Buffers_Acidic and Basic buffer solutions

Buffer Solutions
❖A buffer is a solution of a weak acid and its conjugate base (salt) that resists
changes in pH in both directions—either up or down, when small quantities
of an acid and a base(alkali) are added to it.
❖ A buffer works best in the middle of its range, where the amount of
undissociated acid is about equal to the amount of the conjugate base.
❖Acidic buffer solution:
• pH less than 7.
• are commonly made from a weak acid and one of its salts e.g. ethanoic acid and
sodium ethanoate in solution; pH 4.76

Buffer Solutions
• Acidic buffer solution:
• pH less than 7.
• are commonly made from a weak acid and one of its salts e.g. ethanoic acid and
sodium ethanoate in solution; pH 4.76
• Alkaline buffer solution:
• pH greater than 7.
• are commonly made from a weak base and one of its salts e.g. ammonia solution
and ammonium chloride solution; pH 9.25
• Both acid and conjugate base must be present in comparable amounts.

Buffer Solutions
Alkaline buffer solution:
• pH greater than 7.
• are commonly made from a weak base and one of its salts e.g.
ammonia solution and ammonium chloride solution; pH 9.25
❖There are two requirements for a buffer:
❖ Two substances are needed:
1. an acid capable of reacting with added OH- ions and
2. a base that can consume added H3O + ions.

Buffers
(a) .
• The pH electrode is indicating the pH of water that
contains a trace of acid (and bromphenol blue acid–
base indicator).
• The solution at the left is a buffer solution with a pH of
about 7. (It also contains bromphenol blue dye.)
(b) .
• When 5 mL of 0.10 M HCl is added to each solution,
the pH of the water drops several units, whereas the
pH of the buffer stays essentially constant, as implied
by the fact that the indicator colour does not change.

Uses of buffers
❖Keeping pH constant in a wide variety of chemical applications.
❖Enzyme reactions and cell functions have optimum pH’s for
performance.
❖pH control is important for them to function optimally.
❖Fermentation processes in industry
❖Pharmaceutical industry (buffer medium)
❖Calibration of pH meter.
❖Buffers function effectively over a narrow pH range; pKa ±1.

Uses of buffers
Too acidic?
• Increase respiration rate expelling CO2
• Excretory system – excrete more or less bicarbonate
❖Maintaining a constant blood pH
❖blood pH is about 7.
❖pH <7
❖or >7.8 is fatal.

Natural buffers in the human body
CARBONIC ACID BICARBONATE BUFFER
CO2 + H2O  H2CO3. H2CO3 + H2O  HCO3-
- + H3O+
HCO3
- + H2O  CO3
2- + H3O+
• Cellular respiration produces CO2 as a waste product.
• This is immediately converted to HCO3
- ion in the blood.
• On reaching the lungs it is again converted to and released as CO2.
• In the blood it neutralises acids released due to other metabolic
processes.
• In the stomach and duodenum it also neutralises gastric acids and
stabilises the intra cellular pH of epithelial cells
• by the secretion of bicarbonate ions into the gastric mucosa.

PROTEIN BUFFER (HAEMOGLOBIN) SYSTEM
HHB(aq)  H+(aq) + Hb-(aq) HHbO2(aq)  H+ aq)+ HbO2
-(aq)
HHb(aq) + HCO3(aq)  Hb(aq) + H2CO3(aq)
• Protein buffer system helps to maintain acidity in and around the cells.
• Haemoglobin makes an excellent buffer
• binds to small amounts of acids in the blood, before they can alter the pH of the blood.
Other proteins containing amino acid histidine are also good at buffering.

PHOSPHATE BUFFER SYSTEM
H2PO4
- (aq)  HPO4
2-(aq) + H3O+(aq).
• Phosphate buffer system operates in the internal fluids of all cells.
• It consists of dihydrogen phosphate ions as the H+ donor (acid) and
hydrogen phosphate ion as the H+ ion acceptor (base).
• If additional hydroxide ions enter the cellular fluid, they are
neutralised by the dihydrogen phosphate ion.
• If extra H+ enter the cellular fluid then they are neutralised by the
hydrogen phosphate ion.

