Buffer solutions.

ass. prof. I. R. Bekusass. prof. I. R. Bekus

Plan

1.Ionization of water.2.Acid-base theory.3.Buffer solutions.4.Buffer in blood.

Importance of water1. The most abundant substance in living systems.

– Makes up 70% or more of the weight of most organisms.

2. Living organisms depend on water for their existence.– Physical & chemical properties of water make it fit to support

life: High boiling point water remains liquid in most seasons. Ice less dense than water floats on liquid water and water freezes

from top to bottom. So a good insulator: a frozen layer of ice serves as a blanket that protects creatures below.

3. Ubiquitous solvent in cells.4. Excellent solvent of polar and ionic substances.5. Medium where most cell’s metabolic reactions take

place.6. Ionization of water and its acid-base reactions important

for the functions of proteins and nucleic acids.7. The Shapes of proteins and nucleic acids and structure

of biological membranes are consequence of their interaction with water.

“Water molecule” is polar Water has a simple structure. Oxygen has 6 electrons in the outer shell: 1s2, 2s2,

2p4.

Sp3 hybridization H—O—H bond angle is 104.5.

The net charge of water molecule is zero. But O—H bond is polar because O is more electronegative than H. Sharing of electrons between H and O is unequal. The charge on O = -0.82 and on H = +0.41. This charge separation produces permanent dipoles.

Water is а neutral molecule with а slight tendency to ionize. We usually express this ionization as:Н2О = Н+ + ОН-

There is actually no such thing as а free proton (Н+) in solution. Rather, the proton is associated with а water molecule as а hydronium ion, H3O

+. The association of

а proton with а cluster of water molecules also gives rise to structures with the formulas Н5О2

+, Н7О3+, and

so on. For simplicity, however, we collectively represent these ions by H+.

Because the product of [Н+] and [ОН-] is а constant (10-14 М2), [Н+] and [ОН-] are reciprocally related. Solutions with relatively more Н+ are acidic (рН < 7), solutions with relatively more ОН- are basic (рН >7), and solutions in which [Н+] = [ОН-] = 10 -7 М are neutral (рН = 7). Note the logarithmic scale for ion concentration. K is the dissociation constant (ionization constant)

Кw = [Н+][ОН-] =10 -14 M2 at 25 0C.[Н+] = [ОН-] = 10-7 М [Н+] = 10-7 М are said to be neutral[Н+] > 10-7 М are said to be acidic, [Н+] < 10-7 М are said to be basic. Most physiological solutions have

hydrogen ion concentrations near neutrality.

[H+] = [OH-] Neutral solution[H+] > [OH-] Acidic solution[H+] < [OH-] Basic (alkaline) solution

Ion Product of Water, Kw

The ion product constant, Kw, for water

* is the product of the concentrations of the hydronium and hydroxide ions.

Kw = [ H3O+] [ OH− ]

* can be obtained from the concentrations in pure water.

Kw = [ H3O+] [ OH− ]

Kw = [1.0 x 10− 7 M] x [ 1.0 x 10− 7 M]

= 1.0 x 10− 14

Calculating [H3O+]

What is the [H3O+] of a solution if [OH−] is 5.0 x 10-8 M?

STEP 1: Write the Kw for water.

Kw = [H3O+ ][OH− ] = 1.0 x 10−14

STEP 2: Rearrange the Kw expression.

[H3O+] = 1.0 x 10-14

[OH−]

STEP 3: Substitute [OH−]. [H3O+] = 1.0 x 10-14 = 2.0 x 10-7 M

5.0 x 10- 8

рН = - log[H+] The pH of pure water is 7.0, Acidic solutions have рН < 7.0 Basic solutions have рН > 7.0.1 М NaOH -14Household ammonia -12Seawater – 8Milk - 7Blood  - 7.4Saliva - 6.6Tomato juice - 4.4Vinegar - 3Gastric juice - 1.51 М НСl - 0

The relationship between the pH of а solution and the concentrations of The relationship between the pH of а solution and the concentrations of an acid and its conjugate base is easily derived. an acid and its conjugate base is easily derived.

[НА] [НА] [Н[Н++]= ]= KK ---------- ---------- [[АА--]]Taking the negative log of each term Taking the negative log of each term [А[А--]]рН = - рН = - log log К + К + log log ------------------ [А[А--]]

This relationship known as the Henderson-Hasselbach equation, This relationship known as the Henderson-Hasselbach equation, that is often used to perform the that is often used to perform the calculations required in preparation of buffers for use in the laboratory, or other applications. calculations required in preparation of buffers for use in the laboratory, or other applications.

pH pH is commonly expressed as – log

[H+] Pure water has [H+]=10-7 and thus

pH=7.

