pH
Soil pH
What temperature is pH normally quoted for?
25 degrees Celsius
What is pH dependent on?
Concentrate
Temperature
(Of the substance)
What type of acids can their pH be calculated directly and why?
Strong (if the concentrations of the solutions are known)
They are assumed to be totally dissociated into their ions
The pH of what type of acids can’t be found directly and why?
Weak acids
Only a small fraction of the molecules are dissociated into ions
Under which circumstances can the pH of strong acids be calculated directly?
If the concentrations of the solutions are known
What do we need to remember to do with Diprotic acids?
Multiply the concentration with 2 for calculating pH
How do we work out the concentration of a specific substance from pH?
[H+] = 10^-pH
Symbol equation to represent the dissociation
Molar ratio
Strong cid
One that almost totally dissociates into its ions in solution
Example of a strong acid
Hydrochloric acid
HCl dissociating equation
HCl (aq) —> H+ (aq) + Cl-(aq)
What do acids dissociate into?
A proton and an anion
Weak acid
One that only partially dissociates into its ions in solution
Example of a weak acid
Ethanoic acid (all organic acids tend to be weak)
What type of acids all tend to be weak?
Ethanoic acids
Ethanoic acid dissociation equation
CH3COOH (aq) —><— CH3COO- (aq) + H+ (aq)
Acid dissociation constant
Ka
Ka
Acid dissection constant
What type of process is it when acid dissociates in solution?
An equilibrium process
What does each acid dissociation reaction have?
An equilibrium constant, Ka
How would you write out Ka for this reaction?
HA (aq) —> H+ (aq) + A- (aq)
Ka = [H+][A-]
————
[HA]
How do we write out the expression for Ka (the acid dissociation constant)?
Just as we were with KC with products over reactants
Ka unit
moldm^-3 (every time)
Why do stronger acids dissociate more?
The more dissociated an acid is, the more H+ ions and ions there will be, so the stronger the acid
Value of Ka for a weak acid
Low
Value of Ka for a strong acid
High
What does a high value of Ka indicate?
A strong acid
pKa expression
pKa = -log10Ka
Ka expression
Ka = 10^-pKa
What does a higher value of pKa indicate?
A weaker acid
What does a lower value of pKa indicate?
A strong acid
What type of acids have the highest pKa values?
Weak acids
Explain why HNO3 is a strong acid based on its Ka and pKa values
Very high Ka value
Low pKa value
Proof that HNO3 is a very strong acid
Even silver can react with it, and silver is usually inert
What can water be described as?
A weak electrolyte
Why is water described as a weak electrolyte?
It partially dislocates into its ions in solution
Equation for water dissociating into its ions in solution
H2O (l) —><— H+ (aq) + OH- (aq)
Enthalpy change for the dissociation of water in solution
Positive
Kw expression
Kw = [H+(aq)][OH-(aq)]
How is the Kw expression obtained?
Applying the equilibrium law and assuming that the concentration of water is effectively constant
Kc = [H+(aq)][OH-(aq)]
————————
[H2O (l)]
(Eliminating H2O)
Kw = [H+(aq)][OH-(aq)]
Kw
The ionic product of water
The ionic product of water symbol
Kw
Value of Kw at 25 Celsius
1x10^-12mol2dm^-6 (in db)
Relationship between the H+ (aq) and OH- (aq) ions in pure water
Equal concentrations
When are there equal concentrations of H+ and OH- ions?
When pure water dissociates
Concentration of H+ or OH- ions from dissociated pure water + explanation
1x10^-7 (for both - concentrations are equal)
Sqrt of Kw
What type of process it the self ionisation of water?
Endothermic
What happens to the value of Kw as the temperature increases?
Increases
When is Kw a constant value?
At a particular temperature, even though the values of H+ and OH- concentrations may not be equal
Does altering the temperature alter Kw? How?
Higher temperature = higher Kw
What happens to strong bases in water?
