pH + Indicators - Chp 18

Derive an equation for the ionic product of water

Write an equation for the self ionisation of water

pH

-log₁₀[H⁺] where the square brackets indicate the concentration in moles per litre

or

the pH of a solution is the negative logarithm to the base of 10 of the hydrogen ion concentration measured in moles per litre

What is the value of Kw at 25^oC

Kw = 1 × 10^-14

Show that the pH of pure water is 7 at 25^oC?

What trend is shown when the pH of pure water is measured at temperatures other than 25^oC?

• the pH of water decreases as temp decreases HOWEVER it does not become more acidic

• a solution is acidic if there is an excess of H⁺ ion but in the case of water there is a 1:1 ratio between H⁺ ions and OH- ions

• they are always present in equal concentrations

Define the acid dissociation constat + what does it express?

expresses the strength of an adic/how well it donates protons

when K_a < 0.01 it is generally considered a weak acid

the larger the value of K_a the higher the concentration of H⁺ and the stronger the acid

What is the relative strength of conjugate acid base pairs?

• the stronger the acid the weaker its conjugate base

• the weaker the acid the stronger the base

How do you calculate the pH of strong acids + bases?

• if strong then you can assume that it will fully dissociate in water

• the ratio between the acid and the protons will be 1:1

• enter the conc of protons in moles/per litre into the formula

• if base then subtract your answer from 14

What are the limitations of the pH scale?

• limited to the 0-14 range even though values outside of that range are possible in theory

• scale does not work for very concentrated solutions, because when the conc goes above 1M complete dissociation does not always occut

• limited to aqueous solutions

How do you calculate the pH of weak acids and bases

• if weak you will be given K_a or K_b

• you can use the formula to find the concentration of H⁺ or OH- ions

• then sub into pH formula

Acid-base indicator

a substance that changes colour according to the pH of the solution in which it is placed

Explain the colour change of the indicator when placed in an acid or a base:

Acid

• colour change to red

• equilibrium shifts to the right to use up H+ ions added by the indicator

Base

• colour change to blue

• OH^- ions react with H+ ions in the indicator, so equilibrium shifts to the right to produce more H+ ions

Range of an indicator

the pH interval over which there is a clear change of colour for that indicator

Name 3 indicators + their colour change + their ranges

Describe how you would carry out a pH titration

Draw a strong acid + strong base pH curve

What is a suitable indicator for a strong acid strong base titration + why?

• methyl orange

• litmus

• phenolphthalein

an indicator that has one colour at pH 3 and another at pH 10 is needed, and all three of these indicators fall within this range

Draw a strong acid + weak base pH curve

What is a suitable indicator for a strong acid weak base titration + why?

• methyl orange

an indicator that has one colour at pH 3 and another at pH 7 is needed, and methyl orange falls within this range 3-5

Draw a weak acid + strong base pH curve

What is a suitable indicator for a weak acid strong base titration + why?

• phenolphthalein

an indicator that has one colour at pH 7 and another at pH 10 is needed, and methyl orange falls within this rang 8-10

Draw a weak acid + weak base pH curve

Why can’t the end point of a weak acid weak base titration be detected with any base?

there is no sudden pH change - pH changes gradually so there is no sharp colour change