Acis And Base Flashcards
Bronsted Lowry definitions
Bronstead Lowry base is a substance that can accept a proton
Bronstead Lowry acid is a substance that donates protons
Protons in aqueous solutions
Since the H° ion has no electrons of its own, it can only form a bond with another species that has a lone pair of electrons. This extremely small size and its intense electric field cause it to have unusual properties compared with other positive ions. It is never found isolated, In aqueous solutions it is always bonded to at least one water molecule to form the ion H3O+
The ionisection of water
Water is slightly lonised:
H20(l) <=> H°(aq) + OH-(aq)
This may be written:
H20(l)+ H20(l) <=>H30°(aq) + OH°(aq)
This emphasises that this is an acid-base reaction in which one water molecule donates a proton to another.
This equilibrium is established in water and all aqueous solutions
You can write an equilibrium expression:
Kc : (H+ (aq))(OH-(aq))
(H20(aq))
K, is called the ionic product of water and at 298K it is equal to 1.0 × 10-14 mol dm-*. Each H,O that dissociates (splits up) gives rise to one H+ and one OH
Ph scale
pH = -log10[H+(aq)]
Remember that square brackets, [I. mean the concentration in moldm
The smaller the pH, the greater the concentration of H*(aq).
Ph with alkalis
pil measures alkalinity as well as acidity, because as [H° (a9)] goes up.
¡OH*(aq)) soes down. Al 298K, if a solution contains more H° (a9) than OH” (ag), its pit will be less than 7 and it is called acidic. la solution contains more OH (aq) than H°(aq), its pil will be greater than 7 and it is called alkaline
Strong and weak acids and bases
Acids that completely dissociate into ions in aqueous solutions are called strong acids. The word strong refers only to the extent of dissociation and not in any way to the concentration.
Strong bases are completely dissociated into ions in aqueous solutions.
Many acids and bases are only slightly ionised) when dissolved in water. In fact an equilibrium is set
Acids like this are called weak acids. Weak refers only to the degree of dissociation.
The dissociation of weak acids
Imagine a weak acid HA which dissociates:
HA(aq) <=> H+(aq) + A (aq)
The equilibrium constant for this reaction is called the dissociation constant and is notated by Ka
The larger the value of Ka the further the equilibrium is to the right.
the more the acid is dissociated , and the stronger it is
Pka
For a weak acid pKa is often referred to. This is defined as:
pKa=-log10ka
Ph changes in an acid base titration
In an acid-base titration, an acid of known concentration is added from a burette to a measured amount of a solution of a base until an indicator shows that the base has been neutralised.
Alternatively, the base is added to the acid until the acid is neutralised.
You can then calculate the concentration of the alkali from the volume of acid used.
You can also follow a neutralisation reaction by measuring the pH with a pH meter) in which case you do not need an indicator. The pH meter is calibrated by placing the probe in a buffer solution of a known pH
Titration curves
Strong acid and base Curve further away from ph 7
Strong scid weak base acid curves further away and base curves closer to ph 7
Weak acid strong base acid curves closer to ph 7 and base curves further away
Weak acid and base both curve closer to ph 7
In a titration, the equivalence point is the point at which sufficient base has been added to just neutralise the acid
Indicators in titrations
The end point is the volume of alkali or acid added when the indicator just changes colour. Unless you choose the right indicator, the equivalence point and the end point may not always give the same answer.
A suitable indicator for a particular titration needs the following properties:
• The colour change must be sharp rather than gradual at the end point, that is, no more than one drop of acid (or alkali) is needed to give a complete colour change. An indicator that changes colour gradually over several cubic centimetres would be unsuitable and would not give a sharp end point.
• The end point of the titration given by the indicator must be the same as the equivalence point, otherwise the titration will give the wrong answer.
• The indicator should give a distinct colour change..
Indicator types
Left hand side colours are more acidic and right hand side are more alkali
Methyl orange -changes from red to yellow at Ph3.7
Bromophenol blue- changes from yellow to blue at ph 4
Methyl red - changes from red to yellow at ph5.1
Bromothymol blue - changes from yellow to blue at ph 7
Phenolpthalein changes from colourless to red at ph 9.3
Half neutralisation point
The point half-way between the zero and the equivalence point is the half-neutralisation point. This is significant because the knowledge that you can add acid (or base) up to this point with the certainty that the pll will change very little is relevant to the theory of buffers. It also allows you to find the pA, of the weak acid,
How buffers work
Buffers are designed to keep the concentration of hydrogen ions and hydroxide ions in a solution almost unchanged. They are based on an equilibrium reaction which will move in the direction to remove either additional hydrogen ions or hydroxide ions if these are added.
Acidic buffers
Acidic buffers are made from weak acids. They work because the dissociation of a weak acid is an equilibrium reaction.
If a little alkali is added, the OH ions from the alkali will react with HA to produce water molecules and A:
This removes the added OH so the pH tends to remain almost the same.
If H+ is added, the equilibrium shifts to the left- H* ions combining with A ions to produce undissociated HA. But, since [A~] is small, the supply of A soon runs out and there is no A left to combine the added Ht. So the solution is not a buffer.
However, you can add to the solution a supply of extra A by adding a soluble salt of HA, which fully ionises,. This increases the supply of A so that more H+ can be used up. So, there is a way in which both added H+ and OH~ can be removed.
An acidic buffer is made from a mixture of a weak acid and a soluble salt of that acid. It will maintain a pH of below 7
Buffers
Buffers don’t ensure that no change in pH occurs. The addition of acid or alkali will still change the pH, but only slightly and by far less than the change that adding the same amount to a non-buffer would cause.
