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Chemistry > Acids > Flashcards

Acids Flashcards

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1
Q

Definition of a strong acid

A

Fully dissociates in water

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2
Q

definition of a weak acid

A

Partially dissociates in water

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3
Q

Define pH

A

-log([H+])

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4
Q

Define [H=] in terms of pH

A

10^-pH = [H+]

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5
Q

Define Ka for a weak acid

A

The acid dissociation constant of a weak acid - high Ka value means higher [H+] concentration, so a stronger acid

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6
Q

Define pKA

A

-log(Ka) - a logarithmic scale that makes the use of a large range of orders easier to handle. Low pKa means high Ka, meaning stronger acid

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7
Q

Define pH of a base

A 
[H+] = Kw/[OH-]
-log([H+] = pH
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8
Q

Define buffer solution

A

A solution that opposes changes of pH on addition of acids or bases. It is a mixture of a weak acid and the anion of the salt of the same weak acid, such that the acid and the anion are in high excess

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9
Q

What happens when an acid is added to a buffer solution?

A

The H+ ion concentration increases, moving the equilibrium to the left to minimise the change, causing the H+ ion concentration to decrease

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10
Q

What happens when a base is added to a buffer solution?

A

The oh- ions react with the H+ ions to form water. This decreases the concentration of H+ ions, causing more of the acid to dissociate into H+ and A-. This causes the equilibrium position to shift to the RHS.

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11
Q

What is the pH of water at a given dissociation constant?

A 
[H+] = sqrt(Kw)
-log[H+] = pH
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12
Q

Even if water has a pH different to 7, is it still neutral?

A

Yes, because every H+ ion that has dissociated has an OH- ion that has dissociated, causing the H+ concentration to be the cancelled out by the OH- concentration

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13
Q

Who was the first to give a definition for acid? What did they think give a substance acidity?

A

Lavoisier. He decided oxygen was the key to acids. Oxygen meaning acid giver.

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14
Q

What is the most useful definition we use today?

A

Bronstead-Lowry.

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15
Q

Wive equations for the multi-step dissociation of H3PO4

A

H3PO4 H2PH4- + H+
H2PO4- HPO4(2-) + H+
HPO4(2-) PO4(3-) + H+

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16
Q

What is the ionic equation for the reaction of a carbonate with acid?

A

Ca(2+) + 2H+ + CO3(2-) –> Ca(2+)(aq) + CO2 + H2O

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17
Q

What is the ionic equation for the reaction of an acid with a base?

A

H+ + OH- –> H2O

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18
Q

What is the equation for the reaction of a metal oxide with acid?

A

MgO + H+ –> Mg(2+) + H2O

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19
Q

What are acid-base conjugates and what links them?

A

Two species linked by loosing H+ e.g. H3PO4 –> H2PO4-

H3PO4 is the acid, and H2PO4- is its conjugate base

20
Q

What actually happens when H+ is in solution?

A

It does not immediately form protons. The hydrogen ions only leave if there is a bond to be made. Normally, this is hydroxonium. HCl + H2O –> H3O+ + Cl-

21
Q

What is the advantage of pH as opposed to H+ concentration?

A

pH has a smaller range of values, as it is a log scale. A change of 1 on the pH scale indicates a change by factor 10 on the h+ concentration scale. Easier to deal with these values.

22
Q

What is the equation for [H+] in a weak acid?

A

[H+] = sqrt(Ka*[HA])

23
Q

What assumption can we make about strong acid concentrations?

A

[HA]=[H+]

24
Q

What assumption can we make about weak acids when they dissociate?

A

[H+]=[A-]

25
Q

What do you need to remember when working with di-or tribasic acids?

A

There are 2/3 moles H+ produced for every mole of acid that dissociates. This may mean multiplying values to calculate pH.

26
Q

Where does the equilibrium lie in the dissociation of water?

A

Well to the left hand side. Around 10^-130 molecules actually dissociate.

27
Q

How do you calculate the concentration of water in a solution?