How buffers work
❖Equilibrium between acid and base.
❖Example: Acetate buffer
• CH3COOH  CH3COO- + H+
• If more H+ is added to this solution, it simply shifts the equilibrium
to the left, absorbing H+, so the [H+] remains unchanged.
• If H+ is removed (e.g. by adding OH-) then the equilibrium shifts to
the right, releasing H+ to keep the pH constant

How do buffer solutions work?
Acidic buffer solutions
• We'll take a mixture of ethanoic acid and sodium ethanoate as typical.
• Ethanoic acid is a weak acid, and the position of this equilibrium will be well to the left:
• The solution will therefore contain these important things:
• lots of un-ionised ethanoic acid;
• lots of ethanoate ions from the sodium ethanoate;
• enough hydrogen ions to make the solution acidic.
• Other things (like water and sodium ions) which are present aren't important to the
argument.

• Hydrogen ions combine with the ethanoate ions to make ethanoic
acid.
• Since most of the new hydrogen ions are removed, the pH won't
change very much
• Added hydroxide ions are removed either by reacting with ethanoic
acid or by reacting with hydrogen ions to form water.

• The water formed re-ionises to a very small extent to give a few
hydrogen ions and hydroxide ions. But pH does not change much

Alkaline buffer solutions
• E.g., a mixture of ammonia and ammonium chloride solutions as typical.
• Ammonia is a weak base, and the position of this equilibrium will be well to the
left:
• The solution will therefore contain these important things:
• lots of un-reacted ammonia;
• lots of ammonium ions from the ammonium chloride;
• enough hydroxide ions to make the solution alkaline.
• Other things (like water and chloride ions) which are present aren't important
to the argument.

• Added hydroxide ions are removed by reacting with ammonium ions
• Added hydrogen ions (acid) are removed either by reacting ammonia or
by reacting with hydroxide ions to form water.

Limits to the working range of a buffer
❖Consider the previous example:
• CH3COOH  CH3COO- + H+
❖If too much H+ is added, the equilibrium is shifted all the way to the
left, and there is no longer any more CH3COO- to “absorb” H+.
❖At that point the solution no longer resists change in pH; it is useless
as a buffer.
❖A similar argument applies to the upper end of the working range.

Chemistry of buffers
❖Lets look at a titration curve
❖Titration is used to determine the concentration of an acid or base by
adding the OTHER and finding an equivalency point.
❖Equivalency point- point in titration at which the amount of titrant
added is just enough to completely neutralize the analyte solution.
❖At the equivalence point in an acid-base titration, moles of base =
moles of acid and the solution only contains salt and water.

❖Suppose you have a KOH solution,
and you want to know its
concentration (molarity).
❖Slowly add an acid (HCl) with a
known concentration (0.1 M) and
find the equivalency point…in this
case it will be at pH = 7… and we use
an indicator that changes color at
that pH determine when that point
has been reached.
❖So, suppose it takes 10ml of 0.1 M
HCl to buffer 50 ml of the KOH.

❖So, suppose it takes 10ml of 0.1 M
HCl to buffer 50 ml of the KOH.
❖The original concentration of the
base = Vol. acid x conc. of acid
Volume of base
= 10 ml x 0.1 M
50 ml
= 0.02 M

Equilibrium constant/Dissociation constant (Ka)
❖Ka = equilibrium constant for H+ ion transfer is also described as the dissociation
constant (the tendancy of an acid to dissociate).
AH → A- (base conjugate) + H+
Ka = [A-] [H+]/ [AH] = [base] [H+] / [acid]
❖Weak acids have low values… contribute few H+ ions…
❖Because we are usually dealing with very small concentrations, log values are
used…
❖The log constant =

❖Since pK is the negative log of K, weak acids have high values … (-2 –
12).
❖HCl = -9.3 – very low ~complete dissociation
• First rearrange the first equation and solve for [H+]
• [H+] = Ka x [acid]/[base]
❖Then take the log of both sides
• log10[H+] = log10Ka + log10 [acid]/[base]
-pH -pKa

❖-pH = -pKa + log10 [acid]/[base]
❖Multiply both sides by –1 to get the Henderson-Hasselbach equation
• pH = pKa - log10 [acid]/[base]

❖What happens when the concentration of the acid and base are equal?
• Example: Prepare a buffer with 0.10M acetic acid and 0.10M acetate
• pH = pKa - log10 [acid]/[base]
• pH = pKa - log10 [0.10]/[0.10]
• pH=pKa
• Thus, the pH where equal concentrations of acid and base are
present is defined as the pKa
❖If there is more of the conjugate base in the solution than acid, for
example, then pH > pKa. Conversely, if there is more acid than conjugate
base in solution, then pH < pKa.
❖A buffer works most effectively at pH values that are + 1 pH unit from the
pKa (the buffer range).

