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pH Buffer Solution

• Introduction
  
  
• Theory
  
  
• pH versus temperature
  
  • Charts
    
    • Introduction
      
    • pH 7.00 chart
      
    • pH 6.86 chart
      
    • pH 4.01 chart
      
    • pH 3.56 chart
      
      
  • Graphs
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    

• Standard pH buffers are used to calibrate pH meters.
  
  
• pH buffer solutions are designed to resist changes in their pH values when they are contaminated by the additions of acid or basic substances.
  
  
• pH buffers can be purchased ready to go or made up using either commercially preparations sold as capsules or by measuring your own ingredients.
  
  
• Commercial preparations come colour coded to avoid confusion between the different buffer solutions.
  
  Red is usually chosen for pH 7, blue for pH 4 and yellow for pH 10 buffers.
  
  
• To prevent Buffer solution contamination -
  
  
  • Always pour stock buffer solutions in to testing containers.
    
    
  • Always discard buffer solutions after use.
    
    
  • Never pour the testing buffer solution back into the stock solution.
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    

• To understand fully how an acidic pH buffer solution works we have to go back to basics and build an understanding, step by step.
  
  

• An acid is a substance that is able to liberate hydrogen ions (H⁺), in solution.
  
  
• Two examples of acids are hydrochloric acid (HCl) and acetic acid (CH₃COOH - the acid component of vinegar).
  
  
• Acids, such as hydrochloric acid (HCl) and acetic acid (CH₃COOH), disassociate, to varying degrees, into two parts, when in solution.
  
  The first part is a H⁺ (hydrogen ion)
  
  The second part is called the conjugate base (Cl^- and CH₃COO^- respectively) which becomes a negatively charged ion.
  
  The disassociation is indicated in the expressions below.
  
  (1) HCl ==> H⁺ + Cl^-
  
  (2) CH₃COOH ==> CH₃COO^- + H⁺
  
  
• The extent of this dissociation or the number of free hydrogen ions (H⁺) liberated depends mainly on the strength of the acid, the concentration of the acid and the number of free H⁺ already present in solution.
  
  
• The pH of any solution is a measure of the final [H⁺] (free hydrogen ion concentration).
  
  
• HCl is a strong acid and will almost completely disassociate into H⁺ and its conjugate base (Cl^-).
  
  That is because the bonds between hydrogen (H ) and chlorine (Cl) are very weak.
  
  
• CH₃COOH (acetic acid) is a weak acid and will only slightly disassociate into H⁺ and it conjugate base (CH₃COO^-).
  
  That is because the bonds between H and acetate (CH₃COO) are much stronger.
  
  
• As the pH is a measure of the hydrogen ion concentration ([H⁺]) of a solution, a solution involving a weak acid will result in a solution with a higher pH (lower [H⁺]) then a solution involving a strong acid.
  
  A straight 1.0 M HCl solution has a pH of 0.00
  
  A straight 1.0 M acetic acid solution has a pH of 2.37
  
  
• The term, "conjugate base ", is so called as it was originally joined or associated with the original undisassociated acid form and has the potential to act as a base.
  
  It acts as a base by binding hydrogen ions to itself and reducing the solution's pH.
  
  That is, the reverse reaction of expressions (1) and (2).
  
  (3) HCl _(increase) <== H⁺ _(decrease) + Cl^-_(decrease)
  
  
  (4) CH₃COOH _(increase) <== CH₃COO^- _(decrease) + H⁺ _(decrease)
  
  
  The addition of the conjugate bases (Cl^- or CH₃COO^-) to a solution, via additions of their acid salts (KCl or NaOOCCH₃), will increase the interaction and bonding between the H⁺ and the conjugate bases to form the undissociated acid form, with a resultant shift in the solution's equilibrium to the left as shown in the expressions above.
  
  
  Note the changed direction of the arrows.
  
  Expressions (1) and (3), or (2) and (4), are more accurately written as shown below (5 & 6), with the double headed arrow indicating that the reaction or equilibrium can proceed in either direction.
  