The methods for measuring pH fall roughly into the following four

categories:

1) Indicator methods2) Metal-electrode methods

(including the hydrogen-electrode method, quinhydron-electrode method and antimony-electrode method)

3) Glass-electrode methods4) Semiconductor sensor methods

Ways to measure pH pH meter

Electrode measures H+ concentration Must standardize (calibrate) before

using.

BUFFER SOLUTIONSBUFFER SOLUTIONS

Buffers are solutions which can Buffers are solutions which can resist changes in pH by resist changes in pH by addition of acid or alkali.addition of acid or alkali.

Buffer systems

The solution, whose pH values do not practically change when moderate amount of either a strong acid or strong base are added and also as result of dilution, are called buffer solutions.

According to their composition buffers

divided into such groups:

1) Acidic buffer solution – consist of a weak acid and a salt of this weak acid and a strong base (acetate buffer solution )

2) Basic buffer solution - consist of a weak base and a salt of weak base and a strong acid (ammonium buffer solution )

3) Protein ampholytic buffer solution

Buffers are mainly classified of two Buffers are mainly classified of two types: types: ((аа) mixtures of weak acids with their ) mixtures of weak acids with their salt with salt with аа strong base strong base (b) mixtures of weak bases with their (b) mixtures of weak bases with their salt with salt with аа strong acid. strong acid. АА few examples are given below: few examples are given below:НН22СОСО

33 / N / NаНСОаНСО33 (Bicarbonate buffer; (Bicarbonate buffer;

carbonic acid and sodium bicarbonate)carbonic acid and sodium bicarbonate)СНСН

33СООНСООН / / СНСН33СООСООNaNa (Acetate (Acetate

buffer; acetic acid and sodium acetate)buffer; acetic acid and sodium acetate)NaNa

22HPOHPO44/ NaH/ NaH

22POPO44 (Phosphate(Phosphate buffer)buffer)

The most widely used in laboratory practice buffer solution

Buffer mixture name

Composition pH

Formate Formic acid HCOOH and sodium formate HCOONa

3,8

Benzoate Benzoic acid C6H5COOH and sodium benzoate C6H5COONa

4,2

Acetate Acetic acid CH3COOH and sodium acetate CH3COONa

4,8

Phosphate Sodium dihydrogen phosphate NaH2PO4 and sodium hydrogen phosphate Na2HPO4

6,6

Ammonium Ammonium hydrate NH3*H2O and ammonium chloride NH4Cl

9,2

Chemistry of buffers -pH = -pKa + log10 [acid]/[base] Multiply both sides by –1 to get the

Henderson-Hasselbach equation– pH = pKa - log10 [acid]/[base]

Chemistry of buffers 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

A buffer works most effectively at pH values that are + 1 pH unit from the pKa (the buffer range)

Calculating buffer recipes

Henderson-Hasselbach equation– pH = pKa - log10 [acid]/[base]

Rearrange the equation to get– 10(pKa-pH) = [acid]/[base]

Look up pKa for acid in a table. Substitute this and the desired pH into equation above, and calculate the approximate ratio of acid to base.

Because of the log, you want to pick a buffer with a pKa close to the pH you want.

Example You want to make about 500 mL of

0.2 M acetate buffer (acetic acid + sodium acetate), pH 4.0.

Look up pKa and find it is 4.8. 10(4.8 - 4.0) = 100.8 = 6.3 =

[acid]/[base] If you use 70 mL of base, you will

need 6.3X that amount of acid, or 441 mL. Mix those together and you have 511 mL.

Acid–Base ConceptAcid–Base Concept1)1) The Arrhenius theoryThe Arrhenius theory

ACIDACID - a substance that provides H+ ions in water - a substance that provides H+ ions in water

BASEBASE - - a substance that provides OH- ions in watera substance that provides OH- ions in water

2) 2) The BrønstedThe Brønsted--Lowry TheoryLowry Theory

All Brønsted–Lowry bases have one or more lone pairs of electrons:

3) The Lewis Acids and Base theory3) The Lewis Acids and Base theoryLEWIS LEWIS ACIDACID An electron-pair acceptor An electron-pair acceptor

LEWIS LEWIS BASEBASE An electron-pair donor An electron-pair donor

What is an Acid? An acid is a substance which, when

dissolved in water, releases protons. The extent of dissociation, that is, the

amount of protons released compared to the total amount of compound, is a measure of the strength of the acid.

