1. Buffer solutions. ass. prof. I. R. Bekus

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

3. Importance of water 1. 1. The most abundant substance in living systems. – –Makes up 70% or more of the weight of most organisms. 2. 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. 3. Ubiquitous solvent in cells. 4. 4. Excellent solvent of polar and ionic substances. 5. 5. Medium where most cell’s metabolic reactions take place. 6. 6. Ionization of water and its acid-base reactions important for the functions of proteins and nucleic acids. 7. 7. The Shapes of proteins and nucleic acids and structure of biological membranes are consequence of their interaction with water.

4. “Water molecule” is polar Water has a simple structure. Oxygen has 6 electrons in the outer shell: 1s 2, 2s 2, 2p 4. 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.

5. Water is а neutral molecule with а slight tendency to ionize. We usually express this ionization as: Н 2 О = Н + + ОН -

6. 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, H 3 O +. The association of а proton with а cluster of water molecules also gives rise to structures with the formulas Н 5 О 2 +, Н 7 О 3 +, and so on. For simplicity, however, we collectively represent these ions by H +.

7. 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), 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

8. Ion Product of Water, K w The ion product constant, K w, for water * is the product of the concentrations of the hydronium and hydroxide ions. K w = [ H 3 O + ] [ OH − ] * can be obtained from the concentrations in pure water. K w = [ H 3 O + ] [ OH − ] K w = [1.0 x 10 − 7 M] x [ 1.0 x 10 − 7 M] = 1.0 x 10 − 14

9. Calculating [H 3 O + ] What is the [H 3 O + ] of a solution if [OH − ] is 5.0 x 10 -8 M? STEP 1: Write the K w for water. K w = [H 3 O + ][OH − ] = 1.0 x 10 −14 STEP 2: Rearrange the K w expression. [H 3 O + ] = 1.0 x 10 -14 [OH − ] STEP 3: Substitute [OH − ]. [H 3 O + ] = 1.0 x 10 -14 = 2.0 x 10 -7 M 5.0 x 10 - 8

10. 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 -14 Household ammonia -12 Seawater – 8 Milk - 7 Blood - 7.4 Saliva - 6.6 Tomato juice - 4.4 Vinegar - 3 Gastric juice - 1.5 1 М НСl - 0

11. The relationship between the pH of а solution and the concentrations of an acid and its conjugate base is easily derived. [НА] [НА] [Н + ]= K ---------- [А - ] [А - ] Taking the negative log of each term [А - ] [А - ] рН = - log К + log --------- [А - ] [А - ] This relationship known as the Henderson-Hasselbach equation, that is often used to perform the calculations required in preparation of buffers for use in the laboratory, or other applications.

12. pH pH is commonly expressed as – log [H + ] Pure water has [H + ]=10 -7 and thus pH=7.

13. The methods for measuring pH fall roughly into the following four categories: 1) Indicator methods 2) Metal-electrode methods (including the hydrogen-electrode method, quinhydron-electrode method and antimony-electrode method) 3) Glass-electrode methods 4) Semiconductor sensor methods

14. Ways to measure pH pH meter Electrode measures H + concentration Must standardize (calibrate) before using.

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

16. 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.

17. According to their composition buffers divided into such groups: 1) 1) Acidic buffer solution – consist of a weak acid and a salt of this weak acid and a strong base (acetate buffer solution ) 2) 2) Basic buffer solution - consist of a weak base and a salt of weak base and a strong acid (ammonium buffer solution ) 3) 3) Protein ampholytic buffer solution

18. Buffers are mainly classified of two types: (а) mixtures of weak acids with their salt with а strong base (b) mixtures of weak bases with their salt with а strong acid. А few examples are given below: Н 2 СО 3 / NаНСО 3 (Bicarbonate buffer; carbonic acid and sodium bicarbonate) СН 3 СООН / СН 3 СООNa (Acetate buffer; acetic acid and sodium acetate) Na 2 HPO 4 / NaH 2 PO 4 (Phosphate buffer)

