6 Acids, bases and pH Flashcards
State the Arrhenius definition of an:
a) ‘acid’
b) ‘base’
a) = produces H⁺ when dissolved in water
b) = produces OH- and a cation in aqueous solution
State the Bronstead-Lowry definition of an:
a) ‘acid’
b) ‘base’
a) H⁺ donor
b) H⁺ acceptor
State the Lewis definition of an:
a) ‘acid’
b) ‘base’
a) an electron-pair acceptor
b) an electron-pair donor
Why do acids form hydrogen ions (H⁺)?
the only positively-charged ion when dissolved in water
What are the three properties of hydrogen ions?
proton (H⁺) donor
lone-pair acceptor
pH < 7
Define ‘strong acids.’ Give an example.
= are completely dissociated to ions in solution
i.e. hydrochloric acid
HCl → H⁺ + Cl-
Define ‘weak acids.’ Give an example.
= are incompletely dissociated in solution
i.e. carbonic acid
H₂CO₃ ⇌ H⁺ + HCO₃- ⇌ 2H⁺ + CO₃2-
What do both bases and alkalis both do? Give 2 examples of equations of this.
both form hydroxide ions when dissolved in water
NaOH → Na⁺(aq) + OH- (aq)
NH₃ + H₂O ⇌ NH₄⁺ (aq) + OH-(aq)
What are the four properties of hydroxide ions?
proton (H⁺) acceptor
lone-pair donor
pH > 7
slippery feel
State the 2 definitions of an ‘alkali.’
forms hydroxide ions as the only negatively-charged ions when dissolved in water (e.g. NaOH)
a basic salt alkali metal or alkaline earth metal
Define ‘neutralisation.’
acid + base → salt + water
i. e. HCl + NaOH → H₂O + NaCl
overall: H⁺ + OH- → H₂O
Define ‘strong base.’
= H⁺ ion of the acid combines with the OH- of the base to form water
Define ‘salt.’
= the compound formed by the cation of the base and the anion of the acid (H⁺ of acid replaced by metal ion)
Give 2 examples of an acid-base neutralisation reaction and state the lone pair donor and lone pair acceptor reactions.
NH₃ + HCl → NH₄⁺Cl- lone pair lone pair donor acceptor NH₃ + BF₃ → NH₃⁺BF₃- lone pair lone pair donor acceptor
How are acids related to bases?
Acid ⇌ proton + conjugate base
How are bases related to acids?
Base ⇌ proton + conjugate acid
Label the conjugate acid-base pairs [as pairs (1) and (2)] in the following neutralisation reaction.
H₂CO₃ + OH- →
H₂CO₃ + OH- → HCO₃- + H₂O
(1) acid base(1) conjugate conjugate (2)
base (1) acid
Label the conjugate acid-base pairs [as pairs (1) and (2)] in the following neutralisation reaction.
CH₃COOH + H₂O ⇌
CH₃COOH + H₂O ⇌ CH₃COO- + H₃O⁺
(1)acid (2)base (1)base (2)acid
What is meant by acid strength?
not the same as concentration
strong = full dissociation
weak = partial dissociation
How is acid strength quantified?
by the acid dissociation constant, Ka
State the formula for Ka and how it is derived.
General acid-base reaction: HA + H₂O ⇌ H₃O⁺ + A- Therefore: Ka = [H₃O⁺][A-]/[HA][H₂O] Ignore [H₂O] as it is very large and not changed significantly during the reaction: Ka = [H₃O⁺][A-]/[HA] Or better known as: Ka = [H⁺][A-]/[HA]
State the formula for Kb and how it is derived.
General acid-base reaction: B + H₂O ⇌ BH + OH- Therefore: Ka = [BH][OH-]/[B][H₂O] Ignore [H₂O] as it is very large and not changed significantly during the reaction: Ka = [BH][OH-]/[B] Or better known as: Kb = [B⁺][OH-]/[B]
How does the value of Ka/Kb correspond to the strength of an acid/base?
The larger the value the stronger the acid/base
Define ‘pH.’
a measure of the concentration of H⁺
What is pH usually measured by?
by an indicator or using a pH probe
What is pH usually exist as and how is this converted?
It is a very small value in most cells
The p scale converts this to a whole number
State the formula for pH.
pH = -log10[H⁺]
What is pH of 6 equivalent to?
a H⁺ concentration of 1.00x10^-6M
State the formula for pOH and what it measures.
pOH = -log10[OH-]
It measures basicity
State the formula for the ionic product of water (Kw) for the relationship between pH and pOH and how it is derived.