For example, HCl (hydrochloric acid) is a strong acid, because it dissociates completely in water, generating free [H+] and [Cl-].

Acidity can be measured on a scale called pH (more scarily, “the negative logarithm of the hydrogen ion concentration”).

31

AcidsAcids

Lemon juiceLemon juice contains contains citric acid,citric acid, and and vinegar vinegar containscontains ethanoic acid.ethanoic acid.

Some strong acids Some strong acids areare hydrochloric acid, sulphuric hydrochloric acid, sulphuric acid acid andand nitric acid. nitric acid.

Some weak acids Some weak acids areare ethanoic ethanoic acid, citric acid acid, citric acid andand carbonic acid. carbonic acid.

There are many acids present in our everyday lives.

But What’s a Weak Acid?

Some substances, like acetic acid (vinegar!) dissociate poorly in water.

Thus, they release protons, but only a small fraction of their molecules dissociate (ionize).

Such compounds are considered to be weak acids.

Thus, while 1 M HCl is pH = 0 and 1 M acetic acid is only pH = 2.4…

33

Weak acids thus are in equilibrium with their ionized species:

34

HA H+ + A-

Keq =[H+][A-] [HA]

Governed by the Law of Mass Action, and characterized by an equilibrium constant:

Weak acids, their conjugate bases, and buffers…

Weak acids have only a modest tendency to shed their protons.

When they do, the corresponding negatively charged anion becomes a willing proton acceptor, and is called the conjugate base.

The properties of a buffer rely on a balance between a weak acid and its conjugate base.

35

AlkalisAlkalis

When the When the oxides of some oxides of some metalsmetals dissolve in dissolve in waterwater they they make an make an alkali solution.alkali solution.

Alkalis Alkalis react with react with acidsacids and and neutralise them. them.

Many everyday substances are alkalis. They feel soapy. They are corrosive.

Alkalis Alkalis

AlkalisAlkalis are present in many are present in many cleaning substancescleaning substances in use in in use in our homes. our homes.

Kitchen cleaners are alkalineKitchen cleaners are alkaline because they contain because they contain ammoniaammonia or or sodium hydroxidesodium hydroxide, which , which attack grease.attack grease.

Calcium hydroxide and sodium hydroxide are strong alkalis. The most recognisable and common weak alkali is ammonia.

Important buffers

Buffers in the Blood The pH of blood is 7.35 – 7.45 Changes in pH below 6.8 and above

8.0 may result in death The major buffer system in the body

fluid is H2CO3/HCO3-

Some CO2, the end product of cellular metabolism, is carried to the lungs for elimination, and the rest dissolves in body fluids, forming carbonic acid that dissociates to produce bicarbonate (HCO3

-) and hydronium (H3O+) ions.

More of the HCO3- is supplied by the

kidneys. CO2 + H2O ↔ H2CO3 H2CO3 + H2O ↔ H3O+ + HCO3

-

Bicarbonate buffer system

H2CO3 + H2O ↔ H3O+ + HCO3-

Excess acid (H3O+) in the body is neutralized by HCO3

-

H2CO3 + H2O ← H3O+ + HCO3-

Equilibrium shifts left Excess base (OH-) reacts with the

carbonic acid (H2CO3) H2CO3 + OH- → H2O + HCO3

-

Equilibrium shifts right

pH of the bicarbonate buffer The concentrations in the blood of

H2CO3 and HCO3- are 0.0024M and

0.024 respectively H2CO3/ HCO3

- = 1/10 is needed to maintain the normal blood pH (7.35 – 7.45)

Regulation of blood pH

The lungs and kidneys play important role in regulating blood pH.

The lungs regulate pH through retention or elimination of CO2 by changing the rate and volume of ventilation.

The kidneys regulate pH by excreting acid, primarily in the ammonium ion (NH4

+), and by reclaiming HCO3

- from the glomerular filtrate (and adding it back to the blood).

The concentration of carbonic acid in the body is

associated with the partial pressure of CO2.* When CO2 level rises, producing more H2CO3, the

equilibrium produces more H3O+, which lowers the pH – acidosis.

* Decreasing of CO2 level due to a hyperventilation, which expels large amounts of CO2, leads to a lowering in the partial pressure of CO2 below normal and the shift of the equilibrium from H2CO3 to CO2 and H2O. This shift decreases H3O+ and raises blood pH – alkalosis.

