Topic 12 - Acid-base Equilibria Flashcards
What is the definition of a Bronsted-Lowry Acid?
Proton donor
What is the definition of a Bronsted-Lowry Base?
Proton acceptor
What happens in acid-base reactions?
Protons are transferred between substances
Describe what happens in the forwards reaction and backwards reaction when HCl dissociates in water
HCl + H2O <=> Cl- + H3O+
In the forwards reaction HCl acts as an acid and donates a proton to H2O. H2O acts as a base and accepts that proton.
In the backwards reaction H3O+ acts as an acid and donates a proton to Cl-. Cl- acts as a base and accepts that proton.
What are the conjugate pairs in the dissociation of HCl in water?
HCl and Cl-
H2O and H3O+
Conjugate pairs differ by…?
H+
How would you describe the conjugate pairs in the dissociation of HCl and water?
HCl is the conjugate acid of Cl-
Cl- is the conjugate base of HCl
H3O+ is the conjugate acid of H2O
What are the conjugate pairs in the dissociation of NH3 in water?
NH3 + H2O <=> NH4+ + OH-
NH3 and NH4+
H2O and OH-
NH4+ is the conjugate acid of NH3
OH- is the conjugate base of H2O
What is the definition of a strong acid?
One that fully dissociates into its ions
What is the definition of a weak acid?
One that partially dissociates into its ions
What type of conjugate pairs do strong acids, strong bases, weak acids and weak bases have?
Strong acids - weak conjugate base
Strong bases - weak conjugate acid
Weak acids - strong conjugate base
Weak bases - strong conjugate acid
What is the definition of pH?
-log(base 10)[H+]
What value do you use to calculate the pH of a solution?
Hydrogen ion concentration
How do you calculate the hydrogen ion concentration from pH?
10^-pH
A high concentration of H+ ions means…?
A low pH
A low concentration of H+ ions means…?
A high pH
A decrease in pH of 1 on the pH scale is equivalent to…?
The concentration of H+ ions increasing by 10 fold
What is the assumption made when strong acids dissociate?
They dissociate fully into their ions, therefore:
Concentration of acid = concentration of H+ ions when dissociated
What is the definition of the ionic product of water?
The product of the molar concentrations of H+ and OH- ions at a specified temperature
What is the value of the ionic product of water equal to at 298K?
1 x 10^-14
Write the equation for the ionic product of water
Kw = [H+][OH-]
How will temperature affect the ionic product of water?
At higher temperatures the forwards reaction, H2O dissociating into H+ and OH- is favoured as the forwards reaction is endothermic. Therefore, the POE shifts to the right and the value of Kw increases. [H+] and [OH-] increase and so the pH decreases.
Vice versa for when the temperatures are lower
What are the units for the ionic product of water?
mol^2dm^-6
If you have [OH-] how do you calculate the pH of the solution?
- Use Kw to find [H+]
- -Log(base 10)[H+]
What do you need to watch out for when doing acid-base calculations?
Diprotic/Dibasic compounds
What is the definition of a diprotic/dibasic compound?
One that contains 2 potential H+ ions to donate per molecule
If 0.1 moldm^-3 of sulfuric acid dissociates in water, what concentration of H+ ions do you have?
0.2 moldm^-3
Does pure water dissociate into its ions?
Yes, but only slightly
What is the assumption made when deriving Kw?
The expression for Kc for water is:
Kc = [H+][OH-]/[H2O]
Assumption: Since [H2O] is very large and changes by only a very small amount we can consider it to be a constant
Kc[H2O] = [H+][OH-]
Kc[H2O] = Kw
For the dissociation of water, is the forwards reaction endothermic or exothermic?
Endothermic
At higher temperatures the pH of water is…?
Lower (more acidic)
At lower temperatures the pH of water is…?
Higher (more basic)
Is water always neutral?
Yes
When will a solution be neutral?
When the [H+] = [OH-]
If you add an acid or alkali to water, will that change the value of Kw?
NO, it will only change the [H+] and [OH-] in solution
What is the definition of a strong base?
One that fully ionises in solution
What is the definition of a weak base?