Completely dissociated
Example of a strong base
NaOH
Equation for NaOH (a strong base) completely dissociating in water
NaOH (aq) —> Na+ (aq) + OH- (aq)
Why does water act as both an acid and a weak base?
It’s amphoteric
When does the concentration of OH- not equal the concentration of H+?
For a strong base dissolving in water
Why does the concentration of OH= ions not equal the concentration of H+ ions when a strong base dissociates in water?
Due to the ions from the base
The concentration of which ion is equal to the concentration of the base when a strong base dissociates and why?
The concentration of OH- ions
As the base is fully dissociated
How can we work out pH for a strong base dissociating in water since the concentration of H+ is not equal to the concentration of OH- ions?
Concentration of OH- is equal to the concentration of the base
If we re given this and Kw, we can work out [H+] and thus pH
What might we have to do when attempting to work out the pH of a strong base?
Work out how it dissociates
What’s important to note about conjugate acid-base pairs?
Has to be in the order “acid-base”
The others would be base-acid conjugates
How can we work out [H+] and [OH-] with pure water?
Square root of Kw
What do we do if we asked to work out the pH of a final solution, and we only have some values for the initial volumes and concentrations?
C1V1= C2V2
Why is the ionisation of water and endothermic process?
As the value of Kw increases as the temperature increases
How do you write an expression for Ka the acid dissociation constant?
[H+][anion-]
——————
[acid]
What do you do to get the concentration of a strong base with (OH)2?
Divide the concentration by two (not multiply)
What is Ka specific for?
A specific temperature
What is the assumption made for weak acids?
[H+] = [A-]
What is the assumption [H+] = [A-] made for?
Weak acids
[A-]
Concentration of the anion of the acid
What can the equilibrium law only be applied to?
Weak electrolytes
Example of a weak electrolyte
Ethanoic acids
Why can the equilibrium law only be applied to weak electrolytes?
If the electrolytes were stronger, they would fully dissociate, so they would be no equilibrium symbol
How do we calculate the pH of a weak acid? Explain the process
Ka = [H+][A-]
————
[acid]
Since [H+] = [A-], we can think of it as being [H+]^2
So, [H+] = square root of Ka x [acid]
Then use the normal pH expression
How can we consider [H+] in weak acids and why?
[H+] since [H+] = [A-] on the top of the Ka expression
Steps for calculating the pH of a weak acid with Ka
- Balanced symbol equation
- Equation for Ka
- Insert values
Why can we assume that the concentration of a weak acid is the same at equilibrium as when undissociated?
Because the degree of dissociation is so small
What can we assume about the concentration of a weak acid at equilibrium compared to when undissociated and why?
The same
The degree of dissociation is very small
How would we calculate Ka from pH?
[H+] = 10^-pH
Ka = [H+]^2
———
[acid]
Buffer solution
A solution whose pH does not change to any appreciable extent on addition of small amounts of acid or alkali
What does a typical buffer solution contain?
A weak acid and one of its salts (an acid and its conjugate base)
If our buffers are made up of a weak acid and one of its salts, what is largely dissociated into its ions and what’s not?
Acid largely undissociated into its ions (weak)
Salt dissociated into its ions (ionic)
Example of a weak acid and one of its salts (its conjugate base) that make up a buffer solution
Ethanoic acid
Sodium ethanoate
How do we know if a buffer is acidic?
Negative pH
2 processes at work within a buffer solution
Partial dissociation of the acid that makes the buffer into its ions
Complete dissociation of the salt of the acid
In our example of ethanoic cid and sodium ethanoate making up a buffer solution, what is partially dissociating into its ions and what is completely dissociating into its ions?
Partial dissociation of ethanoic acid
Complete dissociation of sodium ethanoate
Partial dissociation of ethanoic acid into its ions equation
CH3COOH —><— CH3COO- + H+
What type of reactions are partial dissociations and what does this mean?
Equilibrium reactions - can manipulate the position
Equation for the complete dissociation of sodium ethanoate into ions
CH3COONa —> CH3COO- + H+
What is the same about the processes of the partial dissociation of the acid and the complete dissociation of the salt of the acid at work in a buffer solution?
