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  • Created on: 20-01-15 15:13

strong and weak

A Bronsted-Lowry acid = a proton donor

Strong acids fully dissociate. 

e.g. HCl, HNO3, first dissociation of H2SO4.

Weak acids partially dissociate.

e.g. H3PO4, CH3COOH, second dissociation of H2SO4.

Ka shows the extent of acid dissociation. The higher the value,the stronger the acid.

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diprotic acids

E.g. H2SO4


H2SO4 + aq >>> H+ + HSO4-

This is a strong dissociation because the H+ fully dissociates.

HSO4- <> H+ + SO4(2-)

This is a WEAK dissociation because the H+ aren't fully dissociated.

note: reversible reaction sign on second dissociation.

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acid base ionic equation

Ionic equations for acid-base reactions are ALWAYS:

H+ + OH- >>> H2O

Kw is the ionic product of water.

The concentration of water is always 55.6moldm-3.

The Kw is always 1x10^-14 mol2 dm-6

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conjugate acid base pairs

When an acid is dissolved, water acts as a base, accepting H+.

HCl+ H2O <> H3O+ + Cl- 

After the reaction, the protonated water acts as the acid, and the ion acts as the base.

HCl = acid 1, H2O= base 2

H3O+ = acid 2, Cl- = base 2

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Why do we use pH?

It is more manageable than H+ concentration which can be extremely big or small numbers.

pH = -log[H+]

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What is a buffer?

A buffer solution is a mixture that minimises small changes in pH on the addition of acids and bases.

What are they made from?

A weak acid, HA, and a conjugate base, A-. e.g a weak acid, and the salt of a weak acid.

The CH3COOH/CH3COO-Na+ system

  • weak acid partially dissociates
  • CH3COOH <> CH3COO- + H+
  • salt completely dissociates
  • CH3COO-Na+>>>> CH3COO- + Na+
  • The high concentration of the base pushes equilibrium to the left, reducing number of H+.
  • Therefore, we have a large resevoir of the weak acid and conjugate base pair.
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Buffer action

Works on the equilibrium of HA <<>> H+ + A-

The weak acid removes alkali, and the conjugate base removes any protons.

Minimise acidic conditions

  • H+ conc is increased
  • the conjugate base reacts with new ions
  • equilibroum shifts to the left, remobing new protons.

Minimising alkaline conditions

  • OH- conc is increased
  • Weak acid reacts with it, to make water
  • The weak acid dissociates more, to make more H+, reducing the number of OH-
  • The equilibrium shifts to the right.
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pH values of Buffer systems

Depends on

  • the acid dissociation constant, Ka, of the buffer system
  • the concentration ratio of the weak acid and its conjugate base

Method 1

Ka = [H+][A-]/[HA]

therefore [H+] = Ka x (acid]/[product ion]

OR H+ = Ka x [CH3COOH]/[CH3COO-]

Method 2- Henderson-Hasselbalch

pH= pKa +log (product ion/ acid)

pH= pKA + log (CH3CCO-/CH3COOH)

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Carbonic acid-hydrogencarbonate buffer system

pH of bloodmust be kept between 7.35 and 7.45. If not acidosis or alkalosis is produced.

The buffer system used - Carbonic acid, H2CO3, and hydrogencarbonate HCO3-. The acid disscoaition constant is 4.3 x 10-7 moldm-3.

H2CO3 <<<>>> H+ + HCO3-

Any increase in acidity

  • H+ are removed by HCO3-.
  • Equilibrium shifts left, removing most of H+

Any increase in OH- ions

  • The H+ reacts with additonal OH-
  • Equilibroium shifts right to produce more H+.
  • All OH- are reacted

Most chemicals that neter the blood are acidic, so a lot of H2CO3 is produced. This is converted into aqueous CO2  by an enzyme. This is then exhaled by the lungs.

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expressions for acids


[HA] >>> [H+]


 [H+] = \/Ka x [HA]

Ka = [H+]^2/ [HA]

PKa = -logKa

Ka is the acid dissociation constant and shows the extent of acid dissociation.

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Neutralisation: strong acid + strong base


Strong acid + strong base

|                   ___________
|                  |
|                  |
|                  |
|                  |

Starts a VERY low pH, sudden incline, and finishes at a VERY high pH

DONT FORGET: vertical line should go up at volume given to you!

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weak acid +strong base

Weak acid + weak base

|                  ____________
|                  |

Starts at semi-low pH (4-6) and finishes at semi-high pH (8-10)

DONT FORGET: vertical line should go up at volume given to you!

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strong acid weak base/ weak acid strong base

Strong acid and weak base

VERY LOW start point (1-5) , and fairly high end point (8-10)

Strong base and weak acid

Very low start point and very high end point.

NOTE: A titration can begin with either an acid or a base, so make sure the graph starts with a pH to match.

Key features of a titration curve

  • When the base is first added, the pH increases very slightly as the acid is in great excess.
  • Within 1-2cm3 of the equivalence point the pH rises sharply. The equivalence point is within the vertical line of the titration curve.
  • As further base sis acid, the pH goes slightly more alkaline, but the base is in excess anyway.

DONT FORGET: vertical line should go up at volume given to you!

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Indicators for titrations


An acid base indicator is often a weak acid, represented as HIn. It is one colour in its acid form (HIn) and another in its cojugate base form (In-)


When there are equal amounts of weak acid and conjugate base, the indicator is at its end point. So it will become a mixture of the two colours, or colourless, suddenly.


pH value of the end point must as close as possible to the pH value of the titrations equivalence point. It must change colour within the pH region of the vertical line of the equivalence point.

For this reason, indicators cant be used for weak acid-weak base titrations, as the equivalence point is so narrow.

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Enthalpy change of neutralisation

The standard enthalpy change of neutralisation

the energy change that accompanies the the neutralisation of an aqueous acid by an aqueous base to form one mole of H2O under standard conditions.


USUAL Q= M x C x At

M= mass of solution (cm3 =g)

C= heat capacity of solution (generally 4.18)

At= temperature change(shoudl alwayt be negative as neutralisation is exothermic.)

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Strong and weak bases

A bronsted Lowry base is a proton acceptor.

A strong base fully dissociates its OH-

e.g. KOH

A weak base partially dissociates its OH-

e.g. H2O <<>> OH- + H+

pH of base 

H+ = kW/[OH-]

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Ionic equations

Acids with metals

2H+ + Mg >>> Mg2+ + H2

Acids and carbonates

2H+ + CO32- >>> H2O + CO2

Acids and hydroxides

2H+ + OH- >>> H2O

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