One that only slightly ionises in solution
Definition of strength of an acid, its Ka?
Measure of the degree of ionisation/dissociation that occurs in solution
A high Ka value means…?
Strong acid
A low Ka value means…?
Weak acid
Which three strong acids should you know?
- HNO3
- H2SO4
- HCl
Which weak acid should you know?
Carboxylic acids
Why is pKa used sometimes when plotting a graph instead of Ka?
When there is a large range of Ka values
pKa = …?
-log(base 10)Ka
A high pKa value means…?
weak acid
A low pKa value means…?
strong acid
pKw = …?
-log(base 10)Kw
Write the dissociation equation and Ka equation for the weak acid HA
HA <=> H+ + A-
Ka = [H+][A-]/[HA]
What are the assumptions that we make for when doing calculations with the Ka of weak acids?
For Ka = [H+][A-]/[HA]
- Since acid is only slightly ionised/dissociated we assume that the equilibrium concentration of acid [HA] is equal to the original acid concentration
- We assume equilibrium [H+] = [A-] so we can rewrite it as [H+]^2. This is because we ignore any H+ present int he solution due to the dissociation of water molecules, which is small
Write the Ka equation for the weak acid HA after we have followed the assumptions
Ka = [H+]^2/ acid concentration
How would the pH of a strong acid change if you diluted it 10 times?
pH would increase by 1 unit
How would the pH of a strong acid change if you diluted it 100 times?
pH would increase by 2 units
How would the pH of a strong acid change if you diluted it 1000 times?
pH would increase by 3 units
How would the pH of a weak acid change if you diluted it 1000 times?
pH would increase by 1.5 units
How would the pH of a weak acid change if you diluted it 100 times?
pH would increase by less than 1 unit
How would the pH of a weak acid change if you diluted it 10 times?
pH would increase by less than 0.5 units
Why is the change in pH different for weak acids versus strong acids?
Weak acids do not fully dissociate in solution (you have an equilibria) when water is added the equilibrium shifts to the right to oppose the change. Whilst there is a higher fraction of the acid dissociated this is countered by the dilution of the dissociated products.
Can pH be negative?
Yes
What are equimolar solutions and how can they be used?
Equimolar solutions are solutions that contain the same number of moles. They can be used to determine whether a substance is an acid, base or salt and whether it is strong or weak.
How do you determine the pH of solutions?
- Use a pH meter to measure the pH of solutions
- Calibrate it first by placing the probe into deionised water, allowing it to settle, and then adjusting the reading so it reads 7.0. Do the same with a standard solution of pH 4 and pH 10 ensuring that you rinse the probe with deionised water between readings
How do you prove that your weak acid pH decreases by 0.5 as you dilute the acid by factors of 10, 100 and 1000?
You rearrange your Ka equation:
[H+] = root(Ka[acid])
When diluting a weak acid the acid concentration decreases by 10 fold so you sub the values into the equation
How do you produce a titration curve?
Measure the pH change of a solution of acid or alkali as alkali or acid is added to it
Describe the shape of a titration curve for a strong acid-strong base neutralisation reaction
Plot of pH against volume of strong base added (could be volume of strong acid added)
Initial pH = pH of acid
Final pH = pH of base
pH gradually changes at first
pH rapidly changes between pH 3.5 and 10.5 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 7 units
Equivalence point = pH 7
Describe the shape of a titration curve for a strong acid-weak base neutralisation reaction
Plot of pH against volume of weak base added (could be volume of weak acid added)
Initial pH = pH of strong acid
Final pH = pH of weak base -> Determine using Kb (probably would not have to do this), normally between 9 or 10
pH gradually changes at first
pH rapidly changes between pH 3.5 and 7.0 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 3.5 units
Equivalence point < 7
This is because your weak base has a strong conjugate acid
Describe the shape of a titration curve for a weak acid-strong base neutralisation reaction
Plot of pH against volume of strong base added (could be volume of weak acid added)
Initial pH = pH of acid -> you can determine this using Ka and concentration
Final pH = pH of strong base
pH changes more rapidly than with SA-SB
pH rapidly changes between pH 7 and 10.5 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 3.5 units
Equivalence point > 7
This is because your weak acid has a strong conjugate base
Describe the shape of a titration curve for a weak acid-weak base neutralisation reaction
Plot of pH against volume of weak base added (could be volume of weak acid added)
pH initially changes in the same way as with S.B - W.A
pH does not change rapidly at any point so there is no vertical section of the graph
Equivalence point = depends on the relative strength of the acid and base but will be around 7
Definition of equivalence point
Point at which moles of H+ = moles of OH-
What does the volume of acid/alkali needed to be added to reach the equivalence point depend on?