Same anion given out in both processes
What is different about the anion given out in the processes of the partial dissociation of the acid and the complete dissociation of the salt in a buffer solution?
Much more of the anion is given out during the complete dissociation of the salt
Which equilibrium law can we apply with buffers too?
The same equation for the acid dissociation constant
Acid dissociation constant equation
Ka = [A-][H+]
————
[acid]
Which concentrations are assumed to be the same in a buffer made up of ethanoic acid and sodium ethanoate solution?
The concentration of the ethanoic acid and the undissociated ethanoic acid
The concentration of the ethanoate ions and the sodium ethanoate
The concentrations of what are large and the concentrations of what are small in a buffer solution made up of ethanoic acid and sodium ethanoate?
Large concentrations —> ethanoic acid and ethanoate ions
Small concentrations —> hydrogen ions
How is the [CH3COO-]/[CH3COOH] ratio kept steady in a buffer solution of ethanoic acid and sodium ethanoate?
The concentrations of the ethanoic acid and ethanoate ions are large and that of the hydrogen ions is small
Expression for [H+] in a buffer solution and explain exactly which values would be inputted
[H+] = Ka x [acid]/[anion]
[acid] = undissolved
[anion] = assumed to be equal to the concentration of the salt
2 assumptions in buffer solutions
[H+] is not equal to [A-] (much higher [A-])
[A-] = [salt]
Explain what happens when acid is added to a buffer solution
The system wishes to decrease the concentration of hydrogen ions using Le Chatelier’s Principle
the position of equilibrium moves to the left
Thus the pH remains constant
Explain how the concentration of hydrogen ions in a buffer solution is decreased when the position of equilibrium moves
The negative ions combine with the hydrogen ions and the majority of the added hydrogen ions are removed
Why is it possible for a buffer solution to remain at a constant pH when acid is added?
Because the negative ion concentration is large
Which equation do we write out to explain how buffer solutions work? Why?
The partial dissociation of the acid
This is the equilibrium process
Explain how buffer solutions work when alkali is added to the system
The hydroxide ions react with the hydrogen ions to form water
The system wishes to increase the hydrogen ion concentration
The position of equilibrium moves to the right
Why does the concentration of hydrogen ions in a buffer solution decrease when an alkali is added?
The hydroxide ions react with the hydrogen ions to form water
Explain what happens in the example of the ethanoic acid and sodium ethanoate buffer solution when acid is added
The ethanoate ions combine with the hydrogen ions and the majority of the added hydrogen ions are removed
Explain what happens in the example of the ethanoic acid and sodium ethanoate buffer solution when alkali is added
The ethanoic acid molecules dissociate to restore the concentration of the hydrogen ions
Can buffers reach a breaking point? How and why?
Yes
If we keep adding more acid
The ions would eventually run out
What are buffer solutions important for in the real world?
Biological systems
Industrial processes
Example of a biological system that buffer solutions are important in?
Maintaining the pH of blood in the range 7.35 to 7.5 (weak alkali)
how is the pH of blood maintained?
By hydrogencarbonate and phosphate buffers
What would happen if the pH of blood were to fall outside the optimum range? Explain
Serious health implications
E.g - acidosis, where enzymes denature. Can occur during an athsma attack due to the inability to get rid of CO2
What do many processed foods contain and why?
Buffers
So that they can safely be eaten without changing the pH of the blood
Example of an industrial process that buffers are important in + explain
Fermentation
Keeps enzyme pH at optimum
What are basic buffers made up of?
A weak base and a salt
Example of a basic buffer solution that maintains an alkaline pH
Ammonium chloride and ammonia solution
What happens in a basic buffer solution of ammonium chloride and ammonia solution?
