5.1.3 - Acids, Bases and Buffers

• Brønsted-Lowry Acids and Bases

  Brønsted-Lowry acid= proton donor.

  Brønsted-Lowry base= proton acceptor.
• Conjugate Acid-Base Pairs

  Contains two species that can beinterconvertedby transfer of a proton.
• Monobasic Acid

  One mole of acid dissociates to form one mole of protons.
• Monobasic, Dibasic, and Tribasic Acids

  Refer to the total number of hydrogen ions (protons) in the acid that can be replaced per molecule in an acid-base reaction.

  
  

  e.g., HCl - monobasic

  H2CO3 - dibasic

  H3PO4 - tribasic
• Role of Protons in the Reactions of Acids

  Acid + Metal -> Salt + Hydrogen

  
  

  Acid + Carbonate -> Salt + Water + Carbon Dioxide

  
  

  Acid + Base -> Salt + Water

  
  

  Acid + Alkali -> Salt + Water
• Dissociation of a Weak Acid

  HA ⇌ H+ + A-
• Acid Dissociation Constant, Ka

  Ka = [H+][A-]/[HA]
• pH

  pH is a logarithmic scale, meaning a change of one pH number is equal to 10 times the difference in concentration of H+

  
  

  pH = -log[H+]

  
  

  [H+] = 10^-pH
• Ka and pKa

  pKa = -logKa

  
  

  Ka = 10^-pKa

  
  

  This exists as pKa values are much more manageable than Ka values.
• pH of Weak Acids

  HA ⇌ H+ + A-

  A weak acid, HA, partially dissociates.

  [H+] depends upon [HA] and the acid dissociation constant, Ka.

  When HA molecules dissociate, A- and H+ ions are formed in equal quantities.

  Based on the above, two approximations exist which allow Ka to be simplified further:

  
  

  Ka = [H+]^2/[HA]
• Determination of Ka for a Weak Acid


  Prepare a standard solution of the weak acid of known concentration.

  
  

  Measure the pH of the standard solution using a pH meter.

  
  

  The acid concentration ([HA]) and the pH (which can be converted to [H+]) allows for Ka to be calculated.
• Limitations of the Approximations Used

  Assumes that the [H+] from the dissociation of water is negligible.

  
  

  However, at a pH greater than 6 the dissociation of water will be significant compared with dissociation of a weak acid.

  
  

  The second approximation assumes that [HA] is much greater than the [H+] at equilibrium.This approximation will only hold for weak acids with small Ka values.
• Ionic Product of Water, Kw

  Kw = [H+] [OH-]

  Kw has a set value at 298K, setting up a neutral point within the pH scale.

  Kw controls the concentration of H+ and OH- ions in aqueous solution.
• Calculating pH in a Dilute Solution

  Find [H+] in original solution.

  
  

  [H+] in dilute = original [H+] × (old volume/new volume)

  
  

  Then -log the dilute [H+].
• pH of Solutions of Strong Bases

  Can be calculated from the concentration of the base, [OH-], and Kw.

  With a strong base, e.g., NaOH, it completely dissociates, so: [NaOH] = [OH-]

  • Find [OH-]
  • Use Kw = [H+][OH-]to calculate [H+]
  • Use [H+] to find pH.
• Calculating pH of Mixtures

  • Write balanced symbol equation.
  • Calculate initial moles of H+ and OH- ions.
  • H+ and OH- react in 1:1 - find excess.
  • Calculate volume of solution by addition, then converted to dm3.
  • Calculate concentration of excess by using n = cV.

  If excess is H+:

  • pH = -log[H+]

  If excess if OH-:

  • Kw = [H+][OH-]
  • therefore,[H+] = Kw/[OH-] then,
  • pH = -log[H+]
• Calculating Weak Acid + Strong Base

  • Calculate moles of HA
  • Calculate moles of OH-
  • Calculate which is in excess.

