Self Ionisation of Water
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Transcript of Self Ionisation of Water
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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.
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Show how [H+] = 1.0 X 10-7
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• 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)
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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
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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
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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.
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pH of Common Substances
Acidic Neutral Basic
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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)
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The ionic product of water is the product of the hydrogen and hydroxide ion concentration in 1litre of water at 25 °C
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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
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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.
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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
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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.
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Measuring pH• Universal Indicator Paper
• Universal Indicator Solution
• pH meter
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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
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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)
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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
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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
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pH of bases: pOHpOH= -logpOH= -log1010 [OH-] [OH-]
pH + pOH = 14pH + pOH = 14pH= 14 - pOH
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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)
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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
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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 = ?
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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 )
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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++]]
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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
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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
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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
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• Calculate the pH of a 1 molar ethanoic acid solution that is only 1.4% ionised
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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
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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
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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
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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
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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
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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
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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
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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
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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
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Page 261, 262
Question NB to practise