The ammonium chloride dissociated completely, releasing all the ammonium ions
Equation for ammonium chloride dissociating completely to release all the ammonium ions
NH4Cl (aq) —> NH4+ (aq) + Cl- (aq)
Key equilibrium in the basic buffer made up of ammonium chloride and ammonia solution
NH4+ —><— NH3 + H+
Explain what happen when we add an acid to a basic buffer
Causes the position of equilibrium to shift to the left
This removes additional H+ ions provided by the acid
Explain what happens when we add a base to a basic buffer
Adding a base removes the H+ ions which causes the position of equilibrium to shift to the right to make more H+ ions
How do we explain the new pH of a solution when acids/bases are added to ammonia?
Use the key equilibrium of ammonia
NH4 —><— NH3 + H+
What do we use to prove that the pH of water is 7?
The fact that [H+] = [OH-] in pure water
sqrt kw = [H+] and use this in pH equation
What does the pH of a salt solution depend on?
The reactants used to make the salt (i.e - whether the acid or alkali are strong or weak)
What is it due to if the pH of a salt solution is not 7?
Hydrolysis
What does hydrolysis cause in a salt solution?
The pH is no longer 7
Nature of a solution made from a strong acid and a strong base
Neutral, pH 7
Nature of a solution made from weak acid and strong base
Alkaline, pH above 7
Nature of a solution made from strong acid and weak base
Acidic, pH less than 7
Example of a salt consisting of a strong acid and strong base
NaCl
Example of a salt consisting of a weak acid and a strong base
CH3COONa
Example of a salt consisting of a strong acid and a weak base
NH4Cl
In which situation is there no hydrolysis in a salt?
In a salt consisting of a strong acid and a strong base
What happens as a result of there being no hydrolysis when a strong acid and a strong base react?
The number of hydrogen and hydroxide ions remains constant and the pH is 7
In what kind of salts does hydrolysis occur?
Salts made from…
Weak acid + strong base
Or
Strong acid + weak base
Explain with the use of equations the hydrolysis of sodium ethanoate with water (a salt made from a weak acid and a strong base)
When sodium ethanoate is dissolved in water it dissociates into its ions
CH3COONa —> CH3COO- + Na+
The anion of the salt then gets hydrolysed as it reacts with water molecules as follows
CH3COO- + H2O —><— CH3COOH + OH-
Why is the final solution alkaline when hydrolysis of a salt made from a weak acid and a strong base occurs?
From the reaction of the anion with water….
CH3COO- + H2O —><— CH3COOH + OH-
The ethanoate ion is acting as a base and the water as an acid. The solution is alkaline as OH- has been released, which is a strong base. It outweighs the effect of the weak acid.
Explain the hydrolysis of ammonium chloride with water (a salt made from a strong acid and a weak base)
When ammonium chloride dissolves in water, it dissociates fully into its ions
NH4Cl —> NH4+ + Cl-
The ammonium ions are hydrolysed by the after as follows:
NH4+ + H2O —><— NH3 + H3O+
Why is the final solution acidic when hydrolysis of a salt made from a strong acid and a weak base occurs?
From the reaction of the anion with water molecules:
NH4+ + H2O —><— NH3 + H3O+
The ammonium ion is acting as the acid and the water as a base. The solution is acidic due to the formation of hydrogen ions during the hydrolysis. H3O+ being a sting acid outweighs the weak base ammonia
What is an indicator?
A substance which is used to measure the end points in titrations
What are indicators usually?
Weak acids which change their colour depending on the pH of the solution in which they are present
Equilibrium equation of methyl orange
Hme (aq) —><— H+ (aq) + Me- (aq)
Red. Yellow.
Acid solution. Alkaline solution.
What does the equilibrium of the indicator methyl orange tell us?
Undissociated acid = red
Dissociated acid = yellow
When will a methyl orange indicator turn colour from yellow to red?
When the pH is 3.7
Methyl orange indicator colour change when the pH is 3.7
From yellow to red
What colour will methyl orange appear at the exact end point?
Orange
When will the methyl orange indicator appear orange?
At the exact end point
Why does methyl orange appear orange at its exact end point?