19. The most widely used in laboratory practice buffer solution Buffer mixture name CompositionpH FormateFormic acid HCOOH and sodium formate HCOONa 3,8 BenzoateBenzoic acid C 6 H 5 COOH and sodium benzoate C 6 H 5 COONa 4,2 AcetateAcetic acid CH 3 COOH and sodium acetate CH 3 COONa 4,8 PhosphateSodium dihydrogen phosphate NaH 2 PO 4 and sodium hydrogen phosphate Na 2 HPO 4 6,6 AmmoniumAmmonium hydrate NH 3 *H 2 O and ammonium chloride NH 4 Cl 9,2

20. Chemistry of buffers -pH = -pKa + log 10 [acid]/[base] Multiply both sides by –1 to get the Henderson-Hasselbach equation – –pH = pKa - log 10 [acid]/[base]

21. 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 - log 10 [acid]/[base] pH = pKa - log 10 [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)

22. Calculating buffer recipes Henderson-Hasselbach equation – –pH = pKa - log 10 [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.

23. 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) = 10 0.8 = 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.

26. Acid–Base Concept 1) The Arrhenius theory ACID - a substance that provides H+ ions in water BASE - a substance that provides OH- ions in water

28. 2) The Brønsted-Lowry Theory All Brønsted–Lowry bases have one or more lone pairs of electrons:

30. 3) The Lewis Acids and Base theory LEWIS ACID An electron-pair acceptor LEWIS BASE An electron-pair donor

31. 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

32. Acids Lemon juice contains citric acid, and vinegar contains ethanoic acid. Some strong acids are hydrochloric acid, sulphuric acid and nitric acid. Some weak acids are ethanoic acid, citric acid and carbonic acid. There are many acids present in our everyday lives.

33. 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

34. Weak acids thus are in equilibrium with their ionized species: 34 HA H + + A - K eq = [H + ][A - ] [HA] Governed by the Law of Mass Action, and characterized by an equilibrium constant:

35. 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

36. Alkalis When the oxides of some metals dissolve in water they make an alkali solution. When the oxides of some metals dissolve in water they make an alkali solution. Alkalis react with acids and neutralise them. Alkalis react with acids and neutralise them.neutralise Many everyday substances are alkalis. They feel soapy. They are corrosive.

37. Alkalis Alkalis are present in many cleaning substances in use in our homes. Kitchen cleaners are alkaline because they contain ammonia or sodium hydroxide, which attack grease. Calcium hydroxide and sodium hydroxide are strong alkalis. The most recognisable and common weak alkali is ammonia.

39. Important buffers

40. 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 H 2 CO 3 /HCO 3 - Some CO 2, 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 (HCO 3 - ) and hydronium (H 3 O + ) ions. More of the HCO 3 - is supplied by the kidneys. CO 2 + H 2 O ↔ H 2 CO 3 H 2 CO 3 + H 2 O ↔ H 3 O + + HCO 3 -

41. Bicarbonate buffer system H 2 CO 3 + H 2 O ↔ H 3 O + + HCO 3 - Excess acid (H 3 O + ) in the body is neutralized by HCO 3 - H 2 CO 3 + H 2 O ← H 3 O + + HCO 3 - Equilibrium shifts left Excess base (OH - ) reacts with the carbonic acid (H 2 CO 3 ) H 2 CO 3 + OH - → H 2 O + HCO 3 - Equilibrium shifts right

42. pH of the bicarbonate buffer The concentrations in the blood of H 2 CO 3 and HCO 3 - are 0.0024M and 0.024 respectively H 2 CO 3 / HCO 3 - = 1/10 is needed to maintain the normal blood pH (7.35 – 7.45)

44. Regulation of blood pH The lungs and kidneys play important role in regulating blood pH. The lungs regulate pH through retention or elimination of CO 2 by changing the rate and volume of ventilation. The kidneys regulate pH by excreting acid, primarily in the ammonium ion (NH 4 + ), and by reclaiming HCO 3 - from the glomerular filtrate (and adding it back to the blood).

