Self Ionisation of Water

Self Ionisation of WaterWater undergoes Self Ionisation

H2O(l) ⇄ H+(aq) + OH-

(aq)

or

H2O(l) + H2O(l) ⇄ H3O+(aq) + OH-

(aq)

The concentration of H+ ions and OH- ions is extremely small.

Because the equilibrium lies very much on the left hand side.

Show how [H+] = 1.0 X 10-7

• Degree of ionisation is extremely small• Kw = Kc[H2O]= [H+][OH-] = = 1 x 10-14 (at 25ºC)• Kw is the Ionic Product of water/dissociation product

of water• Kw is temperature dependent ( not pressure or

concentration dependent)• Increase temperature will increases the ionic

product ( no effect on pH of water though)• Acidic solution [H+] greater [OH-]• Pure Water is a very very weak electrolyte• ( only 1 in every 600 million water molecules ionise)

Kw is temperature dependentT (°C) Kw (mol2/litre2)

0 0.114 x 10-14

10 0.293 x 10-14

20 0.681 x 10-14

25 1.008 x 10-14

30 1.471 x 10-14

40 2.916 x 10-14

50 5.476 x 10-14

Kw of pure water increases as the temperature increases

The ionic product of water is the product of the hydrogen and hydroxide ion concentration in 1litre of water at 25 °C

Kw = [H+][OH-] = 1 × 10-14 at 25 °C

pH

[H+ ] x [OH- ] = 1 x 10-14 = [1 x 10-7 ] x [1 x 10-7 ]

[H+ ] of water is at 250C is 1 x 10-7 mol/litre

Replacing [H+ ] with pH to indicate acidity of solutions

pH 7 replaces [H+ ] of 1 x 10-7 mol/litre where pH = - Log10 [H+ ]

The ionic product of water is the product of the hydrogen and hydroxide ion concentration in 1litre of water at 25 °C

Kw = [H+][OH-] = 1 × 10-14 at 25 °C

pHAt 250C Kw = 1 x 10-14 mol2/litre2

[H+ ] x [OH- ] = 1 x 10-14 mol2/litre2

This equilibrium constant is very important because it applies to all aqueous solutions - acids, bases, salts, and non-electrolytes - not just to pure water.

pH of Common Substances

Acidic Neutral Basic

The pH Scale• Each pH unit is 10 times as large as the

previous one• A change of 2 pH units means 100 times more

basic or acidic

x10x10 x100x100

9

Limitations

1.Doesn’t cover very HIGH concentration (pH above 10-1) or very low pH values (pH below 10-14)

2.Must be aqueous

3.Affected by temperature ( standard temperature is 25°C)

The ionic product of water is the product of the hydrogen and hydroxide ion concentration in 1litre of water at 25 °C

Acid–Base Concentrations in Solutions

OH-

H+OH-

OH-H+

H+

[H+] = [OH-] [H+] > [OH-] [H+] < [OH-] acidic

solutionneutralsolution

basicsolution

conc

entr

atio

n (m

oles

/L)

10-14

10-7

10-1

pH Scale

Soren Sorensen(1868 - 1939)

The pH scale was invented by the Danish chemist Soren Sorensen to measure the acidity of beer in a brewery. The pH scale measured the concentration of hydrogen ions in solution. The more hydrogen ions, the stronger the acid.

The pH Scale

Neutral Weak Alkali

Strong Alkali

Weak Acid

Strong Acid

7 8 9 10 11 12 133 4 5 62 141 7 8 9 10 11 12 133 4 5 62 141 9 10 11 123 4 5 621

pH ScaleThe quantity of hydrogen ions in solution can affect the color of certain dyes found in nature. These dyes can be used as indicators to test for acids and alkalis. An indicator such as litmus (obtained from lichen) is red in acid. If base is slowly added, the litmus will turn blue when the acid has been neutralized, at about 6-7 on the pH scale. Other indicators will change color at different pH’s. A combination of indicators is used to make a universal indicator.

Measuring pH• Universal Indicator Paper

• Universal Indicator Solution

• pH meter

The pH ScaleThe pH ScalepH scale

[H+] > 10-7M, pH < 7

ACIDIC

[H+] < 10-7M, pH > 7

BASIC

[H+] = 10-7M, pH = 7

NEUTRAL

The larger the hydrogenIon concentration

The smaller the pH,The stronger the acid

The pH Scale• Each pH unit is 10 times as large as the

previous one• A change of 2 pH units means 100 times

more basic or acidic

x10x10 x100x100

17

Limitations

1.Doesn’t cover very HIGH concentration (pH above 10-1) or very low pH values (pH below 10-14)

2.Must be aqueous

3.Affected by temperature ( standard temperature is 25°C)

pH is temperature dependentT (°C) pH

0 7.12

10 7.06

20 7.02

25 7

30 6.99

40 6.97

pH of pure water decreases as the temperature increasesA word of warning!If the pH falls as temperature increases, does this mean that water becomes more acidic at higher temperatures? NO!Remember a solution is acidic if there is an excess of hydrogen ions over hydroxide ions.

