Further aspects of equilibria

Acids

Acids that dissociates completely in solution are strong acids while those who dissociates partially are called weak acids

Base

Bases that dissociates fully in solution are called strong bases while those who dissociates partially are called weak bases

Bronsted- Lowry theory

Proton donor : bronsted acid
Proton acceptor : bronsted base
Conjugate acid and conjugate base is there because there is equilibrium

Equilibrium constant, Kw

Ionic product of water
Concentration of water is always constant
Its value at 298 K is 1.00 x 10^-14 mol^2dm^-6
Temperature change → change in Kw
Change in Kw → change in pH

pH

Defined as the negative algorithm to the base of 10 of the hydrogen ion concentration

Since the concentration of the left hand side is a constant, we ignore it in the calculation of Kw
[OH-] will also have the same concentration as [H+] thus we can substitute it to get [H+]^2

Formula: POH + PH =14

Dissociation constant, Ka

At 298 K, value of Ka for dissociation of ethanoic acid is 1.74 x 10^-5 mol dm^-3
Temperature change → change in Ka
To calculate the value of Ka, make 2 assumptions
1. Ignore the concentration of hydrogen ions produced by the ionisation of the water molecules present in the solution. This is reasonable because the ionic product of water which is 1x10^-14 is negligable compared with most values for most weak acids.
2. Assume that the ionisation of the weak acid is so small that the concentration of undissociated HA molecules present at equilibrium is approximately the same as that of the original acid
3. The higher the Ka value, the smaller the PKa value.
pKa = -log10[Ka]

Acid-base indicators

Is a dye or mixture of dyes that changes colour over a specific pH range
Many indicators can be considered as weak acids
Adding an acid to this indicator will shift the position of the equilibrium to the left thus there are now more molecules of colour A
Adding an alkali to this indicator will shift the position of the equilibrium to the right thus there will be more molecules of colour B
The colour of the indicator depends on the relative concentration of HIn and In- and the colour of the indicator during a titration depends on the concentration of H+ ions present.

Indicators

Equivalence point

Equivalence point is the sharp decrease in the graph where all the OH- ions fully reacted with the H+ ions

End point

End point is the point when the indicator changes colour

Graphs

Strong acid and strong base

A sharp fall in the graph line between pH 10.5 and 3.5, in this region tiny additions of H+ ions result in a rapid charge in pH
The midpoint of the slop is at pH 7.0
The midpoint of the sharp fall corresponds to the point at which the H+ ions in the acid have exactly reacted with the OH- ions in the alkali, this is the end point of the titration
Bromothymol blue indicator changed from blue to yellow over the range of 7.6 to 6.0, We can also use phenolphthalein

Strong acid and weak base

A sharp fall between pH 7.5 to 3.5
The midpoint of the slope is at pH 5
We use methyl red instead of phenolphthalein

Weak acid and strong base

A sharp fall between pH 11 to 7.5
The midpoint of the slope is at pH 9
We use phenolphthalein instead of methyl orange

Weak base and weak acid

No suitable indicator as there is slope observable

Buffer solutions

A solution in which pH does not change significantly when a small amount of acids or alkalis are added
Is used to keep the pH constant • Mixture of weak acid and its salts: ethanoic acid and sodium ethanoate (buffer at pH range of 4-7)
Mixture of weak base and its salts: ammonia and ammonium chloride

Limitations of buffer solution: No buffer solution can cope with the excess addition of acids or alkalis. If a very large amount of acid or alkali are added. the pH will change significantly.

Calculation of buffer solution

Buffer to maintain body pH

The reaction is catalyzed by carbonic anhydrase. If the production of H+ ions are left unchecked, it would lower the blood pH and cause acidosis. This may denature enzymes halting important metabolic reactions.

Blood pH: 7.35-7.45

Connection to Le chatelier's principle

If the H+ ion concentration increases:

The equilibrium shifts to the left
H+ ions combine with hydrogen carbonate ions to form carbon dioxide and water until equilibrium is restored
This decreases the concentration of hydrogen ions in the blood and helps keep the pH constant

If the H+ ion concentration decreases:

The equilibrium shifts to the right
Some carbon dioxide and water combine to form H+ and hydrogen carbonate ions until equilibrium is restored
This decreases the concentration of hydrogen ions in the blood maintaining constant blood pH

Solubility product

Solubility product is the product of concentrations of each ion in a saturated solution of a sparingly soluble salt in 298K, raised to the power of their relative concentrations

An equilibrium is established when an undissolved ionic compound is in contact with a saturated solution of its ions. The ions move from solid to the saturated solution at the same rate as they move from the solution to the solid.

The idea of solubility product is only useful for sparingly soluble salts. The smaller the value of Ksp, the lower the solubility of the salt.

Calculation of Ksp

note:

Moles is raised as powers
solubility times the mole

Predicting ppt

Notes:

Remember to check total volume
If the value of Ksp is greater than the solubility product, then a precipitate will form because we can say that the equilibrium shifts to the left forming more BaCo3

Common ion effect

The common ion effect is the reduction in solubility of a dissolved salt achieved by adding a solution of a compound which has an ion in common with the dissolved salt. This often results in precipitation

If we add a solution of sodium chloride, NaCl to the equilibrium->

The chloride ion is common to both sodium chloride and silver chloride
The added chloride ions shift the position of equilibrium to the left
Thus, silver chloride is precipitated

Partition coefficients, Kpc

Let us consider ammonia dissolved in two immiscible solvents that forms two separate layers. A separating funnel with the organic solvent and a aqueous solution of ammonia. The ammonia is soluble in both solvents so when the mixture is left to settle, a dynamic equilibrium is established. At this point, ammonia molecules are moving from the aqueous layer to the organic layer at the same rate as they are moving form the organic layer to the aqueous layer.

The partition coefficient, Kpc is the equilibrium constant that relates the concentration of a solute partitioned between two immiscible solvents at a particular temperature.

Let us use oil and water as the two immiscible liquids. Ammonia can dissolve in both liquids but ammonia forms hydrogen bonds with water while only VDW is present when ammonia dissolves in oil. Thus, ammonia will be more soluble in water. This is due to the polarity of ammonia. Solubility depends on the strength of the intermolecular bonds between the solvent and the solute.
In conclusion, The polarity of a molecule will affect the Kpc relating to solubility and the stronger the Intermolecular forces are, the more soluble the solute will be

Partition coefficient and paper chromatography

In paper chromatography the different partition coefficients of the components in a mixture correspond to their relative solubilities in the two solvents. In paper chromatography the mobile phase is the solvent chosen. The other solvent is the water trapped in the paper’s structure, which is the stationary phase. The solutes in the mixture being separated are partitioned to different extents between the solvents in the mobile and stationary phase. The greater the relative solubility in the mobile phase, the faster the rate of movement as the mobile phase passes over the stationary phase.

In the figure to the right, the mobile phase moves over the stationary phase, carrying solute particles with it. The filter paper is the inert solid in paper chromatography.

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