5.3, 5.4- Acids, Bases and pH, buffers and neutralisation

bronsted-lowry theory

• acid is a substance that donates protons (H⁺)

• base is substance that accepts protons (H⁺)

how are acids categorised

• based on their capacity to donate protons:

  • monoprotic→ donates 1 proton per molecule

  • diprotic→ donates 2 protons per molecule

  • triprotic→ donates 3 protons per molecule

conjugate acid-base pairs

• consists of 2 species that are interconverted by transfer of H⁺

• in forward reaction, acid donates proton, but in reverse reaction conjugate base accepts proton

• in forward reaction, base accepts proton, but in reverse reaction conjugate acid donates proton

water

• water is amphiprotic→ can behave as both acid or base

• when reacting with acids, water accepts proton to form H₃O⁺ (hydronium ion)

• when reacting with bases, water donates proton to form OH^- (hydroxide ion)

metal and acid

• salt and hydrogen

• redox reaction

metal carbonate and acid

• salt, water and CO₂

• neutralisation

insoluble bases and acid

• salt and water

• neutralisation

acids and alkalis

• salt and water

• neutralisation

development acid-base theories

• lavoisier:

  • proposed that all acids contain oxygen

  • doesn’t apply to acids e.g. HCl

• Arrhenius:

  • acids release H⁺ in water and bases release OH^-

  • does not account for bases e.g. NH₃ that do not release OH^-

the pH scale

• logarithmic scale that measures concentration of H⁺ in a solution

• ranges from 0-14:

  • <7 is acidic solution

  • =7 is neutral

  • >7 is alkaline solution

calculating pH (+ inverse)

• pH=-log₁₀[H⁺]

• [H⁺]=10^-pH

H⁺ concentration in different strong acids

• Monoprotic→ [H⁺]=[acid]

• Diprotic→ [H⁺]= 2[acid]

dissociation of H⁺ in weak acids

• partially dissociate in aqueous solution

• conc of H⁺ less than initial conc. of acid

acid dissociation constant (K_a)

• K_a= [H⁺][A^-]/[HA]

• larger K_a= stronger the weak acid

assumptions made for weak acids

• [HA]_equilibrium≈[HA]_initial → ionisation of weak acid is so small that conc. of undissociated HA at equilibrium is approx. same as initial conc.

• [H⁺]_equilibrium≈[A^-]_equilibrium → ionisation of water is negligible

• assumptions simplify K_a formula to K_a=[H⁺]²/[HA]

when do K_a assumptions become inaccurate

• for stronger acids withe K_a>10^-2, conc H⁺ is significant so [HA]_equilibrium≠[HA]_initial

• for very weak acids or very dilute solutions, conc H⁺ from dissociation of water becomes significant so [H⁺]_equilibrium≠[A^-]_equilibrium

pK_a

pK_a=-log₁₀(K_a)

what is a buffer solution

a solution that minimises change in pH on addition of small amounts of acid or base.

type 1 buffers

• weak acid (HA)

• salt of the weak acid/ conjugate base of weak acid (A^-M⁺)

type 2 buffers

• weak acid in excess (xs HA)

• small amount of strong alkali (M⁺OH^-)

how do buffers minimise pH change on addition of acid

• more acid= increased [H⁺]

• based on Le Chatelier’s principle, equilibrium position shifts to left

• more H⁺ ions react with CH₃COO^- to form CH₃COOH

• causes [H⁺] to decrease

how do buffers minimise pH change on addition of alkali

• when alkali added, increase in [OH^-]

• OH^- react with H⁺ to form water

• causes [H⁺] to decrease

• equilibrium will shift to the right

• more CH₃COOH dissociate, so [H⁺] increases

calculating pH of type 1 buffer solutions

• cannot simplify to [H⁺]² because [A^-]>[H⁺]

• We can assume [HA]_{(equilibrium)}= [HA]_original

calculating pH of type 2 buffers

what range does blood plasma need to have a pH between

• 7.35-7.45

• below= acidosis

• above= alkalosis

components of carbonic acid-hydrogen carbonate buffer system

• carbonic acid (H₂CO₃)- weak acid

• Hydrogen carbonate ion (HCO₃^-)- conjugate base

blood buffer equation

H₂CO_3(aq) ⇌ H⁺_(aq) + HCO₃^-_(aq)

what happens if there is an increase in H⁺ in blood buffer system

• [H⁺] increases

• extra H⁺ reacts with HCO₃^- to form H₂CO₃

• position of equilibrium shifts to left to remove most of the extra H⁺

what happens when there is an increase of OH^- in blood buffer system

• increases [OH^-]

• extra [OH^-] reacts with H⁺ to form H₂O

• decreases H⁺

• equilibrium position shifts to the right to restore most of used up H⁺

understanding titration curves

• starts at pH below 7→ shows that acid in conical and alkali in burette

• ends at pH above 7→ shows there is an excess of alkali at end of titration

sections of a titration curve

1. only slight increase in pH when alkali is added because acid is in excess

2. sharp rise in pH when alkali is added because concentrations of acid and alkali are now similar and proportion of H⁺ used up increases

3. slight increase in pH when alkali added because alkali is in excess and more addition of OH^- has little impact on pH

equivalence point

• the volume of solution that completely reacts (and therefore neutralises) the volume of the other solution

finding equivalence point on titration curve

1. find centre of vertical section

2. draw a vertical line going down x-axis to identify corresponding volume

strong acid strong base titration curve

strong acid weak base titration curve

• end of curve shifted down- weak base has lower pH

weak acid strong base titration curve

• start of graph shifted up- weak acid has higher pH

weak acid weak base titration curve

• start shifted up, end shifted down

what are pH indicators

• weak acids in solution

• conjugate weak acid has a different colour to its conjugate base

what is end point of a titration

• equal concentrations of weak acid and conjugate base forms of the indicator→ colour chnage observed

choosing suitable indicators

• indicator must have an end point colour change that coincides with equivalence point