3.9 Acid-base equilibria Flashcards by Anna-Maria Petter

What is the Lowry-Bronsted theory?

The Lowry-Bronsted theory states that acid-base equilibria involves the transfer of protons between substances and substances can be classified as acids or bases depending on their interaction with protons.

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Define a Lowry-Bronsted acid and give an example

A Lowry-Bronsted acid is a proton donor.
Example: Ammonium ions (NH4^+)

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Define a Lowry-Bronsted base and give an example

A Lowry-Bronsted base is a proton acceptor.
Example: Hydroxide ions (OH-)

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Describe the difference between a strong acid and a weak acid

A strong acid dissociates almost completely in water which means nearly all the H+ ions are released.

A weak acid only partially dissociates in water so only a small number of H+ ions are released.

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Describe the difference between a strong base and a weak base

A strong base dissociates almost completely in water so nearly all the OH- ions are released.

A weak base only partially dissociates in water so only a small number of OH- ions are released.

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Give an expression for pH in terms of [H+]

pH= -log10[H+]
Equivalently:
[H+]=10^-pH

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What is the relationship between pH and hydrogen ion concentration, [H+]?

The pH scale is a measure of hydrogen ion concentration. The lower the pH, the higher the concentration of hydrogen ions.

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What is the hydrogen ion concentration of a solution of hydrochloric acid which has a pH of 2.0?

[H+] = 10^-pH
= 10^-2
=0.01 mol dm^-3

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Give examples of strong acids and state the pH range which indicates a strong acid

Examples:
Hydrochloric acid, Sulfuric acid, Nitric acid

pH range of strong acids: 0-3

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Give examples of weak acids and state the pH range which indicates a weak acid

Examples:
Ethanoic acid, hydrogen sulfide, and organic carboxylic acid

pH range of weak acids: 4 to just below 7

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Give examples of strong bases and state the pH range which indicates a strong base

Examples:
Sodium hydroxide, potassium hydroxide, calcium hydroxide

pH range of strong bases: 12-14

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Give examples of weak bases and state the pH range which indicates a weak base

Examples:
Ammonia, methylamine

pH range for weak bases: just above 7 up to 12

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What is the acid dissociation constant, Ka?

The acid dissociation constant, Ka, is a measure of how strong an acid is in a solution.

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Give the formula used to calculate Ka for a reaction of the form
HA(aq) <—> H+(aq) + A-(aq)

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

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What are the units for Ka?

mol dm^-3

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How does the strength of an acid relate to the value of Ka?

Ka is the equilibrium constant for the dissociation of an acid: HA <—> H+ + A-

The stronger the acid, the further to the right the equilibrium lies so there is a higher concentration of products. This causes Ka to increase.

Why is Ka used to find the pH of weak acid?

Weak acids only partially dissociate in water so the concentration of H+ ions is not the same as the acid concentration (as with strong acids). This means the pH cannot be found using [H+] and so Ka is used instead.

Give the formula used to find Ka of a weak acid

For a weak acid you can assume that all the H+ ions in solution come from the acid so that [H+]=[A-] and you can assume that [HA] equilibrium = [HA] start

So for a weak acid:
Ka = [H+]^2 / [HA]

What is the difference between describing an acid/ base as ‘concentrated’ compared to ‘strong’?

‘Concentrated’ implies there are many moles per dm^3.
‘Strong’ relates to the dissociation of the substance and implies the acid/base almost completely dissociates in water.

Give the formulas used to convert between Ka and pKa

pKa = -log10(Ka)
Ka = 10^-pKa

What is the pKa of an acid which has a Ka value of 1.60 x 10^-2 mol dm^-3?

pKa = -log10(Ka)
= -log10(1.6 x 10^-2)
= 1.80 ( 3.s.f)

Give the equation for the ionic product of water

Kw = [H+][OH-]
Or equivalently:
Kw = [H3O+][OH-]

Derive the ionic product of water using the equation for the ionisation of water

In water, the following equilibrium is set up:
H2O <—> H+ + OH-

So, Kc = ([H+][OH-]) / [H2O]. Since [H2O] is very large compared to [H+] and [OH-], [H2O]Kc can be considered to be constant. Then [H2O]Kc = Kw and so Kw = [H+][OH-].

What is the pH of pure water at room temperature?

How does the pH at which water is neutral change as temperature increases?

In the equilibrium of water, the forwards reaction is endothermic and is therefore favoured when the temperature of the water is increased. So, the pH at which water is neutral (when the concentration of H+ and OH- is the same) decreases when the temperature is increased.

What is a pH curve?

Graphs which plot pH against volume of acid or base added are called pH curves and can be used to help identify the point of neutralisation of a solution.

What is the equivalence point on a pH curve?

The equivalence point is also called the end point and is the point at which the pH curve is vertical. This is the point at which the solution has been neutralised.

What is a buffer solution?

A buffer solution is a solution which is able to resist changes in pH when small volumes of acid or base are added. A buffer solution is commonly formed from a weak acid and its salt or weak base and its salt. This produces a mixture containing H+ ions and a large pool of OH- ions which helps to resist any change in pH.

How is an acidic buffer solution made by mixing sodium ethanoate with ethanoic acid?

Ethanoic acid is a weak acid so will partially dissociate so lots of undissociated ethanoic acid molecules will remain in solution. The sodium ethanoate will fully dissociate, producing lots of ethanoate ions, CH3COO-. Therefore the following equilibrium is set up which will move to counteract changes in pH: CH3COOH(aq) <---> H+(aq) + CH3COO-(aq)

Consider the buffer solution made by mixing sodium ethanoate with ethanoic acid. How does the pH change when a small amount of acid is added?

The following equilibrium is set up within the buffer solution: CH3COOH <---> H+ + CH3COO- If a small amount of acid is added, the concentration of H+ increases. Most of the extra H+ ions combine with CH3COO- ions to form CH3COOH so the equilibrium shifts to the left. This reduces the concentration of H+ to near to its original value so the pH does not change.

Explain the significance of buffer solutions in nature

Buffer solutions are common in nature in order to keep systems regulated. Enzymes in living organisms often require an optimum pH and this can be maintained by a buffer solution.

Why are buffers used in industrial processes?

Industrial processes use buffer solutions to maintain the optimum reaction conditions for large scale manufacturing.

Give an example of where buffer solutions are used in industrial processes

Buffers are used in fermentation. Buffer solutions are added before fermentation begins to prevent the solution becoming too acidic. They are also used in fabric dyeing processes and to perform chemical analysis.

Define salt hydrolysis

A reaction where one of the ions from a salt reacts with water to form an acidic or basic solution.

What needs to be considered when selecting an indicator for titration?

The indicator chosen for a titration must change colour at exactly the end point of titration. The indicator should change colour over a narrow pH range, which coincides with the vertical section of the pH curve.

Give an example of an indicator which can be used in a titration and describe its colour change

Phenolphthalein Colour at low pH: colourless Colour at high pH: Pink pH of colour change: 8.3-10 Methyl Orange Colour at low pH: Red Colour at high pH: Yellow pH of colour change: 3.1-4.4

Why should a pH meter be used instead of an indicator in weak acid/ weak base titrations?

In weak acid/ weak base titrations there is no sharp pH change so it is very difficult to successfully use an indicator to get an accurate point of neutralisation. A pH meter should be used instead since it will be more accurate and precise.