Choice of a buffer system
• It depends on:
• Desired buffer pH
• Ka of the weak acid (buffers work well when desired pH is within pKa
±1).
• Temperature (Ka depends on temperature)
• Reactions of the buffer components with medium toxicity, acidity etc

pH of a buffer solution
• The pH of a buffer system is given by the Henderson-Hasselbalch equation:
• for a weak acid and its salt:
• for a weak base and its salt:
• where [salt], [acid] and [base] are the molar concentrations of salt, acid and
base.

Preparation of a buffer solution of a given pH: e.g. prepare 2 litres of solution
of buffer pH 7.4
• Choose weak acid whose pKa is close to desired pH, for example H2PO4
_; Ka =
6.2 x 10-8, pKa = 7.21
• Use Henderson –Hasselbalch equation to calculate amounts of base and acid.
Acid base (salt)
H2PO4
- + H2O → HPO4
2_ + H3O+
Ka = 6.2 x 10-8. pKa = 7.21
7.4 = 7.21 + log[base]/[acid]
7.4 -7.21 = log[base]/[acid]
0.19 = log[base]/[acid]
[base]/[acid]= 100.19 = 1.5; i.e. base to acid ratio of 1.5

Preparation of 2L of pH 7.4 buffer solution
• Choose [base] to be 1.5M, then [acid] will be 1M to give a ratio of 1.5.
• Base = HPO4
2_; choose cation of Na, i.e. Na2HPO4; molar mass = 143g
• Acid = H2PO4
-; with cation of Na, i.e. NaH2PO4; molar mass = 120g
• For one litre of solution mass of base (Na2HPO4) is 1.5 x 143g = 214.5g and mass of acid
(NaH2PO4) is 1 x 120g = 120g.
• For two litres of buffer solution multiply the masses for 1 litre by 2.
• Mass of base (Na2HPO4) = 2 x 214.5g = 429g; mass of acid (NaH2PO4) = 2 x 120g = 240g
• Weigh 429g of Na2HPO4 and 240g of NaH2PO4 and dissolve them in 2 litres of distilled water and
mix well.

Buffer capacity
• Buffer capacity is a measure of the efficiency of a buffer in resisting changes in
pH.
• Conventionally, the buffer capacity ( ) is expressed as the amount of strong acid
or base, in gram-equivalents, that must be added to 1 liter of the solution to
change its pH by one unit.
= gram equivalent of strong acid/base to change pH of 1 liter of buffer
solution
• = the pH change caused by the addition of strong acid/base

Buffer capacity
• In practice, smaller pH changes are measured and the buffer capacity is
quantitatively expressed as the ratio of acid or base added to the change in pH
produced
• The buffer capacity depends essentially on 2 factors:
• Ratio of the salt to the acid or base: The buffer capacity is optimal when the ratio is
1:1; that is, when pH = pKa
• Total buffer concentration: For example, it will take more acid or base to deplete a
0.5 M buffer than a 0.05 M buffer.
• The relationship between buffer capacity and buffer concentrations is given by the
Van Slyke equation:
where C = the total buffer concentration (i.e. the sum of the molar concentrations
of acid and salt).
• Concentration of acid and base should be >>> [H3O+]

Buffer capacity
• Just as we must often compromise the optimal pH for a product, so must we
compromise on the optimal buffer capacity of our solution.
• On the one hand, buffer capacity must be large enough to maintain the product
pH for a reasonably long shelf-life.
• Changes in product pH may result from interaction of solution components with
one another or with the product package (glass, plastic, rubber closures, etc.).
• On the other hand, the buffer capacity of ophthalmic and parenteral products
must be low enough to allow rapid readjustment of the product to physiologic pH
upon administration.
• The pH, chemical nature, and volume of the solution to be administered must all
be considered.
• Buffer capacities ranging from 0.01 - 0.1 are usually adequate for most
pharmaceutical solutions.

Acid - Base Indicators
•Acid - Base indicators (also known as pH indicators) are
substances (dyes) which change colour with pH.
•They are usually weak acids or bases, with the
conjugate acid-base forms having different colours.