H₂O ⇌ H⁺ + OH-
ionic product of water:
Kw= [H⁺][OH-]
(Ignore H₂O as it doesn’t vary)
What is Kw @ 25℃?
1x10^-14
State the formula of pKw and what it must add up to.
pKw = pH + pOH
must always add up to 14
What is H₃O⁺?
a H⁺ donor and acid
What are the values of Ka, [H₃O⁺] and pH with increasing acid strength?
High Ka and [H₃O⁺]
low pH
What are the values of Kb, [H₃O⁺] and pH with increasing base strength?
High Kb and pH
low [H₃O⁺]
State what indicators are, how they work and state the general equation for this.
indicators = weak acid molecules that change colour with a change of the equilibrium position
HLn ⇌ H⁺ + ln-
acid conjugate base
Give an example of an indicators, its colour change and the pH it maintains.
phenolphthalein
changes from colourless to pink
from pH range of 0-8.2 to a pH range of 8.2-12
Calculate the pH from a [H⁺] of 1 pM.
[H⁺] = 1 pM = 1x10^-12 M
= log10(1x10^-12) = -12
= -log10(1x10^-12) = 12
pH = 12 = HIGH
Calculate the pH from a [H⁺] of 1 mM.
[H⁺] = 1 mM = 1x10^-3 M
= log10(1x10^-3) = -3
= -log10(1x10^-3) = 3
pH = 3 = LOW
For a strong acid, what is the starting [Acid] equivalent to?
the [H⁺] of the acid
What is the equation for calculating the pH of weak acids and how is it derived?
HA ⇌ H⁺ + A- Ka = [H⁺][A-]/[HA] Rearrange: logKa = log ( [H⁺][A-]/[HA]) logKa = log [H⁺] + log ([A-]/[HA]) -pKa = -pH + log ([A-]/[HA]) Into the Henderson-Hasselbach equation: pH = pKa +log ([A-]/[HA])
What is a buffer solution?
resists change in pH upon addition of small amounts of either acid or base
What is a buffer solution composed of?
consists of a solution of a weak acid and the salt of its conjugate base
Or, a weak base and the salt of its conjugate acid
it has a high concentration, as there is an equal amount of each component
What pH does a buffer generally maintain?
a pH of around the pKa of the acid
What happens to this general buffer, if more H⁺ is added?
HA ⇌ H⁺ + A-
H⁺ + A- → HA
the conjugate base accepts it, forming the conjugate acid and no pH change occurs
equilibrium position shifts left to resist the change
What happens to this general buffer, if OH- is added?
HA ⇌ H⁺ + A-
H⁺ + OH- → H₂O
it reacts with H⁺ to form water
equilibrium shifts right
more acid dissocitaes to maintain constant [H⁺]
State the buffer system in intrasystemic fluid and what it is used for. (Hint: the system is the same as the blood)
use bicarbonate
for metabolic acids
State the buffer system in blood and what it is used for.
Hint: what is the alkaline substance that accepts protons ?
use bicarbonate and haemoglobin
for metabolic acids and important for carbon dioxide
State the buffer system in intracellular fluid.
Hint: what is made in cells?
proteins and phoshpates
State the buffer system in urine and what it is used for.
phosphate and ammonia
for the formation of NH₄⁺
State the buffer system in bone and what it is used for.
Calcium carbonate and phosphate
for prolonged metabolic acidosis
Explain how blood pH is maintained by bicarbonate in plasma using a diagram.
see document
Define ‘metabolic acidosis/alkalosis.’
a build-up of acidic metabolites (e.g. oxalic acid, lactic acid), failure of acid-base elimination in the kidneys
Define ‘respiratory acidosis/alkalosis.’
too little/too much CO₂ exchanged in lungs
Explain how blood pH is maintained by bicarbonate in plasma.
H⁺ produced by dissociation of H₂CO₃
protonation reduces affinity of haemoglobin for O₂
O₂ released in peripheral tissues
haemoglobin can also transport CO₂ - also favours oxygen release
Draw out the reactions between:
a) histidine and H⁺
b) haemoglobin and CO₂
a) Hint: Pentagon with N, NH and double bond below N and opposite (N becomes NH⁺)
b) NH₂ becomes NH=OOH
For each of the following names identify the substance involved in the acid/base definitions:
a) Arrhenius
b) Bronstead-Lowry
c) Lewis
a) God of water; H⁺/OH- dissolved in ‘water’ or solution
b) Sounds like a philosopher very positive; H⁺ ions
c) Lewis; feeling ‘low’ - ‘electron-pairs’ are negative