CO2 + H2O ↔ H2CO3 ↔ H3O+ + HCO3

-

Respiratory Acidosis: CO2 ↑ pH ↓ Symptoms: Failue to ventilate, suppression

of breathing, disorientation, weakness, coma Causes: Lung disease blocking gas diffusion

(e.g., emphysema, pneumonia, bronchitis, and asthma); depression of respiratory center by drugs, cardiopulmonary arrest, stroke, poliomyelitis, or nervous system disorders

Treatment: Correction of disorder, infusion of bicarbonate

Respiratory Alkalosis: CO2 ↓ pH ↑

Symptoms: Increased rate and depth of breathing, numbness, light-headedness, tetany

Causes: hyperventilation due to anxiety, hysteria, fever, exercise; reaction to drugs such as salicylate, quinine, and antihistamines; conditions causing hypoxia (e.g., pneumonia, pulmonary edema, and heart disease)

Treatment: Elimination of anxiety producing state, rebreathing into a paper bag

Metabolic (Nonrespiratory) Acidosis: H+ ↑ pH ↓ Symptoms: Increased ventilation, fatigue,

confusion Causes: Renal disease, including

hepatitis and cirrhosis; increased acid production in diabetes mellitus, hyperthyroidism, alcoholism, and starvation; loss of alkali in diarrhea; acid retention in renal failure

Treatment: Sodium bicarbonate given orally, dialysis for renal failure, insulin treatment for diabetic ketosis

Metabolic (Nonrespiratory) Alkalosis: H+ ↓ pH ↑ Symptoms: Depressed breathing,

apathy, confusion Causes: Vomiting, diseases of the

adrenal glands, ingestions of access alkali

Treatment: Infusion of saline solution, treatment of underlying diseases

Other important buffers The phosphate buffer system (HPO4

2-/H2PO4-)

plays a role in plasma and erythrocytes.H2PO4

- + H2O ↔ H3O+ + HPO42-

Any acid reacts with monohydrogen phosphate to form dihydrogen phosphate

dihydrogen phosphate — monohydrogen phosphate

H2PO4- + H2O ← HPO4

2- + H3O+

The base is neutralized by dihydrogen phosphatedihydrogen phosphate — monohydrogen

phosphateH2PO4

- + OH- → HPO42- + H3O+

Proteins act as a third type of blood buffer

Proteins contain – COO- groups, which, like acetate ions (CH3COO-), can act as proton acceptors.

Proteins also contain – NH3+ groups, which, like

ammonium ions (NH4+), can donate protons.

If acid comes into blood, hydronium ions can be neutralized by the – COO- groups

- COO- + H3O+ → - COOH + H2O

If base is added, it can be neutralized by the – NH3+

groups

- NH3+ + OH- → - NH2 + H2O

Buffer CapacityBuffer Capacity Buffer capacity is determined by the actual Buffer capacity is determined by the actual

concentrations of salt and acid present, as concentrations of salt and acid present, as well as by their ratio. well as by their ratio.

Buffering capacity is the number of Buffering capacity is the number of grams of strong acid or alkali which is grams of strong acid or alkali which is necessary for necessary for аа change in pH of one change in pH of one unit of one liter of buffer solution.unit of one liter of buffer solution.

The buffering capacity of The buffering capacity of аа buffer is, definеd buffer is, definеd ааs the ability s the ability of the buffer to resist changes in pH when an acid or of the buffer to resist changes in pH when an acid or base is added.base is added.

Buffer Capacity

A buffer counteracts the change in pH of a solution upon the addition of a strong acid, a strong base, or other agents that tend to alter the hydrogen ion concentration.

Buffer capacity β: buffer efficiency, buffer index or buffer value Is the resistance of a buffer to pH changes

upon the addition of a strong acid or base.Definition:It can be defined as being equal to the amount of strong acid or strong base , expressed as moles of H +or OH- ions, required to change the pH of one litre of the buffer by one pH unite.Maximum buffer capacity (βmax) obtain when ratio of acid to salt = 1 i.e. pKa = pH

DefinitionThe ratio of increase of strong base (or acid) to

small change in pH brought about by this addition.

