STPM Chemistry Form 6 Notes – Ionic Equilibrium (Part 11)

by BerryBerryTeacher

in Berry Reference (Notes)

pH indicators or acid-base indicators are typically weak acids or weak bases.  Linking to the previous part, the indicators is commonly used in titrations during experiments involving analytic chemistry (typically to judge the extent of a chemical reaction). Indicators does not change colours at any specific hydrogen ion concentration, but over a range of hydrogen ion concentrations. This is named as the pH range, aka the colour change interval. Part 11 of Berry Berry Easy notes on Ionic Equilibrium for STPM Chemistry Form 6 students will be about indicators. It’s a topic most oft appear in exams but is always creating problems for students who fail to understand and memorise the colours of the indicators. So do take note of the case examples below.

[Tips: Colours are subjective, as they differ between what is perceived by the person conducting the experiment. As such, the use of pH indicators is subjected to non-precise readings. If you need precised measurements, a pH meter would be a better choice. Nonetheless, for STPM level, pH indicators are used. So you'll need to memorise the colours first and foremost. Make a list of them, their type and the colours involved. Hot topic in exams.]

STPM Chemistry Form 6 Notes – Ionic Equilibrium (Part 11)

Natural Universal Indicator_Purple cabbage

Natural Universal Indicator_Purple cabbage

Indicators

  • an organic compound with extended conjugation
  • weak acid (protonated) form and conjugate weak base (deprotonated)
  • form two distinct colour
  • when the pH of the solution is less than the pKa of the indicator (pH solution < pKa (indicator),
Indicator pH range Colour change (acid) Colour change (alkali)
Litmus 4.5-8.3 red blue
Phenolphthalein 8.3-10.0 colourless fuchsia (pink)
Methyl orange 3.1-4.4 red orange
Methyl red 4.4-6.2 red yellow
Bromothymol blue 6.0-7.6 yellow blue

HIn(aq) + H2O(l) <—-> H3O+(aq) + In-(aq)

HIn = acid (colour A)

In- = conjugate base (colour B)

Case 1:

Environment: < pH values (concentration H3O+ is high)
Le Chatelier’s principle: Equilibrium shifts to the left
Colour: A

Case 2:

Environment: > pH values (concentration H3O+ is low)
Le Chatelier’s principle: Equilibrium shifts to the right
Colour: B

Case 3:

Environment: [HIn] = [In-]
Colour: Intermediate colour

KHIn = ([H3O+] [In]) / [HIn]
KHIn = indicator dissociation constant

Choosing an indicator

I) Strong acid – Strong base Titration

  • pH range from 3.5 to 10.5
  • Suitable indicator: Phenolphthalein, methyl red, bromothymol blue)

II) Strong acid – Weak base Titration

  • pH range from 3.1 to 7.0
  • Suitable indicator: Methyl orange

III) Weak acid – Strong base Titration

  • pH range from 8.3 to 10.0
  • Suitable indicator: Phenolphthalein

IV) Weak acid – Weak base Titration

  • pH range has no sharp
  • Suitable indicator: no indicator

V) Weak Polyprotic Acids

  • There are few equivalence points
  • First equivalence point = pH range at 4.6
  • Suitable indicator: Methyl orange, Methyl red
  • Second equivalence point = pH range at 9.0
  • Suitable indicator: Phenolphthalein

Natural acid-base indicator

Anthocyanin (red in acidic solutions and blue in basic)

  • red cabbage leaves
  • rose petals
  • skin of a lemon
  • berries
  • hydrangea

[Bonus: it is important to understand that the pH range stated might shift very slightly depending on the indicator's concentration level in the solution and also the temperature of the solution. So take special note when you run experiments involving indicators.]

In the next part, Part 12 of Berry Berry Easy notes on Ionic Equilibrium for STPM Chemistry Form 6 will be on biology buffer, Henderson-Hasselbalch equation and examples related to it. So stay tuned.

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