AC18: pH Flashcards
How do you calculate pH when you know the concentration of H+ ions
pH = -log₁₀[H+(aq)]
How do you calculate the concentration of H+ ions when you know the pH
[H+] = 10^-pH
What is Kw, the expression for it, and under what conditions this expression is valid
Kw = ionic product of water
Kw = [H+][OH-] = 10^-14 mol^2 dm^-6 at constant temperature (298K)
How can you use Kw to calculate [H+]
(10^-14) / [OH-]
What is the general Kc expression for acids and alkalis
Kc = [H+][A-]
———-
[HA]
What do you assume when calculating pH of a weak acid
[HA] at equilibrium is the same as the initial [HA]
[H+] from ionisation of water is negligible
What is Ka and when do you use it
Ka is the acid dissociation constant, only for WEAK ACIDS, as each HA ionises to give equal numbers of H+ and A-
Kc = [H+][A-] = [H+]^2 = Ka
———- ———
[HA] [HA]
How can you rearrange Ka to get the [H+]
[H+] = root ( Ka[HA])
How do you calculate pKa, and why is pKa used instead of Ka
pKa = -log10Ka
pKa is used because Ka values can be very small and have a large range, therefore difficult to compare or plot a graph
How do you calculate Ka from pKa
Ka = 10^-pKa
What is a limitation of using Ka for calculating pH of very weak acids
having assumed that ionisation of water is negligible, at 25℃, the dissociation of water is 10^-7 moldm^-3, pH = 7. So if acid is very weak (pH>6.5), then [H+] from the acid would be comparable to that from water, this assumption would be invalid.
How can you tell from an equation, if an acid/alkali is weak or strong
if it is weak then reaction will be reversible, if it is strong then reaction is not reversible
What is a limitation of using Ka for calculating pH of ‘stronger’ weak acids
having assumed that [HA] at equilibrium is the same as [HA] at start, a stronger weak acid would ionise significantly and [HA] at equilibrium would not be the same as [HA] at start
What is the function of a buffer solution
they resist changes in pH when small amounts of acid/alkali are added
What does a buffer consist of and what makes a buffer effective
a weak acid (HA) and its conjugate base (A⁻), or excess weak acid and strong alkali. An effective buffer needs sufficient acid (HA) to react with added OH⁻ and sufficient conjugate base (A⁻) to react with added H⁺, suggesting equal amounts of HA and A⁻
Explain how methanoic acid can act as a buffer and give the equation
COOH ⇌ COO⁻ + H⁺
When a small amount of acid is added, the added H⁺ inos combine with the COO⁻ from the buffer and the buffer equilibrium shifts to the left to remove the added H⁺. When a small amount of alkali is added the added OH⁻ ions combine with H⁺ from the buffer and the buffer equilibrium shifts to the right to replace the removed H⁺. This minimises the change in [H⁺] and therefore the pH
Explain what the Ka is for a buffer
In a buffer, [A⁻] = [HA], therefore Ka = [H⁺], and so pKa = pH for an effective buffer
How do you know what weak acid to choose when deciding which buffer to use
A weak acid with a pKa close to the pH we want the buffer to maintain
What is the effective range of a buffer
± 1 pH unit of the pKa of the acid used
What is the pH range of blood
around 7.35 to 7.45
What is the most important buffer in the blood
the carbonic acid-hydrogencarbonate ion buffer
Give the equation for the buffer in the blood and explain in context of the blood, how this buffer works
H₂CO₃ ⇌ H⁺ + HCO₃⁻
When adding lactic acid, equilibrium shifts to the left. When adding alkalis like amines, equilibrium shifts to the right
What is excess carbonic acid in the blood converted to
excess carbonic acid is converted to CO₂ and released from the body via gas exchange
