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Acids & Bases, pH & Buffers

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Edmund Fenton Fowler

on 21 May 2013

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Transcript of Acids & Bases, pH & Buffers

Acids & Bases A Bronsted-Lowry acid is a proton, H+, donor.
A Bronsted-Lowry base is a proton, H+, acceptor. Compound A proton Acid What you need to know Base Compound A proton Some acids Sodium Hydroxide, Na(OH)2
Ammonium Hydroxide, NH4OH
Potassium hydroxide, KOH Hydrochloric acid, HCl
Nitric acid, HNO3
Sulphuric acid H2SO4
Phosphoric acid, H3PO4 Some alkalis An alkali is a type of base that dissolves in water to form hydroxide , OH-(aq), ions Reactions with carbonates Reactions with Bases Reactions with metals Reactions with alkalis Acid-Base Reactions Aqueous acids react with solid carbonates to form a salt, carbon dioxide gas & water. 2HCl(aq)+CaCO3(s)>>CaCl2(aq)+CO2(g)+H2O(l) Aqueous acids react with bases to form a salt and water. 2HNO3(aq)+MgO(s)>>Mg(NO3)2(aq)+H2O(l) Aqueous acids react with alkalis to form a salt & water. H2SO4(aq)+2KOH(aq)>>K2SO4(aq)+H2O(l) Aqueous acids react with metals to form a salt & hydrogen gas. This isn't actually an acid-base reaction; in fact it is a REDOX reaction. 2HCl(aq)+Mg(s)>>MgCl2(aq)+H2(g) Acid + Metal >> Salt + Hydrogen Acid + Alkali >> Salt + Water Acid + Base >> Salt + Water Acid + Carbonate >> Salt + Carbon Dioxide + Water Mono-, di- and tri-basic acids Mono-basic acids Tri-basic acids Di-basic acids Different types of acids can release different amounts of protons depending on their formula. Acids that can release more than one proton do so in more than one step. If an acid hasn't released all of its protons then the reaction forms an equilibrium. A mono-basic acid can release only one proton. A di-basic acid can release two protons. A tri-basic acid can release three protons. HCl(aq)>>H+(aq)+Cl-(aq) H2SO4(aq)><H+(aq)+HSO4 -(aq)
HSO4 -(aq)><H+(aq)+SO4 2-(aq) H3PO4(aq)><H+(aq)+H2PO4 -(aq)
H2PO4 -(aq)><H+(aq)+HPO4 2-(aq)
HPO4 2-(aq)>>H+(aq)+PO4 3-(aq) In all of these reactions the aqueous acid is neutralised:

Neutralisation is a chemical reaction in which an acid and a base react together to produce a salt and water. Acid-Base Pairs A acid-base pair is a set of two species that transform into each other by gain or loss of a proton. Acid + Base >< Conjugate acid + Conjugate Base HNO2 + H2O >< H3O+ + NO2- H2O + NH3 >< HO- + NH4+ Here are examples of water acting as a base and an acid: CH3COOH + H2O >< H3O+ + CH3COO- The best way to understand what conjugate means is to just look at the equation in the reverse direction. The conjugate acid will donate a proton in the reverse direction and the conjugate base will accept it. The acid dissociation constant Strong and weak acids & bases A strong acid or base is one that completely dissociates in solution.
A weak acid or base is one that only partially dissociates in solution. For a weak acid: HA(aq)><H+(aq) + A-(aq) An equilibrium is established for weak species and its position lies well to the left as there are only small concentrations of dissociated ions. K a [H+][A-] K = a [HA] Units: mol dm-3 A large value for dissociation constant means a lot of dissociation, hence it describes a strong acid.
A small value for dissociation constant means not much dissociation, hence it describes a weak acid. K a p K a p K a p = -log(K ) a K = 10 a - pH What you need to know What is pH? The Ionic Product of Water, Kw Calculating the pH value of Strong Bases It is a measure of hydrogen ion concentration.
pOH also exists & it is a measure of hydroxide ion concentration. pH=-log[H+(aq)] [H+(aq)]=10 -pH Remember that a low pH means a high Hydrogen ion concentration and a low Hydroxide ion concentration. For a strong acid: [H+(aq)] = [HA(aq)] HA(aq)>H+(aq) + A-(aq) [H+(aq)] = (Ka x [HA(aq)])^0.5 Kw=[H+(aq)][OH-(aq)] At 25 degrees, its value is 1.00 x 10 mol dm -14 2 6 The reason the constant is useful is that it allows us to convert between Hydrogen ion concentration and Hydroxide ion concentration. N.B. Do not confuse the terms strong & weak with concentrated or dilute.
Strong & weak represent the dissociation of the acid whereas,
Concentrated & dilute mean how much of the acid there is in the solution. There are 2 methods to work this out- Using Kw Using pOH First find [H+(aq)] from [OH-(aq)] using Kw:
Kw = [H+(aq)] [OH-(aq)] =
:. [H+(aq)] = Kw / [OH-(aq)] = =

Then calculate the pH:
pH = -log[H+(aq)] = -log( ) = 12.70 1.00 x 10 mol dm -14 2 6 1.00 x 10 -14 0.050 2.00 x 10 mol dm -13 2 6 2.00 x 10 -13 E.g. A solution of KOH has a concentration of 0.050 M. What is its pH? Useful background knowledge:
The definition of pOH is:
pH + pOH = 14 Find pOH:
pOH = -log[OH-(aq)] = -log(0.050) = 1.30

Then calculate pH:
pH = 14 - 1.30 = 12.70 Buffers & Neutralisation What you need to learn Buffer Solution ... is a mixture that minimises pH changes on addition of small amounts of an acid or base. We primarily focus on acidic buffers. These are buffer solutions with a pH < 7. They consist of a weak acid, HA, and its conjugate base, A-. What happens when an acid is added to the buffer solution? What happens when an alkali is added to the buffer solution? The equilibrium in the buffer solution Learn this reason! How to calculate the pH of a buffer solution They could also ask to do the same calculation the other way around. For a buffer consisting of a weak acid, HA, and its conjugate base, A-: Dissociation of weak acid
-Negligible amount of ethanoate ions compared to the amount coming from sodium ethanoate with fully dissociates. pH= -log(8.7x10^-6)
=5.060 -0.10 mol dm^-3 ethanoic acid
-0.20 mol dm^-3 sodium ethanoate Dissociation of salt:
CH3COO-Na+>>CH3COO-+Na+ [CH3COO-]=
[CH3COO-Na+] Now a quick example of how to do this by video Left to learn:
-Carbonic acid/hydrogencarbonate buffer system
-Titration curves
-Enthalpy changes
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