Acids and Bases Part 1 PDF
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University of Southampton
Joern Werner
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These notes cover the fundamentals of acid-base chemistry, including definitions, equilibrium concepts, and dissociation constants. The text explains how biologists need to understand these concepts for several important biological processes.
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Acids and Bases Part 1 Joern Werner Why biologists care about acids and bases? The building blocks of both proteins and DNA/RNA are amino acids and nucleic acids respectively. Phospholipids have acidic and basic headgroups. In many cases the catalytic function of enzymes depends on...
Acids and Bases Part 1 Joern Werner Why biologists care about acids and bases? The building blocks of both proteins and DNA/RNA are amino acids and nucleic acids respectively. Phospholipids have acidic and basic headgroups. In many cases the catalytic function of enzymes depends on pH (i.e. acid/base equilibrium) Protein function may depend on pH (e.g. oxygen binding to hemoglobin) Organism tissues and cells typically maintain the pH in a narrow range Therefore, much understanding acids and bases, pH and buffers is critical for understanding the processes that govern life. Acids and basis: definitions Acids and bases are all about the behavior of H + and OH-. One definition of an acid is an H+ donor: a compound that undergoes a chemical change to produce an H + ion. The H+ ion is equivalent to a proton (recall lectures on atomic structure). Acid HA ⇌ H+ + A - Anion Hydrogen ion (proton) The acid dissociates into ions: the process is an equilibrium process. CH3COOH ⇌ ? Ethanoic acid CH3COOH ⇌ CH3COO- + H+ Acids and basis: definitions A base is an H+ acceptor; a compound that undergoes a chemical change to combine with an H+ ion. B + H+ ⇌ BH+ NH3 + H+ ⇌ NH4+ Brønsted-Lowry definition Ammonia Ammonium Acid = H+ donor Base = H+ acceptor In water. Hydronium ion HCl dissociates in water: HCl ⇌ H+ + Cl- In H2O H+ ions do not exist independently. They combine with water: H+ + H2O ⇌ H3O+ So then the overall dissociation reaction is: HCl + H2O ⇌ H3O+ + Cl- But we usually focus on the acid and do not bother to include the water: HCl ⇌ H+ + Cl- Acids and basis: definitions. In water. When a base dissolves in water it accepts a hydrogen ion from a water molecule B + H2O ⇌ BH+ + OH- Hydroxide ion After its reaction with a base, the water molecule becomes a hydroxide ion, OH -. Self test exercise: In each case, (a) and (b) is the molecule on the (a) HBr ⇌ H+ + Br- left-hand side of the equilibrium sign acting as (b) CN - + H+ ⇌ HCN an acid or a base? Pairing up acids and bases: conjugate pairs Acids and bases operate as a pair: an acid must have a base to which it donates a H+. Forward reaction: Back reaction: NH3 + H2O ⇌ NH4+ + OH- NH4+ + OH- ⇌ NH3 + H2O Water acts as Water acts as a an acid. base Ammonia acts Ammonium ion as a base: acts as an acid General case Extent of a dissociation reaction HA + H2O ⇌ H3O+ + A- This is an equilibrium. It will reach a point at which there will be no overall change in the composition as forward and back reactions will occur to same extent. In aqueous solutions this happens almost instantaneously. The extent to which an acid dissociates (the position of the equilibrium) is called its strength. The readiness with which an acid donates H + to generate H3O+ is called the strength of the acid. The readiness with which a base accepts H + from water to generate OH - is called the strength of the base, will go to base H2O ⇌ H+ + OH- CH3COOH + H2O ⇌ CH3COO- + H3O+ Weak acid HNO3 + H2O ⇌ NO3- + H3O+ Strong acid Extent of a dissociation reaction A strong acid readily dissociates to form its conjugate base when dissolved in water, releasing a H+ ion in the process. A weak acid does not dissociate to a very great extent. A strong base readily accepts a H+ ion from solution. A weak base does not easily accept a H+ ion. CH3COOH + H2O ⇌ CH3COO- + H3O+ Weak acids have strong conjugate bases Acid is weak - will not Conjugate base is Strong acids have weak conjugate bases easily give up H+ ion strong as very easily accepts H+ The acid dissociation constant The reactions we have been looking at are equilibrium reactions: HA + H2O ⇌ A- + H3O+ Recall that the equilibrium constant, Kc is: [H3O+][A-] [Products] Kc = [HA] [H2O] [Reactants] Consider the dissociation of acid in a dilute aqueous solution. The concentration of water will be large (55.5 M) and in a very good approximation will stay the same. The acid dissociation constant Consider the dissociation of acid in a dilute aqueous solution. The concentration of water will be large (55.5 M) and will not change. We can remove [H2O] from the right-hand side of the equation by multiplying both sides of the equation by this: [H3O+][A-] Replace by [H3O+][A-] [H3O+] [A-] Kc = Kc x [H2O] = x [H2O] Kc x [H2O] = [HA] [H2O] [HA] [H2O] [HA] We can then define something called the acid dissociation constant, Ka, which is just the left-hand side of the equation. So: [H3O+][A-] Define Ka = Kc x [H2O] then Ka= [HA] Interpretation of the acid dissociation constant [H3O+] [A-] HA + H2O ⇌ A- + H3O+ Ka= [HA] What does the magnitude of Ka tell us? It is equivalent to the equilibrium constant – so provides the same information: the relative