It’s a mixture of red and yellow
What is the pH range in which an indicator changes its colour over?
2 units
pH range and colour range of the methyl orange indicator
2.7 —> 4.7
Red —> yellow
pH range and colour range of Congo red indicator
3.0 —> 5.0
Blue —> red
pH range and colour range of litmus indiciator
5.0 —> 8.0
Red —> blue
pH range and colour range of phenolphthalein indicator
8.4 —> 10.4
Colourless —> red
What happens during a titration as the alkali is added to the acid?
Change in pH
What must an indicator do at the end point of a titration is it is to be effective?
It must change colour quickly
List the reactions that can be followed successfully using an indicator to determine the end point
- Strong acid v.s strong alkali
- Strong acid v.s weak alkali
- Weak acid v.s strong alkali
Weak alkali example
NH3
Weak acid example
CH3COOH
Strong alkali example
NaOH
Strong acid example
HCl
In what type of reaction is it not possible to determine its end point?
Between a weak acid and a weak alkali
Why is it not possible o determine the end point of a reaction between a weak acid and a weak alkali?
The change in pH at the end point is small and gradual
What experiment must we do in order to be able to plot a graph of pH v.s volume?
If we have acid in a conical flask and base in a burette (or vice versa) and measure the pH for every 5cm^3 of base added, a graph of pH v.s volume of alkali added to a fixed acid volume can be plotted
Equation for the neutralisation reaction that occurs when an acid and base react
H+ + OH- —> H2O
Is the general shape of a titration curve when an acid and a base react the same?
Yes
Describe the general shape of a titration curve when an alkali is added to an acid
Slow rise
Vertical rise
Slow rise
Describe the general shape of a titration curve when an acid is added to an alkali
Slow drop
Vertical drop
Slow drop
Describe the pH change around the equivalence point
Large change in pH for only a small amount of titrant being added
Equivalence point
The point when the number of moles of acid added is equal to the number of moles of alkali
End point of a titration
The point in a titration at which the indicator changes colour
What type of end point do we need in a titration?
One that changes quickly at the equivalence point
What do we need for a colour change to occur at the equivalence point?
Only 1 drop of titrant
At which pH will the equivalence point be adjacent to in a strong acid-strong base titration curve?
7
What do different equivalence points depend on?
There being different acids or bases
What happens at the equivalence point in terms of pH?
Sudden change
In the case of adding HCl to NaOH, what is the equivalence point representing?
It’s the vertical point where nHCl = nNaOH
How do we work out the pH of the solution using a titration curve?
We take the mean number (the centre point) of the vertical line
How do we work out which indicator to use for a reaction using a titration curve?
The pH of the indicator must fit perfectly within the vertical region of the titration plot
If we used an indicator that didn’t fit perfectly within the vertical region of the titration plot, what would happen?
Outside of the region would give an undershoot (too far before) or overshoot (too far after the vertical region)
How would we work out the initial pH to plot if we had to draw a titration curve for the opposite reaction?
From the concentrations given
Compare how a titration curve for adding strong acid to strong alkali and adding strong alkali to strong acid would be different
Acid to alkali —>slow drop, vertical drop, slow drop
Alkali to acid —> slow rise, vertical rise, slow rise
Equation for adding hydrochloric acid to ammonium hydroxide
HCl + NH4OH —> NH4Cl + H2O
Describe the titration curve for a strong acid and a weak base
Sharper decrease in pH but then the gradient flattens off
Why does the gradient flatten off on the titration curve of a strong acid and a weak base?
Due to buffering
Explain how buffering is happening in the reaction between HCL and NH4OH
the salt (NH4Cl) and the base (NH4OH) are reacting together to reduce the effect of the change
What is it that causes the difference in the shape of the curve with a strong acid-weak base titration curve?
Buffering
Why is there a really slow change in pH to begin with when reacting a strong acid and a weak base compared to strong acids and bases?
There’s no buffering when reacting strong acids and strong bases
Why is the equivalence point at a pH that is below 7 when reacting a strong acid a weak base?