β = Δ B Δ pH

*ΔB = the small increment in gram equiv./liter of strong base

* added to the buffer solution to produce a pH change

Δ pH = pH change

Definition of buffer capacity

A buffer absorbs strong acid and base through the two reactions shown on the left side of our diagram:

A- + H3O+ => HA + H2O HA + OH- => A- + H2O

The buffer will stop working when either one of its components (HA or A-) is exhausted, and therefore cannot neutralize any more strong acid or strong base. The most effective buffering solutions are those which have similar concentrations of HA and A- because then the buffer has the capacity to absorb both acid and base with the same effectiveness. Although the pH of a buffer depends only on the ratio [HA]/[A-], the ability of the buffer to absorb acid or base depends on the overall value of [HA] and [A-]. For instance, above we found a pH change of -0.02 units (from 7.20 to 7.18) when we added 0.010 moles of HCl to 1L of a buffer in which [HA] = [A-] = 0.50 M.

Supposed that we had a buffer with [HA] = [A-] = 5.0 M. How much HCl would we need to add to get a pH change of -0.02 units? The answer is 10x as much as we found above, or 0.10 moles of HCl. This is summarized in this diagram:

A ten-fold increase in the concentration of our buffering agents increased the ability to absorb acid, i.e. the buffer capacity, ten fold. The buffer capacity is directly proportional to the concentration of our buffering agents.

Buffers Buffers ActAct When hydrochloric acid is added to the acetate buffer, the salt reacts When hydrochloric acid is added to the acetate buffer, the salt reacts

with the acid forming the weak acid, acetic acid and its salt. Similarly with the acid forming the weak acid, acetic acid and its salt. Similarly when when аа base is added, the acid reacts with it forming salt and water. base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised.Thus, changes in the pH are minimised.

СНСН33СООН + NaOH = СНСООН + NaOH = СН33COONa + НCOONa + Н22ОО

СНСН33СООСООNNа + HCI = СНа + HCI = СН33СООН + NaCIСООН + NaCI

The buffer capacity is determined by the absolute concentration of The buffer capacity is determined by the absolute concentration of the salt and acid. But the the salt and acid. But the рНрН of the buffer is dependent on the of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - relative proportion of the salt and acid (see the Henderson - Hasselbach's equation). When the ratio between salt and acid is 10:1, Hasselbach's equation). When the ratio between salt and acid is 10:1, the pH will be one unit higher than the pKa. When the ratio between the pH will be one unit higher than the pKa. When the ratio between salt and acid is 1:10, the pH will be one unit lower than the pKa.salt and acid is 1:10, the pH will be one unit lower than the pKa.

Factors Affecting pH of Factors Affecting pH of аа Buffer Buffer The pH of The pH of аа buffer solution is determined buffer solution is determined

by two factors:by two factors:1. The value of pK: The lower the value of 1. The value of pK: The lower the value of

pK, the lower is the pH of the solution.pK, the lower is the pH of the solution.2. The ratio of salt to acid concentrations: 2. The ratio of salt to acid concentrations:

Actual concentrations of salt and acid in Actual concentrations of salt and acid in аа buffer solution may be varied widely, buffer solution may be varied widely, with with попо change in change in рНрН, so long as the , so long as the ratio of the concentrations remains the ratio of the concentrations remains the same.same.

The Conceptual Problem with pH

Because it’s a logarithmic scale, it doesn’t make “sense” to our brains.

But Paul explains it well—every factor of 10 difference in [H+] represents 1.0 pH units, and

Every factor of 2 difference in [H+] represents 0.3 pH units.

Therefore, even numerically small differences in pH, can have profound biological effects…

62

How Can You Actually Determine the pH of a

Solution? Use a pH meter—read the

number. Use pH paper (color patterns

indicate pH). Titrate the solution with precise

amounts of base or acid in conjunction with a soluble dye, like phenolphthalein, whose color changes when a specific pH is reached.

8

Mechanisms for Regulation of pHMechanisms for Regulation of pH

1.1.      Buffers of body fluids, Buffers of body fluids, 2.2.      Respiratory system, Respiratory system, 3. Renal excretion. 3. Renal excretion.

These mechanisms are These mechanisms are interrelated. interrelated.

Acidic solutions have a high H+ Acidic solutions have a high H+ concentration. Base solutions have a concentration. Base solutions have a low H+ concentration. The pH scale is low H+ concentration. The pH scale is used to indicate the acidity or alkalinity used to indicate the acidity or alkalinity of a solution. Pure water with an equal of a solution. Pure water with an equal number of hydrogen and hydroxide ions number of hydrogen and hydroxide ions has a pH of 7.has a pH of 7.

Thank you for attentionThank you for attention