In the case of pure water, there are always the same number of hydrogen ions and hydroxide ions. This means that the water is always neutral - even if its pH change

pH & Indicators

pH= 7 at 25° CpH = -Log10 [H+]

Defined as the negative log to the base 10 of the molar Hydrogen ion

concentration in an aqueous solution

pH of bases: pOHpOH= -logpOH= -log1010 [OH-] [OH-]

pH + pOH = 14pH + pOH = 14pH= 14 - pOH

pH ExercisespH Exercisesa)pH of 0.02M HCl pH = – log10 [H+]

= – log10 [0.020]= 1.6989

= 1.70

b)pH of 0.0050M NaOH pOH = – log10 [OH–]

= – log10 [0.0050]= 2.3pH = 14 – pOH= 14 – 2.3

=11.7

c) pH of solution where [H +] is 7.2x10-8M

pH = – log10 [H+]= – log10 [7.2x10-8]= 7.14

(slightly basic)

pH Calculations

pH

pOH

[H+]

[OH-]

pH + pOH = 14

pH = -log10[H+]

[H+] = 10-pH

pOH = -log10[OH-]

[OH-] = 10-pOH

[H+] [OH-] = 1 x10-14

pH of dilute aqueous solutions of acids

monoproticmonoprotic

diproticdiprotic

HA(aq) H1+(aq) + A1-(aq) 0.3 M 0.3 M 0.3 M

pH = - log10 [H+]pH = - log10[0.3M]

pH = 0.52e.g. HCl, HNO3

H2A(aq) 2 H1+(aq) + A2-(aq) 0.3 M 0.6 M 0.3 M

pH = - log10[H+]pH = - log10[0.6M]

pH = 0.22e.g. H2SO4

pH = ?

What is the pH of a 0.1 molar soltion of NaOH (careful)

What is the pH of 0.05 molar solution of Co(OH)2 ( assume its fully dissociated )

Solving for [H+]Solving for [H+]• A solution has a pH of 8.5. What is the A solution has a pH of 8.5. What is the

Molarity of hydrogen ions in the solution?Molarity of hydrogen ions in the solution?

pH = - log [HpH = - log [H++]]

8.5 = - log [H8.5 = - log [H++]]

Strong and Weak Acids/Bases

Strong acids/bases – 100% dissociation into ions

HCl NaOHHNO3 KOHH2SO4

Weak acids/bases – partial dissociation, both ions and molecules

CH3COOH NH3

Need to know equilibrium constant

pH calculations for Weak Acids and Weak Bases

For Weak Acids

pH = -Log10

For Weak Bases

pOH = Log10

pH = 14 - pOH

[H+]= √ka×Macid

[OH-]= √kb×Mbase

pH of solutions of weak concentrations

Weak Base

pH of a 0.2M solution of ammonia with a Kb value of 1.8 x 10-5

pH = 11.2681

• Calculate the pH of a 1 molar ethanoic acid solution that is only 1.4% ionised

Acid base indicators• Substances that change colour according to pH

of solution• Most are weak acids or bases so must only be

added in small amounts. The colour of the dissociated molecule is different to the colour of the undissociated molecule

• Some indicators dissociate to form weak bases• InH=In- + H+

• InOH = In+ + OH-• Chemical equilibrium alters whether in

presence of acid or base

Theory of Acid Base IndicatorsAcid-base titration indicators are quite often weak acids.

For the indicator HInThe equilibrium can be simply expressed as

HIn(aq, colour 1) H+

(aq) + In-(aq, colour 2)

Methyl orange•HIn (red, Acid)= H+ + In- (yellow, Base)•In acid: the equilibrium lies to the ______ giving it a ___ colour•In base: the equilibrium lies to the ______ giving it a ___ colour• : dynamic equilibrium: apply a stress by adding or removing H+ ions will shift the equilibrium•The equilibrium will shift depending on whether H+ ions or OH- ions exist. Therefore causing a colour change

Draw rough trend graphName of Indicator

Approx Range

Acid ColourLower pH

Base ColourHigher pH

Methyl Orange

3.1-4.4red

yellow

Litmus 5-8 red blue

Phenolphthalein

8.3-10 colourless pink

Acid Base Titration CurvesStrong Acid – Strong Base Strong Acid – Weak Base

Weak Acid – Strong Base

25 cm3 of 0.1 mol dm-3 acid is titrated with 0.1 mol dm-3 alkaline solution.

Weak Acid – Weak Base

Choice of Indicator for Titration

• Indicator must have a complete colour change in the steepest part of the pH titration curve

• Indicator must have a distinct colour change

• Indicator must have a sharp colour change

Indicators for Strong Acid Strong Base Titration

Both phenolphthalein

and methyl orange

have a complete

colour change in the

vertical section of the

pH titration curve

Indicators for Strong Acid Weak Base Titration

Only methyl orange

has a complete

colour change in the

vertical section of the

pH titration curve

Phenolphthalein has

not a complete colour

change in the vertical

section on the pH

titration curve.

Methyl Orange is

used as indicator for

this titration

Indicators for Weak Acid Strong Base Titration

Only phenolphthalein

has a complete

colour change in the

vertical section of the

pH titration curve

Methyl has not a

complete colour

change in the vertical

section on the pH

titration curve.

Phenolphthalein is

used as indicator for

this titration

Indicators for Weak Acid Weak Base Titration

Neither phenolphthalein

nor methyl orange have

completely change colour

in the vertical section on

the pH titration curve

No indicator suitable

for this titration

because no vertical

section

Page 261, 262

Question NB to practise