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Acid-base Equilibria Flashcards

1
Q

Describe a Bronsted-Lowry acid

A

Proton donor

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2
Q

Define a Bronsted-Lowry base

A

Proton acceptor

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3
Q

Acid base reactions involve the transfer of what

A

Protons

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4
Q

In the reaction HNO3 + HNO2 ⇌ NO3- + H2NO2 + what are the conjugate acid-base pairs?

A

Acid - HNO3
Conjugate base - NO3-
Base - HNO2
Conjugate acid - H2NO2+

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5
Q

Identfiy the conjugate acid-base pairs in HCOOH + CH3(CH2)2COOH ⇌ HCOO- + CH3(CH2)2COOH2+

A

Acid - HCOOH
Conjugate base - HCOO-
Base - CH3(CH2)2COOH
Conjugate acid - CH3(CH2)2COOH2+

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6
Q

What ion causes a solution to be acidic (2 marks)

A

H+ ions release in water combining with H2O to form H3O+ oxonium ions.

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7
Q

What causes a solution to be alkaline

A

OH- (hydroxide ions)

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8
Q

What is the equation for the ionisation of water

A 
2H2O (l) ⇌ H3O+ (aq) + OH- (aq)
Or H2O (l) ⇌ H+ (aq) + OH- (aq)
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9
Q

What us the value of Kw at 298K

A

1.0 x 10^-14

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10
Q

What physical factors affect the value of Kw and how do they affect it

A

Temperature only - if temperature is increased, equilibrium moves to the right so Kw increases and the ph of the pure water decreases

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11
Q

What us pKw

A

Sometimes pKw is used instead of Kw to make numbers more manageable
pKw = -log Kw
Kw = 10^-pKw

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12
Q

Why is pure water still neutral even if pH does not equal 7?

A

[H+] = [-OH]

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13
Q

Give an expression for Ph in terms of H+

A

ph = -log10[H+]

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14
Q

Whay is the relationship between pH and concentration of H+?

A

Lower pH = higher concentration of H+

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15
Q

If two solutions have a pH difference of 1, what is the difference in [H+]?

A

A factor of 10

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16
Q

How do you find [H+] from pH

A

[H+] = 10^-pH

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17
Q

How do you find [-OH] from pH? (298k)

A

Find [H+] and then use Kw = [H+][OH-] which is equal to 1.0 x 10^-14 at 298k to calculate [-OH]

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18
Q

What is the difference when finding [H+] of diprotic and triprotic acids

A

Need to multiply the concentration of acid by the number of protons to find [H+]

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19
Q

How do you calculate the pH of a strong alkaline solution

A

Use Kw to calculate [H+] from [OH-]

Then use pH = -log[H+]

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20
Q

Define the term strong acid

A

One which fully dissociates in water (HX → H+ + X-)

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21
Q

How do you calculate the pH of a strong acid

A

pH = -log [H+]

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22
Q

Define the term strong base

A

One which fully dissociates in water (XOH → X+ + OH-)

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23
Q

What is the difference between concentrated and strong

A

Concentrated means many mol per dm3, strong refers to amount of dissociation

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24
Q

What is a weak acid and a weak base?

A

Weak acids and bases do not fully dissociate in water. They only partially dissociate into their ions