Concentrations of your acid + alkali and the volume of whichever one is getting the solution added to it
What are the differences between the strong acid-strong base titration curve and the strong base-weak acid titration curve?
W.A - S.B
- Higher equivalence point
- Higher starting pH
- Graph is steeper initially as you add alkali
- The steep part is shorter comparative to S.A - S.B
When is the half equivalence point?
When you have added half the volume of acid/alkali compared to the equivalence point
For any weak acid, what is true at the equivalence point? Why is this true?
pH = pKa
Because:
HA <=> [H+][A-]/[HA]
At the half equivalence point only half of the acid has been neutralised by the strong base, half the molecules of HA have dissociated and been neutralised so [A-] = [HA]
Therefore:
Ka = [H+] at half equivalence point
Why do suitable indicators need to be chosen for neutralisation reactions?
As they can only indicate a pH within a certain range
Range of indicator needs to lie completely within the range of the vertical section of the titration curve
How do you determine the pH at the equivalence point?
- Determine moles of your acid and base
- Minus the smaller moles from the larger (which one is in excess)
- Turn moles back to concentration
- If [H+] then -log(base 10)[H+]
- If [OH-] then Kw to find [H+] then -log(base 10)[H+]
In what titration curve and which part can buffer behaviour be observed?
weak acid - strong base
- In the initial section of the curve
What is the definition of a buffer solution?
A solution that resists changes in pH when small amounts of acid or alkali are added to it
What are the two ways that buffer solution can be formed?
- Partial neutralisation
- Weak acid/base and conjugate base/acid salt
How is an acidic buffer solution made?
Mixture of weak acid and salt of its conjugate base
e.g CH3COOH and CH3COONa
How is a basic buffer solution made?
Mixture of weak base and salt of its conjugate acid
e.g NH3 and NH4Cl
Describe how a basic buffer solution works?
1st dissociation:
Weak acid or weak base that dissociates partially into its ions
2nd dissociation:
The conjugate base/acid salt that dissociates fully into its ions
- The buffer solution will contain a high concentration of the conjugate base/acid as the salt from the 2nd dissociation dissociates fully
- This means that the POE will lie very far to the left in the first dissociation
-Therefore, you have a high concentration of your undissociated weak acid/base and a relatively low concentration of H+ ions
For any buffer solution there is:
- A RESERVOIR of UNDISSOCIATED WEAK ACID/BASE MOLECULES
- A RESERVOIR OF CONJUGATE BASE/ACID IONS FROM THE DISSOLVED SALT
For a basic buffer, describe how the buffer resists changes in pH when a small amount of acid (H+) or alkali (OH-) is added?
H+:
Your weak base is partially dissociated in water
The H+ ions react with the OH- ions, low concentration of these, forming water. As you have a reservoir of NH3 and H2O this change is reversed by equilibrium 1 shifting to the right which replaces the OH- ions that reacted, keeping the pH almost constant
OH-:
Your weak base is partially dissociated in water
The OH - ions react with the large reservoir of NH4+ ions reforming your original products, so equilibrium shifts to the left. Your OH- concentration hardly changes so your pH almost remains constant
For an acid buffer, describe how the buffer resists changes in pH when a small amount of acid (H+) or alkali (OH-) is added?
H+:
Your weak acid is partially dissociated in water
The H+ ions react with the reservoir of CH3COO- ions causing equilibrium 1 to shift to the left
Your H+ ions concentration hardly changes and therefore, the pH remains almost constant
OH-:
Your weak acid is partially dissociated in water
The OH- ions react with the low concentration of H+ ions present. Equilibrium 1 then shifts to the right to replace the the H+ ions which reacted, can do this because of the reservoir of CH3COOH, therefore the pH remains almost constant
Definition of an acidic buffer?