  (5) HCl <==> H⁺ + Cl^-
  
  
  (6) CH₃COOH <==> CH₃COO^- + H⁺
  
  
  

Acid contaminations of a strong acid solution


• The addition of an acidic compound to a strong acid solution, such as a hydrochloric acid (HCl) solution, increases the solution's hydrogen ion concentration ([H⁺]).
  
  HCl <==> H^{+ (increase)} + Cl^-
  
  Some of the added H⁺ will be taken up or bound to the conjugate base (Cl^-) to form HCl establishing a new equilibrium between each other.
  
  HCl ^(increase) <== H^{+ (decrease)} + Cl^-(decrease)
  
  That is, adding H⁺ to the right hand side of expression will shift the the equilibrium to the left, increasing the undissociated acid form (HCl).
  
  When a strong acid is involved only a few of the added H⁺ are bound up.
  
  The reluctance of the H⁺ to bond to the conjugate base (Cl^-) of a strong acid, is due to the weak attractions between them. This results in most of the added H⁺ to exists as free H⁺.
  
  However, by increasing the concentration of the conjugate base (Cl^-), through additions of potassium chloride (KCl) to the HCl solution, you increase the number of interactions between H⁺ and Cl^-.
  
  HCl <==> H⁺ + Cl^-(increase)
  
  This then shifts the reaction further to the left and with some reduction in the number of free H⁺ from the added acid, helping to maintain the pH close to its initial value.
  
  HCl ^(increase)<== H^{+ (decrease)} + Cl^-(decrease)
  
  Also, by increasing the hydrochloric acid concentration ([HCL]), increases the [H⁺] and means that the percentage change of the [H⁺], and hence pH change, resulting from acid additions, will be relative small.
  
  HCl ^(increase)<==> H^+(increase) + Cl^-(increase)
  
  Therefore, increasing both the [HCl] and [Cl^-] will help buffer the HCL solution to some extent against pH changes due to added acids.
  
  
  Base contaminations of a strong acid solution
  
  
• The addition of a basic compound to a strong acid solution such as a HCl (hydrochloric acid) solution, increases the solution's hydroxide ion concentration ([OH^-]).
  
  Most, but not all of the OH^- (hydroxide ion) will be taken up or bound to the H⁺ to form water molecules (H₂O).
  
  (7) H⁺ + OH^- ==> H₂O
  
  The removal of H⁺ causes a significant shift to the right in the expression's equilibrium, with HCl disassociating to form H⁺ and Cl^-
  
  HCl ==> H⁺ + Cl^-
  
  However, there is an overall decrease in the final total [H⁺].
  
  In summary, the addition of a basic substance (e.g sodium hydroxide - NaOH) releases OH^-, which bonds to and removes most of the H⁺ from the right hand side of the expression and will shift the the equilibrium to the right, decreasing the undissociated acid form (HCl) and replacing most but not all of the removed H⁺.
  
  This predisposition of a strong acid to release H⁺ means only a very small change in the initial [H⁺] and the pH value will occur, on small additions of basic compounds to strong acid solutions.
  
  As well as this, if a strong acid solution had a high initial concentration, the percentage of the H⁺ bound to the added OH^-, would be relatively small with only a small change in the initial pH values.
  
  These two factors combined, sees a concentrated solution of a strong acid serve as a good pH buffer against small additions of basic compounds.
  
  
• Therefore a concentrated HCl solution to which a high concentration of KCl (potassium chloride) has been added will serve as pH buffer solution (of a low pH value) against changes in [H⁺] and [OH^-] and hence changes in pH in either direction.
  
  Such a buffer preparation is shown below.
  
  80 mL of a pH 1 buffer - 65 mL 0.2 M HCl + 25 mL 0.2 M KCl diluted to 100 mL
  
  
  

• The addition of an acidic compound to a weak acid solution such as an acetic acid (CH₃COOH) solution, increases the solution's hydrogen ion concentration ([H⁺].
  
  CH₃COOH <==> CH₃COO^- + H⁺_(increased)
  
  Most, but not all of the added H⁺ will be taken up or bound to the conjugate base (CH₃COO^-), establishing a new equilibrium.
  