  If excess is HA:

  • Calculate moles of excess HA and A- formed.
  • Calculate [HA] left over and [A-] formed.
  • Use Ka = ([H+][A-])/[HA]to find [H+]
  • Find pH

  If excess OH-:

  • Calculate [OH-]
  • Use Kw = [H+][OH-] to find [H+]
  • Find pH

  If moles OH- = HA:

  • pH = pKa of weak acid.
• Buffer Solutions


  A system that minimises pH changes when small amounts of an acid or a base are added.

  
  

  The weak acid, HA, removes added alkali.

  
  

  The conjugate base, A-, removes added acid.

  
  

  As the buffer works, the pH does change but only by a small amount.
• Preparing Weak Acid Buffer Solutions - Weak Acid and its Salt


  A buffer solution can be prepared by mixing a solution of ethanoic acid with one of its salts, e.g, sodium ethanoate.

  
  

  When ethanoic acid is added to water, the acid partially dissociates - this is the source of the weak acid component.

  
  

  When added to water, the salt completely dissociates into ions - this is the source of the conjugate base component.
• Preparing Weak Acid Buffer Solutions - Neutralisation of the Weak Acid

  Adding aqueous solution, e.g., NaOH, to an excess of weak acid.

  
  

  The weak acid is partially neutralised by the alkali, forming the conjugate base.

  
  

  Some of the weak acid is left over, unreacted.
• Action of the Buffer Solution - Conjugate Base

  Removes added acid.

  
  

  [H+] increases.

  
  

  H+ ions react with the conjugate base, A-

  
  

  The equilibrium position shifts to the left, removing most of the H+ ions.
• Action of the Buffer Solution - Weak Acid

  Removes added alkali.

  
  

  [OH-] increases.

  
  

  The small concentration of H+ ions reacts with the OH- ions.

  
  

  HA dissociates, shifting the equilibrium position to the right to restore most of the H+ ions.
• pH and Buffer Solutions


  A buffer is most effective when there are equal concentrations of the weak acid and its conjugate base:

  [HA] = [A-]

  
  

  When [HA] = [A-]:

  - The pH of the buffer solution is equal to the pKa value of HA.
• Calculating the pH of a Buffer Solution

  HA ⇌ H+ + A-

  
  

  Ka = [H+][A-]/[HA]

  
  

  [H+] = Ka x ([HA]/[A-])

  Provided that Ka and the concentrations of HA and A- are known, [H+] and therefore the pH can be calculated.
• Buffer Solutions in the Body


  Blood plasma pH is maintained by the carbonic acid-hydrogencarbonate buffer system.

  
  

  Acid Added:

  • [H+] increases.
  • H+ ions react with the conjugate base, HCO3 -
  • Equilibrium shifts left, removing most of the H+ ions.

  Alkali Added:

  • [OH-] increases.
  • Small concentration of H+ ions reacts with the OH- ions.
  • H2CO3 dissociates, shifting the equilibrium position to the right to restore most of the H+ ions.
• pH Titration Curve

  This is an acid-base titration, base-acid titrations are flipped.

  
  

  Excess of Acid - when the base is first added, the acid is in great excess, and the pH increases very slightly.

  
  

  Vertical Section - as this section is approached, the pH starts to increase more quickly as the acid is used more quickly.Eventually, the pH increases rapidly during addition of a small volume of base.

  
  

  Excess of Base - pH will rise very slightly as the base is now in great excess.

  
• pH Titration Curve - Equivalence Point

  The volume of one solution that exactly reacts with the volume of the other solution.

  
  

  This is the centre of the vertical section of the pH titration curve.
• Acid-Base Indicators


  It is a weak acid that has a distinctively different colour from its conjugate base, A-

  
  

  E.g., methyl orange.

  - HA is red

  - A- is yellow

  
  

  At the end point, the indicator contains equal concentrations of HA and A- meaning the colour at its end point is orange.
• pH Meters

  Prior to using, it must be calibrated:

  • Place the bulb of the pH meter into distilled water, allow the reading to settle, and adjust the reading so it is 7.0
  • Do the same with a standard solution of pH 4 and pH 10.
  • Rinse the probe with distilled water between each reading.

  Alternatively, a calibration curve can be drawn.