A solution that resists changes in pH to keep the solution below pH 7
Definition of a basic buffer?
A solution that resists changes in pH to keep the solution above pH 7
What two ways can a buffer solution be made by partial neutralisation?
Weak base + strong acid
Weak acid + strong base
How does partial neutralisation work?
Partial neutralisation is when your weak acid/base is partially neutralised by a strong base/acid.
The weak acid/base is only partially neutralised because it is in EXCESS
Therefore, all the strong base/acid reacts to form salt + water but there is STILL SOME WEAK ACID/BASE LEFT IN SOLUTION
Therefore, you have your weak acid/base and their conjugate base/acid salt. The salt fully dissociates forming a reservoir of the conjugate base/acid. This causes the POE of the weak acid/base to shift very far to the left so there is a reservoir of undissociated weak acid/base
How does buffer action occur in titration curves?
Partial neutralisation:
When you have a weak acid/base and strong base/acid initially/finally in the curve you have excess of your weak acid/base and so you have a buffer solution
Do concentrations and volumes of substances matter when making a buffer from weak acid/base and conjugate base/acid salt?
No
Do concentrations and volumes of substances matter when making a buffer from partial neutralisation?
Yes
Describe how to determine pH of a buffer solution formed from partial neutralisation
Partial Neutralisation:
1. Calculate moles of acid/alkali
2. Calculate moles of salt
3. As salt is limiting reagent, moles of salt = conjugate base/acid salt
4. Calculate concentration of acid/base
5. Calculate concentration of conjugate base/acid salt
6. Ka = [H+][A-] -> salt concentration/ [HA] -> acid/base concentration
7. Rearrange for H+
8. pH
Describe how to determine pH of a buffer solution formed from the normal method
- . Ka = [H+][A-] -> salt concentration/ [HA] -> acid/base concentration
- Rearrange for H+
- pH
What are the assumptions we make during buffer calculations?
- [HA] - the original acid concentration since very little is dissociated
- [A-] - the salt concentration since the salt is fully ionised and a very small amount of A- comes from the acid
State the Henderson Hasselbalch equation
pH = pKa + log(base 10) [A-]/[HA]
Definition of enthalpy change of neutralisation?
The enthalpy change when one mole of water is formed when solutions of acid and alkali are reacted together
Are enthalpies of neutralisation exothermic/negative or endothermic/positive?
Exothermic/negative - more bonds being made than broken
For a strong acid - strong base neutralisation what is always the value for the enthalpy change of neutralisation?
-57 kjmol^-1
Will weak acid/base neutralisations be more or less exothermic than strong acid/base neutralisations? Why?
Less exothermic - This is because weak acids/bases do not fully dissociate into their ions this means that energy is used ionising the compounds first
The more weak the acid/base the … the value of the enthalpy of neutralisation?
less negative/exothermic - more energy required to ionise the acid/base
What pH does the blood need to be kept around?
pH 7.4
What buffer system is present in the blood
Carbonic acid - hydrogencarbonate buffer system
How does the carbonic acid - hydrogencarbonate buffer system in the blood behave as a buffer?
H2CO3 ions dissociate following the equation:
H2CO3 <=> H+ + HCO3-
There are reservoirs of H2CO3 and HCO3-
[H+] rises:
When the concentration of H+ ions rises in the blood the reservoir of HCO3- ions reacts with the H+ ions and the equilibrium will shift to the left reducing the H+ ion concentration to almost its original value. This stops the pH of the blood form dropping
[H+] falls:
When [H+] falls the reservoir of H2CO3 molecules oppose the change by dissociating, this means the equilibrium will shift to the right increasing the [H+] to its original value
How are the levels of H2CO3 controlled in the blood?
By respiration: as this involves breathing out CO2
H2CO3 <=> H2O + CO2
How are the levels of HCO3- controlled in the blood?
By the kidneys: excess is excreted out in urine