Acid - Base Indicators
• Consider an indicator which is a weak acid, with the formula HIn.
• At equilibrium, the following equilibrium equation is established with its conjugate base:
HIn(aq) + H2O(l)  In-(aq) + H3O+(aq)
Acid base
colour A colour B
• At low pH values the concentration of H3O+ is high and so the equilibrium position lies to
the left. The equilibrium solution has the colour A. At high pH values, the concentration of
H3O+ is low - the equilibrium position thus lies to the right and the equilibrium solution has
colour B.
• The eye only sees colour of predominant species provided its concentration is at least ten times
more than the other species.

Universal Indicator
• Substance that changes colour in different pH environments.
• Contains a mixture of indicators—that can be added to any substance
to determine its pH.
• Forms a continuous spectrum of colours that give an approximate reading of
the concentration of protons in a sample.
• Litmus paper, for example, appears red at low pH and appears blue or violet
at high pH.

Criterion for choice of indicator for acid-base titration
• Choose an indicator which changes colour at (or close to) the equivalence
point.
• The rapid change in pH value at the equivalence point should match with a
sharp change in colour of the indicator at the end point.
Thus consider:
• pH range at equivalence point for the acid and base in the titration.
• Working pH range (pka ±1) for the indicator.
• If several indicators are possible then choose one with the most sharp (easy
to observe) colour change.

Colour change
• Each pH indicator has its own pH range for its color change.
• The colour change is due to ionisation of the acid-base indicator.
• The unionised form has different colour than the ionised form.
The ionisation of the indicator is largely affected in acids and bases as it is
either a weak acid or a weak base.
HIn + H2O  In- + H3O+.
• If the indicator is a weak acid, its ionisation is practically negligible in acids as
the equilibrium shifts to left hand side due to high concentration of H3O+ ions.

Colour change
• Thus, the solution would have colour of the acid form. while it is fairly ionised in
bases and it has colour of ionised (base) form- OSTWALD’S THEORY
• Similarly if the indicator is a weak base, its ionisation is large in acids and low in bases
due to common OH- ions.
• In presence of an acid the equilibrium shifts to the left.
• On addition of alkali, hydrogen ions are removed by OH- ions in the form of water
molecules and the equilibrium shifts to right hand side.
• Thus, the concentration of ionised form increases in solution and they impart their
colour to the solution

pH range for indicator
• pH = pKa + log10[In−] / [Hin]
• For the reaction mixture to impart colour of HIn with confidence, [In-(aq)]/
[HIn(aq)] has to be ≤ 1/10 .
• For the reaction mixture to impart colour In- with confidence, [In-(aq)]/
[HIn(aq)] has to be ≥ 10.
• In other words, for the reaction mixture to impart colour of HIn with
confidence, pH value of the solution should be pKa - 1 or lower, and
• For the reaction mixture to impart colour of In- with confidence, pH value of
the solution should be pKa +1 or higher.
• Hence pH range of an indicator = pKa ± 1.

Colour and pH range for indicators
Indicator
Colour pKln pH range
Acid Base (actual)
Thymol Blue -
1st change
red yellow 1.5 1.2 - 2.8
Methyl Orange red yellow 3.7 3.2 - 4.4
Bromocresol Green yellow blue 4.7 3.8 - 5.4
Methyl Red yellow red 5.1 4.8 - 6.0
Bromothymol Blue yellow blue 7.0 6.0 - 7.6
Phenol Red yellow red 7.9 6.8 - 8.4
Thymol Blue -
2nd change
yellow blue 8.9 8.0 - 9.6
Phenolphthalein colourless pink 9.4 8.2 - 10.0

pH range at Equivalence point: Titration curves
• pH range at equivalence point is obtained from the pH profile for a titration, a plot of
pH (of titrand solution) against volume of titrant added.
• pH shows a large increase (inflexion) at equivalence point.
• pH and pH range at equivalence point depend on strengths of acid and base used in
the titration.
• Indicator works best when its pKa is equal to pH at equivalence point; i.e. [HIn] = [In-]

pH at equivalence point
Titrand Titrant Equiv pH Equiv pH
range
Remark(s)
Strong Strong 7 3 - 11 Steep inflection, sharp change
in pH, slow initial increase in pH
Weak Strong About 9 7 – 11 Less sharp change in pH, larger
initial increase in Ph
Strong Weak About 5 4 -7 Less sharp change in pH, slow
initial increase in pH
Weak Weak No equiv point Reaction goes to equilibrium

Buffers_Acidic and Basic buffer solutions

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