A

1000g in 1dm^3 H2O

n=m/Mr = 1000/18 = 55.6

c=n/v =55.6/1 = 55.6 moldm^-3

28
Q

What does the pH of a base depend on?

A

The dissociation of OH- ions

29
Q

What can we assume about the concentration of bases?

A

[NaOH] = [OH-]

30
Q

What are the five stages to the perfect buffer solutions long answer?

A

1) Buffers resist change to pH on addition of H+ or OH-
2) They consist of a mixture of a weak acid and its conjugate base. These are in large excess
3) The overall equation of the buffer is:
CH3COOH H+ + CH3COO-
The conjugate base is formed by the dissociation of NaCH3COOH to form Na+ (spectator ion) and CH3COO-, which is included in the buffer mixture.
4)On the addition of H+, it reacts with CH3COO- to produce CHeCOOH,. This pushes the equilibrium to the left hand side.
5) ON addition of OH-, it reacts with H+ to form H2O. THis reduces the concentration of H+ in the buffer, moving the position of equilibrium to RHS to oppose the decrease in H+ concentration.

31
Q

What is the alternate method of formation of a buffer?

A

Massive excess of acid reacts with base.
CH3COOH + NaOH –> CH3COONa + H2O.

The salt then fully dissociates, and forms a buffer with the massive excess of CH3COOH.

32
Q

Why is a buffer solution formed when NaOH is mixed with methanoic acid? What must be in excess?

A

The acid must be in excess. The acid reacts with the base to produce salt. The salt dissociates fully into the conjugate base of the acid, and hence the buffer forms because the acid is still in high concentration as there is an excess.

33
Q

Does a buffer system stop a change to pH or minimise it?

A

Minimises it for addition of a small amount of acid/base

34
Q

To show you are dealing with a weak acid using a pH probe, what pH would you expect?

A

pH>1

35
Q

Where is the equivalence point in an acid-base titration curve?

A

In the middle of the vertical section

36
Q

Describe the change in gradient due to the addition of base to acid in a titration

A

Initially, the pH is low, and increases slowly. As it reaches the equivalence point, the gradient increases until it is vertical at the equivalence point. This sharp change in pH is caused by the addition of only one drop of alkali. After the equiv point, the gradient decreased, but the pH is still increasing as the base is in greater excess.

37
Q

What is the end point of a titration?

A

The point at which the concentrations of weak acid and conjugate base are equal. The colour is half way between the colours in acid and base.

38
Q

What dictates the colour of an indicator?

A

It must form different coloured solutions when in acid form and in conjugate base form. It behaves in an equilibrium like a buffer system responding to acid and base addition.

39
Q

How do you choose an indicator?

A

Make sure the end point is in the range of the equivalence point.

40
Q

Why can you not titrate weak acid and weak base?

A

The equivalence point is gradual - there is no vertical section where pH increases sharply, so the colour change would be gradual, hence there is no good indicator to use.

41
Q

What is the enthalpy change of neutralisation?

A

The enthalpy change that accompanies the neutralisation of one mole of aqueous acid with one mole of aqueous base to form one mole of water under standard conditions.

42
Q

How do you calculate the enthalpy change of neutralisation?

A

Find the energy released, using Q=mct. Use a negative sign if the temperature of the surroundings increases. Then find the moles that react. Then find the enthalpy change by dividing the energy change by the number of moles.

43
Q

What are the two equations for calculating the pH of buffers?

A

[H+] =Ka *[HA]/[A-]

Or Henderson-Hasselbach

pH = pKa + log([A-]/[HA])

44
Q

How do you calculate the pH of 2.35E-3 moldm-3 Ca(OH)2?

A

[H+]=1E-14/2(2.35E-3)

2 moles of OH-

45
Q

What is half neutralisation, what is the equation that is valid at this point?

A

[HA]=[A-]

=>Ka=[H+]
=>pKa=pH

46
Q

Why might the pH of H2SO3 (weak acid) be lower than calculated using methmatical method?

A

Small amount of secondary dissociation.

Chemistry (16 decks)

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