position of the reaction at equilibrium. (recall lecture on equilibria) If Ka is large, this tells us that numerator is large, and the denominator is small. So the reaction lies to the right at equilibrium. This means that the acid is strong as it is mostly dissociated (so many H 3O+ ions are produced). If Ka is small the reaction lies to the left at equilibrium and the acid is weak (relatively little is dissociated). Definition and interpretation of the base dissociation constant We can define a base dissociation constant, Kb, in a similar way [BH+] [OH-] B + H2O ⇌ BH+ + OH- Kb = where Kb = Kc x [H2O] [B] A large value for Kb tells us that the reaction lies to the right at equilibrium. The base is strong. A small value for Kb tells us that the reaction lies to the left at equilibrium. The base is weak. Summary Acid and base dissociation reactions are equilibrium reactions. We can define a Ka and Kb which are equivalent to the equilibrium constants for the reactions. A large value of Ka tells us that the acid readily donates a proton. The reaction below lies to the right at equilibrium. The acid is strong HA + H2O ⇌ A- + H3O+ A small value of Ka indicates a weak acid A large value of Kb tells us that the base readily accepts a proton. The reaction below lies to the right at equilibrium. The base is strong B + H2O ⇌ BH+ + OH- A small value of Kb indicates a weak base Introducing pKa and pKb The value of Ka can vary over a large range (many orders of magnitude). So to make it easier to work with the quantities it makes sense to use a logarithmic scale (brings numbers down to a similar order of magnitude). We define a new parameter, pKa: pKa = -log(Ka) Let’s see how it works: 5.5 and 3.4 are easier to For example a weak acid has a Ka value of 3.0 x 10-6 work with than 3.0 x 10-6 and 3.8 x 10-4 respectively The pKa value is –log(3x10-6) = +5.5 log(0.01)=-2 We can do the same for bases: pKb = -log(Kb) log(0.1)=-1 log(1)=0 For example, a weak base has a Kb value of 3.8 x 10-4 log(10)=1 The pKa value is –log(3.8x10-4) = +3.4 etc Ka, pKa and the strength of acids Weak acid Formula Ka pKa Lactic acid CH3CHOHCOOH 1.4 x 10 -4 3.85 Ethanoic acid CH3COOH 1.8 x 10 -5 4.74 Carbonic acid H2CO3 4.2 x 10 -7 6.38 Ammonium ion NH4+ 5.6 x 10 -10 9.25 The ion product of water: Kw We already know that water acts as an acid and a base. Conjugate acid-base pair H2O + H2O ⇌ H3O+ + OH- - Is water a strong acid or a strong base? Base Acid Acid Base - Where does the reaction lie at equilibrium? Conjugate acid-base pair [H3O+] [OH-] Kc = or [H3O+] [OH-] = [H2O]2 Kc = Kw [H2O]2 It turns out the reaction lies almost totally to the left, so very little water dissociates. Hence, we define a special term, Kw , called the ion product of water. KW = [H3O+] [OH-] = 1 x 10 -14 mol2 dm-6 at 25℃ (Note Kw is a product of two concentrations and hence has the units of mol dm -3 * mol dm-3 ) So how do we use the ion product of water? Ion product of water is surprisingly useful: it shows the relationship between the concentration of H3O+ or OH- in any aqueous solution. The expression KW = [H3O+] [OH-] = 1 x 10 -14 mol2 dm-6 is saying ‘the concentration of hydronium ions [H3O+] multiplied by the concentration of hydroxide ions [OH -] has a value of 1 x 10 -14 mol2 dm-6. So, if we know either the concentration of OH - or H3O+ in a particular solution, we can use the expression to calculate the concentration of the other species. Using Kw: A worked example If we have a solution that contains hydronium ions at a concentration of 1 x 10-3 M. The ion product of water, Kw, tells us that [H3O+] [OH-] = 1 x 10 -14 mol2 dm-6 We already know that [H3O+] = 1 x 10-3 M = 1 x 10-3 mol dm-3. So how do we work out [OH-] ? Let’s rearrange the equation for Kw to get [OH-] by itself. [H3O+] [OH-] = 1 x 10 -14 mol2 dm-6 1 x 10 -14 mol2 dm-6 [OH-] = Isolate [OH-] [H3O+] So [OH-] = 1 x 10 -11 M 1 x 10 -14 mol2 dm-6 = = 1 x 10 -11 mol dm-3 1 x 10-3 mol dm-3 More generally Ka , Kb and Kw are related We just saw that Kw = 1 x 10 -14 mol2 dm-6 It turns out that Kw = Ka x Kb So, if we know the value of Kb for a weak base, we can calculate Ka for its conjugate acid, and similarly, if we know Ka for a weak acid, we can calculate Kb for its conjugate base. What magic is this? Let’s take a look! Let’s remind ourselves of the reaction of ammonia with water Ka , Kb and Kw are related (2) In the forward reaction, ammonia is acting [NH4+] [OH-] NH3 + H2O ⇌ NH4+ + OH- Kb = as a base, so we can write a base [NH3] dissociation constant: In the back reaction, ammonium ion is [NH3] [H3O+] NH4+ + H2O ⇌ NH3 + H3O+ acting as an acid, so we can write an acid Ka = dissociation constant: [NH4+] [NH3] [H3O+] [NH4+] [OH-] Ka x Kb = x = [H3O+] [OH-] = Kw [NH4+] [NH3] Summary Ka can have a very wide range of values. To keep comparisons manageable, we use pKa values instead pKa = -log(Ka). Ka is large for strong acids, pKa is small. We use Ka and pKa values when we refer to the strength of an acid or base. This is because we know that the relationship between conjugate acids and bases: strong acid – weak conjugate base and vice versa. So, we only need to know the value of Ka and pKa to in order to predict the strength of the base too. Ion product of water is defined as: KW = [H3O+] [OH-] We saw that Ka x Kb = Kw Hence, if we know the Ka of a weak acid, we can work out the Kb of its conjugate base and vice versa.