  It is the strong attractions or interactions between H⁺ and the conjugate bases of weak acids that account for the extent of the uptake or bonding of added acids in weak acid solutions.
  
  Since the [H⁺] (free hydrogen ion concentration) determines the solutions pH, and only a small change in the [H⁺] occurs, the pH of a weak acid solution will remain relatively stable on small additions of acidic compounds.
  
  That is weak acid solutions serve as good pH buffers against added acidic compounds .
  
  The interactions and bonding between H⁺ and CH₃COO^- can be increased by increasing the concentration of the conjugate base (CH₃COO^-), through additions of sodium acetate (Na.OOCCH₃) to the acetic acid solution.
  
  Na.OOCCH₃ ==> Na⁺ + CH₃COO^-(increase)
  
  This then shifts the reaction below, further to the left and greatly reduces the number of free H⁺ from the added acid.
  
  CH₃COOH ^(increase)<== H^{+ (decrease)} + CH₃COO^{- (decrease)}
  
  Further to above, by increasing the initial [CH₃COOH] (acetic acid concentration) increases the [H⁺]. Therefore any small quantities of added acidic compounds not bound up, will result in only small percentage changes in the [H⁺], and hence only small changes in pH values.
  
  Therefore, a concentrated solution of a weak acid will enhance its pH buffering capacity in regards to acid contaminations.
  
  
  
• The addition of a basic compound to a weak acid solution such as an acetic acid (CH₃COOH) solution, increases the solution's hydroxide ion concentration ([OH^-]).
  
  Most, but not all of the OH^- will be taken up or bound to the H⁺ to form water molecules (H₂O).
  
  H⁺ + OH^- ==> H₂O
  
  The removal or the binding of the added H⁺ causes a small readjustment/shift to the right in the equilibrium of the expression shown below, with CH₃COOH disassociating to form extra H⁺ (to replace those removed) and CH₃COO^-.
  
  CH₃COOH ^(decrease) ==> H^+(increase) + CH₃COO^-(increase)
  
  However, there is an overall decrease in the initial [H⁺].
  
  The reluctance of a weak acid to release H⁺ means a significant change in the initial [H⁺] and pH will occur on small additions of basic compounds to weak acid solutions.
  
  To summarise, adding OH^- to, and removing H⁺ from the right hand side of the expression will shift the the equilibrium to the right, increasing H⁺ to replace some of the H⁺ that were removed in bonding to the introduced OH^-.
  
  However, by increasing the [CH₃COOH] (acetic acid concentration), increases the [H⁺].
  
  CH₃COOH ^(increase)<==> H^+(increase) + CH₃COO^-(increase)
  
  This means that the percentage change of the [H⁺], and hence pH change, resulting from acid additions, will not be as significant.
  
  Therefore, increasing both the [CH₃COOH] and [CH₃COO^-] will buffer the acetic acid solution against pH changes due to any added acids.
  
  
  
• Therefore a concentrated acetic acid (CH₃COOH) solution to which a high concentration of sodium acetate (NaCH₃COO) has been added will form a good pH buffer solution (of a moderate pH value) against changes in [H⁺] and [OH^-] and hence changes in pH in either direction.
  
  Such a buffer preparation is shown below.
  
  An approximate pH 5 buffer can be made by combining 29.5 mL 0.20M acetic acid (1.1 mL conc. / 100 mL) with 70.5 mL of 0.20 M sodium acetate (27.199 g NaC2H3O2/ liter)
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  

• The commercially prepared pH 7.00 buffer and the self prepared pH 6.86 buffers require different temperature corrections.
  
  
• A pH meter is only accurate to+/- 0.01 pH value.
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  
  • Commercially Prepared -pH 7 Standard Buffer Solution
    
    pH of 7.00 at 25 degrees C^o.
    
    pH Buffer 7 is a mixture of Potassium and Sodium Phosphates with a mold inhibitor.
    
    The pH 7 buffer has a blue color to distinguish it from other buffers.
    
    The following Table gives the pH value of this buffer at various temperatures.
    