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25
Give some examples of strong acids
HCl, H2SO4, H3PO4
26
Give some examples of strong bases
NaOH, CaCO3, Na2CO3
27
Give some examples of weak acids
CH3COOH (ethanoic), any organic acid
28
Give some examples of weak bases
NH3, Zn(OH)2
29
What is Ka? (expression
For acid HA, HA ⇌ H+ + A- | Ka = [H+][A-]/ [HA]
30
How do you work out the pH of a weak acid?
Use the equation for Ka, subbing values for [A-] and [HA] | Use pH = -log[H+] to find pH
31
What is pKa
Used to make Ka values more manageable pKa = -log Ka Ka = 10^-pKa
32
What is a titration
The addition of an acid/base of known titration to acid/base of unknown titration to determine the concentration. An indicator is used to show that neutralisation has occured, as is a pH meter
33
Define the term equivalence point
The point at which the exact volume of base has been added to just neutralise the acid or the other way around
34
What generally happens to the pH of the solution around the equivalence point?
There is a large and rapid change in Ph, except in the weak-weak titration
35
How would you calculate the concentration of a reactant if you know the volume and concentration of the other reactant and the volume of that reactant
- Calculate moles of one reactant - Use balanced equation to work out the moles of the other Use concentration = mol/volume to calculate concentration
36
What is the end point?
The volume of acid or alkali added when the indicator just changes colour. If the right indicator is chosen, equivalence point = end point.
37
What are the properties of a good indicator reaction
Sharp colour change - no more than one drop of acid/alkali needed for change - End point must be the same as the equivalence point or titration gives the wrong answer - Distinct colour change so it is obvious when the end point has been reached
38
What indicator would you use for a strong acid-base titration
Phenolphthalein or methyl orange, but phenolphthalein is usually used as it gives a clearer colour change
39
What indicator would you use for a strong acid-weak base titration
Methyl orange
40
What indicator would you use for a strong base-weak acid titration
Phenolphthalein
41
What indicator would you use from a weak acid-weak base titration?
Neither methyl orange or phenolphthalein is suitable as neither gives a sharp change at the end point
42
What colour is methyl orange in acid? In alkali? At what pH does it change?
Red in acid; yellow in alkali. Changes at about pH = -4.5. Approximately the same as the pKa value
43
What colour is phenolpthalein in acid? In alkali? At what pH does it change?
Colourless in acid, red in alkali. Changes at about pH = 9-10. Approximately the same as pKa value
44
What is the half neutralisation point
When volume = half the volume that has been added at the equivalence point
45
Why is there a difference in enthalpy changes of neutralisation values for strong and weak acids
Enthalpy changes of neutralisation are always exothermic. This value for enthalpy change is similar for strong acids and alkalis becuase the same reaction occurs H+ + OH- → H2O Weak acids have a less exothermic enthalpy change of neutralistation since the energy absorbed to ionise the acid is used to break the bond to hydrogen in the un-dissociated acid
46
pH defintion
The pH of an aqueous solution is defined as the reciprocal to the base 10 of the hydrogen ion concentration measured in moles per cubic decimetre, pH=-log[H+]
47
Ka defintion
The acidic dissociation constant which measures the strength of an acid in solution, Ka = [H+][A-]/[HA] Re-arranging for pH gives, pH = pKa + log[A-]/[HA]
48
Amphoteric
A substance that can act as an acid or a base. | For example HCO3- can accept a proton and form H2O and CO2, or donate a proton and form CO3^2-
49
Monoprotic acid
Can release only one H+ upon dissociation
50
Polyprotic acid
Can release more than one H+ upon dissociation
51
Why are acid-base indicators used?
To detect when a reaction reaches its equivalence point. The indicator should be chosen so that its end point matches the equivalence point
52
Why does a pH probe need to be calibrated
So that for each pH reading the pH value is accurate
53
How do you calibrate a pH probe
Submerge pH probe in buffer solutions of three different pHs including pH7 and usually pHs around 4 and 10 pressing the calibrate button each time
54
What is accuracy
The more accurate the data the closer it is to the actual value
55
What equipment is used to carry out a titration
- A pipette and pipette filler are used to accurately measure out the volume of a reactant before transferring it into the conical flask - A burette is a controlled way to add small volumes of one reaction into the other reactant
56
How do you carry out a titration
- Once the pipette has been used to place one reactant into the conical flask, fill the burette with the other reactant and record the intial volume - Add a few drops of indicator to the conical flask - Open the burette tap and allow the reactant to flow into the conical flask, swiriling it - Close the burette tap once the expected colour change occurs, use white tile to make the colour change easier to identify - Record final burette volume - Repeat until you get concordant results and then calculate the mean titre
57
How do you carry out a titration to calculate pH
- Add 25cm3 of 0.1moldm-3 ethanoic acid solution into a conical flask with a few drops of phenolphthalein - Sodium hydroxide goes into burette - Titrate the solutions together until the mixture turns pale pink - Add another 25cm3 of 0.1 moldm-3 ethanoic acid solution into the conical flask using a pipette - Record the pH of the resulting solution
58
How to calculate Ka from titration
Using the restulting solution, you know only half the acid has been titrated so [HA] = [A-] - You can then cancel [A-] and [HA] in the Ka equation as they equal the same and so Ka = [H+] - Convert the resulting solutions pH to [H+] to give a value for the acid dissociation constant, Ka ([H+] = 10^-pH)
59
What are some sources of uncertainty in CP9
- Innacuracy of burette readings | - Difficulty identifying the exact end point
60
What are ways to overcome uncertainties in CP9
- Read the value from the bottom of the mensicus | - Use a white tile to see colour change more clearly
61
Define a neutral solution
A solution where the [H+] concentration is equal to the [OH-] and this is the case for pure water. Bear in mind this can occur at any pH
62
Conjugate acid
When a base accepts a proton, the species formed becomes the conjugate acid
63
Conjugate base
When an acid donates a proton, the species formed becomes the conjugate base
64
Can pure water conduct electricty?
Yes, it has slight electrical conductivity as it self ionises, the ionic product of water given as Kw = [H+(aq)][OH-(aq)]
65
The dissacotation constant of water, Kw, increases with water. When temperature increases what happens to water?
It remains neutral
66
An aqueous solution of ethanoic acid is gradually diluted, what does this mean?
The pH increases
67
What is the conjugate base of the acid HCO3-?
CO3^2-
68
Why are aqeous solutions of sodium ethanoate slightly alkaline?
The ethanoate ions react with water to give OH- ions
69
A solution of HCl has pH3. When it is made ten times more dilute, pH is?
4

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