Topic 12- Part 2

Topic 12- Part 2

1. pH Indicators and Half Equivalence

   • pH INDICATOR- Compound added in small amounts to a solution so pH of a solution can be determined visually
     • All indicators are weak acids
   • Litmus is a weak acid. There will be an equilibrium established when this acid dissolved in water
     • Hlit(aq) (acid) ⇌ H+(aq) + Lit-(aq) (alkali)
       • Adding H+ ions causes equilibrium to shift left so colour change would be from blue to red
       • Hydroxide ions react with and removes H+ ions. Equilibrium shifts right to replace lost H+ ions. Litmus turns blue
   • Using indicators: Indicator range must be smaller and within vertical range of titration curve. pKin value similar to pH at equivalence point (end point)
     • Kin=[H+] pKin=pH
   • ENDPOINT- Good approximation for the equivalence point as long as a suitable indicator has been selected. Point reached when indicator changes colour
     • EQUIVALENCE POINT- Point at which chemically equivalent quantities of bases and acids mixed
   • HALF EQUIVALENCE- Point at which half the acid has been turned to salt [HA]=[A-]. It is half the concentration hence half volume
     • Steps to finding Ka with half equivalence:
       • 1) 1/2 equivalence volume
         • 2) Draw line to pH at 1/2 equivalence volume
           • 3) pH=pKa
             • 4) Ka= 10^(-pKa)

2. Buffers

   • BUFFER- solution that resists in pH when small amounts of acid or alkali are added
   • USES:
     • Swimming pools- prevents chemicals from irritating skin.
       • Soda- prevents acidic flavouring from damaging teeth
         • Rivers- protects against effect of acid rain
   • ACIDIC BUFFERS FORMATION
     • Acidic buffers have a pH lower than 7. Made of a weak acid and conjugate base (salt)
       • Example of ethnic acid and sodium/potassium ethanoate
         • CH3COOH ⇌ CH3COO- + H+
           • Barely dissociates, causing equilibrium to lie to the left as weak acids barely dissociate
         • CH3COONa --> CH3COO- + Na+
           • [CH3COO-] remains high. Equilibrium would lie to the right, so lots of soluble salt as it fully dissociates
     • ACIDIC BUFFERS- FUNCTION
       • Small amounts of H+ added: CH3COOH ⇌ CH3COO- + H+ : Equilibrium shifts left. Concentration of CH3COOH increases. The H+ would react with CH3COO- hence larger amounts of CH3COOH to be formed
       • Small amounts of OH- added: H+ + OH- --> H2O CH3COOH ⇌ CH3COO- + H+ : Reacts with H+ and equilibrium shifts right to replace H+. Reacts with CH3COOH forming salt and water
         • OR CH3COOH + OH- --> CH3COO- + H2O : Reacts with CH3COOH forming salt and water [HA] : [A-] remains constant
   • ALKALINE BUFFERS FORMATION
     • Alkali buffers have a pH greater than 7. Made from a weak base and conjugate acid (salt)
       • Example of ammonia and ammonium chloride
         • NH3 + H2O ⇌ NH4+ + OH-
           • Equilibrium shits to the left as weak bases barely dissociate
         • NH4Cl --> NH4+ + Cl-
           • Lots of soluble salt as fully dissociates
     • ALKALINE BUFFERS- FUNCTION
       • Small amounts of H+ added: H+ + OH- --> H2O NH4 + H2O ⇌ NH4+ + OH-. Reacts with OH- and equilibrium shifts right to replace OH-
         • OR H+ + NH3 --> NH4+. Reacts with NH3 forming NH4+. [NH3] : [NH4+] remains constant.
       • Small amounts of OH- added: NH3 + H2O ⇌ NH4+ + OH- Reacts with NH4+ equilibrium shifts left.
   • HOW DOES A MIXTURE ACT AS A BUFFER SOLUTION?
     • Buffers are made up of weak acid and its conjugate base. The contain large reservoirs of [HA] and [A-]
       • When small amounts of H+] is added H+ + A- --> HA so equilibrium shifts to the let to counter the change
         • When small amounts of [OH-] is added OH- + H+ --> H2O, OH- + HA --> H2O + A-. The equilibrium shifts to the right to replace lost H+
           • Concentration of [HA] is almost equal as [A-] and remains unchanged
   • EXPLAINING pH CURVE
     • BUFFERING CAPACITY- Quantity of H+/OH- ions it can absorb without a significant change in its pH
     • Initially, since [ethanoic] acid is a weak acid, there is a relatively low concentration of hydrogen ions.
       • As these react with the alkali, pH changes rapidly
         • At the same time, however [sodium ethanoate] is formed
           • Once a moderate amount of this is present, solution acts as a buffer
             • Until such time as its buffering capacity is exhausted [A-] > [HA]
               • Which is near to the endpoint of the titration
   • When pH at equivalence of the titration is greater than 7 HCOO- + H2O --> HCOOH + OH-. Since OH- produced pH will be >7
     • When pH at equivalence is less than 7 NH4+ +H2O --> NH3 + H2O. Since H3O+ produced concentration of H+ ions increases hence lower pH

3. Bicarbonate Buffer System

   • pH of blood is controlled by the carbonic acid and hydrogen carbonate buffer mixture, where hydrogen carbonate HCO3- is conjugate base.
     • Dissociation of carbonic acid: H2CO3 ⇌ H+ + HCO3 -
       • Buffer deals with production of. lactic acid through the addition of H+ causing equilibrium to shift left, pH would decrease.
       • Buffer deals with production of excess urine via equilibrium shifting to the right since H+ is removed
   • Respiratory System
     • H2O + CO2 ⇌ H2CO3 (Carbonic acid) ⇌ HCO3 - (bicarbonate ion) + H+
     • Carbonic acid buffer is coupled to the respiratory system. Carbonic acid is unstable and decomposes to form CO2 and H2O. More carbonic acid in blood, the more CO2 dissolves in blood sombre CO2 exhaled.
     • Increasing exercise increases the amount of CO2 hence equilibrium shifts right hence more carbonic acid is formed causing second equilibrium to shift right hence HCO3- and H+ formed so pH decreases.

4. Enthalpy of Neutralisation

   • STANDARD ENTHALPY OF NEUTRALISATION- Enthalpy change when one mole of water is produced by the neutralisation of an acid with an alkali at 100kPa and 298K
   • Ionic equation is always H+ + OH- --> H2O
   • Confirming reaction is exothermic: Measure initial temperature, mix solutions in. polystyrene cup, measure temperature each minute. Extrapolate graph and if temperature increases, it is an exothermic reaction.
   • Weak Acids
     • Ionic equation: HA(l) + OH-(aq) --> A-(aq) + H2O(l)
     • Don't fully dissociate. Energy is required to be fully dissociated to overcome electrostatic forces of attraction between oppositely charged ions
     • delta Hn (weak acid) = delta Ht (strong acid) + delta Hn (strong acid)