    

    DEGREES-Co	DEGREES-Fo	pH
    0	32	7.12
    5	41	7.09
    10	50	7.06
    15	59	7.04
    20	68	7.02
    25	77	7.00
    30	86	6.99
    35	95	6.98
    40	104	6.97
    45	113	6.97
    50	122	6.98
    55	131	6.98
    60	140	6.98
    70	158	6.99
    80	176	7.00
    90	194	7.02

    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
  • Commercially Prepared -pH 6.86 Standard Buffer Solution
    
    pH of 6.86 at 25 degrees C^o.
    
    pH Buffer 6.86 is a mixture of Potassium and Sodium Phosphates.
    
    The pH 6.86 buffer has a blue color to distinguish it from other buffers.
    
    The following Table gives the pH value of this buffer at various temperatures.
    
    

    DEGREES-Co	DEGREES-Fo	pH
    0	32	6.984
    5	41	6.951
    10	50	6.923
    15	59	6.900
    20	68	6.881
    25	77	6.865
    30	86	6.853
    35	95	6.844
    40	104	6.838
    45	113	6.834
    50	122	6.833

    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
  • Commercially Prepared -pH 4 Standard Buffer Solution
    
    pH of 4.01 at 25 degrees C^o.
    
    pH Buffer 4.01 is a mixture of Potassium Acid Phthalate, Primary Standard.
    
    The pH 4.01 buffer has a red color to distinguish it from other buffers.
    
    The following Table gives the pH value of this buffer at various temperatures.
    
    

    DEGREES-Co	DEGREES-Fo	pH
    0	32	4.01
    5	41	4.00
    10	50	4.00
    15	59	4.00
    20	68	4.00
    25	77	4.01
    30	86	4.01
    35	95	4.02
    40	104	4.03
    45	113	4.04
    50	122	4.06
    55	131	4.07
    60	140	4.09
    70	158	4.12
    80	176	4.16
    90	194	4.20

    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    
    • Commercially Prepared -pH 3.56 Standard Buffer Solution
      
      pH of 3.56 at 25 degrees C^o.
      
      pH Buffer 3.56 is a concentrated solution of potassium bitartrate.
      
      The following Table gives the pH value of this buffer at various temperatures.
      
      

      DEGREES-Co	DEGREES-Fo	pH
      10	50	-
      15	59	-
      20	68	3.56
      25	77	3.56
      30	86	3.58

      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      
      

    • Note that the final [H⁺] (free hydrogen ion concentration) determines the solution's pH.
      
      
      • An acidic buffer solution consists of a weak acid and its conjugate (associated) base.
        
        
      • Acetic acid- CH₃COOH - like all acids are able to release hydrogen ions (H⁺) in solution.
        
        (1) CH₃COOH <==> CH₃COO^-+H⁺
        
        Acetic Acid - CH₃COOH
        
        and its conjugate base CH₃COO^-
        
        CH₃COO^- is a conjugate base, as it is able to accept hydrogen ions (H ⁺)
        
        (2) CH₃COO^-+H⁺ <==> CH₃COOH
        
        Acids and their conjugate base can be represented by AH and A^- respectively.
        
        
      • The species AH, A^- and H⁺ all exists in equilibrium with each other.
        
        
      • The relative concentrations of AH and A^- determine the solution's hydrogen ion concentration ([H⁺]) and hence its pH.
        
        
      • Acetic acid is a weak acid and will predominantly exist in its undissociated form, AH (CH₃COOH) in solution.
        
        
      • An acidic compound, added to a weak acid solution such as an acetic acid (CH₃COOH) solution, increases the [H⁺] (concentration), most of which, but not all of which, will be taken up or bound to the conjugate base and a new equilibrium will be reestablished between the species involved - as in expression (1).
        
        The degree to which this occurs depends on the acid strength of the added acid and more importantly on the concentration of the undissociated form, A^-, of the weak acid.
        
        
      • A basic or alkaline substance added to a acid solution increases the hydroxyl ion concentration ([OH^-]) concentration, some of which will bind to the free H⁺ present in solution.
        
        The degree to which this occurs depends on the base's strength of the added base and more importantly on the free H⁺ concentration, which is dependent on the concentration of the of the acidic compound (AH) present in solution.