Since the salt formed is acidic
Describe the change in pH when reaching a strong acid and a weak base?
Rapid fall of pH to start with but the rate of fall slows down
Describe the titration curve when adding a weak acid to a strong base
pH rises quite quickly and then flattens off due to buffering
Sharp rise and then the gradient decreases markedly
When is the buffering when reacting a strong base and weak acid and why?
When the solution is acidic as it is weak
Where is the equivalence point when reacting a strong base and a weak acid and why?
pH above 7 since this is where neutralisation has occurred and the salt formed is an alkaline salt
Why sit here term “neutral point” not appropriate and the term neutralisation isn’t always helpful?
Because, for example, when adding a weak acid to a strong base, the salt formed will be alkaline and this is where neutralisation has occurred
Describe the titration curve of adding a weak acid to a weak base
The titration is rarely performed as there is no clear end-point, it is more of an inflexion at the equivalence point
There is, however, indications of buffering at either side of the equivalence point
pH of the equivalence point when adding a weak base and a weak acid
7
What is often used when reacting a weak acid and a weak base?
A pH probe
Example of a reaction where there is more than one equivalence point
The reaction between ethanedioic acid and sodium hydroxide
Equation for the reaction between ethanedioic acid and sodium hydroxide
H2C2O2 + 2NaOH —> Na2C2O2 + 2H2O
When we have a reaction where there’s more than one equivalence point, what do the different equivalence points represent?
When the first and second reactions have completely finished
Why does a reaction with ethanedioic acid have 2 equivalence points?
Ethanedioic acid is dibasic (or Diprotic) - it has 2 acidic hydrogens
How does ethanedioic acid give its different equivalence points?
It loses the acidic groups individually
What do we have to do if we have 2 different end-points for a reaction?
Would have to switch indicators - a double titration
Henderson-Hasselbalch equation
pH(buffer) = pKa + log[salt]/[acid]
What is the Henderson-Hasselbalch equation used for?
It’s an alternative way to work out buffer pHs
Which radio do we use with the Henderson-Hasselbalch equation?
The ratio of salt to acid concentrations
What is true if the concentration of the salt is equal to the concentration of the acid (Henderson-Hasselbalch equation)?
Since [salt]/[acid] = 1 and log(1) = 0
pHbuffer = pKa
When is pHbuffer = pKa true? Why?
Half way towards the equivalence point
Half of the acid will have turned into salt
Half-equivalence point
Where half of the acid has been neutralised
Which two things are true at the half-equivalence point?
Amount of salt formed = amount of acid remaining
pH = pKa
How could we work out Ka using the half-equivalence point?
The half-equivalence point would be halfway towards the equivalence point, so half of the volume at the equivalence point
We then extrapolate at this volume to find pH
Here, pH = pKa
Ka = 10^-pKa
How do we explain why buffering occurs in the buffer zone?
-write the equilibrium between the conjugate acid and base
-explain what happens to the position of equilibrium when acids or bases are added
How is the pH of a buffer solution kept constant when acid is added?
The anions react with the H+ ions to remove them from the solution to keep the pH constant
How do we get the reversible reaction for a basic buffer solution?
The salt dissociates completely, releasing ammonium ions which set up the reversible reaction
Ka in a titration plot
Use the pH at the half value for neutralisation (not half way of the vertical region) in ka = 10^-pH
Explain why we get 2 vertical regions on a titration curve of a dibasic acid
2 acidic protons = the Ka of each proton is different = 2 vertical regions
Explain the location of the 2 vertical regions for a dibasic acid
The first vertical region occurs at half the volume of the second vertical region as each occurs after removing the same number of protons
What happens when we have an indicator that doesn’t change colour within the vertical region of a titration curve?
The indicator will change colour gradually when the acid/base is added
Important use for buffer solutions
Storing enzymes at constant pH
What has to be true about an indicator for it to be suitable to use?
Its colour change needs to be complete in the vertical region
