Acid-Base Equilibria Flashcards
State Arrhenius Theory of acids and bases
An acid is a substance that releases H+ ions when dissolved in water, a base is a substance that releases OH- ions in the presence of water.
Limitation of Arrhenius Theory
Applies only to reactions which take place in aqueous medium.
Certain compounds, Ammonia behaves like bases although it does not have a OH- in its structure.
Explain what is meant by the terms Bronsted-Lowry acid, Bronsted-Lowry base and conjugate acid-base pair
Bronsted-Lowry acid is a proton, h+, donor
Bronsted-Lowry base is a proton, h+, acceptor
For every Bronsted acid, it has a conjugate base which has one H+ less than the acid.
For every Bronsted base, it has a conjugate acid which has one H+ more than the base.
Under which conditions can Bronsted-Lowry reactions occur?
They can occur in aqueous, non-aqueous and gaseous systems as emphasis is on proton transfer.
Factor affecting strength of Base
Extent of ionisation in aqueous solution. Strong base completely…
Strong base:
concentration of acid = concentration of OH-
Weak base:
concentration of base /= concentration of OH-
Limitation of Bronsted-Lowry reaction
A Bronsted-Lowry acid has to contain an ionisable Hydrogen atom.
Factor affecting strength of Acid
Extent of dissociation in aqueous solution. Strong acid completely…
Strong acid:
concentration of acid = concentration of H+/ H3O+
Weak acid:
concentration of acid /= concentration of H+/H30+
Auto ionisation of water ionic equation
H2O + H2O -> H3O+ + OH-. Arrow is reversible.
Ionic Product of water, Kw value and formula
Kw = (H3O+)(OH-) = 1.0 x 10^-14 mol^2dm^-6 at 25degrees celsius
pH scale and pOH scale formula. Formula to find concentration of H+ and OH- in relation
pH = -log(H30+)
pOH = -log(OH-)
(H+) = 10^-pH
(OH-) = 10^-pOH
Relationship of pOH and pH
pH + pOH = -log( 1.0 x 10^-14) = 14
When is auto-ionisation of water considered negligible and when is it not?
Negligible when concentrations of strong acid and strong base are high.
Not negligible when acids or bases are very diluted;
(H3O+) and (OH-) is lesser or equal to 1 x 10^-7 mol/dm3
This will result in (H3O+) and (OH-) contributed by auto-ionisation of water to be considered significant and hence needed to be included.
Formula to find H3O+ concentration using Ka, acid dissociation constant.
(H+) = (Ka x c)^1/2. Square rooted.
Limitations of using pH as indicator for strength of acidity or basicity
pH varies with its concentration. pH of a stronger acid need not be lower than pH of a weaker acid if concentrations are different.
To use pH as an indicator of acid strength between 2 acids, concentration of 2 acids must be the same.
Definition of a Buffer solution
A buffer solution is a solution in which its pH does not change significantly on the addition of a small amount acid or base.
Suitable indicator for titration between:
SA and SB
WA and SB
SA and WB
WA and WB
SA and SB -
Methyl orange (3.1 - 4.4)
Phenolphthalein (8.3 - 10.0)
WA and SB -
Phenolphthalein (8.3 - 10.0)
SA and WB -
Methyl orange (3.1 - 4.4)
WA and WB -
End point cannot be detected accurately by any indicator
Hence, pH>7 use phenolphthalein
pH<7 use methyl orange
Which indicator is the best indicator to test for strength of acid and why?
Acid dissociation constant, Ka, as it is a constant at constant temperature and does not vary with concentration. At the same temperature, Ka, of stronger acid is always higher than the Ka of a weaker acid, regardless of their concentrations.
Relative strength of conjugate Acid-Base Pair
The more readily an acid donates a proton, the less readily its conjugate base would accept a proton. Vice Versa.
The stronger an acid, the weaker its conjugate base; the stronger a base, the weaker its conjugate acid.