45. The concentration of carbonic acid in the body is associated with the partial pressure of CO 2. * When CO 2 level rises, producing more H 2 CO 3, the equilibrium produces more H 3 O +, which lowers the pH – acidosis. * Decreasing of CO 2 level due to a hyperventilation, which expels large amounts of CO 2, leads to a lowering in the partial pressure of CO 2 below normal and the shift of the equilibrium from H 2 CO 3 to CO 2 and H 2 O. This shift decreases H 3 O + and raises blood pH – alkalosis. CO 2 + H 2 O ↔ H 2 CO 3 ↔ H 3 O + + HCO 3 -

47. Respiratory Acidosis: CO 2 ↑ 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

48. Respiratory Alkalosis: CO 2 ↓ 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

49. 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

50. 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

51. Other important buffers The phosphate buffer system (HPO 4 2- /H 2 PO 4 - ) plays a role in plasma and erythrocytes. H 2 PO 4 - + H 2 O ↔ H 3 O + + HPO 4 2- Any acid reacts with monohydrogen phosphate to form dihydrogen phosphate dihydrogen phosphate — monohydrogen phosphate H 2 PO 4 - + H 2 O ← HPO 4 2- + H 3 O + The base is neutralized by dihydrogen phosphate dihydrogen phosphate — monohydrogen phosphate H 2 PO 4 - + OH - → HPO 4 2- + H 3 O +

52. Proteins act as a third type of blood buffer Proteins contain – COO - groups, which, like acetate ions (CH 3 COO - ), can act as proton acceptors. Proteins also contain – NH 3 + groups, which, like ammonium ions (NH 4 + ), can donate protons. If acid comes into blood, hydronium ions can be neutralized by the – COO - groups - COO - + H 3 O + → - COOH + H 2 O If base is added, it can be neutralized by the – NH 3 + groups - NH 3 + + OH - → - NH 2 + H 2 O

55. Buffer Capacity Buffer capacity is determined by the actual concentrations of salt and acid present, as well as by their ratio. Buffering capacity is the number of grams of strong acid or alkali which is necessary for а change in pH of one unit of one liter of buffer solution. The buffering capacity of а buffer is, definеd аs the ability of the buffer to resist changes in pH when an acid or base is added.

56. 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 ( β max ) obtain when ratio of acid to salt = 1 i.e. pKa = pH

57. Definition The 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

58. 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.

59. 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.

60. Buffers Act 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 when а base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised. 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 when а base is added, the acid reacts with it forming salt and water. Thus, changes in the pH are minimised. СН 3 СООН + NaOH = СН 3 COONa + Н 2 О СН 3 СООNа + HCI = СН 3 СООН + NaCI The buffer capacity is determined by the absolute concentration of the salt and acid. But the рН of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - 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 salt and acid is 1:10, the pH will be one unit lower than the pKa. The buffer capacity is determined by the absolute concentration of the salt and acid. But the рН of the buffer is dependent on the relative proportion of the salt and acid (see the Henderson - 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 salt and acid is 1:10, the pH will be one unit lower than the pKa.

61. Factors Affecting pH of а Buffer The pH of а buffer solution is determined by two factors: 1. The value of pK: The lower the value of pK, the lower is the pH of the solution. 2. The ratio of salt to acid concentrations: Actual concentrations of salt and acid in а buffer solution may be varied widely, with по change in рН, so long as the ratio of the concentrations remains the same.

62. 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

63. 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

64. Mechanisms for Regulation of pH 1. Buffers of body fluids, 1. Buffers of body fluids, 2. Respiratory system, 2. Respiratory system, 3. Renal excretion. 3. Renal excretion. These mechanisms are interrelated. Acidic solutions have a high H+ concentration. Base solutions have a low H+ concentration. The pH scale is used to indicate the acidity or alkalinity of a solution. Pure water with an equal number of hydrogen and hydroxide ions has a pH of 7.

65. Thank you for attention